General Chemistry Kinetics Titration Acid Bases & Concentrations Lab Reports

Personal Protective Equipment Requirements




Safety Glasses


Gloves Nitrile


Fume Hood





Lab Coat


Gloves Heavy


Bio Hood



Experiment 11


Equilibrium—Shifting reactions

A lesson on Le Chatelier’s Principle



word image 352 word image 353

Experiment adapted from one used at Colby College




Students will be able to…

  • Describe Le Chatelier’s Principle
  • Determine the shift in reaction that occurs with various changes in experimental conditions: concentrations, temperature, etc.




No chemical reaction goes to completion. When a reaction stops, some amount of reactants remain. For example, although we write


2 CO2 (g)  2 CO (g) + O2 (g) (1)


as though it goes entirely to products, at 2000K only 2% of the CO2 decomposes. A chemical reaction reaches equilibrium when the concentrations of the reactants and products no longer change with time. The position of equilibrium describes the relative amounts of reactants and products that remain at the end of a chemical reaction. The position of equilibrium for reaction (1) is said to lie with the reactants, or to the left, because at equilibrium very little of the carbondioxide has reacted. On the other hand, in the reaction

H2 (g) + ½ O2 (g) H2O (g) (2)


the equilibrium position lies very far to the right since only very small amounts of H2 and O2 remain after the reaction reaches equilibrium. Since chemists often wish to maximize the yield of a reaction, it is vital to determine how to control the position of the equilibrium.


The equilibrium position of a reaction may shift if an external stress is applied. The stress may be in the form of a change in temperature, pressure, or the concentration of one of the reactants or products. For example, consider a flask with an equilibrium mixture of CO2, CO, and O2, as in reaction (1). If a small amount of CO is then injected into the flask, the concentration of CO2 increases. Here the external stress is the increase in concentration of CO. The system responds by reacting some of the added CO with O2 to yield an increased amount of CO2. That is, the position of equilibrium shifts to the left, yielding more reactant and less CO.


Reaction (1) also shifts with changes in pressure. Starting with reaction (1) at equilibrium, an increase in pressure causes the position of equilibrium to shift to the side of the reaction with the smaller number of moles of gas. That is, by shifting the equilibrium position to the left, the reaction decreases the number of moles of gas, thereby decreasing the pressure in the flask. In so doing, some of the applied stress is relieved. On the other hand, an increase in pressure for reaction (2) shifts the equilibrium position to the right to decrease the number of moles of gas.


The response of a reaction at equilibrium to changes in conditions is summarized by LeChâtelier’s Principle: A system perturbed from equilibrium shifts its equilibrium position to relieve the applied stress.


For an increase in temperature, the reaction shifts in the endothermic direction to relieve the stress. The decomposition of CO2, reaction (1), is endothermic in the forward direction. Upon an increase in temperature, the equilibrium position shifts in the forward direction to minimize the temperature increase. The formation of ammonia is exothermic:

N2 (g) + 3 H2 (g)  2 NH3 (g) (3)


Upon an increase in temperature, the equilibrium position shifts to the left, which is the endothermic direction.


The Iron-Thiocyanate Equilibrium


When potassium thiocyanate, KNCS, is mixed with iron(III) nitrate, Fe(NO3)3, in solution, an equilibrium mixture of Fe 3+ , NCS , and the complex ion FeNCS 2+ is formed:


Fe3+ (aq) + NCS (aq)  FeNCS2+ (aq) (4) yellow colorless red


The solution also contains K + and NO3 ions, but these are spectator ions and do not participate in the reaction. The relative amounts of the various ions participating in the reaction can be judged from the color of the solution. In neutral or slightly acidic solutions, Fe3+ is light yellow, NCSis colorless, and FeNCS2+ is red. If the solution is initially reddish, and the equilibrium shifts to the right (more FeNCS2+ the solution becomes darker red, while if the equilibrium shifts to the left (less FeNCS2+ the solution becomes lighter red or perhaps straw-yellow.




