Electrolytes vs Nonelectrolytes Different Solutions and A Soil Sample Lab Report


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Experiment 5

How Can Geckos Walk Up Walls?:

A lesson on Intermolecular Forces

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Students will be able to…

  • explain how dominant forces (gravity and IMF) change with scale. At the macroscale, gravity is dominant, while at the nanoscale intermolecular forces are dominant.
  • compare and contrast attractive versus repulsive intermolecular forces.
  • describe, compare, and contrast three major types of intermolecular forces.
  • predict physical/macroscopic properties such as vapor pressure and boiling points based on identification of a substance’s intermolecular forces.





You may have seen geckos (and some insects) climbing up walls and scurrying across ceilings. Why may this be the case? How are these creatures able to do this, whereas humans are unable to climb up a flat wall or cling to a flat ceiling (let alone move across it)? Some people have thought maybe they have a glue they are emitting or that their feet are like suction cups; however, geckos are able to run very quickly and so these theories do not support much of the data that has been collected. One current thought is that it all has to do with intermolecular forces, the attractions between the small structures within the geckos feet and the wall.








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Intramolecular forces are “bonding” forces, such as covalent bonds, which exist within each molecule. These forces influence the chemical properties of the substance. Intermolecular forces are “nonbonding” forces that exist between molecules. They influence the physical properties of the substance.


In this experiment, you will first investigate three types of intermolecular forces: dispersion forces (van der Waals), dipole-dipole forces, and hydrogen bonding. You will then use your knowledge gained from Part A to design an experiment to determine the identity of 5 unknown liquids (hexane, acetone, butanol, water, and glycerin).





  1. For each of the intramolecular forces below, explain which each one is, which is the strongest, and give an example of a molecule that has that force.
    1. Dipole-dipole interaction
    2. Hydrogen-bonding
    3. Van der Waal’s force


  1. For each of the molecules below, explain which of the three intermolecular forces is involved (may be more than one).
  1. Methanol (CH3OH)
  2. Chlorine gas (Cl2)
  3. Acetic acid (CH3COOH)
  1. Water (H2O)
  2. Propane (C3H8)
  3. Phosphorus (P4)


  1. Define each of the properties that you will be looking at in this experiment. For each, would you expect Methanol (CH3OH) to have a high or low value and why?


    1. Capillary Action e. Melting point
    2. Surface Tension f. Freezing point
    3. Solubility in water g. Boiling point
    4. Solubility in hexane (C6H14) h. Vapor Pressure


  1. Complete the table below. Some columns have been completed for you. As you are searching for these values, some helpful search terms will include:
  2. Physical properties of … (insert each chemical here) or SDS of … (insert each chemical here)
  3. Include a list of your references. You can use Wikipedia for this assignment.







Butanol (n-butanol)

Glycerin (glycerol)

Molecular Formula








Molecular Structure


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Molecular Weight (g/mol)









Tension at

20oC (dyn/cm)









Pressure (mm Hg at 20 or








Melting Point (oC)








Boiling Point (oC)








Solubility in water






Slightly soluble






  • Water (distilled)
  • Ethanol (on each bench)
  • Cyclohexane (on each bench)
  • 4 unknown liquids (A, B, C, D) o hexane, acetone, butanol, and


  • Glass plate
  • Stopwatch


  • Test tubes (6 per group) and rack
  • Wax paper
  • Penny
  • Capillary tubes
  • Ruler
  • Waste beaker

Personal Protective Equipment Requirements


Instructor Demo


Safety Glasses


Gloves Nitrile


Fume Hood


Instructor Always


Lab Coat


Gloves Heavy


Bio Hood






A. Discovery: How do intermolecular forces relate to the physical properties of substances?

You will examine several physical properties of three liquids: water, ethanol, and cyclohexane.


1. Volatility

Place a drop of each liquid on a glass plate. Record the time it takes for the liquid drop to evaporate. Record your observations.


2. Solubility

If two liquids mix together (no layer forms), they are soluble. Try different combinations of the three liquids in a small test tubes to determine which are soluble in each other. Use approximately 2 mL of each liquid. Record your observations. Dispose of your combinations in a waste beaker.


