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"; Write The Procedure- Determination Of Ka For A Weak Acid - Chem Homework Help
write the Procedure- Determination of Ka for a Weak Acid

Determination of Ka for a Weak Acid

Determination of Ka for a
Weak Acid
Hands-On Labs, Inc.
Version 42-0151-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will learn about pH titration curves and the
interactions that occur between a weak acid and a
strong base. You will also learn the equations that
allow the Ka to be determined from a titration curve
and will then perform a pH titration curve to find the
Ka for an unknown weak acid. You will identify the
weak acid by comparing the experimental Ka to a
variety of known weak acid Ka values.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Identify the defining characteristics of acids and bases.

● Define acidity, and explain how the pH scale relates to the acidity of a substance.

● Distinguish between strong and weak acids in relationship to the acid ionization constant (Ka).

● Discuss the purpose of a pH titration curve, and define the half-equivalence point and
equivalence point.

● Perform pH analysis utilizing pH indicator strips.

● Perform a titration of a weak acid against a strong base.

● Construct an accurate titration curve using experimental results.

● Determine Ka for an unknown weak acid utilizing the data collected during the experiment.

Time Allocation: 2.5 hours

www.HOLscience.com 2 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

Materials
Student Supplied Materials

Quantity Item Description
1 Bottle of distilled water
1 Dish soap
1 Drinking glass or cup
1 Pair of scissors
1 Roll of paper towels
1 Sheet of white paper
1 Source of tap water

HOL Supplied Materials

Quantity Item Description
1 Glass beaker, 100 mL
1 Graduated cylinder, 10 mL
1 Pair of gloves
1 Pair of safety goggles
1 pH test strips, wide range
1 Experiment Bag: Determination of Ka for a Weak Acid

1 – Phenolphthalein solution, 1% – 0.5 mL in pipet
1 – Pipet, empty short stem
1 – Sodium hydroxide, 1 M – 30 mL in dropper bottle
1 – Unknown #108 – 25 mL in dropper bottle

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

www.HOLscience.com 3 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

Background
Acids and Bases

Scientists categorize substances as either acids or bases depending on their properties and the
way they interact with other chemicals. Classification of an acid or a base is useful when predicting
how a chemical will react with another chemical. For example, researchers must consider acid-
base reactions when designing antacids to help treat heartburn or when creating cleaning
products. Because acids are proton donors and bases are proton acceptors, the interaction of an
acid and a base will result in a chemical reaction that is called neutralization.

The generalized definition of an acid is as a substance that has a sour taste and will react with a
base. Strong acids are corrosive to metals and will cause severe burns on contact with skin. While
strong acids, such as concentrated sulfuric acid, can be very dangerous, other acids, such as citric
acid, are weak and are found in everyday food items such as lemons, limes, and oranges. See Figure
1. Likewise, the generalized definition of a base is a substance that tastes bitter, feels slippery, and
interacts with acids. The tendency of bases to taste bitter is believed to be an evolutionary trait
to prevent early humans from eating toxic plants. Bases feel slippery to the skin because they
degrade the fatty acids and oils in skin on contact. Bases can also be considered strong (and very
dangerous) or weak. An example of a strong base is pure bleach, while an example of a weak base
is soap which feels slippery on the skin. Chemically, acids react with water to result in an increased
concentration of H+ ions, while bases react with water to result in an increased concentration of
OH- ions. Do not personally test an acid’s sour taste, or a base’s bitter taste and slippery feel.

Figure 1. Acids. A. Sulfuric Acid (H2SO4) is a strong acid © WH Chow. B. Citric acid, found in oranges,
lemons, and limes is a weak acid. © Nattika.

www.HOLscience.com 4 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

The bubbling volcano that one
typically sees at a student science

fair is the result of a chemical
reaction between an acid and a

base. The acid in a volcano is vinegar
(acetic acid) and the base in the
reaction is baking soda (sodium

bicarbonate).

pH Scale

Scientists use the pH scale as a convenient way to measure the acidity of a solution. See Figure
2. Acidity is a measure of the amount of dissolved hydrogen ions, H+, in a solution. The greater
the number of hydrogen ions present in a solution, the more acidic the solution. For example, a
solution with a pH of 1 has more hydrogen ions and is more acidic than a solution with a pH of 6.
The pH scale is a logarithmic system and is calculated using the following formula:

A reading of 0 to less than 7 is acidic, over 7 to 14 is basic, and 7 is neutral. A neutral substance is
a substance that is neither an acid nor a base. Because the pH scale is logarithmic, each number
on scale has a 10-fold difference in acidity compared to the next number. For instance, a pH of 6
is 10 times more acidic than a pH of 7.

Figure 2. pH scale

Bases can be measured by the presence of an ion called a hydroxide ion (abbreviated as OH–). The
pH value of a base increases as the amount of hydroxide ions in the solution increases.

A pH indicator changes color at a specific pH, allowing scientists to qualitatively measure the pH
of a substance. Each indicator changes color within well-defined pH ranges, allowing researchers
to select a specific indicator for a specific pH range. As the pH of the solution changes within the
specific range, so does the color of the indicator. For example, below a pH of 8.2, the indicator
phenolphthalein is colorless. At a pH of 8.2, the phenolphthalein indicator turns pink and will
darken to a dark fuchsia as the pH reaches 10.

www.HOLscience.com 5 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

Ionization Constants

In addition to the strength (or weakness) of an acid or base being described by pH, the strength
(or weakness) of acids and bases is further quantitatively defined by their acid or base dissociation
(ionization) constant. An acid ionization constant (Ka) is the equilibrium constant, describing the
interaction between water and an acid to form a hydronium ion (H3O

+) and a conjugate base. A
base ionization constant (Kb) is the equilibrium constant, describing the interactions between
water and a base, which form a conjugate acid and a hydroxide ion (OH-). See Figure 3. The larger
the value of Ka, the stronger the acid, and the smaller the value of Ka, the weaker the acid. See
Table 1.

Table 1. Relationship between Ka and acid strength.

Acid Ka Value
HSO4

– 1.2 x 10-2 Stronger Acidity
HC2H2ClO2 1.35 x 10

-3

HF 7.2 x 10-4

HC2H3O2 1.8 x 10
-5

HOCl 3.5 x 10-8

NH4
+ 5.6 x 10-10 Weaker Acidity

A strong acid completely dissociates in water, forming a conjugate base that is a weaker base
than water, while a weak acid only partially dissociates in water, forming a conjugate base that
is a stronger base than water. While the acid ionization constant (Ka) of an acid can be calculated
using the concentrations of the acid, base, hydronium ion (conjugate acid), and conjugate base,
as shown in Figure 3, the Ka can also be calculated from a pH titration curve.

Figure 3. Acid and Base Dissociation Constants. An example reaction for Ka is shown in red and an
example reaction for Kb is shown in blue.

www.HOLscience.com 6 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

Equivalence Point

A pH titration curve plots the change in pH that occurs as a strong base is titrated into a weak
acid. A weak acid will act as a buffer to the addition of each drop of the strong base, maintaining a
primarily constant pH until the equivalence point occurs. The equivalence point is the moment in
the titration where exactly enough strong base (titrant) has been added to completely react with
the weak acid (analyte). The equivalence point is identified by a dramatic increase in pH, resulting
in a sharp vertical increase in the titration curve. Following the equivalence point, the titration
curve will maintain the pH of the strong base, as all of the weak acid in the titration has reacted
with the strong base. See Figure 4.

Figure 4. pH titration curve.

At the point halfway to the equivalence point is the half-equivalence point. The half-equivalence
point of a pH titration curve is the moment when the exact volume of strong base has been titrated
into the weak acid, to convert half of the weak acid into its conjugate base. More specifically, the
concentrations of weak acid and conjugate base in the reaction are equal. Thus, when referring
back to the equations in Figure 3, at the half-equivalence point [A-] = [HA], so Ka = [H3O

+]. When
the equation is further derived:

www.HOLscience.com 7 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

To determine the half-equivalence point of a pH titration curve, one would use a pH meter to
measure and record the change in pH that occurs as each drop of strong base is titrated into the
weak acid. A typical pH meter is able to monitor the precise pH of a solution to the nearest tenth,
hundredth, or thousandth value. See Figure 5.

Figure 5. pH Meter. A typical laboratory pH meter is shown. © Yenyu Shih

As a pH meter is priced above a reasonable cost for one-time student use, it is not feasible to use
for this experiment. In the first portion of this exercise, the Ka of an unknown weak acid will be
determined through the use of a pH indicator. You will identify the equivalence point of a titration
between the weak acid and the strong base NaOH. Then you will use the pH indicator strips to
determine the pH at intervals throughout the titration, including the equivalence point and the
half-equivalence point. You will use this data to create a pH titration curve and to determine the
Ka of the unknown weak acid. Finally, you will use the Ka value to identify the weak acid from a
list of known Ka values.

www.HOLscience.com 8 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

Exercise 1: Determination of Ka
In this experiment you will determine the volume of strong base (NaOH) required to reach the
equivalent point in a titration with an unknown weak acid.

Procedure

Completely read all instructions and assemble all equipment and supplies before beginning
work on this experiment.

Note: It is extremely important to deliver drops from your pipet accurately. Thus, you MUST hold
the pipet vertically and squeeze the pipet bulb consistently to make sure you deliver the same size
drop each time.

Part 1: Determination of Equivalence Point

1. Gather the empty short-stem pipet, glass or cup, and distilled water.

2. Pour some distilled water into the glass or cup.

3. Fill the short-stem pipet with water by placing the tip of the pipet in the water, pressing down
on the bulb, and then releasing.

4. Practice releasing single drops of water from the pipet until you are able to consistently
release evenly-sized drops of water.

5. When you are able to fill the pipet and also release even-sized drops of water, empty the pipet
of all water, filling and releasing air from the pipet to ensure all water is released.

6. Put on the safety glasses and a pair of safety gloves.

7. Measure 4 mL of the unknown weak acid from the dropper bottle into the graduated cylinder.
It is important to measure exactly 4 mL of the acid.

8. Carefully pour the 4 mL of unknown weak acid from the graduated cylinder into a clean, dry,
100-mL glass beaker.

9. Place the beaker onto a clean sheet of white paper.

10. Use scissors to carefully cut the tip of the phenolphthalein pipet and add exactly 2 drops of
phenolphthalein solution to the unknown weak acid in the beaker.

11. Carefully swirl the beaker to completely mix the unknown weak acid and phenolphthalein
solution. The solution will be colorless and clear.

12. Gather the bottle of Sodium Hydroxide (NaOH) from the kit.

a. Carefully remove the cap from the bottle and remove the dropper insert from the top
of the bottle.

b. Gently bend the dropper tip to the side and move back and forth to separate the tip
from the bottle.

www.HOLscience.com 9 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

c. Place the tip to side of the bottle, leaving the bottle of NaOH open at the top. See
Figure 6.

Figure 6. Sodium hydroxide with dropper tip removed from bottle.

13. Place the tip of the empty short stem pipet into the Sodium Hydroxide bottle and fill the pipet
with the NaOH.

Note: You will need to count the total number of drops of NaOH added to the beaker. Read all steps
before moving on.

14. Add NaOH from the pipet into the beaker containing the 4 mL of unknown weak acid 1 drop
at a time, swirling and observing the solution in the beaker after each drop until the color
changes to a pale-pink color for at least 5 seconds. Count each drop of NaOH added to the
beaker, so that you will know exactly how many drops of NaOH were required to cause the
solution to change from clear and colorless to pale-pink in color.

Note: It is important to stop adding NaOH at a pale-pink color and before a dark-pink color forms.

15. Record the number of NaOH drops added to the beaker in Data Table 1 of your Lab Report
Assistant.

Note: The data in Data Table 1 will be used to help you set up the titration in Part 2.

16. Carefully pour the solution down the sink drain and use the dish soap, tap water, and paper
towels to wash and dry the beaker.

17. Repeat steps 7 through 16 an additional time to collect data for Trial 2.

18. Average the total number of drops required to reach the equivalence point from Trial 1 and
Trial 2 and record it in Data Table 1.

www.HOLscience.com 10 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

Part 2: Construction of pH Titration Curve

19. Measure 4 mL of the unknown weak acid from the dropper bottle into the graduated cylinder.
It is important to measure exactly 4 mL of the acid.

20. Carefully pour the 4 mL of unknown weak acid from the graduated cylinder into a clean, dry,
100-mL glass beaker.

21. Gather a pH indicator strip from the kit and place it into the unknown weak acid in the beaker
so that all 4 indicator squares come into contact with the weak acid. See Figure 7.

Figure 7. pH indicator strip in beaker. Notice that in either direction, all 4 indicator squares are
immersed in the solution. A. Strip facing down. B. Strip facing up.

22. As soon as all 4 squares are in contact with the solution, remove the strip from the beaker.
Place the strip next to the provided pH scale. Determine the pH of the solution, to the closest
0.5 pH, and record it in Data Table 2 in your Lab Report Assistant next to “NaOH Drops Added
– 0” in the “Trial 1” column. See Figure 8.

23. To accurately read the pH indicator strip, place the strip between the two closest pH indicator
color matches and determine if the strip is closest in color to a whole number value or in
between two color matches, providing a 0.5 value. See Figure 8.

Figure 8. pH scale reading. A. pH value for this indicator strip would be 4.5. Note that the
bottom indicator square is closer in color to a pH of 5.0, while the indicator square 2nd from
the bottom is closer in color to a pH of 4.0. B. pH value for this indicator strip would be 4.0.

When the strip is between the pH 3.0 and pH 4.0, the colors of the indicator strip align almost
perfectly with the indicator colors of pH 4.0.

www.HOLscience.com 11 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

24. From Data Table 1, copy the average number of drops required to reach the equivalent point
and record it in Data Table 2 in the space provided next to “Equivalent Point = ( Drops).”

25. Divide the number of drops required to reach the equivalent point by 2 and record the value
in Data Table 2 in the space provided next to “Half-Equivalent Point = ( Drops).” If this
value ends in 0.5, record down to the closest value. (Example: 41.5 drops would round to 41.0
drops.)

26. As done in Part 1, place the tip of the empty short-stem pipet into the sodium hydroxide
bottle and fill the pipet with the NaOH.

Note: You will need to count the total number of drops of NaOH added to the beaker. Read all steps
before moving on.

27. Add exactly 10 drops of NaOH from the pipet into the beaker containing the 4 mL of unknown.

28. Swirl the beaker to fully mix the NaOH with the unknown weak acid. After swirling, wait 10
seconds before moving on to the next step.

29. Gather a pH indicator strip from the kit and place it into the solution in the beaker so that all
4 indicator squares come into contact with the weak acid. See Figure 7.

30. As soon as all 4 squares are in contact with the solution, remove the strip from the beaker.
Place the strip next to the provided pH scale. Determine the pH of the solution, to the closest
0.5 pH, and record it in Data Table 2 next to “NaOH Drops Added – 10.” See Figure 8.

31. To accurately read the pH indicator strip, place the strip between the 2 closest pH indicator
color matches and determine if the strip is closest in color to a whole number value, or in
between 2 color matches, providing a 0.5 value. See Figure 8.

32. Repeat steps 26 through 31 for the remaining 13 titration steps in Data Table 2. It is important
to collect the pH at exactly the number of drops required to reach the half-equivalence point
and the equivalence point.

Note: To gather and record the values for Data Table 2, while not missing a collection point, consider
the following example: If the half-equivalence point was 43.0 drops, you would collect the pH value
at 40 drops. Add 3 additional drops of NaOH, swirl the beaker, wait 10 seconds, and collect the pH
value at 43 drops. Add 7 additional drops of NaOH, swirl the beaker, wait 10 seconds, and collect the
pH value at 50 drops.

33. When all pH values (up to 120 drops of NaOH) have been determined and recorded in Data
Table 2, carefully pour the solution down the sink drain and use the dish soap, tap water, and
paper towels to wash and dry the beaker.

34. Repeat steps 19 through 33 for Trial 2. Record all values in Data Table 2 in the “Trial 2” column.

35. Average the pH values in Data Table 2 for Trials 1 and 2 and record in pH value (Average) in
Data Table 2.

www.HOLscience.com 12 ©Hands-On Labs, Inc.

Experiment Determination of Ka for a Weak Acid

36. The average pH at half-equivalence point is equal to the pKa of the unknown weak acid. Record
this value in Data Table 3 in your Lab Report Assistant next to pKa.

37. Determine the Ka of the unknown weak acid using the following series of equations:

38. Record the Ka of the unknown weak acid in Data Table 3.

39. Use the values for Ka and pKa for a variety of acid in Table 2 to determine the identity of the
unknown weak acid. Record the identity of the unknown weak acid in Data Table 3.

Table 2. Ka and pKa Values of Acids

Acid Name Acid Formula Ka Value pKa Value
Chlorous acid HClO2 1.2 x 10

-2 2.92
Monochloroacetic acid HC2H2ClO2 1.35 x 10

-3 3.87
Hydrofluoric acid HF 7.2 x10-4 4.14

Acetic acid HC2H3O2 1.8 x 10
-5 4.75

Carbonic acid H2CO3 4.3 x 10
-7 6.37

Hypochlorous acid HOCl 3.5 x 10-8 7.46
Hydrocyanic acid HCN 6.2 x 10-10 9.21

40. Calculate the percent error of the Ka between the experiment value of Ka (determined in the
experiment) and the known value of Ka (provided in Table 2) and record in Data Table 3. Use
the equation below to calculate percent error:

41. Calculate the percent error of the pKa between the experiment value of pKa (determined in
the experiment) and the known value of pKa (provided in Table 2) and record it in Data Table
3.

42. When you are finished uploading photos and data into your Lab Report Assistant, save and
zip your file to send to your instructor. Refer to the appendix entitled “Saving Correctly,” and
the appendix entitled “Zipping Files,” for guidance with saving the Lab Report Assistant in the
correct format.

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Experiment Determination of Ka for a Weak Acid

Cleanup:

43. Properly dispose of used chemical and pipets.

44. Clean all items with soap and water, and return to the lab kit for future use.

Questions
A. Using the data collected in Data Table 2, create a graph of “Drops of NaOH added” vs “pH”

where “drops of NaOH added” is on the x-axis and “pH” is on the y-axis.

B. pH indicator strips work by changing color in the presence of solutions with varied pH values.
Thinking about your procedure steps and results in Part 1, why do you think the pH of the
unknown weak acid was not determined with pH indicator strips until Part 2?

C. Discuss possible causes of error in the experimental procedure. Why do you think there was
a percent of error in the pKa and Ka values of the unknown weak acid, in comparison to the
values presented in Table 2?

D. Do you think using a pH meter instead of pH indicator strips would have created a larger or
smaller percent error? Explain your answer.

E. Why was phenolphthalein a good indicator to use for determining the equivalence point
between the unknown weak acid and strong base?

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Experiment Determination of Ka for a Weak Acid

Equilibrium and Le
Châtelier’s Principle
Hands-On Labs, Inc.
Version 42-0166-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in which
to complete the exercise.

Experiment Summary:

In this experiment, you will describe the components
of a reaction at chemical equilibrium and use Le
Châtelier’s principle to predict the direction a chemical
system will shift upon changes in concentration,
temperature, and pressure. You will perform
equilibrium reactions and test Le Châtelier’s principle
by manipulating concentration and temperature of
the reactions.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Define chemical equilibrium and chemical system.

● Define equilibrium constant and reaction quotient.

● Explain how changes in temperature, pressure, and concentration affect the chemical system
of a reaction.

● State Le Châtelier’s Principle.

● Perform chemical equilibrium reactions and manipulate chemical systems through concentration
and temperature.

● Perform calculations to determine the equilibrium constant (K) and reaction quotient (Q) of
reactions.

● Apply Le Châtelier’s principle to predict changes and explain observed changes in a chemical
system.

Time Allocation: 2.5 hours

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Experiment Equilibrium and Le Châtelier’s Principle

Materials
Student Supplied Materials

Quantity Item Description
1 Dish soap
1 Hot water bath (hot water, cup)
1 Ice bath (ice, water, cup)
1 Pair of scissors, 4 in
1 Roll of paper towels
1 Source of tap water
1 Tape: clear, duct, or masking

HOL Supplied Materials

Quantity Item Description
2 Pairs of gloves
1 Pair of safety goggles
1 Test tube cleaning brush
1 Well plate – 24
1 Experiment Bag: Equilibrium and Le Châtelier’s Principle:

1 – Potassium chromate (K2CrO4), 1 M, 2 mL in pipet
1 – Potassium ferrocyanide (K4Fe(CN)6), 0.2 M, 2 mL in pipet
1 – Iron(III) nitrate (Fe(NO3)3), 0.1 M, 2 mL in pipet
1 – Hydrochloric acid (HCl), 2 M, 10 mL in dropper bottle
1 – Sodium hydroxide (NaOH), 2 M, 10 mL in dropper bottle
3 – Short, thin-stem pipets

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

www.HOLscience.com 3 ©Hands-On Labs, Inc.