  1. Consider the reaction: NO (g) + O3 (g)  NO2 (g) + O2 (g) . This is an exothermic reaction. What would you expect to happen (shift left, shift right, no change) if
    1. NO was added
    2. CO was removed
    3. N2 (Nitrogen gas) was added
    4. Temperature is increased


  1. What is a reasonable redox reaction of Tin (II) with Iron (III)? Make sure that the reaction is balanced.


  1. What is the reaction of Ag+ with FeSCN 2+ Write out the complete balanced reaction.


MATERIALS—Virtual Lab (materials for recording only)


  • Netbooks
  • Water
  • Beakers
  • Graduated cylinder
  • Spatula
  • Spot well plate
  • Rack with 5-10 mL test tubes  Ice bath (one per bench)

Hot water bath (one per bench)


 50 mL of each solution per lab bench) o 1 M Fe(NO3) 3 o 1 M KNCS o 0.1M SnCl2 o 0.1 M AgNO3 o 0.1 M Na2HPO4 o 1 M NH4OH



Personal Protective Equipment Requirements

Instructor Demo


Safety Glasses


Gloves Nitrile


Fume Hood


Instructor Always


Lab Coat


Gloves Heavy


Bio Hood




PROCEDURE—Virtual—Data is based on procedure below


For each of the external stresses described below, necessary information is provided regarding the manner in which one or more of the chemical species is affected. You will use a spot plate containing multiple wells and use a different well for each of the operations described, recording your observations of the color change of the solution.


A. Operations to Introduce an External Stress– Record your observations in your data table


  1. Add one drop each of 1 M Fe(NO3)3 and 1 M KNCS to 25 mL of distilled water. Mix well.


  1. Add a few drops of this solution to each of seven wells of a spot plate. One well serves as a color standard against which to judge color changes in the other wells. The other six wells are for performing your operations to introduce an external stress.


  1. Add one drop of 1 M Fe(NO3)3 to one of the wells, mix, and observe.


  1. Add one drop of 1 M KNCS to a second well, mix, and observe.


  1. Add one drop of 0.1M SnCl2 to a third well, mix, and observe. Tin(II) ions are involved in a redox reaction with the Iron (III).


  1. Add one drop of 0.1 M AgNO3 to a fourth well, mix, and observe.


  1. Add one drop of 0.1 M Na2HPO4 to a fifth well, mix, and observe.


  1. Add one drop of 1 M NH4OH to a sixth well, mix, and observe.


B. Effect of Temperature on the Equilibrium– Also record your observations in the data table.


  1. Pour about 4-5 mL of the iron-thiocyanate solution made above into three test tubes. Set one tube aside as a color standard against which to judge color changes in the other tubes.


  1. Gently warm the second tube in a hot water bath on a hot plate. Do not boil the solution.


  1. Cool the third tube in a beaker of ice water. Observe.





  1. In a table, summarize your observations for each of the reactions that you perform on the ironthiocyanate equilibrium. As an example, if you added a drop of concentrated HCl to the standard solution, the blood-red color lightens or perhaps disappears altogether. This change in color indicates that the FeNCS2+ concentration decreases.


To explain this result, it is necessary to know that in the presence of a large excess of Cl–, Fe3+ forms complex ions: Fe 3+(aq) + 6 Cl (aq) FeCl6 3- (aq)


The increase in Clreduces the Fe3+ concentration in accord with Le Chatelier’s Principle and some FeNCS2+ dissociates to replace some of the Fe3+ removed by reaction with Cl–.


  1. Write a complete balanced equation for each of the added materials (Procedure steps 4-8)


  1. Based on the results, is the reaction of iron with thiocyanate exothermic or endothermic? Explain


  1. Be sure to submit your completed data table with you’re A&D.





  1. State 2 possible errors in this procedure (human error is not acceptable) and what impact each error (separately) would have on the final results.


