CHEM 130 Lab Manual Fall 2011

CHEM 130 Laboratory Manual

A Green Approach to Introductory Chemistry

Fall 2011

CHEM 130 Lab 1 Revised 5/2011

Table of Contents

12 Principles of Green Chemistry…………………………………………… 3

Experiments

1. Measurement of Mass, Volume, and Density ……………………………. 4

2. Properties of Matter, Mixtures, and Solutions …………………………... 10

3. Acids and Bases: The Effects of Acid Rain ...…………………………….. 23

4. Titrations: How Much Antacid is Needed to Neutralize Soda? ..……….. 30

5. Chemical Reactions Part I. An Iodine “Clock” ………………………….. 37

6. Chemical Reactions Part II. Metals and Oxidation ……………………... 42

7. Stoichiometry and the Determination of Molar Mass …………………... 47

8. Reaction Rates and Catalysis: Hydrogen Cars and Other Stuff ..……… 48

9. Liquids, Solids, and Intermolecular Forces ……………………………… 74

10. The Ideal Gas Law: Can We Make a Greener Airbag? ..……………… 81

Appendix

Chem 130 Laboratory Technique Guide ………………………………...…. 89


Twelve Principles of Green Chemistry *1. Prevention

It is better to prevent waste than to treat or clean up waste after it has been created.

2. Atom EconomySynthetic methods should be designed to maximize the incorporation of all materials used in the process into the final product.

3. Less Hazardous Chemical SynthesesWherever practicable, synthetic methods should be designed to use and generate substances that possess little or no toxicity to human health and the environment.

4. Designing Safer ChemicalsChemical products should be designed to effect their desired function while minimizing their toxicity.

5. Safer Solvents and AuxiliariesThe use of auxiliary substances (e.g., solvents, separation agents, etc.) should be made unnecessary wherever possible and innocuous when used.

6. Design for Energy EfficiencyEnergy requirements of chemical processes should be recognized for their environmental and economic impacts and should be minimized. If possible, synthetic methods should be conducted at ambient temperature and pressure.

7. Use of Renewable FeedstocksA raw material or feedstock should be renewable rather than depleting whenever technically and economically practicable.

8. Reduce DerivativesUnnecessary derivatization (use of blocking groups, protection/ deprotection, temporary modification of physical/chemical processes) should be minimized or avoided if possible, because such steps require additional reagents and can generate waste.

9. CatalysisCatalytic reagents (as selective as possible) are superior to stoichiometric reagents.

10. Design for DegradationChemical products should be designed so that at the end of their function they break down into innocuous degradation products and do not persist in the environment.

11. Real-time analysis for Pollution PreventionAnalytical methodologies need to be further developed to allow for real-time, in-process monitoring and control prior to the formation of hazardous substances.

12. Inherently Safer Chemistry for Accident PreventionSubstances and the form of a substance used in a chemical process should be chosen to minimize the potential for chemical accidents, including releases, explosions, and fires.


* Anastas, P. T.; Warner, J. C.; Green Chemistry: Theory and Practice, Oxford University Press: New York, 1998, p.30. By permission of Oxford University Press.

Experiment 1: Measurement of Mass, Volume, and Density

IntroductionHave you ever wondered what happens to fish in

the winter time? When you go to a lake in the winter you can see how the top is frozen but if you’re like me you wonder what it looks like down below. Thankfully for fish and other creatures that live in the water, ice floats on top of water so the fish can continue to live deeper in the water under the ice. But this brings up an important question. Why does ice float on top of the water?

In the lab this week you will be learning about several common measurements in the laboratory and you will be using them to measure density. Density is a fundamental property of materials that differs from one substance to another. The density of a given substance may change with temperature. The units of density are those of mass divided by volume. So essentially, density is how much stuff you have in a given space. Mass is just a measure of how much of something you have and volume is how much space it takes up.

So this week’s lab is going to focus on studying density. We are going to look at how the density of water changes as a function of temperature. We are also going to be measuring the density of several different liquids. You will then identify an unknown liquid based on whether it sinks or floats on top of a liquid of known identity.

SafetyThe materials in this lab are not harmful to the environment. Do not eat or drink

in the lab. Be sure to wear safety glasses while performing the experiments. Isopropyl alcohol is flammable. Do not pour any of the materials down the sink since most of them can be reused in the future. Do not put oil down the sink (it is bad for the pipes).

Experimental ProcedureThis lab is going to consist of three parts. Part A of the experiment is going to

explore the density of water as a function of temperature. Part B will look at the density of several common liquids. Part C will involve identifying an unknown liquid based on what your experimental data from part B and an experiment you design yourself.

Part AYour instructor or T.A. will give you instructions on how to use an electronic balance to weigh out your samples for this experiment. Also take a look at “Appendix A” before coming to lab. It tells you how to use a balance as well.