Experiment Equilibrium and Le Châtelier’s Principle

Background
Equilibrium

Chemical reactions may go to completion, that is, the reactants produce products until the one
of the reactants (the limiting reactant) is used up. However, most reactions are reversible and
products react to reform the reactants. The reactants, products, and energy associated with a
chemical reaction are referred to as a chemical system. The reaction reaches a state of chemical
equilibrium when the rate of the forward reaction is the same as the rate of the reverse reaction,
which means that the concentrations of the reactants and products are constant. The term
dynamic equilibrium is used to emphasize that the reactions are still occurring, even though the
reaction appears to have stopped changing. See Figure 1.

Figure 1. (Top) This reaction proceeds to completion, as noted by the pink arrow. (Bottom) This
reaction is reversible, as noted by the blue arrows which move in two directions.

The situation at equilibrium varies greatly from reaction to reaction. Some chemical reactions reach
equilibrium with mainly reactants present, other reactions reach equilibrium with appreciable
amounts of both reactants and products present, while still others do not reach equilibrium until
mostly products are present. In addition, it makes no difference if a reaction is started with 100%
reactants or with 100% products, or whether a reaction is started with 1 mole of reactants or 10
moles of products, it will reach the same equilibrium point where free energy is equal to zero.
See Figure 2.

Figure 2. Reaction at chemical equilibrium. In this reaction, the reactants (9 green balls and 9
pink balls) react to form product (9 purple balls). As the reaction is at chemical equilibrium, the
reaction does not go to 100% completion. Rather, at chemical equilibrium, there is a mixture of

both products and reactants: 8 purple balls, 1 green ball, and 1 pink ball.

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Experiment Equilibrium and Le Châtelier’s Principle

Equilibrium Constant

When a reaction is at equilibrium (and at a specific temperature), the concentrations of reactants
and product remain constant, which makes it possible to define an equilibrium constant, K. (See
below) The equilibrium constant, K, is equal to the molar concentrations of the products, each
raised to a power equal to the coefficient in the balanced chemical equation, divided by the molar
concentrations of the reactants, each raised to a power equal to the coefficient in the balanced
chemical equation.

If the products concentrations are much higher than the reactant concentration at equilibrium,
then K will be a large number. K values close to 1 mean that both reactants and products are
present in similar amounts at equilibrium. Small K values (less than 1) indicate that reactant
concentrations are higher than product concentrations at equilibrium.

Generally, if the value of K is greater than 1, we say that the reaction favors the products and if the
value of K is less than 1, the reaction favors the reactants. The larger the value of K, the more the
reaction favors the products. For example, consider the sample reaction and equilibrium constant
calculations below:

The concentrations of the reactants and products at equilibrium are: [N2] = 1.4 x 10
-3 M, [Cl2] =

4.3 x 10-4 M, and [NCl3] = 1.9 x 10
-1 M. The equilibrium constant is calculated as shown below:

As the value of K is much larger than 1, the reaction favors the formation of product. See Figure 3.

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Experiment Equilibrium and Le Châtelier’s Principle

Figure 3. Relationship between K and the formation of products vs. reactants.

Le Châtelier’s Principle

Once a reaction is at equilibrium, any change in concentration (or pressure of a gaseous reactant or
product) will disrupt the chemical system. The reaction will spontaneously return to equilibrium,
but at a slightly different position than the original one. For example, adding more reactant to a
system at equilibrium will speed up the forward reaction and produce more products. The resulting
increase in product concentration will cause the reverse reaction to speed up also until it equals
the forward rate, re-establishing the equilibrium at a higher product concentration than before.
We say that this reaction “shifted to the right” to again achieve equilibrium. Adding product to
a reaction at equilibrium results in the opposite situation. Le Châtelier’s principle states that a
chemical system will adjust in response to in order change to return to equilibrium. See Table 1.

Table 1. Chemical system responses to change.

Type of Change
Chemical System Shift

(Right shifts toward product.
Left shifts toward reactant.)

Increase concentration of reactant
OR

Decrease concentration of product
Right

Decrease concentration of reactant
OR

Increase concentration of product
Left

Increase temperature of an exothermic reaction Left
Increase temperature of an endothermic reaction Right
Decrease temperature of an exothermic reaction Right

Decrease Temperature of an endothermic reaction Left
Decrease pressure (gases only) More gas molecules
Increase Pressure (gases only) Less gas molecules

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Experiment Equilibrium and Le Châtelier’s Principle

Reaction Quotient

The value of the equilibrium constant for a reaction varies with temperature. In a chemical
system at equilibrium, if the forward reaction is exothermic, then the reverse reaction must be
endothermic, and vice versa. Increases in temperature cause both reactions to speed up but will
have a greater effect on the endothermic reaction. Increases in temperature are said to “favor”
the endothermic process while decreases in temperature “favor” the exothermic process. See
example below:

If additional heat were applied to the exothermic reaction, the system would shift to the left,
reducing the overall temperature of the reaction. Likewise, if additional AB3 was added to the
reaction, the system would again shift to the left to create the appropriate balance of products
and reactants.

The chemical system can also be shifted toward the right upon addition of A2 or B2, as both
additions would cause an increased concentration of product. Likewise, as both A2 and B2 are
gaseous substances, increasing or decreasing the pressure of these two gases would cause a shift
of the chemical system to re-achieve the equilibrium position.

The above examples discussed predicting how a reaction at equilibrium will respond to changes.
It is also possible to predict how a newly combined mixture of reactants and products will proceed
to reach equilibrium. We do this by determining the reaction quotient, Q, and comparing it
to the value of the equilibrium constant. Q is calculated the same way as K, but using initial
concentrations rather than equilibrium concentrations. See Table 2.

Table 2. Relationship between reaction quotient (Q) and equilibrium constant (K).

Relationship Chemical System Response
Q is equal to K System remains at equilibrium

Q is greater than K The reverse reaction will predominate as the system approaches equilibrium.

Q is less than K The forward reaction will predominate as the system approaches equilibrium.

Example:

For example, examine the reaction for the synthesis of ammonia, which has a K = 6.0 x 10-2. Note
that the initial concentrations are differentiated from equilibrium concentrations by the subscripted
“0.”

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Experiment Equilibrium and Le Châtelier’s Principle

The following concentrations of gases are placed in a container: [H2] = 1.0 x 10
-2M, [N2] 5.0 M and

[NH3] = 1 x 10
-4 M. Calculate Q to determine which reaction (forward or reverse) will proceed to

reach equilibrium.

As Q is less than K (Q = 0.0020, K = 0.060), the forward reaction will predominate as the system
approaches equilibrium.

Hemoglobin is a protein
in your red blood cells that

can combine with oxygen to allow
your blood to carry that oxygen
to cells in your body. Hemoglobin
combines with oxygen in the lungs

and releases oxygen to your cells. The
partial pressure of oxygen in these two

locations differs: high in the lungs,
low in the cells throughout the

body.

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Experiment Equilibrium and Le Châtelier’s Principle

Exercise 1: Equilibrium of Chromate and Dichromate
In this exercise, you will investigate Le Châtelier’s principle on chromate-dichromate equilibrium.

Note: The potassium chromate used in this experiment is potentially toxic and contact with both skin
and eyes must be avoided. Ensure that gloves and goggles are worn when working with this chemical.

1. Put on your safety goggles and gloves.

2. Set the 24-well plate on the table and use scissors to carefully snip off the tip of the potassium
chromate (K2CrO4) chemical pipet. Using the well plate as a pipet holder, set the pipet upright
in a well. See Figure 4.

Figure 4. Well plate with chemical pipet.

3. Place 8 drops of potassium chromate in the well directly in front of the pipet. See Figure 5.

Figure 5. Placing 8 drops of potassium chromate into a well.

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Experiment Equilibrium and Le Châtelier’s Principle

In the remainder of this exercise, you will investigate the equilibrium reaction between chromate
and dichromate (the potassium is a spectator ion, and does not participate in this equilibrium
reaction):

4. Record the color of the chromate in the well plate in Data Table 1 of your Lab Report Assistant.

Note: The indication for which way the chemical system shifts is based on color.

5. Add four drops of hydrochloric acid (HCl) to the potassium chromate.

6. Add drops of sodium hydroxide (NaOH) to the chemical system until the color change is
complete. Record the number of drops of NaOH added and the resulting color change in Data
Table 1.

7. Set up a cold water bath (cup of ice mixed with cold water) and a hot water bath (cup of very
hot water). See Figure 6.

Figure 6. Water baths. A. Cold water bath. B. Hot water bath. Note that the water is not boiling,
but is hot enough to create steam along the side of the cup (red arrow).

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Experiment Equilibrium and Le Châtelier’s Principle

8. Place 8 drops of potassium chromate into an empty well.

9. Add 4 drops of HCl to the potassium chromate in the well.

10. Use an empty, short-stem pipet to draw-up all of the potassium chromate/HCl from the well
and turn the pipet over so all of the reaction is pooled at the bottom of the pipet. See Figure 7.

Figure 7. Drawing-up the K2CrO4/HCl reaction into an empty pipet.

11. Observe the color of the reaction and record in Data Table 2 of your Lab Report Assistant.

12. From the color of the reaction, determine if the reactants or products are favored. Record in
Data Table 2 and explain your answer.

13. Place the reaction in the pipet into the cold water bath, submerge the liquid in the water bath,
and tape the end of the pipet to the cup to keep the pipet upright. Allow the pipet to remain
in the cold water bath for 2-3 minutes. See Figure 8.

Figure 8. Pipet in cold water bath.

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Experiment Equilibrium and Le Châtelier’s Principle

14. After the pipet has been in the cold water bath for 2-3 minutes, remove the pipet from the
cold water bath and observe the reaction, including both color and additional observations.
To completely observe the color of the reaction, swirl the contents in the pipet.

15. From the color and observations of the reaction in the pipet, determine if the reactants or
products were favored as the result of the cold-water bath. Record in Data Table 2 and explain
your answer.

16. Place the reaction in the pipet into the hot water bath, submerge the liquid in the water bath,
and tape the end of the pipet to the cup to keep the pipet upright. Allow the pipet to remain
in the hot water bath for 2-3 minutes. See Figure 9.

Figure 9. Pipet in hot water bath.

17. After the pipet has been in the hot water bath for 2-3 minutes, remove the pipet from the
hot water bath and observe the reaction, including both color and additional observations. To
completely observe the color of the reaction, swirl the contents in the pipet.

18. From the color and observations of the reaction in the pipet, determine if the reactants or the
products were favored as the result of the hot-water bath. Record in Data Table 2 and explain
your answer.

19. Repeat steps 13-18 as necessary, to complete Data Table 2. You may not need to repeat the
steps.

20. Place the caps on the HCl and NaOH bottles and leave the well-plate out, as they will be used

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Experiment Equilibrium and Le Châtelier’s Principle

in the next exercise.

Questions
A. Use your results to determine if the forward reaction in the potassium chromate/HCl reaction is

endothermic or exothermic. Explain your answer, using Table 1 to help construct your thoughts.

B. Write the equation for the equilibrium constant (K) of the reaction studied in this exercise.

Use the information below to answer Questions C, D, and E:

The equilibrium constant (K) of the reaction below is K = 6.0 x 10-2, with initial concentrations as
follows: [H2] = 1.0 x 10

-2 M, [N2] = 4.0 M, and [NH3] = 1.0 x 10
-4M.

C. If the concentration of the reactant H2 was increased from 1.0 x 10
-2 M to 2.5 x 10-1M, calculate

the reaction quotient (Q) and determine which way the chemical system would shift.

D. If the concentration of the reactant H2 was decreased from 1.0 x 10
-2 M to 2.7 x 10-4M, calculate

the reaction quotient (Q) and determine which way the chemical system would shift.

E. If the concentration of the product NH3 was decreased from 1.0 x 10
-4 M to 5.6 x 10-3M, calculate

the reaction quotient (Q) and determine which way the chemical system would shift.

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Experiment Equilibrium and Le Châtelier’s Principle

Exercise 2: Equilibrium of Ferrocyanide and Ferric
Ferrocyanide
In this exercise, you will investigate Le Châtelier’s principle on ferrocyanide and ferric ferrocyanide.

Note: Never mix cyanide containing compounds with any kind of acid, as toxic fumes will be
produced. Only add cyanide compounds to bases as instructed in this exercise.

1. Put on your safety goggles and gloves.

2. Set the 24-well plate on the table and use scissors to carefully snip off the tip of the potassium
ferrocyanide (K4Fe(CN6)) chemical pipet. Wipe the scissors with a damp paper towel and then
carefully snip off the tip of the iron(III) nitrate (Fe(NO3)3) chemical pipet. Using the well plate
as a pipet holder, set the pipets upright in a well. See Figure 10.

Figure 10. Pipets in well plate.

3. Place 8 drops of potassium ferrocyanide in an empty well of the well plate.

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Experiment Equilibrium and Le Châtelier’s Principle

In the remainder of this exercise, you will investigate the equilibrium reaction between ferrocyanide
and ferric ferrocyanide (also called iron(III) ferrocyanide):

4. Record the color of the potassium ferrocyanide in the well plate in Data Table 3 of your Lab
Report Assistant.

Note: The indication for which way the chemical system of the reaction shifts is based on color.

5. Add 1 drop of iron(III) nitrate to the potassium ferrocyanide. Record the color of the ferric
ferrocyanide in Data Table 3.

6. Add drops of NaOH to the reaction until the chemical system shifts, as noted by a color change.
Record the number of drops of NaOH added in Data Table 3.

7. Observe the reaction after the chemical system has shifted and record observations in Data
Table 3.

8. Repeat steps 3 through 7 as needed to complete Data Table 3. Note that you may not need to
perform the steps additional times.

9. Clean all equipment and return to the kit for future use.

10. Dispose of used chemical pipets properly.

11. When you are finished uploading photos and data into your Lab Report Assistant, save and
zip your file to send to your instructor. Refer to the appendix entitled “Saving Correctly,” and
the appendix entitled “Zipping Files,” for guidance with saving the Lab Report Assistant in the
correct format

Questions
A. From your observations and data collected in Data Table 3, describe the direction of the

chemical system shift upon addition of NaOH.

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Experiment Equilibrium and Le Châtelier’s Principle

The Properties of Water
Hands-On Labs, Inc.
Version 42-0129-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will explore unique characteristics of water.
You will investigate surface tension, cohesion
and adhesion, density, specific gravity, and phase
changes. You will experiment with capillary action,
create a heating curve, and calculate the density and
specific gravity for a variety of substances.

© Hands-On Labs, Inc. www.HOLscience.com 1

EXPERIMENT

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Describe and list some of the defining characteristics of water.

● Define adhesion, cohesion, and capillary action.

● Describe melting, vaporization, sublimation, condensation, freezing, and deposition.

● Produce a heating curve of water.

● Prepare a graph and interpret results.

● Calculate density and specific gravity.

Time Allocation: 4.5 hours

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Experiment The Properties of Water

Materials
Student Supplied Materials

Quantity Item Description
1 Clear tape
1 Coin (Penny, nickel, dime, or quarter)
1 Crushed ice*
1 Cup or drinking glass, 2 oz.– 8 oz., flat top
1 Distilled water
1 Dish soap
1 Food coloring (blue or green)
1 Isopropyl alcohol (10 mL)
1 Matches
2 Needle or straight pin
1 Olive oil (10 mL)

50-100 Paper clips
1 Roll of paper towels
1 Source of tap water
1 Timer (watch, cell phone, etc)

HOL Supplied Materials

Quantity Item Description
1 Beaker, 100 mL, glass
1 Beaker, 250 mL, glass
1 Burner fuel
1 Burner stand
1 Capillary tubes (3 tubes of varying diameter)
1 Graduated cylinder, 25 mL
1 Digital scale
1 Metric ruler
1 Pair of safety goggles
3 Pipets, empty, short stem
1 Test tube, 25 x 150 mm
1 Thermometer, analog

* The experiment requires crushed ice. Ensure that you have ample time to prepare ice before
starting.

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Experiment The Properties of Water

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software such as Microsoft® Word or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

www.HOLscience.com 4 ©Hands-On Labs, Inc.

Experiment The Properties of Water

Background
Unique Properties of Water

Water is one of the most important substances on Earth, as it is an essential resource for all living
things. Approximately 70% of Earth’s surface is covered in water and the human body is composed
of approximately 60% water. Water is the only substance that people regularly experience in all
three states of matter: solid, liquid, and gas. As water is essential to living organisms, it is critical
to gain an understanding of the properties of water and how it facilitates daily life.

Many of the unique properties of water are the result of its polar covalent molecular structure,
and ability to form intermolecular hydrogen bonds with other water molecules. Hydrogen bonding
is the result of an attraction between molecules that contain a partially positive hydrogen atom
(H+) and a highly electronegative atom such as oxygen (O), nitrogen (N), or fluorine (F). Hydrogen
bonding is exhibited between water molecules (H2O) in its liquid and solid phases. When multiple
water molecules interact, the partially positively charged hydrogen end of one molecule is
attracted to the partially negatively charged oxygen end of another molecule. See Figure 1.

Figure 1. Molecular structure of water. A. Polarity of water: a center of negative charge (δ-) near
the oxygen atom and centers of positive charge (δ+) near the hydrogen atoms. ©molekuul.be B.

Hydrogen bonding in liquid H2O.

Water, as a liquid, does not exist as separate water molecules, but as a loose network of molecules
linked by hydrogen bonds. These hydrogen bonds, while much weaker than the covalent bonds
within the water molecule, are responsible for the tendency of water molecules to “stick” together.
The term used to describe the tendency of like molecules to “stick” together is cohesion. Water
has a high cohesion compared to many other substances. The cohesive nature of water is evident
in a water droplet; without cohesion the water would not hold its shape. See Figure 2. When
liquids with strong intermolecular hydrogen bonds, such as water, have strong cohesive forces
they also exhibit high surface tension. Surface tension is the resistance of the object to increase
its surface area. Surface tension creates a “skin” or “shell” around the liquid, strong enough to
support light objects. Surface tension is the property that allows small insects to walk on water
without breaking through the surface tension (skin) of the liquid.

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Experiment The Properties of Water

Figure 2. Cohesive nature of water. As a result of the intermolecular hydrogen bonding in water,
the water prefers to stick together rather than interact with its surrounding surface, creating

surface tension around the water droplet. © koh sze kiat

In addition to cohesive forces, water also exhibits adhesive forces. While cohesive forces are the
result of the intermolecular hydrogen bonds within the water, adhesive forces are the result of
attractive forces between the water and its container, or surrounding surface. Water molecules
have a polar covalent molecular structure; therefore, they have a desire to interact with surfaces
containing polar bonds, such as glass. See Figure 3.

Figure 3. Adhesive nature of water. (Left) Water droplet on nonpolar surface, (Right) Water
droplet on a polar surface. As the result of the polar water molecule’s desire to interact with the

polar glass surface, the water droplet will spread out along the surface of the glass.

A third property of water results from the interplay of the cohesive and adhesive forces and is
called capillary action. Capillary action is the spontaneous upward rising of a liquid through a
narrow tube, or opening.

Capillary action is how water and
the nutrients that water carries are
brought up from the roots to the

leaves in a plant.

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Experiment The Properties of Water

Density and Specific Gravity

Density is defined as mass per unit of measure, which is most often volume. Density is a way to
describe how heavy an object is for its size. The density of a liquid is usually reported as grams
per milliliter (g/mL), while the density of a solid is usually reported as grams per cubic centimeter
(g/cm3). The density of water is accepted at 1 g/mL. Substances with a density greater than 1 g/
mL will sink when placed in water, while objects with a density less than 1 g/mL will float when
placed in water. The density of an object is determined with the following equation:

Specific gravity is the ratio of the density of a substance to the density of a reference substance.
For liquids, this reference substance is usually water. When calculating the specific gravity, the

densities must be calculated at the same temperature and pressure.

Substances with a specific gravity less than 1 are less dense than water, and thus tend to float
when placed in water. Conversely, substances with a specific gravity greater than 1 are denser
than water, and thus tend to sink when placed in water.

Phase Changes

Adding energy to, or removing energy from, a substance (such as water) may cause a change in
its state of matter. Energy may be added or removed by increasing or decreasing temperature.
When energy is added to a solid or a liquid substance, the molecules move more quickly, and
may eventually break the weak intermolecular forces that hold them together. Breaking these
intermolecular forces can result in a change in the state of matter. On the other hand, when energy
is removed from a liquid or a gas substance, the molecules move more slowly and eventually
form intermolecular bonds. Because intermolecular forces are different for each substance,
temperatures in which a phase change occurs differ for each substance, and become unique
properties of those substances.

As shown in Figure 4, changes among states of matter may fit into six categories. Energy is required
to overcome molecular attraction. Changes in states of matter moving from a solid to liquid phase
(melting), a liquid to gas phase (vaporization), or a solid to gas phase (sublimation), require
energy. When moving in the opposite direction, from a gas to liquid phase (condensation), a
liquid to solid phase (freezing), or a gas to solid phase (deposition), energy is released. Notice
that during the processes sublimation and deposition, the phase change is between the solid and
gas state, and the particles do not enter an intermediary liquid phase.