3. Surface Tension

Place a drop of each liquid on a piece of wax paper. Describe each drop and measure the diameter of each drop. How many drops of each liquid can you place on the surface of a penny? Record your observations


4. Capillary Action

Obtain capillary tubes that are labeled with the different liquids. Place one end of the tube into the appropriate liquid. Record your observations of the liquid level in the tubes



B. Determination of Mystery Liquids (will be developed during lab)

Using your findings from Part A, develop a procedure to determine the correct identity of the four unknown liquids. Perform at least three tests on each liquid to verify your results. You will have everything from Part A available to you. If you need something special, see your instructor. You may NOT use taste, smell, or touch as one of your tests. You must show your procedure to your instructor before proceeding.









  1. For Part A, determine the polarity and intermolecular force(s) involved in each liquid (water, ethanol, cyclohexane). Use your data (volatility, solubility, surface tension, and capillary action) to support your answer.


  1. Identify each unknown liquid and the intermolecular force(s) involved in each liquid. Provide evidence from your experiments that confirms your conclusions.


  1. Fill in the blanks: The stronger the IMF the
    1. ____(higher/lower) the melting/freezing/boiling points because ______.
    2. ____(higher/lower)the vapor pressure because _______________.
    3. ____(higher/lower) the viscosity because _________.
    4. ____(stronger/weaker) the surface tension because __________.
    5. ____(stronger/weaker)the capillary action because ________.


  1. How can a gecko walk upside-down on a ceiling? Why can geckos walk upsidedown on ceilings, while humans cannot?





  1. Explain, using intermolecular forces, why at room temperature, chlorine (Cl2) is a gas, bromine (Br2) is a liquid and iodine (l2) is a solid.


  1. Indicate the intermolecular forces for isopropyl alcohol (structure below)

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      1. How would you expect isopropyl alcohol to behave in water? Why?
      2. Would you expect isopropyl alcohol to evaporate faster or slower than water? Why?


  1. For the following, determine the best answer AND justify your decision:
    1. Determine which substance has the lowest boiling point:

a. HF b. H2 c. CO2 d. Ne

    1. Determine which substance has the lowest vapor pressure at 25°C:

a. HF b. Br2 c. CO d. Ar

    1. Determine which substance has the highest surface tension:

a. HCl b. H2O c. CH3OH d. CH3CH2OH

  1. Molecular iodine would be most soluble in ____ because ____.
    1. Hexane
    2. Water
    3. Gasoline (octane C8H18)
    4. Methanol



http://static.abbottnutrition.com/cms/PEDIALYTE_2011/IMAGES/pdl-liters-product-20130904.jpg https://dejavuiv.com/wp-content/uploads/2013/03/gatorade.jpg Experiment 7

Electrolytes vs Nonelectrolytes

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…

  • Measure the conductivity of different solutions and a soil sample
  • Create diluted samples and discuss how dilution influences conductivity
  • Determine if each sample, including unknowns, is an electrolyte or nonelectrolyte
  • Evaluate which drink they believe athletes should consume based upon conductance and if a soil sample will be good for plant growth


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Conduction of electricity requires movement of charged particles, such as ions. Substances that form ions when they dissolve in water can conduct electricity and are classified as electrolytes. Substances such as molecular compounds that do not form ions when dissolved in water do not conduct electricity and are classified as nonelectrolytes.

word image 1197 Conduction occurs in electrolytic conductors as both positive ions (cations) and negative ions (anions) migrate toward electrodes. Electrons do not flow in a solution of an electrolyte. Electrolytic conduction involves a migration of ions from one part of the conductor to another.

One way to monitor the levels of electrolytes in a solution is by testing the conductivity. Conductance values are often reported in units of microsiemens (1 μS/cm = 10-6S/cm).We will measure the conductance of several solutions using a conductivity probe similar to that shown in the figure. When the probe is in a solution with ions, an electrical circuit is completed across the electrodes which are on either side of the hole in the probe. A voltage is applied to the two electrodes and the resulting current is proportional to the conductance of the solution.

Electrolytic conduction plays an important role in the function of electrochemical cells, batteries, electrolysis, and electroplating. A compounds ability to ionize and conduct electricity also plays a significant role in biological systems. For example, potassium ions are essential for photosynthesis and respiration in plants. Potassium ions and sodium ions are necessary for electrical impulses, which govern nerve and muscle function. In both plants and animals, a specific concentration of the electrolytes is required for physiological systems to function properly and avoid damage to cells. For humans, there are many products on the market that claim to remedy the loss of electrolytes whether due to excessive exercise or diarrhea. You will measure the conductance of some of these products.