1. You are going to pre-weigh a 25 mL plastic graduated cylinder on the top loading balance. Be sure that you are comfortable with the process of weighing a liquid. Please check the “laboratory techniques” section for addition instruction in how to do this.

2. Then you are going to add 10 mL of water to the graduated cylinder and measure the mass again.

3. By subtracting the mass of the cylinder from the mass of the cylinder + water you can get the mass of the water contained in the cylinder.

4. Once you have the mass of the water inside the cylinder measured, you are going to place the graduated cylinder in the freezer.

5. Note what time you put the cylinder with water in the freezer. 6. You will finish up part A after you have completed parts B and C.

Part BYou now have some practice with the process of measuring volumes and masses of liquids. Now you are going to do it with three liquids in addition to water.Density of Common Liquids

1. Measure the mass of your second graduated cylinder. 2. Acquire 5 mL of Karo syrup by pouring from the container into your graduated

cylinder. It is a very thick liquid so the pouring process may be a little tricky.3. Measure the mass of the cylinder with the liquid and subtract the difference.4. Pour the karo syrup into the “used karo syrup” container. Do not pour the karo

syrup into the container that you got it out of. This is a standard practice in chemistry. Never return chemicals to the container that you got them out of to prevent contamination.

5. Rinse out your graduated cylinder with soap and water. Use a cloth or paper towel to wipe out the inside of the cylinder so it is dry.

6. Now measure the mass of your empty cylinder again. This time acquire 5 mL of cooking oil and follow the same procedure as before.

7. Pour the cooking oil into the “used cooking oil” container. Wash out the cylinder with soap and water and then dry as before.

8. Repeat the process finally with rubbing alcohol to determine the mass of 5 mL of rubbing alcohol. Pour the rubbing alcohol into the “used rubbing alcohol” container when you are finished.

9. Calculate the densities of all of the different liquids before continuing with part C.

Part CThis is where you get to design your own part of the experiment. You are now armed with information about the density of four different liquids. Design an experiment that would allow you to identify an unknown liquid. You will need to have two different methods that verify that your data support your proposed identification of the liquid.

Experimental HypothesisWhen you are designing your own experiment you need to have a hypothesis. Every experiment that a scientist performs has a scientific question behind it. For example, an atmospheric chemist might ask: why is the sky blue? Your hypothesis is your predicted


answer to the scientific question behind your experiment. So the atmospheric chemist in our example might say “ It is hypothesized that the sky is blue because of the way the sun interacts with gas molecules in the air”. The scientist would then proceed to test this hypothesis by performing experiments to see whether the data support or reject their hypothesis. It is totally OK if your hypothesis is not supported by the experiment. This happens all the time in science! Developing a hypothesis and then performing experiments is a standard practice in scientific research, so it’s a good idea for us to start doing this now with our experiments.

Experimental Design Hints-You have already collected quite a bit of data on the density of liquids. For one of your procedures you could use the experience you already have to figure out the density of your unknown liquid. (Don’t forget to write a hypothesis for this experiment!)-What happens if I take two liquids of different densities and put them in the same container? Which one will be on top and which one will be on the bottom? You may be able to use this information to design an experiment to help identify your unknown. If you go this route, you may use any liquid you like from the four that we used in the lab. ONLY USE 2 mL of each liquid in any mixture experiment. Your hypothesis for this experiment might be something like: “It is hypothesized that when unknown liquid A and water are mixed, liquid A will float on the surface of water due to its lower density”).*Note: Physical characteristics of the liquid can be helpful in the identification process. However, just saying that it has the right viscosity to be liquid X is not sufficient evidence by itself but may be helpful in addition to your other two experiments.Once you are done with your experiments pour any pure unknown liquid into the reuse bottle for that unknown liquid. Any mixtures of liquids you may have made can be poured into the “mixtures” bottle.

The rest of Part ANow that you have completed the rest of the experiment, go and check on your graduated cylinder with 10 mL of water.

1. Read the current volume of the water (now ice) in the graduate cylinder. Note any differences on your data sheet.

2. Measure the mass of the graduated cylinder after freezing. Note any difference in the mass on the data sheet.

3. Complete the rest of the calculations on the data sheet as well as the Post Lab questions.

***What to do with your Materials when you are doneWe want to reuse as much of the material as possible from this lab. DO NOT just pour your materials down the sink after you are done with individual parts of the experiment. Instead, put the liquids in the reuse containers labeled for each liquid (including the unknowns.) References and Additional Reading

1. scifun.chem.wisc.edu/HomeExpts/layeredliquids.htm2. www.stevespanglerscience.com