MassDensity =
Volume

2

(substance)

(H O)

densitySpecific Gravity =
density

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Experiment The Properties of Water

Figure 4. Changes among the phases of matter. Red arrows represent phase changes that
require energy (heat) and blue arrows represent phase changes that release energy (cooling).

Heating Curve

A heating curve, as shown in Figure 5, is a graph that plots the temperature of a substance over
time as heat is applied to the substance. The heating curve describes temperature changes over
states of matter, and can be used to determine the melting point and boiling point of substances.
There is no increase in the temperature of the substance at the melting point during the melting
process or at the boiling point during the boiling process because all of the energy applied to the
substance is used to weaken intermolecular forces until they are all broken.

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Experiment The Properties of Water

Figure 5. Model heating curve. Melting and boiling points occur when temperature does not
increase.

Pressure may also affect the state of matter. When increasing pressure, molecules move more
quickly and bump into each other more often. This increases the temperature of the solution and
also causes more of the molecules to get excited. This increases the number of intermolecular
bonds that are broken. Therefore, with more pressure, substances are more likely to vaporize, or
sublime. When pressure is decreased, molecules are less likely to bump into each other, and fewer
intermolecular bonds are broken. Substances are more likely to condense or undergo deposition
with less pressure. However, pressure has very little effect on the liquid/solid transition. This
emphasizes the fact that liquids and solids are already in a condensed state and are considered
incompressible. Therefore, increasing pressure does not increase the frequency of liquid and
solid molecules bumping into each other and does little to raise the melting point.

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Experiment The Properties of Water

Exercise 1: Water and its Unique Properties
In this exercise, you will investigate the properties of cohesion and adhesion in water. You will
demonstrate how these properties work together to create surface tension and capillary action.

Part I

1. Fill the 100-mL glass beaker with tap water so that the beaker is almost full.

2. Hold a needle vertically, as close to the water in the beaker as possible without actually
touching the needle to the water.

Note: vertically is up and down in direction, imagine the post of a stop sign.

3. Gently drop the needle onto the water.

4. Record the observations in Data Table 1 of your Lab Report Assistant in the “Vertical” Column.

5. Carefully retrieve the needle from the water and dry it.

Note: The needle must be completely dry between each step. If you would like to repeat any steps,
make sure to dry the needle before you begin.

6. Hold the needle horizontally, as close to the water in the beaker as possible without actually
touching the water.

7. Gently place the needle onto the water.

Note: Do not drop the needle at an angle.

8. Record the observations in the “Horizontal” column of Data Table 1.

9. Drain the water from the beaker and dry both the beaker and needle.

Part II

1. Place a paper towel on a flat surface.

2. Place the 2 oz.– 8 oz. cup or drinking glass on the paper towel.

3. Fill the cup or drinking glass just above the brim with tap water without allowing the water to
overflow. See Figure 6.

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Experiment The Properties of Water

Figure 6. Water filled just above brim.

4. Estimate how many paper clips you think can be added to the cup before the water overflows,
breaking the surface tension. Record the estimation in Data Table 2 of your Lab Report
Assistant.

5. Add paper clips to the cup 1 at a time while counting the number added. The paper clips will
sink to the bottom of the cup. Add paper clips until the surface tension breaks and water spills
over the cup.

6. Record the results in Data Table 2.

7. Drain the water from the cup or glass and then dry both the cup or glass and paper clips.

Part III

1. Place a paper towel on a flat surface.

2. Place a coin of any type on the paper towel. Ensure that the coin is level.

3. Fill a pipet with tap water.

4. Make a guess as to how many drops of water you think can be placed on the coin before the
surface tension of the water will break and the water will flow off the face of the coin. Record
your hypothesis in Data Table 3 of your Lab Report Assistant.

5. Using the pipet, place 1 drop of water at a time onto the coin until the water flows over the
face of the coin, see Figure 7. Count the drops as they are added and record the final number
of drops in Data Table 3.

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Experiment The Properties of Water

Figure 7. Dropping water onto a coin and proper hand placement.

6. Dry the coin and remove the water from the pipet. Keep the pipet for use in the Exercise 3.

Questions
A. How did the experiment in Part I demonstrate surface tension? Use your experiment

observations when answering this question.

B. In Part I, when adding the needle to the water, which approach worked best to balance the
needle on the water — the vertical or horizontal placement? Explain your answer.

C. In Part II, how did your paper clip estimation compare to your paper clip results?

D. In Part III, how were the properties of adhesion and cohesion demonstrated in the experiment?

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Experiment The Properties of Water

Exercise 2: Capillary Action
In this exercise, you will further investigate the properties of cohesion and adhesion in water. You
will determine how the internal diameter of a capillary tube affects the rate the water travels up
the height of the tube.

1. Remove the small cardboard piece that contains the 3 capillary tubes of varying internal
diameters. The internal diameters of the capillary tubes are 0.25 mm, 0.58 mm, and 1.0 mm.
Think about adhesion and cohesion and develop a hypothesis as to which capillary tube would
be the first to draw water through the height of the tube. Which tube would be the last to
draw water though the height of the tube? Record your hypothesis in Data Table 4 of your
Lab Report Assistant.

2. Carefully remove the 3 capillary tubes from the cardboard piece. See Figure 8.

Figure 8. Capillary tubes on cardboard piece.

3. Invert the 25 x 150 mm test tube so that it is able to stand on a table.

4. Use tape to carefully affix each of the capillary tubes to the outside of the 25 x 150 mm test
tube. Ensure that the bottom of the capillary tubes are evenly aligned, 2–5 mm above the
bottom of the test tube, see Figure 9. It is important that the tape is placed near the center of
the capillary tubes and not blocking the openings at the top and bottom of the tubes.

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Experiment The Properties of Water

Figure 9. Capillary tubes affixed to the outside of the 25 x 150 mm test tube. The bottoms of the
tubes are evenly aligned. Note that the opening of the test tube is flush against the table.

5. Fill the 100-mL glass beaker with approximately ½ inch of tap water.

6. Add 2–3 drops of blue or green food coloring to the water and swirl the beaker to mix the
solution.

7. Start the timer and immediately place the test tube into the beaker with the colored water.
See Figure 10. The water will immediately begin to travel up the capillary tubes. Allow the
capillary tubes to remain in the colored water for 30 minutes. Do not disturb or bump the
setup during the 30 minutes.

Figure 10. Capillary tubes in colored water.

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Experiment The Properties of Water

8. At the end of the 30 minutes, use the ruler to measure the distance (cm) the water traveled in
each capillary tube. Record the data in Data Table 5 of your Lab Report Assistant.

9. Create a line graph of the relationship between the internal diameter of the capillary tube and
the distance water traveled. Plot the internal diameter on the independent axis (x-axis), and
plot the distance traveled on the dependent axis (y-axis).

10. Resize the graph and insert it into Data Table 6 of your Lab Report Assistant. Refer to the
appendix entitled, “Resizing an Image” for guidance with resizing an image.

11. Wash and dry the beaker. Return all items to the kit for future use.

Questions
A. How did this activity demonstrate capillary action? Explain your answer using your experiment

results and observations.

B. Did your experimental results support or refute your hypothesis? Use the data collected in
Data Table 5 to support your answer.

C. Describe the relationship between the internal diameter and the distance the water traveled.
Was there any correlation between the two? Use the graph created in Data Table 6 to support
your answer.

D. How would you expect your results to change if the capillary tubes were made of plastic
instead of glass?

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Experiment The Properties of Water

Exercise 3: Density and Specific Gravity
In this experiment, you will determine the density of water, olive oil, and isopropyl alcohol. You
will then calculate the specific gravity of each liquid.

1. Gather the graduated cylinder, distilled water, short stem pipets, isopropyl alcohol, and olive
oil.

2. Place the clean, dry, 25-mL graduated cylinder on the tared scale. Record the mass of the
graduated cylinder (g), in Data Table 7 of your Lab Report Assistant under “Mass A” column
for water.

3. Fill the graduated cylinder with 5.0 mL of distilled water; use the short stem pipet to help
measure exactly 5.0 mL of water. Record the volume in Data Table 7.

4. Place the 25-mL graduated cylinder with 5.0 mL distilled water on the tared scale. Record the
mass of the graduated cylinder + liquid (g) in Data Table 7 under “Mass B”.

5. Calculate the mass of the water by subtracting “Mass A” from “Mass B.” Record the mass of
the water in Data Table 7.

6. Pour the water down the drain and fully dry the graduated cylinder.

7. Repeat steps 2–6 for the isopropyl alcohol and olive oil.

8. Calculate the densities of the water, isopropyl alcohol, and olive oil. Record the density of
each liquid in Data Table 7.

9. Calculate the specific gravity for each of the 3 liquids and record in Data Table 7.

10. Wash and dry all items and return them to the kit for future use.

Questions
A. If the three substances tested in this exercise were poured into the same graduated cylinder,

in what order would you expect them to layer? Use your data in Data Table 7 to support your
answer.

B. Ice floats in water. Use your experience with density and specific gravity in this exercise to
explain how ice is able to float in water. What happens to water as it freezes that allows it to
float in water?

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Experiment The Properties of Water

Exercise 4: Temperature and Phase Changes
In this exercise, you will make observations of the phase changes of water (H2O). You will measure
temperature and create a heating curve to determine the melting point and boiling point of water.

1. Gather the 250-mL beaker, approximately 150 mL of crushed ice, a watch or timer, the
thermometer, burner stand, burner fuel, and matches.

Note: Large ice cubes may be crushed by placing them in a large plastic bag, placing the bag on a
durable surface, and breaking the pieces apart with a hammer or other heavy object.

2. Fill the beaker to about the 150-mL line with crushed ice.

3. Place the thermometer in the center of the ice. Do not allow the thermometer to touch the
sides or bottom of the beaker.

4. After holding the thermometer in the ice for about a minute, note the time and record
temperature at 0 minutes in Data Table 8 of your Lab Report Assistant. Additionally, record
your observations about the state of matter (solid, liquid, or gas) of the water in Data Table 8.

5. Uncap the burner fuel, light the wick with a match or lighter, and place the fuel under the
stand on a pie pan as shown in Figure 11.

Figure 11. Burner setup. Note that the flame is blue which is sometimes difficult to see.

6. Place the beaker on the burner stand. Keep holding the thermometer in the middle of the ice.

7. Start the timer and begin taking temperature and observation readings every minute,
recording your findings in Data Table 8.

Note: It is important that you record both the temperature AND the state or states of matter present
every minute throughout the experiment.

8. Gently stir the ice with the thermometer as it heats.

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Experiment The Properties of Water

9. Continue to stir the ice or water and record temperature and observations every minute until
the water has boiled for 5 minutes. Do not allow the thermometer to rest on the glass of the
beaker.

10. Extinguish the burner fuel by lightly placing its cap over the flame; do not tighten cap until the
burner fuel container has fully cooled.

11. Thoroughly wash and rinse the equipment for future use.

12. Using the temperature data recorded in Data Table 8, create a heating curve.

● Plot time (minutes) on the x-axis (horizontal axis) and temperature (°C) on the y-axis
(vertical axis). Connect the plotted points with a line.

● Label the heating curve to show each phase of matter (solid, solid + liquid, liquid, liquid +
gas).

● Label the melting point and boiling point on the heating curve.

● Resize the graph and insert it into Data Table 9 of your Lab Report Assistant.

Note: An example heating curve is given in Figure 5 of the Background.

13. When you are finished uploading photos and data into your Lab Report Assistant, save your
file correctly and zip the file so you can send it to your instructor as a smaller file. Refer to the
appendix entitled “Saving Correctly” and the appendix entitled “Zipping Files” for guidance
with saving the Lab Report Assistant correctly and zipping the file.

Questions
A. Are there parts of the curve with positive slopes and parts that are flat (slope of zero)? What

states of matter are present when the slope of the heating curve is positive and what states
of matter are present when the slope is zero or close to zero?

B. Describe the key characteristics for the three states of matter.

C. Define the melting point. What was the observed melting point of water?

D. Define boiling point. What was the observed boiling point of water?

E. What happens to heat energy when it is not increasing the temperature of the substance in
the beaker? Use your heating curve to explain your answer.

F. Was temperature perfectly constant during your test while the water was melting and while
it was boiling? Explain why or why not.

G. The published melting point of H2O is 0°C, and the published boiling point is 100°C. Why may
you have found different values?

H. Use the following information to determine if the intermolecular forces of isopropyl alcohol
are greater or weaker than the intermolecular forces of water. Explain your answer. The
melting point of isopropyl alcohol (rubbing alcohol, C3H8O) is about -90°C and the boiling
point is about 82°C.

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Experiment The Properties of Water

The Properties of Water
Hands-On Labs, Inc.
Version 42-0129-00-02

Lab Report Assistant
This document is not meant to be a substitute for a formal laboratory report. The Lab Report
Assistant is simply a summary of the experiment’s questions, diagrams if needed, and data tables
that should be addressed in a formal lab report. The intent is to facilitate students’ writing of lab
reports by providing this information in an editable file which can be sent to an instructor.

Exercise 1: Water and its Unique Properties
Data Table 1. Needle Observations.

Vertical Observations Horizontal Observations

Data Table 2. Paper Clips Needed to Break Surface Tension.

Estimation Result

_____ Paper clips _____ Paper clips

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Experiment The Properties of Water

Data Table 3. Drops of Water.

Estimation Result

Questions
A. How did the experiment in Part I demonstrate surface tension? Use your experiment

observations when answering this question.

B. In Part I, when adding the needle to the water, which approach worked best to balance the
needle on the water — the vertical or horizontal placement? Explain your answer.

C. In Part II, how did your paper clip estimation compare to your paper clip results?

_____ Drops of water _____ Drops of water

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Experiment The Properties of Water

D. In Part III, how were the properties of adhesion and cohesion demonstrated in the experiment?

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Experiment The Properties of Water

Exercise 2: Capillary Action
Data Table 4. Capillary Tube Hypothesis.

Hypothesis

Data Table 5. Distance Traveled.

Capillary Tube Distance (cm)
0.25 mm

0.58 mm

1.0 mm

Data Table 6. Internal Diameter versus Distance Traveled.

Graph

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Experiment The Properties of Water

Questions
A. How did this activity demonstrate capillary action? Explain your answer using your experiment

results and observations.

B. Did your experimental results support or refute your hypothesis? Use the data collected in
Data Table 5 to support your answer.

C. Describe the relationship between the internal diameter and the distance the water traveled.
Was there any correlation between the two? Use the graph created in Data Table 6 to support
your answer.

D. How would you expect your results to change if the capillary tubes were made of plastic
instead of glass?

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Experiment The Properties of Water

Exercise 3: Density and Specific Gravity
Data Table 7. Density and Specific Gravity.

Mass A Mass B Mass B – A

Liquid Volume (mL)
Graduated

Cylinder (g)

Graduated

Cylinder with
liquid (g)

Liquid (g)
Density

g/mL
Specific
Gravity

Water

Isopropyl
alcohol

Olive oil

Questions
A. If the three substances tested in this exercise were poured into the same graduated cylinder,

in what order would you expect them to layer? Use your data in Data Table 7 to support your
answer.

B. Ice floats in water. Use your experience with density and specific gravity in this exercise to
explain how ice is able to float in water. What happens to water as it freezes that allows it to
float in water?

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Experiment The Properties of Water

Exercise 4: Temperature and Phase Changes
Data Table 8. Temperature and Observations for Heating Curve.

Time (Min) Temperature (°C) Observations
0
1
2
3
4
5
6
7
8
9

10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30

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Experiment The Properties of Water

Data Table 9. Internal Diameter versus Distance Traveled.

Graph

Questions
A. Are there parts of the curve with positive slopes and parts that are flat (slope of zero)? What

states of matter are present when the slope of the heating curve is positive and what states
of matter are present when the slope is zero or close to zero?

B. Describe the key characteristics for the three states of matter.

C. Define the melting point. What was the observed melting point of water?

D. Define boiling point. What was the observed boiling point of water?

Reaction Order and Rate
Laws
Hands-On Labs, Inc.
Version 42-0195-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will determine how chemical concentration
affects the rate of a reaction, calculate the order of
the reactants, and define the rate law for a reaction
between hydrochloric acid (HCl) and sodium
thiosulfate (Na2S2O3). Additionally, you will define
zero, first, and second order reactions.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Define chemical kinetics and reaction rates, and discuss why they are important for
understanding chemical reactions.

● Differentiate between independent and dependent reaction rates and how they are graphically
represented.

● Define a rate law and how it can be used to determine reaction orders.

● Examine the effects of varying reactant concentrations in a chemical reaction.

● Analyze data to determine the order of a reaction rate.

● Summarize the rate law for an observed reaction.

Time Allocation: 1.5 hours

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Experiment Reaction Order and Rate Laws

Materials
Student Supplied Materials

Quantity Item Description
1 Bottle of distilled water
1 Dish soap
1 Sheet of white paper
1 Source of tap water
1 Timer, clock, or watch with second hand

HOL Supplied Materials

Quantity Item Description
1 Pair of gloves
1 Pair of safety goggles
1 Permanent marker
1 Test tube cleaning brush
1 Well plate – 24
1 Experiment Bag: Reaction Order and Rate Laws

1 – Dropper bottle, empty 15 mL
1 – Hydrochloric acid, 1 M, 15 mL in dropper bottle
2 – Short, thin-stem pipets
1 – Sodium thiosulfate, 0.3 M, 15 mL in dropper bottle

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Reaction Order and Rate Laws

Background
Chemical Kinetics

Chemical kinetics is the study of how quickly chemical reactions occur, and the factors that affect
this speed. The speed of a chemical reaction is called the reaction rate. The reaction rate is a
measure of the change in reactant concentration per unit of time.

Reaction rates may be almost instantaneous, such as the chemical reaction that occurs when our
eyes adjust for night vision. In order for our eyes to see in dark conditions or at night, a chemical
called rhodopsin absorbs a photon and isomerizes, then breaks down into metarhodopsin II and
sends a signal to the brain in less than a second. Alternatively, reaction rates may be billions of years
long, such as in the formation of diamonds. Determining the reaction rate of a chemical reaction
is useful to scientists in many ways, including the development of drugs and the manufacturing
of chemicals. See Figure 1.

Figure 1. An almost-instantaneous reaction rate of hydrogen and oxygen causes an explosion
which gives the space shuttle enough energy to propel into the outer orbit of Earth. © Jose

Antonio Perez

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Experiment Reaction Order and Rate Laws

Reaction Rates may be Independent or Dependent of the Concentration of Reactant

Reaction rates may be constant with the changing concentration of the reactants, or they may be
dependent on the reactant concentration. If the reaction rate is independent of the concentration,
then the relationship between the amount of reactant consumed and time is linear. If the reaction
rate is dependent on the concentration of the reactant, then the relationship is curvilinear, as it is
dependent on the changing amount of reactant. See Figures 2A and 2B. In Figure 2B, the reaction
rate is changing (getting slower) as the amount of reactant is consumed.

Figure 2. Reaction rate examples. A. A hypothetical example of the reaction rate that is
independent of the reactant concentration. In this hypothetical example, the amount of

reactant is used at a rate of 0.2M every 50 seconds. B. A hypothetical example of the reaction
rate that is dependent on the reactant concentration. The amount of reactant used per time is

constantly changing with concentration.

When chemists process
food, they use knowledge of

chemical kinetics to understand
how foods decompose and what
increases or decreases the rate

of decomposition.

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Experiment Reaction Order and Rate Laws

The dependence of the reaction rate upon reactant concentrations is expressed in the “Rate Law”
equation.

Rate Law Equation

Where k is the rate constant, [A] and [B] are the concentrations of reactants; m and n are exponents
that reflect how the rate depends on the reactant concentrations. The exponents m and n give the
“order of the reaction.” Thus, A is said to be of mth order and B to be of the nth order. For example,
if m is two, chemical A would be a second order reactant, and if n was four, B would be a fourth
order reactant. The rate constant and the exponents m and n cannot be determined merely by
looking at the balanced chemical equation; rather, they must be determined experimentally.

The sum of m plus n is the overall reaction order. The exponents m and n are usually positive
whole numbers, but they may be fractions or may even be negative numbers. In the example
given previously, if m was two and n was four in the rate law equation, then the overall reaction
order would be six.

Orders of Reactants

The order of the reactant gives the following information about how the concentration of the
reactant changes the rate:

● Zero Order: If a reactant has an order of zero, the rate is independent of the reactant’s
concentration and that reactant does not appear in the rate law (because anything raised to
the zero power is one). Refer to Figure 2; view the graph that shows when the reaction rate is
independent of the concentration rate. For example, consider “Reaction Example 1” as shown
below. The order of CO is zero and therefore [CO] does not appear in the rate law.