  1. Explain how spectroscopy (think cranberry lab) could be used in this experiment. Give some thought and detail about what that protocol might be. Write a short paragraph explaining what you might do in lab.


  1. If I wanted to maximize the amount of FeNCS 2+ (aq)—what three variables would you use and how would you use each (add or remove this, change that). Would you have 100% product? Explain.


  1. You are the plant manager making Ammonia (NH3) by the Haber process starting with pure nitrogen and hydrogen as reactants (reaction 3 above). What are four (4) ways to improve the amount of product (and thereby guaranteeing you a promotion and a raise from the owners)? 8

Analysis of Cranberry Juice Blends

A lesson on Concentrations

Personal Protective Equipment Requirements

Instructor Demo


Safety Glasses


Gloves Nitrile


Fume Hood


Instructor Always


Lab Coat


Gloves Heavy


Bio Hood



Students will be able to…

  • Prepare a series of cranberry juice standards
  • Create a calibration curve and use the curve to determine the concentration of cranberry juice in cranapple juice


Cranapple juice is made by combining cranberry juice with apple juice. Different brands have different ratios (concentrations) of the two juices in their drink. In this experiment you will be asked to compare the spectrum of apple juice with the spectrum of cranberry juice and decide the concentration of cranberry juice in the unknown cranapple juice.

The UV-Vis spectrum is additive, meaning that if two compounds/juices absorb at the same wavelength the spectrum that you see will be the sum of the two individual absorbances. The first thing you will need to do is to take the spectrum of pure apple juice and pure cranberry juice to determine what wavelength to investigate. Ideally you will find a peak where only one compound/juice absorbs light.

In order to determine the concentration of a substance, you must first create a calibration plot which is a graph of known data (absorbance vs concentration) that is used to determine the unknown concentration of a substance in a solution. Several solutions are prepared with known concentrations of the substance of interest and are called standards and are the independent variable. The measured property, absorbance at a selected wavelength, is the dependent variable.

Cranberry juice is a red juice who’s colored increases as concentration increases. This colored complex obeys Beer’s law; as the concentration of the solution increases, the amount of light absorbed by the solution increases (which is what was discussed above). In this lab, you will create a set of standards to develop a calibration curve and you will then determine the concentration of ethanol in a wine sample.


  1. Given the following graph; answer the following questions:
    1. What does “y” stand for in the equation?
    2. What does “x” stand for in the equation?
    3. Use the equation for the trendline below to find the concentration of Red #40 in Kool-Aid if a solution has an absorbance of 0.425. Show your work.
  2. Suppose you create a standard with an initial concentration of 20 mg/L of Red #40 dye. You dilute 5.00 mL of this to a final volume of 100.0 mL. What is the concentration of dye? Show your work.
  3. What two assumptions are you making when using spectroscopy for this analysis? Explain why those assumptions may be valid for this experiment.
  4. In cranberry juice—what compound is giving it the red color?
  5. Another common red food coloring is Red #3–  Erythrosine. What changes might have to be made to the lab to perform this experiment?


  • 10.00 volumetric flask
  • 1 mL volumetric pipet
  • 2 mL graduated pipet
  • Cranberry Juice
  • Apple Juice
  • conical tubes
  • cuvets
  • Vernier spectrophotometer (SpectroVis)
  • Unknown