3. Picture Credit: http://www.nsf.gov/discoveries/disc_images.jsp?cntn_id=100659&org=NSF

Experimental Data (Example of how you might enter it in your lab notebook)

Part A

a.) Mass of Empty Cylinder: ____________________________b.) Volume of H2O ____________________________c.) Mass of Cylinder + H2O: ____________________________d.) Mass of H2O: ____________________________e.) Time Placed in Freezer: ____________________________

Sample in freezer------------------------------------------------------------------------------------------------------------

f.) Volume of ice ____________________________g.) Mass of Cylinder + Ice ____________________________h.) Mass of Ice ____________________________

Part B

Liquid Mass of Cylinder (g)

Volume of Liquid (mL)

Mass of Cylinder +Liquid (g)

Mass of liquid (g)

Density (g/mL)

Karo Syrup

Oil

Rubbing Alcohol

Part C (Include your experimental data for part C here!)


Unknown Liquid Label _________ Identity_________________


Experiment 1. Post Lab Questions

1. Calculate the expansion in volume of water as it goes from liquid water to ice. The equation to do this is as follows:

The difference in volume here is the difference in volume between water and ice.

2. Why does ice float on top of water? Explain this in your own words from what you learned in this experiment.

3. What were your hypotheses for the experiments in Part C?

4. Describe how you determined the identity of your unknown liquid.

5. Where your hypotheses supported by your experiments? Explain.


Experiment 2: Properties of Matter, Mixtures and Solutions

IntroductionHave you ever seen the show CSI? You probably

have. It is a very exciting show which has become extremely popular. Why is it so cool? Whatever reason people like it, it is one of the ways many people have been exposed to the science of forensics (at least the Hollywood version of it!). Many of the tools used in forensics to identify who committed a crime involve the use of chemistry. You are going to learn about some of the processes used by many scientists (including real forensic scientists) to separate mixtures of chemicals.

Many things in the world around us are made up of chemicals. Chem 130 lab is not your first experience with chemistry, you have been dealing with chemistry all of your life! By the way, your life itself exists in part due to chemistry. To begin to talk about chemistry in a scientific way, we should talk about some terminology. It will be very helpful for you to know and remember that chemistry is a science that involves a lot of new language for you to learn. Treat it like a mixture of a science class and a language class and this will help a lot. So let’s start with some terms:

Matter: Anything that occupies space and has mass. No need to be tricky here. It basically just means stuff. Matter is just something that has characteristics that we can observe. Can you think of some examples of this? You should be able to think of a lot!Substance: Webster defines this as matter of particular or definite chemical composition. OK… maybe that doesn’t help too much. In fact, a substance is something that is made up of all of the same stuff. Chemically it is all the same. An example would be distilled water. This is just water, there isn’t any other stuff in it. Compound: Your textbook defines this as: A substance composed of two or more elements in fixed, definite proportions. So in other words, this has different elements which are bonded together in specific ratios (example, water is H2O, it has two hydrogens for every one oxygen and it is always this way). A compound cannot be separated by physical means (like boiling or filtering or something).Mixture: A mixture is two or more different kinds of substances combined in variable proportions. In other words, now instead of being bonded together like we saw in a compound the different substances just occupy the same physical space. Think about a bowl of lucky charms. You have cereal pieces, marshmallows (definitely the best part), and milk. They are all in the bowl together mixed up. But, you could separate them (say you filter the cereal and marshmallows through a screen or something and then pick out the marshmallows). In the lab it’s the same way. Mixtures can be separated by physical means (in our cereal example we separated them based on their ability to go through a screen). Physical Properties: Those properties that a substance displays without changing its composition.We have already hinted at this. Physical properties are things like size, and


physical state (solid, liquid, or gas). There are other physical properties that we can use like magnetism and whether something will dissolve in water or not. It is very important to note though that exploring a physical property does not change the chemical composition of the substance. Chemical Properties: Properties that a substance can display only through changing its composition. So when we explore chemical properties we are actually changing the composition of the substance we are looking at. Solution: A mixture of two or more substances that is completely uniform (the word we often use is homogeneous, which just means that it is perfectly mixed without any variation in the mixture (think creamy peanut butter). (The opposite of homogeneous is heterogeneous which would be like chunky peanut butter. The mixture is not uniform but variable.). Most of the time when we talk about solutions, one of the substances in the mixture is a liquid. The solutions that you will be dealing with in this lab are water based solutions.Soluble: This means that a solid will dissolve in a liquid forming a homogeneous solution.

Alright! Enough with the definitions already! Let’s talk about the experiment now. You are going to be given a mixture of 4 different substances. This is very

similar to a sample that a technician may have collected at a crime scene. Your task will be to separate the mixture based on physical properties of the components and then try to identify a scenario that may fit with what you found in your sample. The district attorney and the director of the crime laboratory have several ideas about what happened. Once you have separated your mixture and identified the components, you will be able to tell them which scenario fits with your data.