Reaction Example 1:

● First Order: When a reactant is first order the reactant will have an exponent of 1. For example,
consider “Reaction Example 2,” in which both F2 and ClO2 have an order of 1. Increasing the
concentration of F2 or ClO2 increases the rate and decreasing the concentration of F2 or ClO2
decreases the rate. For instance, if the concentration of F2 is doubled, the rate will also double.
If the concentration of F2 is halved, the rate will also be cut in half. Therefore, the rate is
directly proportional to any reactant that is first order. Refer to Figure 2; view the graph that
shows when the reaction rate is dependent of the concentration rate.

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Experiment Reaction Order and Rate Laws

Reaction Example 2:

● Second Order: Reactants that are second order will appear in the rate law as the concentration
of the reactant squared. For example, in the previous example, “Reaction Example 1,” [NO2]
appears in the rate law and is squared. Therefore, NO2 has an order of 2. As a result, if the
concentration of NO2 is doubled, this will cause the rate to increase by a factor of 4.

● Orders above Second Order: Any order above second order will increase the rate to the power
as indicated by the exponent. For example, an order of 3 would increase the rate to the 3rd
power. If a chemical was third order and if it were doubled, the rate would increase by 23, or 8.

Experimentally Determining a Rate Law

Use the following reaction as an example of how a rate law can be determined from experimental
data:

From the following sample data set, determine the rate law:

Initial Concentration Initial Concentration
[NO] [O2] Initial Rate

Experiment 1 0.010 M 0.020 M 0.025 M/s
Experiment 2 0.020 M 0.020 M 0.105 M/s
Experiment 3 0.010 M 0.040 M 0.050 M/s

Notice that when the concentration of one chemical is changed, the other chemical’s concentration
is held constant. When testing for the order of [NO], [O2] is held constant. When testing for [O2],
[NO] is held constant. In doing so, the variable can be computed by using a ratio to determine the
variables m and n since we have isolated the effect of the concentrations of one of the reactants
on the outcome.

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Experiment Reaction Order and Rate Laws

Using a ratio, you can isolate the variable m:

Step 1: Determine the initial rate of two experiments in which the concentration of NO ([NO])
varied (from Table 1), but the concentration of O2 ([O2]) stayed the same.

Step 2: Use known values to compute unknown values. Since you have at least two experiments
in which the chemicals are the same but the concentrations are different, use a ratio of the two
experiments to determine the unknown variables.

Step 3: In order to find the value of the exponential variable, take the natural log (ln) of both sides
of the equation.

Step 4: Determine the initial rate of two experiments in which the concentration of O2 ([O2])
varied (from Table 1), but the concentration of NO ([NO]) stayed the same.

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Experiment Reaction Order and Rate Laws

Step 5: Use known values to compute unknown values. Since you have at least two experiments
in which the chemicals are the same but the concentrations are different, use a ratio of the two
experiments to determine the unknown variables.

Step 6: In order to find the value of the exponential variable, take the natural log (ln) of both sides
of the equation.

Therefore, from this data set, we see that the rate law is:

The reaction is first order in O2, second order in NO, and third order overall.

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Experiment Reaction Order and Rate Laws

Step 7: In order to find “k,” plug in all known numbers using experimental data. The following is
data from Table 1, Experiment 1.

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Experiment Reaction Order and Rate Laws

Exercise 1: Determining the Rate Laws
In this exercise, you will be determining the rate law for the reaction between HCl and Na2S2O3.
Reaction rate data will be generated and used to determine the rate law for the reaction between
hydrochloric acid, HCI, and sodium thiosulfate, Na2S2O3.

From the general form for a rate law given above the general rate law for the reaction between
HCl and Na2S2O3 is written as:

By determining how the reaction rate varies as the concentration of the reactant is varied in the
orders m and n, you can determine the rate law.

Note: Completely read all instructions and assemble all equipment and supplies before beginning
work on this experiment.

Procedure

Part 1: Varying the Concentration of 1.0 M HCl

1. Gather the materials needed for this laboratory found in the materials list. Wear gloves and
safety goggles throughout this experiment.

2. Draw a black “X” on the white sheet of paper with the permanent marker. Make the X just
large enough to show under one well of the 24-well plate. See Figure 3.

Figure 3. Draw a black “X” on the white sheet of paper just large enough to show under one
well of the 24-well plate.

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Experiment Reaction Order and Rate Laws

3. Pour distilled water into the empty dropper bottle labeled “For Distilled Water.” Push the tip
of the dropper bottle into the bottle so that it snaps.

4. From the distilled water dropper bottle, carefully add:

● 6 drops of distilled water to wells C2 and D2.

● 8 drops of distilled water to wells C3 and D3.

Note: No water will be added to wells C1 and D1.

5. From the HCl dropper bottle, carefully add:

● 12 drops of the HCl to wells C1 and D1.

● 6 drops of the HCl to wells C2 and D2.

● 4 drops of the HCl to well C3 and D3.

6. Carefully add 8 drops of Na2S2O3, sodium thiosulfate, into each of the following wells of the
24-well plate: A1, A2, A3, B1, B2, and B3. See Figure 4 for setup.

Figure 4. Setup in the 24-well plate. For this portion of the experiment, you will be using
columns 1, 2, and 3 of the well plate.

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Experiment Reaction Order and Rate Laws

7. Use the permanent marker to label the bulb of an empty short-stem pipet “Na2S2O3.”

8. Slide the well plate so that well C1 is over the “X” on the sheet of white paper.

9. Slightly tip the 24-well plate forward so that the drops of the wells containing Na2S2O3 puddle
together along the bottom rim. Suck all of the contents of one of these wells into the Na2S2O3
pipet. See Figure 5.

Figure 5. Slightly tip the 24-well plate forward so that the drops of the wells containing Na2S2O3
puddle together along the bottom rim and then suck all of the Na2S2O3 from one well into the

empty pipet.

10. When you set the well plate back on the paper (with well C1 over the “X”), lightly shake the
plate back and forth so the liquid covers the entire bottom of the wells again.

11. Zero the timer and hold it in one hand, ready to begin timing.

12. Take the pipet of Na2S2O3 in your other hand. Squeeze all of the contents from the bulb into
well C1 and immediately begin timing.

13. Carefully observe the reaction in Well C1. The moment that the “X” is no longer visible stop
the timer. Record the exact time in seconds under the “Reaction Time (sec): Trial 1” column in
the Data Table 1 in your Lab Report Assistant in the row entitled “C1, D1.”

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Experiment Reaction Order and Rate Laws

14. Repeat steps 8 through 13 for well D1 and record the reaction time data under the “Trial 2”
column in the Data Table 1 in the row entitled “C1, D1.”

15. Repeat steps 8 through 14 for wells C2 and D2 and then for wells C3 and D3.

16. Determine the reaction concentrations of HCl and Na2S2O3, using the following equations.

For Example:

Note: Because you are reaching the same amount of moles of the product (the solid that blocks
the view of the ”X”), the same amount of moles of the reactants are being used for each trial (M).
Therefore, this number should be the same for each trial. Since you are using a ratio to determine the
exponents m and n, when you divide one rate by the other, that number of moles will cancel. This is
why you only have the rate in sec-1 instead of in the units of M sec-1.

17. Calculate the average reaction time of trials 1 and 2. Record this value in Data Table 1 in the
column labeled “Reaction Time: Average.”

18. Calculate the reaction rate for each row in Data Table 1 by taking the inverse of the average
reaction time (1 divided by the average reaction time). Record this number in Data Table 1 in
the column labeled “Reaction Rate (sec-1).”

Part 2: Varying the Concentration of 0.30 M Na2S2O3
19. Use the same sheet of white paper with an “X”on it as you used in Part I.

20. From the distilled water dropper bottle, carefully add:

● 6 drops of distilled water to wells C5 and D5.

● 8 drops of distilled water to wells C6 and D6.

Note: No water will be added to wells C3 and D3.

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Experiment Reaction Order and Rate Laws

21. From the Na2S2O3 dropper bottle carefully add:

● 12 drops of the Na2S2O3 to wells C4 and D4.

● 6 drops of the Na2S2O3 to wells C5 and D5.

● 4 drops of the Na2S2O3 to wells C6 and D6.

22. Carefully add 8 drops of HCl, hydrochloric acid, into each of the following wells of the 24-well
plate: A4, A5, A6, B4, B5, and B6. See Figure 6 for setup.

Figure 6. Setup in the 24-well plate. For this portion of the experiment, you will be using
columns 4, 5, and 6 of the well plate.

23. Use the permanent marker to label the bulb of an empty short-stem pipet “HCl.”

24. Put well C4 over the “X” on the sheet of white paper.

25. Slightly tip the 24-well plate forward so that the drops of the wells containing HCl puddle
together along the bottom rim and you can suck all of the contents of one of these wells into
the “HCl” pipet. Refer to Figure 5.

26. When you set the well plate back on the paper, lightly shake the plate back and forth so the
liquid covers the entire bottom of the wells again.

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Experiment Reaction Order and Rate Laws

27. Reset the timer.

28. Squeeze the contents of the HCl pipet into well C4 and immediately begin timing the reaction.

29. Carefully observe the reaction in well C4. The moment that the “X” is no longer visible stop
the timer. Record the exact time in seconds under the “Reaction Time (sec): Trial 1” column in
the Data Table 2 in your Lab Report Assistant in the row entitled “C4, D4.”

30. Repeat steps 24 through 29 for well D4 and record your data under the “Trial 2” column in the
Data Table 2 in the row entitled “C4, D4.”

31. Repeat steps 24 through 29 for wells C5 and D5 and then for wells C6 and D6.

32. Determine the reaction concentrations of HCl and Na2S2O3. Refer to Part I, step 16.

33. Calculate the average reaction time for the 2 trials. Record this value in Data Table 2 in the
column labeled “Reaction Time: Average.”

34. Calculate the reaction rate for each row in Data Table 2 by taking the inverse of the average
reaction time (1 divided by the average reaction time). Record this number in Data Table 2 in
the column labeled “Reaction Rate (sec-1).”

Note: Because you are reaching the same amount of moles of the product (the solid that blocks
the view of the ”X”), the same amount of moles of the reactants are being used for each trial (M).
Therefore, this number should be the same for each trial. Since you are using a ratio to determine the
exponents m and n, when you divide one rate by the other, that number of moles will cancel. This is
why you only have the rate in sec-1 instead of in the units of M sec-1.

Cleanup:

35. Properly dispose of used chemicals.

36. Clean the equipment with soap and water and allow to thoroughly dry.

37. Return cleaned equipment to the lab it for future use.

Questions
A. Determine the reaction order for HCl using calculations described in the Background section.

Show your work. Note that your answer will probably not be an even whole number as it is in
the examples, so round to the nearest whole number.

B. Determine the reaction order for Na2S2O3 using calculations described in the Background
section. Show your work. Note that your answer will probably not be an even whole number
as it is in the examples.

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Experiment Reaction Order and Rate Laws

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Experiment Reaction Order and Rate Laws

C. Write the rate law for the reaction between HCl and Na2S2O3.

D. Using the following rate law, and the experimental values given, calculate k:

Experiment [F2] (M) [ClO2] (M) Initial rate (M/s)
1 0.5 0.5 0.300 M/s
2 0.8 0.8 0.768 M/s
3 0.5 0.8 0.480 M/s

E. Describe sources of error in this experiment.

Hess’s Law
Hands-On Labs, Inc.
Version 42-0158-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will compare exothermic and endothermic
reactions and learn about the enthalpy of a reaction
(H) and how it relates to Hess’s Law. Using a student-
made basic calorimeter, you will calculate the amount
of heat gained or lost from three reactions (qrxn)
and calculate the change in enthalpy (ΔH) of three
different reactions. Two of the chemical reactions are
intermediate reactions for a third reaction. You will
then compare the percent difference by measuring
an experimental value directly and applying Hess’s
Law by using the data taken from the intermediate
reactions.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● State the difference between an exothermic and an endothermic reaction.

● Define enthalpy (H) and describe how it is related to exothermic and endothermic reactions.

● Describe the change in enthalpy (ΔH) of a reaction and the standard enthalpy of formation
(ΔH°f).

● Define Hess’s law.

● Create a cooling trend based on data taken in a student-made calorimeter.

● Determine the amount of reactants and the amount of products that are in a solution based
on molarity of the reactants.

● Calculate the amount of heat gained or lost (qrxn) from three reactions and the change in
enthalpy (ΔH) of the reactions.

● Evaluate the percent difference between change in enthalpy (ΔH) as calculated by Hess’s law
and measured temperature changes of chemical reactions.

Time Allocation: 2.5 hours

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Experiment Hess’s Law

Materials
Student Supplied Materials

Quantity Item Description
1 Bottle of distilled water
2 Coffee mugs
1 Pair of scissors
1 Roll of paper towels
1 Source of tap water
1 Timer, clock, or watch with second hand

HOL Supplied Materials

Quantity Item Description
1 Digital thermometer
4 Foam cups, 8 oz.
1 Graduated cylinder, 25 mL
2 Pairs of gloves
1 Pair of safety goggles
1 Experiment Bag: Hess’s Law

2 – Ammonium, NH3 (labeled as aqueous Ammonia: NH4OH), 2 M, 10 mL
2 – Ammonium chloride, NH4Cl – 2 M – 10 mL
2 – Hydrochloric acid, HCl – 2 M – 20 mL
2 – Sodium hydroxide, NaOH – 2 M – 20 mL
1 – Long thin stem pipet, 4.5 mL

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Hess’s Law

Background
Thermochemistry

Thermochemistry is the study of heat (thermal energy) that accompanies chemical reactions.
Nearly all chemical reactions involve the release or absorption of heat energy which spontaneously
transfers from warmer to cooler matter. If a chemical reaction occurs and heat is released to the
surrounding environment, the reaction is considered exothermic. Therefore, in an exothermic
reaction, the surrounding temperature increases. Alternatively, if a chemical reaction absorbs
heat energy from its surrounding environment, the reaction is endothermic. In an endothermic
reaction, the surrounding temperature decreases. For reactions that are happening in an aqueous
solution, the water is surrounding the reactants and products. That means that the reactants and
products will release or absorb energy to or from the water. See Figure 1.

Figure 1. Qualities of endothermic and exothermic reactions. ©Aspect3D

The amount of energy that flows as heat during a chemical reaction can be measured. Enthalpy
(H) is the sum of the internal energy of the system (E), which includes heat, and the product of the
pressure (P) and volume (V). When a chemical reaction is at a constant pressure, the amount of
heat given off or absorbed is called the change in enthalpy (ΔH) or the heat of reaction. The unit
of measure for change in enthalpy, as identified by the International System of Units, is the joule
(J). One thousand joules is equal to one kilojoule (1 kJ). When a change in enthalpy accompanies
the formation of 1 mole of a compound under standard conditions from its pure elements, the
enthalpy is denoted as (ΔH°f), which represents the standard enthalpy of formation. The standard
enthalpy of formation is known for a vast number of substances.

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Experiment Hess’s Law

Thermodynamics

Thermodynamics is the study of how heat, energy, and work are interrelated. The First Law of
Thermodynamics is the Law of Conservation of Energy which states that energy can neither be
created nor destroyed but it can be converted from one form to another. Therefore, all energy
transferred between a system and its surroundings during the conversion of reactants to products
must be accounted for as heat or work. When the standard enthalpy of formation is known for all
products and reactants in a reaction, then the change in the enthalpy of the reaction (ΔH) is equal
to the enthalpy of the products (Hproducts) minus the enthalpy of the reactants (Hreactants).

When a reaction releases heat energy and is exothermic, the enthalpy of the products is less than
the reactants, and the ΔH will be a negative number. When the reaction absorbs energy from the
surroundings and is endothermic, the products have greater enthalpy than the reactants, and ΔH
is a positive number. For example, water (H2O) is formed from the pure elements hydrogen (H2)
and oxygen (O2). The reaction is exothermic, and 286 kJ of energy are released:

Hess’s Law

Hess’s law states that the total change in enthalpy is independent of the pathway. This means
that if a reaction takes place in one step or in multiple steps, the ΔH for the overall process must
be the sum of the ΔH values of the constituent reactions.

For example, the decomposition of liquid water into its component elements, H2(g) and O2(g),
occurs in two reactions. In the first reaction, liquid water [H2O(l)] becomes gaseous water vapor
[H2O(g)]. In the second reaction, water vapor [H2O(g)] decomposes into its component elements;
H2(g) and O2(g). Each reaction has its own enthalpy, which are combined to obtain the total change
in enthalpy. Note that as this reaction was reversed from formation to decomposition so the sign
of the enthalpy changed; going from negative to positive.

Notice that this reaction confirms Hess’s law: the heat absorbed in the first reaction (+44 kJ) and
the heat absorbed in the second reaction (+242 kJ) equals the total heat absorbed in the overall
reaction (+286 kJ).

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Experiment Hess’s Law

Dr. Germain Hess
(1802 – 1850) was a Swiss-

born Russian chemist who studied heat
in chemical reactions which formed the

foundation of thermochemistry. After leaving
his career in medicine, Hess focused on chemistry
through his early investigations of minerals and
natural gas. In 1840, Hess announced the law of
constant heat summation, which later became
known as Hess’s law. The experiments that led

to Hess’s law joined thermodynamics and
chemistry, demonstrating that the law of
conservation energy applied to chemical

and physical changes.

Hess’ Law Laboratory

In this laboratory experience, you will measure ΔH for two reactions and use your findings to
predict ΔH for a third reaction. More specifically, you will measure the temperature of a reaction
that has the products NH3 + NaCl + H2O and a reaction that has the products H2O + NaCl. You will
then use your findings to predict the temperature of a reaction that produces NH4Cl. You will then
perform the reaction for NH4Cl and compare your predicted ΔH to temperatures you measure.
The following chemical equations describe the three reactions:

You will perform the third reaction as a tool to validate Hess’s law; the ΔH for the third reaction
should equal the sum of the ΔH values of the two constituent reactions.

Remember that the intermediate reactions can be reversed (and therefore, the sign will be
reversed) when calculating the enthalpy change. For these intermediate reactions, they can be
added together to get the overall reaction (when adding equations, you can cross out the effects
of the same chemicals on the opposite side of the equation). You will reverse the second reaction:

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Experiment Hess’s Law

The overall intermediate reactions will be added together to equal a third reaction:

Comparison of the calculated results will verify the generalization known as Hess’s law. In this case,
the target reaction NH3 + HCl → NH4Cl can also be performed directly and the results compared
to the second two reactions.

Heat energy exchanges do not occur instantaneously and it takes time for energy to move from
a hot object to a cold one. An acceptable solution to this problem is to obtain a cooling curve for
the heat energy exchange and then extrapolate the data back to the exact time that the exchange
began. You will create a cooling trend by taking the temperature of the solution every 20 seconds.
Figure 2 shows a sample graph from hypothetical data.

Figure 2. Example of a cooling trend. The green dot represents the temperature of the system
before the reaction has begun.

You will create a cooling trend by taking data every 20 seconds. You cannot get the immediate
change in heat from mixing the solutions, so you will extrapolate back to point “0” from the
cooling curve. Using the same example data as in Figure 2, the extrapolation is done by extending
the trend line from the cooling curve back to time “0.” See Figure 3.

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Experiment Hess’s Law

Figure 3. Example data for a cooling curve. The trend line has been extended to the “0” time
point to predict the initial temperature as the reaction began. The yellow dot represents

temperature at time “0”; the green dot represents temperature before the reaction began.

Calculations

Molarity (M or mol/L) is a measure of concentration that is calculated as number of moles of
solute (mol) per volume of solution (L). The equation is stated as:

The stoichiometry (relative quantities of reactants and products) of a balanced chemical equation
allows for the calculation of the amount of unknown reactants or products. In this experiment,
you will calculate the amount of moles in the reactants to quantify the amount of moles produced
in the product. To calculate the quantitative amount of product expected through a chemical
reaction, one needs a balanced chemical equation and a known mass or molar amount for one
of the reactants.

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Experiment Hess’s Law

Step 1) Check to ensure that the chemical equation is balanced.

Step 2) Determine the number of moles of one of the reactants.

Step 3) Evaluate the molar ratio of the reactant and the product.

For example, consider the following equation:

Given 0.4 moles of H2, and 0.2 moles of O2 (half the moles of H2), the reaction produces 0.4 moles
of H2O (the same number of moles as H2, and double the number of moles of O2).

After the moles of product are determined, the heat gained or lost from the reaction can be cal-
culated. The equation used to calculate heat gained or lost (qrxn) is:

Note: For all solutions in this experiment, use density to determine mass.

Note: The heat of the reaction, qrxn, could be determined by taking the negative of the heat gained
by the solution, qsoln, plus that gained by the calorimeter, qcal: qrxn = -(qsoln + qcal). A small amount
of heat is absorbed by the calorimeter, which can be measured as the calorimetry constant, qcal.
However, the amount of heat lost to the calorimeter is so insignificant that it is often left off or simply
assumed to be 1J x ΔT. (qcal = c x ΔT).