  1. Initializing Equipment:
    1. Plug in spectrophotometer to computer and open Logger Pro Program. You should see a colored spectrum appear.
    2. To Calibrate:
  2. From Experiment tab on top, select calibrate → spectrophotometer 1
  3. You will see a message that spectrophotometer needs to warm up – 90 sec.
  4. Place cuvet of distilled water (your blank) in holder – clear side by white arrow
  5. Click Finish Calibration
  6. When finished, click OK
  7. Determining Wavelength:
    1. Pipette cranberry juice into the cuvette. Place the cuvette into the spectrovis and click on collect. After the values have stabilized (around 1 minute) click stop.
    2. Note the color of the data and sketch the graph in your notebook. Note any peaks in the data that you see.
    3. Pipette apple juice into the cuvette. Place the cuvette into the spectrovis and click on collect. In the popup box click on store latest run. After the values have stabilized (around 1 minute) click on stop.
    4. Note the color of your new spectrum and sketch the graph overtop of the cranberry juice. In your notebook describe the two spectra and which wavelength you think will be the best for determining the amount of cranberry juice in your sample.
  8. Making your standards:
    1. Create a set of 5–10 mL cranberry standards ranging from 0% cranberry juice to 30% cranberry juice by volume in apple juice. Verify your calculations and chosen wavelength with your instructor prior to making your samples.
    2. After creating your standards, pour each standard into separate conical tubes and label your tubes.
  9. Absorbance measurements: Take the absorbance measurement of the standards and samples.
    1. Determining absorbance vs concentration at particular wavelength
  10. Click configure spectrophotometer button ( )
  11. Select absorbance vs concentration
  12. Change “Single 10 nm Band” to “Individual Wavelengths” and select the wavelength from step 2
  13. Click OK
  14. A window will pop up asking you if you want to save the data – select Yes
  15. Two new columns will appear on left (concentration and absorbance) as well as a plot window
  16. Put in first sample (most dilute sample)
  17. Click Collect and then Keep
  18. A window will open with concentration. Input your calculated concentration. You should also record this data in your laboratory notebook as you do each sample.
  19. Remove your sample and place the next most dilute sample. Once the absorption has stabilized click keep and input your calculated concentration.
  20. Repeat for your remaining samples.
  21. After your last sample click on Stop
  22. Click on insert linear fit and write down the equation in your lab notebook.
  23. Insert your unknown sample and record the absorbance once it has stabilized.

Waste: All solutions can be disposed down the drain.


  1. Create your standard curve (absorbance vs. concentration).
    1. Concentration on x-axis and absorbance on y-axis. Remember labels, trendline, and R2 value.
  2. From your curve what is the concentration of cranberry juice in each cranapple juice//juice blend?
  3. How does the percent cranapple juice you calculated compare with the actual percentage? Why might these numbers differ? (Note: human error is not acceptable.)


  1. The density of pure ethanol (alcohol) is 0.789 g/mL. The percent, by volume, of alcohol in a sample of white wine is 12.5%, which has density of 0.98167 g/mL. What is the percent, by mass, of alcohol in the sample?
  2. Given the following information, determine the percent, by volume, of alcohol in a particular mixed drink (cocktail). (Show your work and include graph!)


Concentration (%)













A 1.0 mL sample of the mixed drink was diluted to 10.00 mL and the absorbance was found to be 0.136.

  1. Windshield washer fluid, composed of water and methanol, comes in two different formulas, winter and summer. Each has as different water to methanol ratio. In an experiment, a student collected data to determine the density of different percent methanol solutions to compare to the windshield washer fluid.
    1. Given the following information:
      1. Prepare a calibration curve
      2. Determine percent of methanol in winter and summer windshield washer fluid. (Show your work and include graph!):


Percent Methanol

Density (g/mL)

























10 mL of washer fluid

Density (g/mL)

Winter fluid (-20 °F)


Summer fluid (+32 °F)


    1. It is known that summer formula (+32 °F) is 1.25% methanol by volume and winter formula (-20 °F) is 37.09% methanol by volume. Why might these numbers vary compared to what you calculated (not human error)?