Although different plants can tolerate different levels of electrolytes, a general guideline for conductivities is given in the table below.


Soil Quality

< 200 μS/cm

Not enough electrolytes

200 μS – 1200 μS/cm

Healthy range for many crops

> 1200 μS/cm

Too many electrolytes

The electrical conductivity of a soil sample indicates the presence or absence of salts, but does not indicate which salts might be present. Salinity is total concentration of all dissolved salts in water. The salinity of soil is not entirely from sodium chloride. There are other mineral contributing to salinity such as, calcium carbonate (CaCO3), sodium sulfate (Na2SO4), and magnesium chloride (MgCl2). There are also salts in fertilizers that contribute to the soluble salts in soil, hence increasing the conductivity of a soil sample. Ammonium nitrate (NH4NO3) and potassium nitrate (KNO3) are salts found in fertilizers that would add to the salinity of the soil.

In this experiment, you will examine the conduction of several solutions, including multiple replenishing drinks to determine which is the most effective and a soil sample to determine its suitability for sustaining plant growth, by measuring their conductance (the conductivity of these solutions) and categorize them as electrolytes or nonelectrolytes. Then, you will identify two solutions of unknown composition by measuring their conductance and deciding whether the solution contain electrolytes or nonelectrolytes. You will also see how conductivity changes as the concentration of the electrolytes changes.


  1. Explain how you could make 10.0 mL of a 0.010 M C6H12O6 solution from a 0.500 M C6H12O6 solution. (Show work)
  2. Complete the table below:


Ionic or Not?

Conduct or Not in water?










  1. After a student dissolved 0.7 g of white solid in water, she tested the conductivity and determined that the solution did not conduct electricity.

The student then concluded: “when the white solid dissolved it ‘just disappeared’ so some sort of error must have been made because all solids are ionic compounds and, therefore, are electrolytes.

Comment on the student’s conclusion. Is her statement a valid conclusion? If so, identify what is correct. If not, write a conclusion that would be correct.

  1. After reading that fertilizers contain potassium ions, an electrolyte essential for plant growth, your friend tells you that his plants will be the healthiest around because he is going to give his plants double the recommended amount of fertilizer. Will his plants win at the state fair? Why or why not?
  2. Even though this is a virtual lab, complete the safety section for your pre-lab as if you did the lab in person.

MATERIALS (NONE FOR FA20—VIRTUAL LAB)—Still comment on safety for pre-lab

  • Netbook
  • 0. 50 M ammonium nitrate (NH4NO3)
  • 0. 50 M glucose (C6H12O6)
  • 0.50 M potassium nitrate (KNO3)
  • 0.50 M sodium chloride (NaCl)
  • 0.50 M calcium chloride (CaCl2)


  1. Preparing Soil Sample
  2. Measure approximately 3 grams of soil noting which sample you take. Record the exact mass.
  3. Carefully pour the soil into a large test tube. For every 1 g of soil, add 5 mL of deionized water. Record the amount of water added.
  4. Stretch a piece of parafilm over the top of the tube. While holding your fingers over the tube, gently invert the tube several times. Allow the soil and water to sit for at least 30 minutes, gently inverting the tube periodically.
  5. Prior to taking a measurement, allow the soil to settle to the bottom of the test tube.
  6. Measuring Conductance
  7. Measure conductance for all solutions.




( S)




( S)

Glucose (0.02 M)



DI water


Glucose (0.01 M)



Tap Water


KNO3 (0.02 M)





KNO3 (0.01 M)





NaCl (0.02 M)





NaCl (0.01 M)



Soil solution


CaCl2 (0.02 M)



Unknown X


CaCl2 (0.01 M)



Unknown Y



  1. Create a results table
    1. Column 1: List the 0.02 M solutions, water, drinks and soil sample in order of increasing (lowest to highest) conductivity.
    2. Column 2: Indicate the conductivity value of each.
    3. Column 3: Classify each sample as a nonelectrolyte or electrolyte.
  2. Summarize what happens to conductivity when a solution is diluted. Describe any numerical proportion between dilution and conductivity.
  3. Based on your data, would you expect unknown X to be a 0.02 M solution of table sugar (C12H22O11) or KCl? Why?
  4. Based on your data, could unknown Y be a 0.02 M solution of CH3OH or NaNO3? Why?
  5. Based on your data, determine the soil’s suitability for sustaining plant growth. Make sure you support your reasoning.
  6. What drink would you recommend for athletes who exercise for extended periods, and why? Use your data and the ingredient labels of the drinks to support your reasoning.