Once the total thermal energy transfer is known, the enthalpy of reaction can be determined
using the following equation:

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Experiment Hess’s Law

To find the percent difference, use the following equation (use the value collected directly from
NH3 + HCl → NH4Cl as the “theoretical value” and the value obtained from the sum of the other
two reactions as the “experimental value”):

Example:

Consider the following reaction:

Note: The numbers in this example are not intended to be the correct ΔH for this equation.

Using the example data shown in Figures 2 and 3, assume that 30 mL of NH3 and 30 mL of HCl
was used, and they were each a 3 M solution. Because you know the molarity and the amount of
solution used, you can determine the amount of each chemical in moles:

Considering the stoichiometry of the chemical equation, the reactants and the product are each
in a 1:1 ratio. Therefore, when the heat of the reaction is determined, it is for 0.09 moles of NH4.

When performing this experiment, you extrapolated the data to get an increase in temperature
of 15 degrees within the calorimeter. You can determine the heat of the reaction:

Finally, assume that when the ΔH of the intermediate equations were added together, they
summed to -48.22 kJ/mol.

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Experiment Hess’s Law

Exercise 1: Testing Hess’s Law
In this laboratory experience, you will measure ΔH for two reactions and use your findings to
predict ΔH for a third reaction.

Important: Wear safety goggles and gloves throughout this laboratory exercise.

Procedure

Part 1: First Reaction (NaOH + HCl)

1. Construct a calorimeter from 2 foam cups:

a. Set a foam cup in a coffee mug. This will act as the base of the calorimeter and will
hold the solution as the reaction proceeds.

b. Trim the lip of a second foam cup, about 3 cm from the top using a pair of scissors.
This will act as the lid of the calorimeter. Place the trimmed foam cup upside down
in the top of the other foam cup, providing an insulated environment for monitoring
temperature as shown in Figure 4.

c. Using the digital thermometer, poke a hole in the top cup so a thermometer can be
inserted. See Figure 4 again for complete setup.

Note: Use care when inserting the thermometer into the calorimeter since it has a pointed tip that
could puncture the lower cup if inserted too forcefully. However, ensure that the thermometer can
touch the bottom of the bottom cup. If not, trim a little more of the top “lid” of the calorimeter.

Figure 4. Calorimeter. A. Cutting off 3 cm of the top of a foam cup. B. Complete setup for
calorimeter.

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Experiment Hess’s Law

2. Use the clean 25 mL graduated cylinder to accurately measure 10.0 mL of 2M NaOH from the
dropper bottle. When nearing 10.0 mL mark, add drops slowly to ensure a measurement of
exactly 10.0 mL.

3. Remove the lid from the calorimeter and pour the 10.0 mL of NaOH into the foam base.
Replace the lid with the digital thermometer in it.

4. Rinse and dry the graduated cylinder and accurately measure 10.0 mL of 2M HCl.

5. Pour the 10.0 mL of HCl into an unused foam cup. Set the foam cup in an empty coffee cup to
prevent it from tipping over.

6. Turn on the digital thermometer and set it to record in Celsius (°C). Wait at least 1 minute to
allow the temperature to equilibrate. Record the temperature of the NaOH solution in Data
Table 1 of your Lab Report Assistant.

7. Remove the thermometer, rinse the tip with distilled water, dry it with a paper towel, and
place the thermometer in the HCl solution. Wait at least 1 minute and record the temperature
of the solution in Data Table 1.

8. Calculate the average starting temperature by averaging the initial temperatures of NaOH and
HCl and record the average in Data Table 1.

9. Remove the thermometer, rinse the tip with distilled water, and dry the thermometer with a
paper towel.

10. Pour the contents of the foam cup containing HCl into the base of the calorimeter containing
NaOH, combining the 2 solutions. Quickly place the foam lid on top of the cup containing the
combined solutions and insert the thermometer through the hole in the lid. Be careful when
inserting the thermometer to ensure its pointed tip does not puncture the lower foam cup.
Immediately record the initial temperature of the mixed solutions in Data Table 2 of your Lab
Report Assistant.

11. Record temperature every 20 seconds for 5 minutes in Data Table 2.

12. Create a scatter plot using the data in Data Table 2. Plot time on the independent axis (x-axis)
and temperature on the dependent axis (y-axis). Add a linear trendline that represents
temperature change. Extend the trendline to the y-axis, time “0”. The extrapolation at the
“0-second time” most likely represents the “highest change in temperature of the mixture.”
Compare the extrapolated temperature to the initial recorded temperature of the mixture and
record the greater temperature in Data Table 1 as the highest temperature of the mixture.
The graph should look similar to the sample cooling curve in Figure 3. Resize and insert an
image of the graph into Graph 1 of your Lab Report Assistant. Refer to the appendix entitled,
“Resizing an Image” for guidance.

13. Subtract the greatest difference in temperature from the initial average temperature of the 2
separate solutions to get “ΔT,” and record it in Data Table 1.

14. Properly dispose of the solution in the calorimeter. Recall that the reaction of an acid (HCl)
and base (NaOH) is a neutralization reaction that produces salt and water.

15. Rinse all equipment with distilled water in preparation for the second reaction. This includes
the foam cups and the 25 mL graduated cylinder.

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Experiment Hess’s Law

Part 2: Second Reaction (NaOH + NH4Cl)

16. Place the calorimeter assembly into an empty coffee mug again to prevent it from tipping
over. Make sure you are still wearing your safety goggles and gloves.

17. Use the clean 25 mL graduated cylinder to accurately measure 10.0 mL of 2M NaOH from the
dropper bottle. When nearing 10.0 mL mark, add drops slowly to get the exact amount of
10.0 mL.

18. Pour the 10.0 mL of NaOH into the bottom cup of the foam calorimeter.

19. Rinse and dry the graduated cylinder and accurately measure 10.0 mL of 2M NH4Cl.

20. Pour the 10.0 mL of NH4Cl into another foam cup and place the cup into a second empty
coffee cup to prevent it from tipping over.

21. Turn on the digital thermometer and place it into the NaOH solution. Wait at least 1 minute
and record the temperature of the solution in Data Table 3 of your Lab Report Assistant.

22. Remove the thermometer, rinse the tip with distilled water, dry it with a paper towel, and
place the thermometer into the NH4Cl solution. Wait at least 1 minute and record the initial
temperature of the solution in Data Table 3.

23. Remove the thermometer, rinse the tip with distilled water, and dry it with a paper towel.

24. Pour the contents of one foam cup into the second one, combining the 2 solutions. Quickly place
the foam lid on top of the cup containing the combined solutions and insert the thermometer
through the hole in the lid. Be careful when inserting the thermometer to ensure its pointed
tip does not puncture the lower foam cup. Record the initial temperature of the combined
solutions as quickly as you see it in Data Table 4 of your Lab Report Assistant.

25. Record the temperature every 20 seconds for 5 minutes in Data Table 4.

26. Create a scatter plot using the data in Data Table 4. Plot time on the independent axis (x-axis)
and temperature on the dependent axis (y-axis). Add a linear trendline that represents
temperature change. Extend the trendline to the y-axis, time 0. The extrapolation at the
“0-second time” most likely represents the “highest temperature of the mixture.” Compare
the extrapolated temperature to the initial recorded temperature of the mixture and record
the greater temperature in Data Table 3 as the highest temperature of the mixture. The graph
should look similar to the sample cooling curve in Figure 3. Resize and insert an image of the
graph into Graph 2 of your Lab Report Assistant.

27. Subtract the greatest difference in temperature of the combined solutions from the initial
average temperature of the 2 separate solutions to get “ΔT,” and record it in Data Table 3.

28. Properly dispose of the solution in the calorimeter.

29. Rinse all equipment with distilled water in preparation for the third reaction. This includes the
foam cups, and the 25 mL graduated cylinder.

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Experiment Hess’s Law

Part 3: Third Reaction (NH3 + HCl)

3(aq) (aq) 4 (aq) NH + HCl NH Cl→

30. Place the calorimeter assembly into an empty coffee mug again to prevent it from tipping
over.

31. Use the clean 25 mL graduated cylinder to accurately measure 10.0 mL of 2M NH3 from the
dropper bottle. When nearing the 10.0 mL mark, add drops slowly to get the exact amount
of 10.0 mL.

32. Pour the 10.0 mL of NH3 into the bottom cup of the foam calorimeter.

33. Rinse and dry the graduated cylinder and accurately measure 10.0 mL of 2M HCl.

34. Pour the 10.0 mL of HCl into another foam cup and place the cup into a second empty coffee
cup to prevent it from tipping over.

35. Turn on the digital thermometer and place it into the NH3 solution. Wait at least 1 minute and
record the temperature of the solution in Data Table 5 of your Lab Report Assistant.

36. Remove the thermometer, rinse the tip with distilled water, dry it with a paper towel, and place
the thermometer into the HCl solution. Wait at least 1 minute and record the temperature of
the solution in Data Table 5.

37. Remove the thermometer, rinse the tip with distilled water, and dry it with a paper towel.

38. Pour the contents of one foam cup into the second one, combining the 2 solutions. Quickly
place the foam lid on top of the cup containing the combined solutions and insert the
thermometer through the hole in the lid. Be careful when inserting the thermometer to
ensure its pointed tip does not puncture the lower foam cup. Record the initial temperature
as quickly as you see it in Data Table 6 of your Lab Report Assistant.

39. Record the temperature that you see every 20 seconds for 5 minutes in Data Table 6.

40. Create a scatter plot using the data in Data Table 6. Plot time on the independent axis (x-axis)
and temperature on the dependent axis (y-axis). Add a linear trendline that represents
temperature change. Extend the trendline to the y-axis, time 0. The extrapolation at the
“0-second time” most likely represents the “highest temperature of the mixture.” Compare
the extrapolated temperature to the initial recorded temperature of the mixture and record
the greater temperature in Data Table 5 as the highest temperature of the mixture. The graph
should look similar to the sample cooling curve in Figure 3. Resize and insert an image of the
graph into Graph 3 of your Lab Report Assistant.

41. Subtract the greatest difference in temperature of the combined solutions from the initial
average temperature of the 2 separate solutions to get “ΔT,” and record it in Data Table 5.

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Experiment Hess’s Law

Cleanup:

42. Properly dispose of the solution in the calorimeter.

43. Rinse all equipment with distilled water. This includes the foam cups, and the 25 mL graduated
cylinder.

44. Dry and return all equipment to the lab kit for future use.

45. When you are finished uploading photos and data into your Lab Report Assistant, save and
zip your file to send to your instructor. Refer to the appendix entitled “Saving Correctly,” and
the appendix entitled “Zipping Files,” for guidance with saving the Lab Report Assistant in the
correct format.

Questions
A. What were the starting temperatures for each of the reactions? Use the scatter plots in Graph

1, Graph 2, and Graph 3 to support your answers.

B. Calculate the heat loss or heat gain of the 3 solution mixtures (qrxn). Then calculate the ΔH for
each reaction. Show your calculations.

C. Use Hess’s law to determine ΔH for the first 2 reactions and then add them together to
determine ΔH for the third reaction: NH3 + HCl → NH4Cl.

D. Compare the results of question “C” with the experimental results of the reaction:

NH3 + HCl → NH4Cl (calculate the percent difference).

E. Use the thermodynamic quantities given below to calculate the theoretical ΔH for this
reaction: NH3 + HCl → NH4Cl

• ΔH°f for NH3 (aq) = – 80.29 kJ/mol

• ΔH°f for HCl (aq) = – 167.2 kJ/mol

• ΔH°f for NH4
+

(aq) = – 132.5 kJ/mol

• ΔH°f for Cl- (aq) = – 167.2 kJ/mol

F. What was the percent difference of the various methods used when comparing the results of
Hess’s law method and the experimental results to the theoretical value?

G. Were the reactions performed in this laboratory exercise exothermic or endothermic? How
could you determine this?

H. Define Hess’s law. Did your experimental results support Hess’s law? Use your data to explain
your answer.

I. What are some possible sources of error in this experiment?

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Experiment Hess’s Law

Oxidation-Reduction
Activity Series
Hands-On Labs, Inc.
Version 42-0186-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will perform redox reactions and identify
oxidizing and reducing agents. You will investigate
the reactivity of metals and develop an activity
series.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Discuss oxidation-reduction reactions and identify oxidizing and reducing agents.

● Summarize the rules for assigning oxidation numbers.

● Describe single displacement reactions, spectators ions and activity series.

● Perform single displacement reactions on metals to develop an activity series.

● Write redox reactions based on experimental results.

● Apply the appropriate rules for assigning oxidation numbers.

Time Allocation: 2.5 hours

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Experiment Oxidation-Reduction Activity Series

Materials
Student Supplied Materials

Quantity Item Description
1 Pair of scissors
1 Pencil
1 Pocket knife, disposable nail file, or sandpaper
1 Roll of paper towels
1 Sheet of white paper
1 Timer or clock

HOL Supplied Materials

Quantity Item Description
1 Pair of gloves
1 Pair of safety goggles
1 Plastic tweezers
1 Test tube, 13 x 100 mm
1 Test tube cleaning brush
1 Well plate – 24
1 Experiment Bag: Oxidation/Reduction Activity Series:

1 – Copper (II) sulfate, 1 M – 3 mL in pipet
1 – Copper foil, 4 pieces in bag 2” x 3”
1 – Lead metal, 4 pieces in bag 2” x 3”
1 – Lead (II) nitrate, 0.2 M – 2 mL in pipet
1 – Mossy zinc, 6-8 pieces in bag 2” x 3”
1 – Silver nitrate, 0.1 M – 6 mL in dropper bottle
1 – Zinc nitrate, 2 M – 3 mL in pipet

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

www.HOLscience.com 3 ©Hands-On Labs, Inc.

Experiment Oxidation-Reduction Activity Series

Background
Chemical Reactions

An oxidation-reduction reaction, or redox reaction, is a chemical reaction that occurs when
electrons are transferred from one reactant to another. For example, hydrogen (H2) and fluorine
(F2) react to produce hydrogen fluoride (HF). In the reaction, there is a partial transfer of one
electron from the hydrogen atom to the fluorine atom, causing a chemical bond. The chemical
equation for the redox reaction is shown in Equation 1.

The oxidation number, often referred to as the oxidation state, represents the charge an atom
would have if electrons were completely transferred. In truth, hydrogen only partially donates an
electron to fluorine, but for purposes of understanding the movement of the electron, charges
can be assigned to each element. In hydrogen fluoride, the hydrogen atom has an oxidation
number of +1 because one negative electron was lost, and the fluorine atom has an oxidation
number of -1 because one electron was gained. It may be said that the hydrogen atom is in
the +1 oxidation state, while the fluorine atom is in the -1 oxidation state. For the purpose of
visualization, oxidation numbers can be written under each atom of the chemical equation, as
shown in Equation 2. Notice that the charge for each of the reactants, elemental hydrogen and
elemental fluorine, is 0.

Oxidation and Reduction

Oxidation is the loss of electrons by a substance undergoing a chemical reaction. During
oxidation, the oxidation number of the element increases and becomes more positive. Reduction
is the gain of electrons by a substance undergoing a chemical reaction. During reduction, the
oxidation number of the element decreases and becomes more negative. In Equation 2, hydrogen
is oxidized, it loses an electron, and its oxidation number increases; fluorine is reduced, it gains an
electron, and its oxidation number decreases. Hydrogen may be referred to as the reducing agent
because it donates electrons and reduces fluorine. Conversely, fluorine may be referred to as the
oxidizing agent because it accepts electrons and oxidizes hydrogen.

A mnemonic device for remembering whether electrons are lost or gained is “OIL RIG”:

Oxidation Is Loss, Reduction Is Gain.

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Experiment Oxidation-Reduction Activity Series

During a reaction, electrons are neither created nor destroyed; they are merely transferred
among atoms. Oxidation and reduction always occur together: oxidation does not occur without
reduction, and reduction does not occur without oxidation. Therefore, the sum of the oxidation
numbers left of the reaction arrow equals the sum of the oxidation numbers right of the reaction
arrow. In Equation 2, the oxidation numbers of the reactants total 0, and the oxidation numbers
of the products also total 0.

Figure 1. Walnuts are one of the most antioxidant-rich food sources available. © Pakhnyushcha

As a natural
consequence of cellular

metabolism, oxidative processes
cause damage to the cells and

tissues of the human body. Oxidation
is widely believed to be one of the

primary causes of aging. Foods such
as walnuts, artichokes, and blueberries

contain antioxidants, chemical substances
that inhibit oxidation. See Figure 1.

Scientists are currently working to fully
understand the role of oxidation in the

aging process and the effectiveness
of consuming antioxidant-rich

foods.

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Experiment Oxidation-Reduction Activity Series

Assigning Oxidation Numbers

Many chemical compounds are composed of more than one of the same type of atom, and the
oxidation numbers of each atom can be summed to find the total contribution to charge. For
example, the compound magnesium chloride (MgCl2) is composed of one magnesium ion and
two chloride ions. The chlorine ions each have an oxidation number equal to -1, which results in a
total contribution of -2. The magnesium atom has an oxidation number equal to +2. Therefore, the
MgCl2 formula unit has a total charge of zero. The oxidation numbers and the total contribution
to charge can be designated simultaneously. In the following, the formula is listed in the first line,
the oxidation numbers are listed in the second line, and the total contribution to charge is listed
in the last line:

Molten magnesium chloride (MgCl2) can be decomposed into the pure elements magnesium
(Mg) and chlorine (Cl) through electrolysis. Equation 3 shows the redox reaction. Notice that the
sum of the total contribution of charge for the reactants is 0 and the total contribution of charge
for the products is also 0.

The rules for assigning oxidation numbers follow. Although there are some exceptions, the rules
generally work well and should be applied in the order they are listed.

Rules for Assigning Oxidation Numbers

1. When a substance is in its elemental form (existing alone without bonds to other elements),
the oxidation number is 0.

2. A monoatomic ion has an oxidation number equal to its ionic charge.

3. A hydrogen atom is in the +1 oxidation state when it is combined with a nonmetal.

4. A hydrogen atom is in the -1 oxidation state when it is combined with a metal.

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Experiment Oxidation-Reduction Activity Series

5. The alkali metals of Group IA on the periodic table have the +1 oxidation state in a compound.

6. The alkaline earth metals of Group IIA on the periodic table have the +2 oxidation state in a
compound.

7. Nonmetals in Group VIIA on the periodic table often have the -1 oxidation state in a compound.

8. Oxygen usually has an oxidation number of -2 in compounds. Exceptions are rare but can
occur in compounds that contain more than one oxygen atom.

9. The sum of the oxidation numbers within a formula is equal to the overall charge of the
formula. (In a neutral molecule, such as CH4, no charge is written next to the molecule, and
the sum of the oxidation numbers is zero.)

10. The most electronegative element in a compound (the element that most attracts electrons)
will have a negative oxidation number.

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Experiment Oxidation-Reduction Activity Series

Redox and Single Displacement Reactions

In this lab, single displacement reactions will be performed to demonstrate the process of
oxidation and reduction. In a single displacement reaction, an element reacts with a compound
and part of the compound is released to become a free molecule. For example, when iron (Fe)
is placed in a solution of copper (II) sulfate (CuSO4), iron takes the place of copper to form iron
(II) sulfate (FeSO4) solution; pure copper (Cu) is released as a free element. Fe is oxidized and
acts as the reducing agent; Cu2+ is reduced and acts as the oxidizing agent. The sulfate ion (SO4)
is considered a spectator ion because it is neither oxidized nor reduced. Equation 4 shows the
chemical reaction with oxidation numbers and total charge contributions.

When an iron nail is placed in a beaker of CuSO4 solution, iron displaces copper, producing FeSO4
and pure Cu. The Cu accumulates on the outside of the nail, as shown in Figure 2.

Figure 2. Iron nail and copper (II) sulfate (CuSO4) solution. Iron reacts to displace copper in
CuSO4. Pure copper accumulates on the outside of the nail.

The reaction between iron and copper (II) sulfate occurs because iron metal is more easily oxidized
than copper. In other words, iron atoms are more apt to lose electrons, thus iron is more “active”
than copper and replaces copper ions in solution. Electrons move from iron atoms to copper ions
forming iron ions and copper atoms. These atoms have moved from a higher energy, less stable
location (the active iron atoms) to a lower energy, more stable location (the less active copper
atoms).

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Experiment Oxidation-Reduction Activity Series

Metals can be ordered from most active (easily oxidized) to least active (not easily oxidized) in a
list called an activity series. Figure 3 is an activity series for a small selection of metals.

Figure 3. Activity series of five metals.

Note that iron is listed higher than copper because it is more active and more easily oxidized.
Figure 3 confirms that a reaction occurs between solid iron and copper ions in solution. On the
other hand, iron is listed below chromium in the activity series. What would you expect if solid
iron was placed in a solution containing chromium ions? Based on what you know, will the iron
displace the chromium ions? That is, will iron atoms transfer electrons to chromium ions? In this
laboratory experiment, you will create an activity series based on your observations of chemical
reactions.