HIS.Q 205.01.02 2020 Review Sheet Exam #3


Enduring Vision, Chs. 9-15

Terms to Know:

Both lecture/films and Enduring Vision

Book Only

**Terms in the book but not on this sheet, will NOT be on the exam**

Chapter 9

Market Economy

Erie Canal

Five Civilized Tribes/Five Tribes

Indian Removal Act

Trail of Tears

Eli Whitney and Cotton Gin

Lowell, MA/Waltham, MA mills

Chapter 10

Political Democratization

Henry Clay

South Carolina Exposition and Protest (1828)

Andrew Jackson

Tariff of 1832

Compromise Tariff of 1833

Force Bill

Second Bank of the United States/Bank Crisis

Nullification/Nullification Crisis

Whig Party

Panic of 1837

Second Great Awakening

The Burned Over District

Charles Grandison Finney


William Lloyd Garrison

Elizabeth Cady Stanton

Seneca Falls Convention

Sojourner Truth

Chapter 11

McCormick Reaper

New York Stock Exchange



Penny press

Minstrel Shows

P.T. Barnum

James Fenimore Cooper

Ralph Waldo Emerson

Henry David Thoreau

Margaret Fuller

Walt Whitman

Nathaniel Hawthorne

Edgar Allen Poe

Herman Melville

Hudson River School

George Catlin

Frederick Law Olmstead

Chapter 12

Nat Turner

Plantation Agriculture

Virginia emancipation legislation (1831-1832)

George Fitzhugh, Cannibals All! (1856)

The Impending Crisis of the South

Task System and Gang Labor

Frederick Douglass

Harriet Tubman

The Underground Railroad

Denmark Vescey

Spirituals (also in Northup)

Chapter 13

German and Irish Immigrants


Mexican independence

Empresario Program

Stephen F. Austin

Sam Houston

The Alamo (1836)

Battle of San Jacinto (1836)

James K. Polk

Mexican War 1846-48

Treaty of Guadalupe Hidalgo (1848)

Free Soil Party

California Gold Rush

Chapter 14

Compromise of 1850

Fugitive Slave Act

Kansas-Nebraska Act (1854)

Popular Sovereignty

Bleeding Kansas

Lecompton Constitution

Uncle Tom’s Cabin

Slave Power/Slaveocracy

Know Nothing Party

Republican Party

Dred Scott v. Sandford (1857)/Dred Scott Decision

Caning of Charles Sumner


Chapter 15

Jefferson Davis (also in Burns’s Civil War)

Fort Sumter (1860)

First Battle of Bull Run (1861)

Antietam (1862)

Emancipation Proclamation (1862/3)

Ulysses S. Grant

William T. Sherman

Robert E. Lee

Vicksburg (1863)

Gettysburg (1863)

Freedmen’s Bureau

Appomattox Courthouse (1865)

New York City Draft Riots


Spectrum of Cranberry Juice

Spectrum of Apple Juice

Cranberry Juice Standards

Concentration (Vol %)

Trial 1 Absorbance

Trial 2 Absorbance





















Trial 1 absorbance

Trial 2 absorbance

Ocean Spray-Cranapple



Wegmans Cranapple



Apple Juice



White Cranberry Juice



Ocean Spray Cranraspberry


0.874 Experiment 9

Designing Hot and Cold Packs:

A lesson on Exothermic and Endothermic Reactions



Students will be able to…

  • Determine the most effective chemical for use in a cold pack
  • Determine the most effective chemical for use in a hot pack
  • Design an effective hot pack or cold pack
  • Describe the differences in exothermic and endothermic reactions based upon the reactions they conducted in the lab


Cold packs are used to treat sprained ankles and similar injuries. A cold pack is typically made of a thin plastic inner bag containing water. That bag, in turn, is surrounded by a heavier plastic bag containing a solid substance. When the pack is twisted, the inner bag breaks and releases the water. As the solid substance dissolves in the water, energy is absorbed and the resulting mixture gets colder.

Some hot packs are similar in design; however, they are used to warm up the body (or keep it warm), prevent hypothermia, and soothe sore muscles. One hot pack of similar design requires you to empty the contents of a chemical pack into the hot pack container and then simply add water. The reaction will then reach a temperature of 160° F and will keep that heat for an hour and a half assuming the individual is well insulated from the cold.

In this experiment, you will first determine temperature changes when several different solid substances are dissolved in water. Using graphical methods, you will then determine which substance you would choose in developing your own cold pack or hot pack and explain why this would be the case. You will then create your own design for the most effective cold or hot pack.