  1. A gardener noticed that some of his plants showed signs of a potassium deficiency. After testing the soil conductivity, he found it was in an acceptable range for those plants. Assuming he carried out the test correctly, can the gardener rule out potassium deficiency? Why or why not?
  2. True or False: 12.0 g of all nitrate (NO3) salts (ie. NaNO3 and Ba(NO3)2) have the same conductivity. Explain.
  3. True or False: 8.0 g of all sugars (sucrose, glucose, lactose, etc) have the same conductivity. Explain.
  4. A soil sample from my front lawn has a higher than normal salinity. What might be two logical reasons for this?

C:\Users\wlammel8\AppData\Local\Microsoft\Windows\Temporary Internet Files\Content.IE5\5O6IUPCX\Birthday1[1].jpg Experiment 6 C:\Users\wlammel8\AppData\Local\Microsoft\Windows\Temporary Internet Files\Content.IE5\5O6IUPCX\cartoon-1294877_960_720[1].png

A lesson on gas laws

We’ve all heard of Boyle’s Law and Charles’s Law, but have you ever heard of Cole’s Law?

It’s thinly sliced cabbage.

Experiment adapted from “Target Gas Law Lab” by Flinn Scientific (Pub # 91654-061616)

Personal Protective Equipment Requirements

Safety guidelines due to Covid-19 have to be followed at all times!

Instructor Demo


Safety Glasses


Gloves Nitrile


Fume Snorkle


Instructor Always


Lab Coat


Gloves Heavy


Bio Hood



Students will be able to:

  • Explain the assumptions made when describing an ideal gas
  • Evaluate laboratory data to calculate ideal gas law constant
  • Critique laboratory protocols to identify potential issues with experiment
  • Analyze laboratory data to determine appropriate quantitative results
  • Synthesize information (descriptive and quantitative) from laboratory experiment to propose novel alternative protocols.


Our understanding of the behavior of gases started with Galileo in the 1600s and his development of the thermometer. This invention looked at the expansion of liquids as they warmed as a way to quantify the amount of heat in a substance. His student, Torricelli, took this concept one step further and invented the barometer in 1644. This was the first attempt to measure the pressure caused by the earth’s atmosphere. (Fig. 1)

Barometer Barostar of Torricelli - Catawiki

Original thermometer invented by Fahrenheit offered at auction
Fig. 1—Torricelli’s barometer Fig. 2—Fahrenheit’s Thermometer

Pascal took this one step further and used this new barometer to study the changes in atmospheric pressure as a function of altitude. He discovered that air pressure decreases as the height above sea level increases.

Daniel Farenheit developed an improved thermometer with mercury as the liquid (see Fig. 2) which allowed a greater range of temperatures to be investigated. He developed a uniform temperature scale based on a mixture of salt/water (0 oF) and body temperature (he assumed 100 oF, actually 98.6 oF).

In 1787, Jacque Charles looked at the relationship of temperature and volume and discovered the inverse connection of the two variables: . Just after the turn of the century, John Dalton’s interest in weather motivated him to study gases. He discovered that the pressure of a mixture of gases is the sum of the individual pressures (remember, he is also the one that came up with an early theory of atoms).

Joseph Louis Gay-Lussac in 1808 formulated two gas laws: one relating pressure and temperature and the second looking at the volumes of gas reactants and products in a chemical equation. Soon (1811), Avogadro extended this work proposing a theory that gases are composed of small molecules. Therefore, equal volumes of gases at the same temperature and pressure must contain the same number of molecules.

Finally, Emile Clapeyron combined all of these laws into what we know as the Ideal Gas Law.

(PV = nRT). R is a constant whose numerical value will depend on the units used for the other variables.