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Experiment Oxidation-Reduction Activity Series

Exercise 1: Describing an Oxidation-Reduction
Reaction
In this exercise, you will add a solution that contains silver ions to elemental copper to elicit
a redox reaction. You will observe the reaction and then write an equation that describes the
movement of electrons.

CAUTION! You must wear your gloves and goggles during both exercises.

Procedure

1. Gather a test tube, the silver nitrate (AgNO3) solution, the copper foil pieces (Cu), and the
plastic forceps (tweezers).

Note: Use plastic forceps, as metal forceps may react with the other metals in the experiment.

2. Record initial observations of the appearance of the AgNO3 and Cu in Data Table 1 in your Lab
Report Assistant.

Note: To observe the AgNO3, place a small drop on a folded piece of paper towel, plastic wrap, or a
plastic baggie. Then, discard in the trash. Be careful, as AgNO3 causes stains.

3. Use the plastic forceps to add 2 pieces of Cu to the test tube.

4. Add the entire contents of the AgNO3 bottle to the test tube. Discard the empty bottle in a
trash bin.

Note: Take care to avoid any skin contact with the silver nitrate.

5. Observe the reaction for about 1 minute. Describe the appearance of any solids, liquids, and
gases in the test tube. Record observations in Data Table 1.

6. Allow the reaction to continue for 30 minutes, and again record observations of the appearance
of liquids, gases, and solids in Data Table 1.

Note: As you wait, continue with the final steps of the exercise and complete the questions at the end
of the exercise.

7. In Data Table 1, write a balanced chemical equation for the redox reaction. Include oxidation
numbers and total charge contributions for the elements involved in the reaction.

8. Identify the element that is oxidized and the element that is reduced. Identify the spectator
ion, the ion that exists in the same form in both the reactants and the products. Record each
in Data Table 1.

9. Identify which element acted as the oxidizing agent, and identify which element acted as the
reducing agent. Record each in Data Table 1.

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Experiment Oxidation-Reduction Activity Series

Cleanup:

10. Properly dispose of the reactants.

11. Clean the test tube with soap and water, dry, and place the test tube back in your kit for future
use.

Questions
A. Define oxidation, reduction, and oxidation number. Describe how oxidation and reduction

affect the oxidation number of an element.

B. Define oxidizing agent, reducing agent, and spectator ion.

C. In the reaction of copper and silver nitrate, a new substance appeared in the test tube.
Describe the physical appearance of the substance and identify its chemical formula.

D. Given an activity series in which the most active metals are at the top of the list and the least
active metals are at the bottom of the list, would copper be listed above silver or would silver
be listed above copper? Support your answer with data from Data Table 1.

E. Solid copper sulfide and silver nitrate react to form copper (II) nitrate and solid silver sulfide.
Write a balanced chemical equation that describes the reaction. Identify the oxidation number
of each element in the reaction. (You do not need to include the total contribution of charge.)
Is this reaction a redox reaction or a non-redox reaction? Explain your answer.

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Experiment Oxidation-Reduction Activity Series

Exercise 2. Creating an Activity Series
In this exercise, you will make observations of copper, lead, and zinc and determine if a successful
chemical reaction has occurred. Use your observations to order the metals in an activity series.

CAUTION! Make sure that you are wearing gloves and goggles. In this lab, you will handle solid
lead, which can be hazardous. Handle with caution, and keep children and pets away from this and
any experiment.

Procedure

1. Locate a 24 well plate. See Figures 4 and 5 for examples of labeling.

Note: Row and column labels for each well are etched in the plastic of the well plate, as shown in
Figure 4. Labeling a white piece of paper and placing it under the well plate can help you accurately
add the chemicals to each well. See Figure 5 for an example of labeling. In this lab, you will only need
to label wells A1, A2, B1, B2, C1, and C2.

Figure 4. Labeled 24-well plate.

Figure 5. Examples of well plate labels. A. Drawing rows and columns of the well plate. B. Well
plate over a hand-labeled sheet of white paper.

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Experiment Oxidation-Reduction Activity Series

2. Locate the copper foil pieces (Cu) and plastic forceps.

3. Place a piece of copper in well A1 and well A2 of the well plate.

4. Locate the lead (II) nitrate (Pb(NO3)2) pipet, and snip the tip off with scissors. Add 15 drops of
Pb(NO3)2 to well A1.

5. Observe well A1 for about 1 minute and record any immediate signs of a chemical reaction in
Data Table 2 in your Lab Report Assistant.

Note: Some examples of indicators of a chemical reaction include darkening of the metal and
formation of bubbles. A chemical reaction may or may not occur.

6. Note the time, as observations will again be made after 30 minutes.

7. Locate the zinc nitrate (Zn(NO3)2) pipet, and snip the tip off with scissors. Add 15 drops of
Zn(NO3)2 to well A2.

8. Observe well A2 for about 1 minute and record any immediate signs of a chemical reaction in
Data Table 2.

9. Gather 2 lead squares (Pb) and a paper towel. Use the forceps to place the 2 pieces of lead
on the paper towel.

10. Locate a pocket knife, nail file, or piece of sandpaper. Carefully scrape the surface of the lead,
removing any rust that may have accumulated.

CAUTION! Always keep the lead on top of the paper towel, as lead shavings can be hazardous.
Scrape the surface of the lead until it looks shiny, as shown in Figure 6. Gently wipe each square on
the paper towel to remove any lead shavings.

Figure 6. Lead square (left) and prepared lead that has been cleaned of rust (right).

11. Place the prepared lead squares in wells B1 and B2 of the well plate.

12. Carefully dispose of the paper towel and all lead particles. If using a nail file or piece of
sandpaper, dispose of the item in a trash bin. If using a pocket knife, thoroughly wash the
knife with soap and tap water. Alternatively, you may take the lead to a waste disposal station.

www.HOLscience.com 13 ©Hands-On Labs, Inc.

Experiment Oxidation-Reduction Activity Series

13. Remove your gloves, carefully turning them inside out and placing them in the trash. Wash
your hands with soap and water; dry them with paper towels. Dispose of the paper towels in
the trash bin.

14. Put on a new pair of gloves before continuing the exercise.

15. Locate the copper (II) sulfate (CuSO4) pipet and snip the tip off with scissors. Add 15 drops of
CuSO4 to well B1. Record observations after about 1 minute in Data Table 2.

Note: Forceps may be used to lift the lead from the copper solution for closer examination. Rinse the
forceps with tap water. See Figure 7.

Figure 7. Using the forceps to examine lead.

16. Add about 15 drops of Zn(NO3)2 to well B2. Record observations after about 1 minute in Data
Table 2.

17. Locate 2 pieces of mossy zinc (Zn). Place a piece of zinc in well C1 and well C2.

18. Add about 15 drops of CuSO4 to well C1 and record observations after about 1 minute in Data
Table 2.

19. Add about 15 drops of Pb(NO3)2 to well C2 and record observations after about 1 minute in
Data Table 2.

20. After 30 minutes have passed, again record observations for all reactions.

21. Use the observations recorded in Data Table 2 to determine whether a chemical reaction
occurred in each instance. Use Data Table 3 of your Lab Report Assistant to record “yes” or
“no” to indicate whether or not a reaction occurred.

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Experiment Oxidation-Reduction Activity Series

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Experiment Oxidation-Reduction Activity Series

22. In Data Table 3, for each instance that you recorded “yes,” write a balanced chemical equation
that represents the reaction. Include oxidation numbers and the total contribution of charge
for the elements involved in the reaction underneath each element or compound (as
demonstrated in the Background). For each instance you recorded “no,” write “no reaction”
in place of the chemical equation.

Cleanup:

23. Properly dispose of remaining reactants and pipets.

24. Wash and dry the equipment and return to the kit for future use.

Questions
A. List each of the metals tested in Exercise 2. Indicate the oxidation number when each element

is pure and the oxidation number when each element is in a compound.

B. Which of the metals in Exercise 2 was the strongest oxidizing agent? Was there an instance
when this metal also acted as a reducing agent? Explain your answer using data from Data
Table 3.

C. Which of the metals in Exercise 2 was the strongest reducing agent? Was there an instance
when this metal also acted as an oxidizing agent? Explain your answer using data from Data
Table 3.

D. How does ease of oxidation correlate with activity? Do highly active metals tend to donate
electrons or accept electrons from other metals?

E. Create an activity series for copper, lead, and zinc. Place the most active metal at the top of
the list.

Colligative Properties and
Osmotic Pressure
Hands-On Labs, Inc.
Version 42-0149-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will explore the colligative properties of freezing
point, boiling point, and osmotic pressure in solutions.
You will observe changes to these colligative
properties by adding a controlled amount of solute
to a solution. You will define colligative properties as
well as discuss membrane permeability and osmotic
pressure.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Explain the four colligative properties of a solution: vapor pressure, freezing point depression,
boiling point elevation, and osmotic pressure.

● Describe the process of osmosis and define osmotic pressure.

● Describe the mathematical relationship between osmotic pressure, molarity, and temperature.

● Observe and describe the process of osmosis through a semipermeable membrane.

● Determine the molecular mass of a compound using osmotic pressure data.

● Examine how the freezing and boiling points of solutions change as a result of the amount of
solute present.

Time Allocation: 2.5 hours

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Experiment Colligative Properties and Osmotic Pressure

Materials
Student Supplied Materials

Quantity Item Description
1 Aluminium pie pan
1 Bottle of distilled water
1 Crushed ice
1 Glass bowl
1 Light corn syrup, 2 oz.
1 Matches or lighter
1 Measuring spoon, 0.5 tsp
1 Plastic drinking cup
1 Roll of paper towels
1 Salt
2 Small rubber bands
1 Source of tap water
1 Timer, clock, or watch with second hand

HOL Supplied Materials

Quantity Item Description
1 Aluminum cup, 2 oz
1 Burner fuel
1 Burner stand
1 Dialysis tubing, 6 in
1 Digital scale
1 Funnel
1 Glass beaker, 250 mL
1 Graduated cylinder, 25 mL
1 Pair of safety gloves
1 Pair of safety goggles
1 Thermometer

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

www.HOLscience.com 3 ©Hands-On Labs, Inc.

Experiment Colligative Properties and Osmotic Pressure

Background
Solutions and Vapor Pressure

A solution consists of a solvent (the dissolving medium) and one or more solutes (substances
dissolved into the solution). Most of the properties of a solution are dependent on the properties
of the solute. However, colligative properties are four properties that are not dependent on the
identity of the solute, but dependent on the number (concentration) of solute particles that are
present in the solution. The four colligative properties of solutions are:

1. Vapor pressure lowering

2. Boiling point elevation

3. Freezing point depression

4. Osmotic pressure

Consider a liquid that is put into an open container at room temperature, such as water in a beaker.
After a period of time, all the liquid will evaporate. Evaporation occurs because the molecules in a
liquid escape the surface of the liquid to form a gas, or vaporize due to the pressure of the moving
molecules. Any volatile substance, or substance that will evaporate, has a specific vapor pressure.
The rate of evaporation is dependent upon environmental conditions such as air temperature, air
pressure, and humidity.

Solutions that contain a solvent and non-volatile solute particles have lower vapor pressures than
the pure solvents under the same conditions. See Figure 1. The presence of the non-volatile
solute particles reduces the likelihood that the solvent will evaporate. Each component of the
solution, the non-volatile solute and the solvent, lowers the vapor pressure of the other.

Figure 1. Adding a solute to a solution decreases the vapor pressure.

Boiling point is defined as the temperature at which the vapor pressure of a liquid is equal to the
atmospheric pressure on the liquid. As solutions (with non-volatile solutes) have lower vapor
pressures than pure solvents at the same temperature, a higher temperature is required to boil
the solution than the pure solvent. The elevated boiling points are directly related to both the
number (concentration) of solute particles and the nature of the solvent.

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Experiment Colligative Properties and Osmotic Pressure

The presence of solute also interferes with the formation of a solid as the solution cools. As
a result, solutions with non-volatile solutes have lower freezing points than pure solvents. Just
as for the boiling point, the lowered freezing points are directly related to both the number
(concentration) of solute particles present and the nature of the solvent. See Figure 2.

Figure 2. Adding a non-volatile solute to a volatile solvent will decrease the vapor pressure of
the solvent. © Volker Rauch

An antifreeze is an additive which

lowers the freezing point of a water-
based liquid. An antifreeze mixture is

used to achieve freezing-point depression
for cold environments and also achieves
boiling-point elevation (“anti-boil”) to

allow higher coolant temperature.

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Experiment Colligative Properties and Osmotic Pressure

Osmosis and Osmotic Pressure

If two solutions of different concentrations are separated by a semi-permeable membrane, solvent
from the less concentrated solution will pass through the membrane into the more concentrated
solution, diluting it. This process will continue until the concentration of solute is the same on
each side of the membrane. This phenomenon is called osmosis.

In biological systems, if a cell is placed into a salt solution in which the salt concentration in the
solution is lower than in the cell, the solution is said to be hypotonic. Water will move from the
solution into the cell and the cell will expand to the point where it may burst. On the other hand,
if a cell is placed into a salt solution in which the salt concentration in the solution is higher than
in the cell, the solution is said to be hypertonic. Water will move out of the cell and into the
solution, causing the cell to shrink.

The pressure that must be applied to stop the movement of solvent through the membrane is
the osmotic pressure of the solution. The amount of pressure needed to stop the movement
of solvent is related to the concentration of solute particles in the two solutions. It is important
to note that a semipermeable membrane can be specifically selected to allow passage of some
small particles. Likewise, biological membranes may allow passage of certain small molecules and
not others. See Figure 3.

Figure 3. Osmotic pressure, due to a semipermeable membrane, will cause molecules to move
one direction in an attempt to maintain equilibrium on both sides of the membrane. © Dreamy

Girl

Osmosis is the basis
for hemodialysis, where

the blood of patients with
kidney malfunction is filtered

to remove waste products
normally removed by the

kidneys.

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Experiment Colligative Properties and Osmotic Pressure

Colligative properties and molar mass calculations

Historically, certain colligative properties: freezing point depression, boiling point elevation,
and osmotic pressure have been used to determine the molar mass of a solute. Now there are
instrumental methods to determine this. Of these three, osmotic pressure is the most sensitive
and gives the best results. The molar mass of a solute in a solution can be found through the
following equation:

Sample Problem:

0.125 grams of a protein were dissolved in 100 mL of water at 25°C. The solution has an osmotic
pressure of 5.15 mm Hg. What is the molar mass of the protein?

As the gas constant, R, requires atmospheres as pressure units, we first have to convert 5.15 mm
Hg to atmospheres, and the temperature from Celsius to Kelvin:

The values can then be placed into the equation:

Solve for M (Molarity = mol/ L):

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Experiment Colligative Properties and Osmotic Pressure

0.125 grams of a protein were dissolved in 100 mL of water. Therefore, the mass of the protein
per liter of solution is:

We now know that the concentration of the solution (2.77 x 10-4 M) was produced by adding 1.25
grams of protein to 1 liter of water. Therefore, 2.77 x 10-4 mol of the protein has a mass of 1.25
grams. The molar mass of the protein is calculated as shown below:

In the following exercises, you will explore three of the four colligative properties of a solution:
freezing point, boiling point, and osmosis.

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Experiment Colligative Properties and Osmotic Pressure

Exercise 1: Colligative Properties – Osmosis
In this exercise, you will observe the results of osmosis through a semipermeable membrane.

Note: Completely read all instructions and assemble all equipment and supplies before beginning
work on this experiment.

Procedure

1. Gather the bowl, distilled water, and the piece of dialysis tubing.

Note: Treat the dialysis tubing gently as you only have 1 piece available.

2. Soak the dialysis tubing for 5 minutes in a bowl filled with distilled water. Ensure the tubing is
completely submerged.

3. While the dialysis tubing is soaking, gather the funnel, 2 rubber bands, the light corn syrup,
and some paper towels.

4. After the tubing has soaked 5 minutes, remove it from the water and put it on a paper towel.

5. Pour the water down the drain and rinse the bowl with distilled water.

6. Refill the bowl halfway with distilled water.

7. To open the dialysis tubing, carefully rub the dialysis tubing between your fingers until the
middle of the tubing opens. See Figure 4.

Figure 4. Carefully rubbing the dialysis tubing until the middle of the tubing opens.

8. Carefully close off 1 end of the dialysis tubing by folding the end and then tying it off using a
cut rubber band. See Figure 5.

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Experiment Colligative Properties and Osmotic Pressure

Note: You may want to cut the rubber band and tie off the tubing or you may want to just wrap the
rubber band around the tubing several times. Test the closure by adding distilled water to the tube.
Repeat the tying process if the closure leaks. Empty the tubing of distilled water after testing.

9. Insert the funnel into the open end of the dialysis tubing and hold the tubing snugly to the
funnel during the next step. See Figure 5.

Figure 5. Inserting the funnel into the open end of the dialysis tubing.

10. Use the funnel to carefully fill the dialysis tubing ⅓ full with the light corn syrup. Remember
that the corn syrup is viscous, so it will move into the tubing slowly. Try to avoid getting any
light corn syrup on the outside of the tubing.

11. Carefully push out most of the air that is inside the tubing above the corn syrup. Then close
off the open end with the rubber band (as done previously). If any corn syrup does get onto
the tubing, after folding over the other side of the dialysis tubing and tying it off with a rubber
band, gently rinse the outside of the tubing with distilled water.

12. Gently pat the tubing dry with a paper towel. Place the plastic cup on the digital scale and
tare the scale. Then, place the dialysis tubing with the light corn syrup inside the plastic cup
to determine the mass of the dialysis tubing/corn syrup. Record the mass in Data Table 1 of
your Lab Report Assistant at time “0.”

13. Place the tubing filled with light corn syrup into the distilled water in the cup or bowl. Ensure
that the tubing is completely submerged. Allow the tubing to remain in the water for 30
minutes.

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Experiment Colligative Properties and Osmotic Pressure

14. After 30 minutes, remove the tubing from the water. Gently pat the tubing dry with a paper
towel and determine the mass of the dialysis tubing/corn syrup as done previously in Step 12.
Record the mass in Data Table 1.

15. Place the tubing filled with light corn syrup back into the distilled water in the cup or bowl.
Ensure that the tubing is completely submerged. Allow the tubing to remain in the water for
another 30 minutes.

Note: Use the same distilled water in the bowl, DO NOT replace the water.

16. After 30 minutes, remove the tubing from the water. Gently pat the tubing dry with a paper
towel and determine the mass of the dialysis tubing/corn syrup as done previously in Step 12.
Record the mass in Data Table 1.

Cleanup:

17. Pour the liquids down the drain, wash and rinse the bowl, and place the used dialysis tubing
into the garbage.

Questions
A. In your experiment, is the light corn syrup in the dialysis tubing hypertonic or hypotonic to

the water?

B. 0.302 grams of an antibiotic was dissolved in 500 mL of water at 23.6°C. The solution has an
osmotic pressure of 8.34 mm Hg. What is the molar mass of the antibiotic? Show your work.

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Experiment Colligative Properties and Osmotic Pressure

Exercise 2: Colligative Properties – Freezing Point
In this exercise, you will determine the freezing point of tap water and 2 more solutions with
varying amounts of salt (a nonvolatile solute) added to them.

Note: Completely read all instructions and assemble all equipment and supplies before beginning
work on this experiment.

Procedure

1. Gather the following equipment and set it up near a source of tap water: the large plastic cup,
crushed ice (made from tap water), measuring spoon for a 0.5 teaspoon (~2.5 mL), salt, digital
scale, thermometer, stopwatch, and 25 mL graduated cylinder.

2. Place 100 g of ice in the plastic cup.

● Place a cup on the digital scale and tare the cup.

● Carefully place enough ice in the cup, by adding small pieces, until you get 100.0 g. See
Figure 6.

Note: If the digital scale in your kit measures a maximum of 100 g, you will need to weigh the ice in
two batches of 50 g. Remember to tare the scale with the empty cup before each 50 g measurement.

Figure 6. Ice in cup

3. Ensure that the thermometer and stopwatch are nearby.

4. Pour 100 mL of tap water into the cup, using the graduated cylinder to pour the water in 25
mL at a time.

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Experiment Colligative Properties and Osmotic Pressure

Note: Though some minerals are in the tap water, if you are using ice, the tap water from the ice will
get into the water. Since colligative properties are relative to the amount of solute in the solution, you
can use tap water and just compare the data from the increase in solute relative to the data from tap
water alone.

5. Begin timing for this exercise.

6. Every 30 seconds, record the temperature of the water near the ice.

a. Stir the solution continuously so the temperature is even throughout the cup.

b. When measuring the temperature, keep the thermometer near the ice in the cup,
remembering that ice floats in water.

7. Record the temperature at each time point (every 30 seconds) in Data Table 2 of your Lab
Report Assistant.

8. After recording is completed, empty the water and ice into a sink.

9. Fully dry the plastic cup and repeat Step 2, weighing 100 g of ice in the cup.

10. Pour 0.5 tsp (2.5 mL) of salt into the plastic cup over the ice and stir. Ensure that when you
stir the mixture, you also stir near the bottom of the cup to dissolve all of the salt into the
solution.