  1. Write the sentence correctly: Energy is (released/required) to break a bond and (released/required) when a bond forms.
  2. Determine the following for both an exothermic and endothermic process:
    1. What happens to the temperature of system?
    2. What happens to the temperature of the surroundings?
  3. Categorize a cold pack and hot pack as either exothermic or endothermic. Explain your decisions.


  • Netbooks
  • Vernier computer interface
  • Temperature probe
  • Water
  • Beakers
  • Graduated cylinder
  • Spatula
  • Ammonium chloride, NH4Cl
  • Calcium chloride, CaCl2
  • Citric acid, H3C6H5O7
  • Magnesium sulfate, MgSO4
  • Potassium chloride, KCl
  • Sodium bicarbonate, NaHCO3
  • Sodium carbonate, Na2CO3
  • Sodium chloride, NaCl

Personal Protective Equipment Requirements

Instructor Demo


Safety Glasses


Gloves Nitrile


Fume Hood


Instructor Always


Lab Coat


Gloves Heavy


Bio Hood



Reusable Heat Packs–

PROCEDURE—Virtual—Data is based on procedure below

  1. The temperature of 25 mL of DI water was recorded. This amount was the same for all of the experiments.
  2. 1 g of salt was then added and the water stirred while the temperature was recorded. Time was stopped once the temperature changed directions. If the flask heated up, once the temperature started to decrease. Time at Final T was how long the flask stayed at Tfinal. See data table.
  3. Using the data how could you make the best hotpack? You may mix as many of the solids as you want, and use as much water as you want. You will need to justify your answers in the A&D section.
    1. Using the graph, determine the time between the minimum and maximum temperatures.
    2. Repeat for the other solids, remembering to rinse the probe with DI water.
  4. After your initial testing with all the solids, you need a plan for creating the most effective hot OR cold pack. You can only use a maximum of 3 grams of solid in your pack. Water is unlimited. You can combine solids if desired.


Hot and Cold Packs Experiment




Total Time (sec)

Time to Reach Tf (sec)

Time at Final (sec)

(Total – Time to Tf)

Ammonium Chloride






Calcium Chloride






Citric Acid






Magnesium Sulfate






Potassium Chloride






Sodium Bicarbonate






Sodium Carbonate






Sodium Chloride







  1. Explain what you saw in the demonstration and why this occurred. How can this reaction be used as a reusable hot pack?
  2. Determine your temperature change and rate of change (∆T/ ∆t) for each of the chemicals above.
  3. Categorize each chemical substance, including the demonstration substance, as having an exothermic or endothermic reaction when combined with water. Explain.
  4. In your initial discovery of determining which substances were effective for use in a hot or cold pack, you found that depending upon the substance, the temperature either increased or decreased, but the degree of temperature change also varied. Discuss some possible reasons why you believe this may have been the case.
  5. Which substance would you use for a cold pack? Explain your decision.
  6. Which substance would you use for a hot pack? Explain your decision.
  7. Explain your final design for the most effective cold or hot pack (how much water did you use, how much solid substance did you use). How did you arrive at this final design? Discuss your findings based upon your design (how cold or hot did it get, how long did it take to reach this temperature, how long did it stay at temperature).


  1. Exothermic Reaction
    1. In your own words, describe an exothermic reaction.
    2. Give two examples of exothermic reactions/processes and explain why these are examples of an exothermic reaction/process. (not hot/ cold pack)
  2. Endothermic Reaction
    1. In your own words, describe an endothermic reaction.
    2. Give two examples of endothermic reactions/processes and explain why these are examples of an endothermic reaction/process. (not hot/ cold pack)
  3. Calcium chloride is an ice-melting compound used for sidewalks and streets. This is an exothermic reaction in which calcium chloride reacts with water which releases heat to the surroundings causing the ice to melt. Using the language of breaking and making bonds and the energy required/released when that happens, how can you describe the temperature change observed when calcium chloride is dissolved in water?
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