Your task will be to investigate a reaction that produces gas and from the volume measured, determine the gas constant and compare these values to literature.


  1. Perform the following conversions to the proper number of significant figures
    1. Pressure—The pressure in Rochester last Thursday was 30.06 inches of Mercury. Look up the appropriate conversions and give the answer in:
      1. Torr
      2. Atmospheres
      3. mm Hg
      4. Pascals
      5. Psi (pounds per square inch)
    2. Temperature—At the same time, the temperature was 62 oF. Convert to:
      1. oC
      2. K
      3. oR (Rankine)
    3. Volume—A trenta beverage at Starbucks © is 31 ounces. Convert this to:
      1. mL
      2. cc (cubic centimeters)
      3. cups
      4. L
      5. Dram (term used to measure whiskey)
  2. What are three assumptions that you make when using ideal gas law?
  3. In this experiment, you are reacting zinc metal with hydrochloric acid to form hydrogen gas and zinc chloride. Write and balance the chemical equation.
  4. Using your reaction of #3, if you react 1.0 g of Zinc metal, how much volume of gas will you generate at STP.


2.0 M HCl (20 mL per student pair)

Zinc metal strip (1 gram per student pair)

100 mL beaker

250 mL Erlenmeyer flask

Gas collecting set-up

100 mL graduated cylinder

Ring stand

Clamp for ring stand to hold cylinder

Bent glass tubing and stoppers (see below)

Balance (good to 0.01 g)


Plastic wrap

SET-UP (approximate)

C:\Users\wlammel8\AppData\Local\Microsoft\Windows\Temporary Internet Files\Content.Word\IMG_6028.jpg


  1. Fill trough at least ½ full with water
  2. Fill your graduated cylinder totally with water. Cover with a piece of plastic wrap, invert and put end into trough under water level. You should be able to remove the plastic wrap and the water stay in the cylinder.
  3. Mass out about 0.1 grams of zinc. Set aside
  4. Set up the rest of your apparatus making sure the end of the tubing goes into the graduated cylinder.
  5. Pour 5 mL of acid into your Erlenmeyer.
  6. Make sure your snorkel is over your flask.
  7. Put the zinc in the flask and immediately stopper the flask. The gas should be bubbling into the graduated cylinder.
  8. Take the temperature of the water in the trough.
  9. When reaction is complete, read the gas level on the graduated cylinder. Be sure to read to 0.1 mL.
  10. Clean out the flask and repeat 2 more times (refilling graduated cylinder with water each time and using new acid).


  1. Look up the atmosphere pressure for Rochester for the day you do the lab. Convert to the proper units (atm).
  2. Convert your volume of gas to L.
  3. Convert your temperature to Kelvin.
  4. Using the ideal gas law, use the known value of “R” to calculate the moles of gas produced.
  5. From the amount of zinc used (you may assume excess acid), calculate the moles of gas that should have been produced.
  6. Determine % error for each trial
  7. Determine the average error for this experiment.
  8. You will likely be off. You are assuming at the gas in the graduated cylinder is pure hydrogen. This is an error—why? (hint—it relates to “muggy air” in summer)
  9. See if you can find a way to correct for this error. Detail your calculations. Recalculate your error.
  10. There is at least one other assumption that you are making in this experiment. Think about it and how it might change your results.


  1. If the zinc had an oxide coating, how would it impact the results (would your moles of hydrogen (calculated or actual) be too high or too low and why?
  2. Based on your results, what is the maximum amount of zinc that you could have used and still have the gas fit in a 100.0 mL cylinder?
  3. There was air in your flask (and not hydrogen) when you conducted this experiment. Why does this not impact your results?
  4. To impress your instructor, you repeated the experiment with 0.1 g of aluminum.
    1. Write down the balanced reaction
    2. Calculate the amount of hydrogen (in mL) that should have been generated at 30.13 in Hg (pressure) and 21.9 oC.
  5. If I had used 1.0 M HCl by mistake, would there have been enough acid to react with 0.1 g of zinc? Show calculation.
  6. If, by mistake, I had grabbed a cylinder that had not graduation markings, could the experiment be saved? Explain how.

CHM 140L-Fa 20

Ideal Gas Lab



Zinc Powder(g)

mL of Gas










Room Temp


Atmospheric Pressure

30.16 in Hg

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