11. Repeat Steps 4-9. When the water is poured into the cup, stir the water and salt mixture.
Record all data in Data Table 2.

12. Rinse the cup well with tap water a few times, pouring out the rinse water each time.

13. Fully dry the plastic cup and repeat step 2, weighing 100 g of ice in the cup.

14. Pour 1.0 tsp (5 mL) of salt into the plastic cup over the ice and stir. Ensure that when you stir
the mixture, you also stir near the bottom of the cup to dissolve all of the salt into the solution.

15. Repeat Steps 4-9. When the water is poured into the cup, stir the water and salt mixture.
Record all data in Data Table 2.

16. Make a graph of your data for all 3 columns of Data Table 2. Insert the graph in Data Table 3
of your Lab Report Assistant.

● Plot temperature on the y-axis and time on the x-axis

● Input your data from the tap water and saltwater solutions

● The 5 consecutive readings indicate the freezing point for the solution

Questions
A. Describe the three freezing points. Is there a relationship between the amount of solute in the

solution and the freezing temperature?

B. What are some practical applications of freezing point depression?

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Experiment Colligative Properties and Osmotic Pressure

Exercise 3: Colligative Properties: Boiling Point
In this exercise, you will determine the boiling point of tap water, and 2 solutions with tap water
and varying amounts of salt added to them.

Procedure

1. Gather the following materials and set up near a source of tap water: 250 mL beaker, burner
stand, burner fuel, matches, aluminum cup, pie pan, salt, measuring spoon for a 0.5 teaspoon
(2.5 mL), thermometer, stopwatch, and 25 mL graduated cylinder.

2. Using the 25 mL graduated cylinder, measure 100 mL of tap water (in 25 mL increments), and
pour the water into the 250 mL beaker.

3. Assemble the burner setup and light the fuel, as shown in Figure 7.

● Place an aluminum pie plate on a solid work surface away from flammable objects. Set the
burner stand towards the back of the pie plate.

● Uncap the burner fuel and set cap aside. Place the burner fuel on the pie plate just in front
of the stand.

● Use matches or a lighter to ignite the fuel. BE CAREFUL the flame may be nearly invisible.

● Gently slide the fuel under the stand without disturbing the beaker.

● The small, 2 oz. aluminum cup will be placed over the fuel to extinguish the flame. Set the
aluminum cup next to the burner setup so you are ready to extinguish the flame at any
point.

Figure 7. Burner fuel setup.

4. When the water is at a rolling boil, stir the water with the thermometer and take the
temperature near the middle of the 250 mL beaker.

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Experiment Colligative Properties and Osmotic Pressure

5. Record the temperature in Data Table 4 of your Lab Report Assistant, in the Temp (°C) of
Control column.

6. Use the small, 2 oz. aluminum cup to extinguish the burner fuel flame. See Figure 8.

● Do not touch the metal stand or the beaker; they may be hot.

● Carefully slide the burner fuel canister out from underneath the burner stand. The sides
of the burner fuel canister will be warm, but not hot.

● Place the aluminum cup directly over the flame to smother it. The cup should rest on top
of the fuel canister, with little or no smoke escaping. Do not disturb the burner stand and
beaker; allow everything to cool completely.

Figure 8. Extinguishing burner.

7. Once the beaker has cooled, pour the water from the beaker down the drain.

8. Fully dry the beaker and repeat Step 2.

9. Pour 0.5 tsp (2.5 mL) of salt into the beaker with water in it, and stir the mixture well with the
measuring spoon or thermometer until the salt has dissolved into the mixture.

10. Repeat Steps 3-7. Record data in Data Table 4.

11. Rinse the beaker well with tap water a few times, pouring out the rinse water each time.

12. Fully dry the beaker and repeat Step 2. Do not light the burner yet.

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Experiment Colligative Properties and Osmotic Pressure

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Experiment Colligative Properties and Osmotic Pressure

13. Pour 1.0 tsp (5 mL) of salt into the beaker with water in it, and stir the mixture well with the
measuring spoon or thermometer until the salt has dissolved into the mixture.

14. Repeat steps 3-7. Record data in Data Table 4.

Cleanup:

15. Wash the beaker with soap and water.

16. Return all items to the kit for future use.

Questions
A. Compare the three boiling points. Is there a relationship between the amount of solute in the

solution and the boiling temperature?

B. What are some practical applications of boiling point elevation?

Antacid Analysis and
Titration
Hands-On Labs, Inc.
42-0139-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will use a back-titration technique to determine
the amount of acid that a commercial antacid is
capable of neutralizing. You will be introduced to
experimental controls, and use a control to validate
the antacid neutralization analysis.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Identify and explore the causes of acid reflux disease.

● Investigate the relationship between antacid and gastric acid and define how antacids
neutralize gastric acid.

● Define titration, equivalence point, and pH indicator.

● Compare and contrast titrations and back titrations.

● Review back titration calculations and explain how control experiments are used to support
experimental results.

● Perform a titration, back titration, and control experiment.

● Determine how much acid an antacid is able to neutralize.

Time Allocation: 2.5 hours

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Experiment Antacid Analysis and Titration

Materials
Student Supplied Materials

Quantity Item Description
1 Bottle of distilled water
1 Dish soap
1 Metal spoon
1 Pair of scissors
1 Roll of paper towels
2 Sheets of white paper
1 Source of tap water

2-6 Thick textbooks

HOL Supplied Materials

Quantity Item Description
1 Digital scale
1 Glass Beaker, 100 mL
1 Graduated cylinder, 10 mL
1 Pair of gloves
1 Pair of safety goggles
1 Short stem pipet
1 Syringe, 10 mL
1 Stopcock
1 Test tube cleaning brush
1 Test tube clamp
1 Experiment Bag: Antacid Analysis and Titration

2- HCl, 1 M, 30 mL in dropper bottle
1- Phenolphthalein solution, 1% – 0.5 mL in pipet
2- Sodium hydroxide, 1 M – 30 mL in dropper bottle
2- Antacid tablets

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Antacid Analysis and Titration

Background
Acid Reflux Disease

If you have ever wondered what antacids are, or more specifically the chemistry behind how
an antacid works, you are not alone. Antacids are used to neutralize gastric acid, a substance
secreted in the stomach to promote the digestion of food. Gastric acid is produced and secreted
by specialized glands in the stomach, where it functions to break down the food we consume into
smaller nutrient particles so they can be absorbed by the small intestine. Gastric acid is composed
primarily of hydrochloric acid (HCl), glycoproteins, and enzymes, and has a pH close to 2.0. The
stomach is lined with mucus, a natural secretion that withstands and protects the stomach from
direct contact with the otherwise corrosive HCl. In a healthy digestive system, the gastric acid
remains in the stomach. See Figure 1.

Figure 1. Digestive System. © Leonello Calvetti

The lower end of the esophagus, the muscular tube that transports food from the mouth to
the stomach, is surrounded by a ring of muscles known as the lower esophageal sphincter. The
lower esophageal sphincter acts to prevent the stomach contents from moving upward into the
esophagus. When this sphincter malfunctions or is otherwise compromised, gastric acid refluxes
(moves back) into the esophagus resulting in acid reflux disease. Acid reflux disease causes
inflammation and irritation of the esophageal lining. See Figure 2.

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Experiment Antacid Analysis and Titration

Figure 2. Gastric acid refluxes past the lower esophageal sphincter (as noted by the two arrows
in the top image) and enters the esophagus, causing what is commonly known as acid reflux

disease. © Alila Sao Mai

Antacid Neutralization and Titration

While there are many ways to treat acid reflux disease, including prescription drugs, surgery, and
diet modifications, the initial treatment for controlling the symptoms of acid reflux disease is
through the use of over-the-counter medications, including antacids. Antacids are basic substances
that neutralize, or raise the pH, of gastric acid (primarily HCl). In a neutralization reaction, the
acid and the base first dissociate in solution, producing hydrogen (H+) and hydroxide (OH-) ions
respectively, which then react to produce a salt and water. Commercial antacids contain a wide
variety of basic substances as their active ingredient, including aluminum hydroxide (Al(OH)3),
magnesium hydroxide (Mg(OH)2), sodium bicarbonate (NaHCO3), and calcium carbonate (CaCO3).
See Figure 3.

Figure 3. Antacid neutralization reactions. (Top Reaction) Antacid neutralization reaction with
aluminum hydroxide (Al(OH)3) as the active ingredient. (Bottom Reaction) Antacid neutralization
reaction with calcium carbonate (CaCO3) as the active ingredient. Note that both reactions react

with the HCl to form a salt and water, and in the case of calcium carbonate, a gas (CO2).

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Experiment Antacid Analysis and Titration

An antacid’s effectiveness is based both qualitatively, by the physical relief it provides, and
quantitatively, by calculating the amount gastric acid neutralized by the antacid. The technique
used to calculate the amount of gastric acid neutralized by an antacid is titration, or more
specifically, back titration. Titration is a direct, quantitative, volumetric technique, where a
solution of a known concentration (titrant) is added to a solution of an unknown concentration
(analyte) until the equivalence point is reached. The equivalence point of a titration, also known
as a stoichiometric point, is the moment in a titration where exactly enough titrant has been
added to completely react with the analyte.

A back titration is an indirect, quantitative volumetric technique where a known quantity
of reagent is added to a known volume and concentration of analyte, and allowed to react. It
is expected that the reaction is not complete and some analyte remains in the solution. The
amount of analyte remaining is determined in a second step, by a titration reaction. A solution
of known concentration (titrant) is added to the solution until the equivalence point is reached,
which is indicated by a change in color. Consider, for example, a back titration to determine the
amount of gastric acid neutralized by an antacid with the active ingredient aluminum hydroxide
(Al(OH)3). A known mass of the Al(OH)3 containing antacid is mixed with an excess known volume
and concentration of HCl and allowed to react. The remaining HCl, which was not neutralized by
the antacid, is then titrated with a known concentration of NaOH until the equivalence point is
reached. In this example, the antacid is the substance, the HCl is the analyte, and the NaOH is
the titrant. The most effective antacid will leave the fewest HCl molecules after the reaction, and
require less of the NaOH in the back titration. This back titration method can easily be used to
compare the effectiveness of different antacids. The most effective antacid will leave the fewest
HCl molecules after the reaction and require less of the NaOH solution in the back titration step.

In either a direct titration or indirect titration, the equivalence point can be identified through
use of a pH indicator. A pH indicator is a substance that changes color when the pH of a
solution changes, allowing scientists to qualitatively measure the moment when the analyte has
completely reacted with the titrant. A common indicator for a titration between a weak acid
and a strong base is phenolphthalein. Phenolphthalein is a pH indicator, which turns bright-pink
in solutions with a pH of 8.2 or higher. Thus, equivalence points in titrations are marked by the
analyte changing color from colorless to bright pink. See Figure 4 for a schematic representation
of the back titration process.

There are many commercial
advertisements for both

prescription and over-the-counter
drugs to help with symptoms of

gastroesophageal reflux disease (GERD).
GERD is simply acid reflux disease that
occurs chronically, resulting in similar

symptoms and treatment as acid
reflux disease.

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Experiment Antacid Analysis and Titration

Figure 4. Schematic representation of the back titration process.

Calculating Results

The quantity of HCl neutralized by the antacid is calculated indirectly by 1) calculating the amount
of HCl present in the initial sample and 2) calculating the amount of HCl neutralized by the NaOH
in the back-titration step. The difference between these two is the amount of HCl neutralized by
the antacid.

For example, calculate a back titration with 0.5 g of antacid (Al(OH)3), 20 mL of 1.5M HCl as the
analyte, and NaOH with a concentration of 1.0M as the titrant. To reach the equivalence point, 27
mL of NaOH was required.

Step 1) Calculate the initial amount of HCl available for neutralization by the antacid.

Note: The molecular weight of HCl is calculated by adding the molecular weights of the two elements
in the compound: H + Cl (1.008g + 35.45g), thus 1 mole of HCl is equal to 36.46g.

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Experiment Antacid Analysis and Titration

When an HCl solution is neutralized to the stoichiometric point with NaOH, the number of moles
of NaOH required to reach the stoichiometric point is equal to the amount of moles of HCl present.

Step 2) Calculate the number of moles of NaOH required to reach the stoichiometric point
(neutralize the excess HCl) after initial neutralization with the 0.5 g of Al(OH)3.

Step 3) Calculate the amount of HCl neutralized by the 0.027 moles of NaOH.

Step 4) Calculate the amount of HCl neutralized by the antacid.

Step 5) Calculate the gram per gram neutralization of HCl by the antacid.

3
3

0.11g HCl = 0.22 g HCl neutralized / 1g Al(OH)
0.5g Al(OH)

To double check the results of the calculations and to confirm the results of the back titration,
a control experiment may be run. In a control experiment the variable tested (the antacid) is
removed from the experiment as a tool to quantitatively confirm that results of the experimental
design were set up to evaluate a single variable. A control experiment for the back titration of
an antacid is to perform a titration between 20 mL of 1.5M HCl and 1.0 M NaOH, removing the
antacid from the experiment. A positive confirmation in the control experiment would be for the
moles of NaOH required to reach stoichiometric quantities to be equal to the initial number of
moles of HCl present in the experiment. This result in the control titration would confirm that
the antacid was indeed neutralizing some of the HCl in the test, since without the antacid, more
NaOH is needed for the neutralization.

For example, assume that in a control experiment, 30 mL of 1.0M NaOH was required to reach
stoichiometric quantities when titrated into 20 mL of 1.5M HCl.

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Experiment Antacid Analysis and Titration

As shown in the above equations, in this example control experiment, the moles of NaOH required
to reach stoichiometric quantities and neutralize the HCl is equal to the number of moles of HCl
present in 20 mL of 1.5M HCl. Additionally, the number of grams of HCl neutralized by 30 mL of
1.0M NaOH is equal to the total number of grams of HCl initially present in the back titration
experiment. Furthermore, as the back titration required 27 mL of 1.0M NaOH to neutralize the
excess HCl, following the initial neutralization with the antacid, the difference in NaOH volume
(30 mL – 27 mL) between the back titration and control experiment should equal a neutralization
of 0.11 g HCl.

As the calculations show, the control experiment verifies that the back titration was successful in
quantitatively determining the amount of HCl neutralized by 0.5 g of antacid (Al(OH)3).

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Experiment Antacid Analysis and Titration

Exercise 1: Back Titration of Antacid Neutralization
In this exercise, you will perform a back titration to determine the amount of acid that a commercial
antacid is able to neutralize.

Procedure

Note: Please read all steps and safety information before starting the procedure.

1. Gather the test tube holder, small stopcock, 10-mL syringe (titrator), and either 2 thick
textbooks and the lab kit box or 5-6 thick textbooks. See Figure 5.

Figure 5. Titrator and small stopcock.

2. Remove the plunger from the titrator and place it back in your kit.

3. Attach the stopcock to the tip of the titrator by placing the larger, clear, plastic end of the
stopcock into the tip of the titrator and then twist the stopcock into place. The stopcock
should fit tightly into the titrator so that the liquid will not leak. See Figure 6.

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Experiment Antacid Analysis and Titration

Figure 6. Fitting the stopcock into the titrator.

4. Stack the 5 textbooks or stack 2 textbooks on top of the lab kit box.

5. Clamp the test tube holder around the middle of the titrator and slide the long end under the
top 2 books in the stack. Place a sheet of white paper next to the bottom of the stack and set
the 100-mL beaker on the sheet of white paper. The end of the stopcock should be located
near the top of the beaker, approximately 1 cm above to 1 cm below the top of the beaker.
See Figure 7.

Figure 7. Titration setup. Note the location of the end of the stopcock. It is important that the
placement of the titrator allows for the white knob to be easily adjusted. If this is not the case,
then either adjust the location of the books in the stack or slightly adjust where in the test tube

clamp the titrator is located.

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Experiment Antacid Analysis and Titration

6. Use the pipet to fill the titrator with 7 – 9 mL of distilled water.

Note: You must use distilled water for this step and not tap water.

7. Using both hands, one on the titrator and one on the stopcock, practice releasing water from
the titrator into the beaker. The goal is to be comfortable releasing only one drop at a time
from the titrator. See Figure 8.

Figure 8. Proper hand positioning for titration. When the open circle is facing you, the titrator is
closed, when the open circle is directly under the titrator spout, the titrator is open and liquid

will flow.

8. When you are comfortable using the titrator, pour the water in the beaker down the drain,
remove the titrator from the test tube clamp, and remove the stopcock from the titrator.
Thoroughly dry each of these 3 items with paper towels.

9. When all items are completely dry, reassemble the titration setup, as shown in Figure 7.

10. Put on your safety gloves and goggles.

11. With the stopcock in the closed position, fill the titrator with 9 – 10 mL of the 1.0M NaOH.

12. Move the beaker away from the titrator and place a crumpled paper towel directly below the
titrator.

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Experiment Antacid Analysis and Titration

13. Using the stopcock, allow a few drops of the NaOH to flow through the titrator into the paper
towel. This will fill the tip of the titrator with NaOH solution and remove any air bubbles from
the titrator.

14. Place the paper towel with the NaOH drops into the trash.

15. On a sheet of clean, white paper, use a clean spoon to crush an antacid tablet into a fine
powder. See Figure 9.

Figure 9. Crushing and antacid tablet. A. Use the back of a spoon to carefully crush the tablet. B.
Tablet should be crushed into a fine powder, as shown.

16. Place the clean, dry beaker on the scale and tare the scale so it reads 0.0 g. See Figure 10.

Figure 10. Tared scale with 100-mL glass beaker.

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Experiment Antacid Analysis and Titration

17. Use the clean spoon to carefully weigh 0.5 g of the crushed antacid tablet into the glass
beaker on the tared scale. Record this mass of the crushed antacid in Data Table 1 in your Lab
Report Assistant.

Note: Save the remaining crushed antacid powder and second tablet for the second trial.

Note: The glass beaker is acting as the weigh boat in this step.

18. Put on both the safety gloves and goggles.

19. Use the graduated cylinder to measure exactly 5 mL of 1M HCl.

20. Pour the 5 mL of HCl into the 100-mL beaker containing the 0.5 g of crushed antacid.

21. When the HCl is added to the antacid a fizzing reaction (formation of gas) will occur. While
carefully holding the beaker, swirl the beaker to thoroughly mix the antacid into the HCl until
the fizzing subsides.

Note: There may be a few granules of antacid that do not incorporate into the HCl, this is normal.

22. Cut off the tip of the phenolphthalein pipet with scissors and add 1 drop of phenolphthalein
to the antacid/HCl mixture in the beaker.

23. Carefully swirl the mixture in the beaker to ensure that the indicator is incorporated into the
antacid/HCl mixture; the solution will be colorless and clear.

24. Place the beaker containing the antacid/HCl/indicator solution back in the titration setup,
under the titrator.

25. Read the volume of NaOH in the titrator and record in Data Table 1 next to “Initial NaOH
Volume (mL)” under Trial 1.

26. Open the stopcock and add 1 drop of NaOH to the colorless and clear antacid/HCl/indicator
sample in the beaker. After the drop is added, gently swirl the beaker and observe the color
for 5 seconds.

Note: It is important to add the NaOH 1 drop at a time to accurately identify the equivalence point
and avoid overshooting the titration.

27. Continue adding NaOH to the beaker, 1 drop at a time, swirling and observing after each drop
until the color changes to a pale-pink color for at least 5 seconds. See Figure 11.

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Experiment Antacid Analysis and Titration

Figure 11. Endpoint of titration. The titration is complete when the color changes to a bright
pink for at least 5 seconds.

28. Read the volume of the NaOH solution remaining in the titrator and record this volume in
Data Table 1 next to “Final NaOH Volume (mL)” under Trial 1.

29. Calculate the total volume of NaOH used by subtracting the final NaOH volume from the
initial NaOH volume and record the total volume in Data Table 1.

30. Leave the titrator assembly intact. You will need it for future titrations in this experiment.

31. Pour the contents of the beaker down the drain and flush the drain with water. Thoroughly
wash the beaker with soap and water to remove all of the antacid/HCl/NaOH/indicator
solution from the beaker. When the beaker is clean, rinse the beaker with distilled water and
then thoroughly dry.

32. If necessary, add more NaOH to the titrator.

Note: It is only necessary to add more NaOH to the titrator if there is less than 1 mL more than the
total volume of NaOH used in the previous trial. For example, if the total volume of NaOH used in
Trial 1 was 2.1 mL, then needs to be at least 3.1 mL of NaOH in the titrator.

33. Repeat steps 16 through 31 one additional time (Trial 2).

34. Average the results from the 2 trials (mass of antacid and total volume of NaOH used) and
record the two averages in Data Table 1.

Note: Leave the titration setup intact (including the remaining NaOH in the titrator) for use in
Exercise 2.

Using the averages from Data Table 1 and the sample calculations provided in the Introduction
section:

35. Calculate the initial amount (grams) of HCl available for neutralization by the antacid. Record
in Data Table 2 in your Lab Report Assistant.

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Experiment Antacid Analysis and Titration

36. Calculate the number of moles of NaOH required to reach the stoichiometric point (neutralize
the excess HCl), after initial neutralization with the antacid. Record in Data Table 2.

37. Calculate the amount (grams) of HCl neutralized by the antacid. Record in Data Table 2.

38. Calculate the gram per gram neutralization of HCl by the antacid. Record in Data Table 2.

39. Leave all materials set up for Exercise 2.

Questions
A. If an antacid tablet weighed 1.6 grams, how many moles of gastric acid (HCl) would it neutralize?

Use the results obtained in Data Tables 1 and 2 to explain and quantify your answer.

B. If you performed this experiment with a titrant of 0.5M NaOH, would you expect your results
in Data Table 2 to change or stay the same? Explain your answer.

C. The reaction that occurred when the antacid mixed with the HCl resulted in an additional
product (besides a salt and water). Did you see evidence of this product? Describe the
experimental evidence you witnessed that supports the formation of the additional product.

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Experiment Antacid Analysis and Titration

Exercise 2: Control Experiment
In this exercise, a control experiment is performed to validate or contradict the results obtained
in Exercise 1.

Procedure

Note: Please read all steps and safety information before starting the procedure.

1. Put on your safety goggles and glasses.

2. Ensure that the 100-mL glass beaker is clean and dry.

3. Check the titrator setup to make sure that all components are as shown in Figure 7.

4. Use the graduated cylinder to measure 5 mL of the 1.0M HCl.

5. Carefully pour the 5 mL of 1.0M HCl into the clean, 100-mL glass beaker.

6. Add 1 drop of phenolphthalein to the HCl in the beaker. Carefully swirl the mixture to ensure
that the indicator is incorporated into the HCl. The solution will be colorless and clear.

7. With the stopcock in the closed position, ensure the titrator is filled with 9 – 10 mL of the 1.0M
NaOH. Add additional 1.0M NaOH as necessary.

8. Place the beaker containing the HCl/indicator solution back in the titration setup under the
titrator.

9. Read the volume of NaOH in the titrator and record in Data Table 3 in your Lab Report
Assistant next to “Initial NaOH Volume (mL)” under Trial 1.

10. Open the stopcock and add 1 drop of NaOH to the colorless and clear HCl/indicator sample
in the beaker. After the drop is added, gently swirl the beaker and observe the color for 5
seconds.

Note: It is important to add the NaOH 1 drop at a time to avoid overshooting the titration.

11. Continue adding NaOH to the beaker, 1 drop at a time, swirling and observing after each drop
until the color changes to a bright-pink color for at least 5 seconds. See Figure 11.

12. Read the volume of the NaOH solution remaining in the titrator and record this volume in
Data Table 3 next to “Final NaOH Volume (mL)” under Trial 1.

13. Calculate the total volume of NaOH used by subtracting the final NaOH volume from the
initial NaOH volume and record the total volume in Data Table 3.

14. Pour the contents of the beaker down the drain and flush the drain with water. Thoroughly
wash the beaker with soap and water to remove all of the HCl/NaOH/indicator solution
from the beaker. When the beaker is clean, rinse the beaker with distilled water and then
thoroughly dry.

15. If necessary, add more NaOH to the titrator.

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Experiment Antacid Analysis and Titration

Note: It is only necessary to add more NaOH to the titrator if there is less than 1 mL more than the
total volume of NaOH used in the previous trial. For example, if the total volume of NaOH used in
Trial 1 was 2.1 mL, then there needs to be at least 3.1 mL of NaOH in the titrator.

16. Repeat steps 2 through 14 one additional time (Trial 2).

17. Average the results from the 2 trials and record in Data Table 3.

Using the average from Data Table 3 and the sample calculations provided in the Introduction
section:

18. Calculate the number of moles of NaOH used to neutralize the 10 mL of 1.0M HCl. Record in
Data Table 4 in your Lab Report Assistant.

19. Calculate the amount of HCl (grams) neutralized by the 1.0M NaOH in the control experiment.
Record in Data Table 4.

20. Calculate the difference in NaOH volume between the back titration average and the control
experiment average. Record in Data Table 4.

21. Calculate the amount of HCl (grams) neutralized by the difference in NaOH volume. Record in
Data Table 4.

22. Clean all materials and return to the lab kit for future use.

23. When you are finished uploading photos and data into your Lab Report Assistant, save and
zip your file to send to your instructor. Refer to the appendix entitled “Saving Correctly,” and
the appendix entitled “Zipping Files,” for guidance with saving the Lab Report Assistant in the
correct format.

Questions
A. Did the control experiment verify or refute the results from Exercise 1? Use your results from

Exercises 1 and 2 to validate your answer.

B. Why is it important to perform multiple trials of experiments?

C. If a commercial was aired that claimed a new antacid was able to neutralize 25 times more
acid as the antacid tablet investigated in this experiment, how would you test this claim?

D. Describe possible sources of error in both Exercise 1 and Exercise 2. Describe possible ways to
reduce this error in future experiments.

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Experiment Antacid Analysis and Titration

Electrochemical Cells and
Cell Potentials
Hands-On Labs, Inc.
Version 42-0153-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

You will learn about galvanic cells and how cell
potential is calculated. You will prepare a copper/
zinc galvanic cell and measure the cell potential of
the reaction. You will monitor the potential of the
cell as the reaction proceeds.

EXPERIMENT

© Hands-On Labs, Inc. www.HOLscience.com 1

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Define electrochemistry and compare redox, oxidation, and reduction reactions.

● Describe electrochemical cells including the flow of electricity through a galvanic cell.

● Predict the anode and cathode of a redox reaction using the standard reduction potentials.

● Construct a galvanic cell.

● Operate a multimeter and interpret voltage data.

● Calculate the standard cell potential for a redox reaction.

Time Allocation: 4 hours

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Experiment Electrochemical Cells and Cell Potentials

Materials
Student Supplied Materials

Quantity Item Description
1 Camera, digital or smartphone
1 Pair of scissors
1 Roll of paper towels

HOL Supplied Materials

Quantity Item Description
1 Digital multimeter
1 Filter paper, 20 cm x 20 cm
2 Glass beakers, 100 mL
2 Jumper cables
1 Pair of gloves
1 Pair of safety goggles
1 Plastic cup, 9 oz
1 Experiment Bag: Electrochemical Cells and Cell Potentials

1 – Copper sulfate (CuSO4), 1.0 M, 75 mL
1 – Potassium chloride (KCl), 1.0 M, 30 mL
3 – Strips of copper metal, 2 in. x ¼ in.
3 – Strips of zinc metal, 2 in. x ¼ in.
1 – Zinc sulfate (ZnSO4), 1.0 M, 75 mL

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Electrochemical Cells and Cell Potentials

Background
Electrochemistry

Electrochemistry is the study of the electrical aspects of chemical reactions, concerned with two
processes: the generation of an electrical current resulting from a spontaneous chemical reaction,
and the use of an electrical current to produce a chemical reaction. These two processes describe
oxidation-reduction (redox) reactions. A redox reaction is a chemical reaction in which there is
a transfer of electrons (change in oxidation state) from one substance to another. The reaction
is termed “redox” because it is composed of two half-reactions: an oxidation reaction in which
electrons are lost and a reduction reaction during which electrons are gained. In the oxidation
reaction the loss of electrons causes an increase in the oxidation number. Likewise, in a reduction
reaction the gain of electrons causes a decrease in the oxidation number. See Figure 1.

Figure 1. Redox reaction between zinc and copper. The full reaction is shown in the top line. In
the middle line is the oxidation reaction; notice that zinc loses two electrons to form the zinc
ion. In the bottom line is the reduction reaction; notice that copper ion gains two electrons to

form the copper atom. The electrons gained and lost in the half-reactions cancel each other out
in the full redox reaction.

Electrochemical and Galvanic Cells

A device that uses redox reactions to either use or produce electricity is called an electrochemical
cell. There are two types of electrochemical cells: electrolytic cells, which use electrical energy,
and galvanic cells, which produce electrical energy from a spontaneous redox reaction. A
spontaneous reaction occurs naturally and does not require external influence (such as electrical
energy). The focus of this experiment will be on galvanic cells. See Figure 2.

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Experiment Electrochemical Cells and Cell Potentials

Figure 2. A simple galvanic cell for the redox reaction between zinc and copper. Oxidation
occurs at the anode end, as copper gains electrons. Reduction occurs at the cathode end, as

zinc donates electrons. The voltmeter measures the amount of electrical energy produced by
the cell.

In a galvanic cell the oxidation and the reduction portions of the redox reaction occur in separate
locations (such as glass beakers), with a wire to facilitate the transfer of electrons between the
locations. As shown in Figure 2, the wire may be attached to a voltmeter that measures the
potential difference of electrical charge between the two locations. If a light bulb were hooked
up to the wire, the light would burn dimly when a small potential difference exists and brightly
when a large potential difference exists. In each of the two locations, an electrode is placed in a
solution containing the same ion as the electrode. For example, in Figure 2, a copper electrode
is placed in the copper solution and a zinc electrode is placed in the zinc solution. The electrode
where oxidation occurs is called the anode, and the electrode where reduction occurs is called the
cathode. To complete the cell (electrical circuit), the two locations are connected with a medium
that facilitates the transfer of the ions (zinc ions and copper ions) from one location to another.
This connection between the two half-cells is called the salt bridge, and it contains an inert
electrolyte solution. A solution is inert if it does not react with the ions of either the electrodes
or the solutions holding the electrodes. When the galvanic cell is complete, the electrons flow
through the cell, from the anode to the cathode.

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Experiment Electrochemical Cells and Cell Potentials

Reduction Potentials

A galvanic cell produces electrical energy that can be measured by a voltmeter. The cell voltage is
the difference in electric potential between the cathode and the anode. The total amount of electric
energy that a cell is expected to produce is called the standard cell potential (E°cell). Standard cell
potential is calculated based on the assumption that the cell is in standard state conditions: the
concentration of anode solution and cathode solution is 1M, the pressure is 1 atmosphere, and
the temperature is 25°C. The standard cell potential is the contribution of standard reduction
potential from the reduction half-reaction (E°cathode) and the standard reduction potential from
the oxidation half-reaction (E°anode), as shown in the equation below:

The standard reduction potentials of half-reactions are constants. See Table 1 for a list of standard
reduction potentials for a number of half-cell reactions. All half-reactions are shown as reduction
reactions, hence standard reduction potentials.

Table 1. Standard Reduction Potentials.

Half-Reaction E°(Volts)
F2(g) + 2e

– → 2F-(aq) +2.87
Cl2(g) + 2e

– → 2Cl-(aq) +1.36
Br2(l) + 2e

– → 2Br-(aq) +1.07
Ag+(aq) + e- → Ag(s) +0.80

Fe3+(aq) + e- → Fe2+(aq) +0.77
Cu2+(aq) + 2e- → Cu(s) +0.34

One of the most common galvanic cells is the
battery. A battery contains a positive electrode (the

cathode) and negative electrode (the anode). These are denoted
by “+” and “-“ symbols on the side of the battery. The electrodes
take up most of the internal space inside the battery and access
areas where chemical reactions occur. The anode experiences an

oxidation reaction in which charged ions interact with the anode to
produce and release electrons. The cathode experiences a reduction
reaction, whereby electrons are absorbed. The reactions result in
the production of electricity, energy that travels in a circuit to

power cell phones, flashlights, and cars.
© Eric Strand, © ekler

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Experiment Electrochemical Cells and Cell Potentials

Half-Reaction E°(Volts)
2H+(aq) + 2e- → H2(g) 0.00
Fe2+(aq) + 2e- → Fe(s) -0.44
Zn2+(aq) + 2e- → Zn(s) -0.76
Al3+(aq) + 3e- → Al(s) -1.66

Mg2+(aq) + 2e- → Mg(s) -2.37
Ca2+(aq) + 2e- → Ca(s) -2.87

K+(aq) + e- → K(s) -2.93

The more positive the reduction potential, the larger the ability of the half-reaction to behave as
the oxidizing agent. Likewise, the more negative the potential, the larger the ability of the half-
reaction to behave as the reducing agent. Given two half reactions, the one with more negative
potential value will be the oxidizer. For example, consider the role of zinc as a reducer in Equation
1 below and as an oxidizer in Equation 2 below:

In equation 1, the cell potential of the half-reaction of zinc is -0.76V and the cell potential of
the half-reaction of copper is +0.34V. In this reaction, the cell potential of the zinc is much more
negative than the copper, and thus the zinc acts as the reducing agent (anode) in the reaction.

In equation 2, the cell potential of the half-reaction of zinc is -0.76V and the cell potential of the
half-reaction of calcium is -2.87V. The cell potential of the calcium is much more negative than the
zinc, and thus the calcium acts as the reducing agent (anode) in the reaction. The driving force of
a reaction, pulling electrons from the anode in one location to the cathode in the other location,
is dependent on the difference between the cell potentials of the half-reactions. The larger the
difference, the more electrical energy the redox reaction will create.

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Experiment Electrochemical Cells and Cell Potentials

The standard cell potentials for Equation 1 and Equation 2 are calculated below:

From the calculations, more electrical energy will be produced from the reaction occurring in
Equation 2 (2.11V) than the reaction occurring in Equation 1 (1.00V). As a redox reaction proceeds,
and the electrons travel from the anode to the cathode, the total cell potential for the reaction
will decrease.

In the experiment, a galvanic cell for the redox reaction between copper and zinc will be prepared.
In a galvanic cell, the Ecell must be positive for a spontaneous reaction to occur. The zinc solution
will be zinc sulfate (ZnSO4) and the copper solution will be copper sulfate (CuSO4). The direction
of electron transfer in the redox reaction will be tested by dipping the copper electrode directly
into the zinc solution and the zinc electrode directly in the copper solution to see which electrode
becomes plated with the ion of the solution. The total potential of the cell will be calculated and
compared to the total amount of electrical energy produced in the galvanic cell, as measured with
a multimeter.

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Experiment Electrochemical Cells and Cell Potentials

Exercise 1: Construction of a Galvanic Cell
In this exercise, you will create and experiment with a galvanic cell.

Procedure

1. Gather all of the supplies listed in the materials list.

2. Use the scissors to cut a strip of the filter paper approximately 1.5 inches in width (1/4 the size
of the sheet of filter paper). See Figure 3.

Figure 3. Cutting a strip of filter paper.

3. Fold the strip of filter paper in half (widthwise) and then in half again. See Figure 4.

Figure 4. Folding the filter paper in half and then in half again.

4. Put on the safety gloves and goggles.

5. Create the salt bridge by carefully winding the folded filter paper into a circle so that it fits into
the bottom of the 9 oz plastic cup. Add the potassium chloride to the cup with the filter paper
until the paper is completely covered with the potassium chloride. See Figure 5.

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Experiment Electrochemical Cells and Cell Potentials

Figure 5. Folding filter paper in cup. The potassium chloride is added to the cup to over the
filter paper.

6. Allow the paper to soak up the potassium chloride for a minimum of 10 minutes or until you
are ready to add it to the galvanic cell, as described later in the experiment.

7. Place the 2 glass beakers on a table. Add approximately 45 mL of zinc sulfate (approximately ½
of the bottle) to one of the beakers. To the second beaker, add approximately 45 mL of copper
sulfate.

8. Pick up a fresh strip of zinc and insert one end of it into the copper sulfate solution. After
approximately 5 seconds, remove the zinc from the copper sulfate and place it on a piece of
paper towel. See Figure 6.

9. Pick up a fresh strip of copper and insert one end of it into the zinc sulfate solution. After
approximately 5 seconds, remove the copper from the zinc sulfate and place it on the piece
of paper towel. See Figure 6.

Figure 6. Metal in solutions. A. Zinc being inserted into copper sulfate. B. Copper being inserted
into zinc sulfate.

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Experiment Electrochemical Cells and Cell Potentials

10. Observe the 2 metal strips and record observations in Data Table 1 in your Lab Report
Assistant.

11. From the observations, determine which of the 2 reactions is spontaneous. Record this in the
observations section of Data Table 1.

12. Set up the multimeter as follows and see Figure 7:

a. Make sure the on/off switch of the multimeter is in the “off” position.

b. Place the end of the black probe into the bottom right hole of the multimeter.

c. Place the end of the red probe into the hole directly above the location of the black
probe. Ensure that the probes are pushed all the way into the multimeter.

d. Turn the voltage dial so that the arrow end of the dial is pointing to 20 DCV.

e. Add 1 jumper cable clip to each end of the probes. It does not matter what color
jumper cable clips are provided in your kit, or which color is attached to either probe.

Figure 7. Multimeter setup.

13. Put the salt bridge into place by submerging 1 end on the copper sulfate and the other end in
the zinc sulfate. Adjust the beakers as necessary so that the salt bridge does not sink between
the beakers. See Figure 8.

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Experiment Electrochemical Cells and Cell Potentials

Figure 8. Salt bridge. Notice that either end of the salt bridge is fully submerged in solution.

14. Clip a fresh piece of zinc onto one of the jumper cable clips and clip a fresh piece of copper
onto the other jumper cable clip.

15. Place the zinc into the zinc sulfate solution, so that the metal is submerged in the solution,
but the jumper cable clip is above, and not touching, the solution or salt bridge. See Figure 9.

16. Place the copper into the copper sulfate solution, so that the metal is submerged in the
solution, but the jumper cable clip is above, and not touching the solution or salt bridge. See
Figure 9.

Note: It may take a few minutes to find the correct placement of the copper and zinc into the
solutions to keep the jumper cable clip above the solution. Adjust the jumper cable clips as necessary
to find the correct placement.

Figure 9. Metals placed into their solutions. Notice the placement of the metal and the jumper
cable clips relative to the solution.

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Experiment Electrochemical Cells and Cell Potentials

17. Turn the multimeter on, and observe whether the total voltage is positive or negative. If
the voltage reads positive, the galvanic cell was prepared correctly and can be allowed to
progress. If the voltage is negative, quickly turn off the multimeter and swap the jumper cable
clips from one metal to the other. For example, if a negative voltage was measured with the
setup in Figure 9, the black jumper cable clip would be switched to hold the zinc, and the
yellow jumper cable clip would be switched to hold the copper.

18. When the metals and jumper cable clips are arranged so that the multimeter has a positive
reading, allow approximately 5 minutes for the multimeter reading to stabilize. When the
multimeter reading has stabilized record the voltmeter reading in Data Table 2 in your Lab
Report Assistant, under 0 minutes.

19. Look at a clock or watch and record the multimeter reading for the galvanic cell every 15
minutes for 2.5 hours.

20. While the reaction in the galvanic cell is progressing, use Table 1 in the Background section
to determine the 2 half-reactions and standard reduction potentials for the redox reaction
occurring in your galvanic cell. Record the half reactions, identifying which is the oxidation
and which is the reduction half-reaction. Also record the corresponding reduction potentials
in Data Table 3 in your Lab Report Assistant.

21. Record the equation for the complete redox reaction occurring in the galvanic cell in Data
Table 3.

22. Calculate the standard cell potential for the redox reaction occurring in the galvanic cell, and
record in Data Table 3.

23. When all multimeter readings have been taken and recorded in Data Table 2, take a photograph
of your galvanic cell. In the photograph, include a small piece of paper that displays your name
and the date. Resize and insert the photograph in Data Table 4 in your Lab Report Assistant.
Refer to the appendix entitled, “Resizing an Image” for guidance.

24. When you are finished uploading photos and data into your Lab Report Assistant, save and
zip your file to send to your instructor. Refer to the appendix entitled “Saving Correctly,” and
the appendix entitled “Zipping Files,” for guidance with saving the Lab Report Assistant in the
correct format.

Cleanup:

25. Turn the multimeter off and carefully take apart the galvanic cell.

26. Properly dispose of solutions, metal pieces, and the salt bridge.

27. Wash lab equipment with soap and water and thoroughly dry.

28. Return cleaned equipment to the lab kit for future use.

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Experiment Electrochemical Cells and Cell Potentials

Questions
A. What were the concentrations of the solutions (zinc solution, copper solution, and salt

bridge)? Were the concentrations consistent with those of standard state conditions? Explain
your answer.

B. Was the amount of electric energy produced in your galvanic cell consistent with the standard
cell potential of the reaction (as calculated in Data Table 3)? Hypothesize why it was or was
not consistent.

C. Was there evidence of electron transfer from the anode to the cathode? Use your data in
Data Table 2 to explain your answer.

D. For the following redox reaction in a galvanic cell, write the oxidation half-reaction and the
reduction-half reaction, and calculate the standard cell potential of the reaction. Use Table 1
in the Background as needed. Explain how you identified which half-reaction is the oxidizer
and which is the reducer. Show all of your work.

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Experiment Electrochemical Cells and Cell Potentials

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