SafetyEven though the materials used in this lab are not toxic you are still working in a chemistry laboratory and therefore need to be careful. Make sure that you wear safety glasses when performing the experiment. Do not eat or drink in the laboratory. Only heat water by itself and add to the mixture, do not heat any solutions of the mixture.

Experimental ProcedureYou are going to use about 1 gram of mixture. It does not need to be exactly a

gram, but you need to know the exact mass of your sample. Separation of the Mixture and Determination of Components1. To start your separation scheme take a look at the mixture. What does it look like you have in there? 2. Take a look at the tools you have in front of you. You should have a small sieve (essentially a piece of window screen with a handle), a magnet, a glass filter funnel, and funnel paper. You should also have a beaker and an Erlenmeyer flask.3. Notice if any of the materials seem large enough that you could screen them out of the mixture. If you do have things that are large enough to separate with a screen, do this before you do anything else to the mixture. Make observations about the material(s) that you were able to separate with the screen.


4. Now that you have separated out any larger material, consider your next step. You have a magnet in front of you and also a funnel and filter paper. Before we get our mixture wet, we should try to see if anything in the mixture is magnetic. Wrap your magnet in a plastic bag and pass it through the mixture. Does it pick anything up? Make observations of the material(s) that you were able to separate with a magnet.5. You have now done some separations with your mixture using both size and magnetism as separation factors. Now consider a new factor, solubility. Do you think there is anything in there that is soluble in cold water? Place a filter paper in a glass funnel and put your mixture in the paper. Pour cold water through the mixture. (WARNING!!!! Don’t add too much water! If you do end up dissolving your material you want it in a minimum of water so that the water can easily be evaporated later!!!) If there are no longer any crystals in the mixture, then your compound is soluble in cold water. If your compound is soluble in cold water, you don’t need to worry about adding hot water to your mixture.6. We know that sand will not dissolve in water, so if you have sand present in your mixture it will certainly not dissolve in cold water. Did your other substances dissolve? If not, the next step is to try boiling hot water and see if this will dissolve the other substance. BUT!!! Don’t use too much water (more than about 20 mL) otherwise you will have a hard time getting the crystals back out of the water.NOTE: You will NOT be given something that dissolves in cold water and something else that dissolves in hot water. It will be only one or the other.7. If your material dissolved in cold water, you will have to evaporate all of the water off using an evaporating dish and a hotplate. Once it gets close (there is only a puddle of water in the bottom of the dish about the size of a dime) pull the dish off of the hotplate and allow the heat of the plate to evaporate the rest of the water. Otherwise it will splatter your component all over the place. 8. If your compound was not soluble in cold water but soluble in hot water, place the solution in an ice bath until your component precipitates (crystalizes) and then filter it out of the water.9. Once you are done with your separation, allow all of the components to dry and then get their masses. Make sure you include all of the data from your experiment in your laboratory manual. See the experimental data page for an example of the type of data you want to collect and how you might want to organize it.10. Once you have identified all of your components, select the appropriate crime scene scenario from the options on the next page.

*Once you are done with the experiment, place all components in containers so that they can be reused.

Physical properties of Possible ComponentsSand: Do I really need to describe this one? It is small solid crystals that are not soluble in hot or cold water. Commonly found in the desert.Steel Shot: Small round metal pieces that are shiny and silver colored. These are fairly large compared with many of the other components and may be separated through the use of a screen.


Lead Shot: Small round metal pieces that have a dark grey color. These are fairly large compared with many of the other components and may be separated through the use of a screen.Steel Filings: These are very small shiny silver colored pieces of metal that would go through a screen easily. These are magnetic, and so they could be easily separated with a magnet. Copper Filings: These are very small shiny copper colored pieces of metal that would go through a screen easily. These are not magnetic, and so they could not be separated with a magnet. Iron Filings: These are very dark black pieces of metal that would go through a screen easily. These are magnetic, and so they could be easily separated with a magnet. Benzoic Acid: This is a white powder that is used as a food additive. These crystals are not very soluble in water. In cold water they do not dissolve at all, but they will dissolve in boiling water. On cooling the boiling water in an ice bath they will form a solid crystal again.Powdered Sugar: This as you well know is a white crystalline solid. Sugar is soluble in cold and hot water. The only way to get it back out of the solution is to boil off all of the water.

Possible Crime Scene Scenarios

1. A drug deal was going on in the desert. One of the suspects used a shotgun which contained lead shot and shot a vehicle leaving traces of copper metal. The “drugs” were actually a fake. Instead of being cocaine, the drug dealer substituted sugar which is a white powder that is soluble in cold water.

2. A drug deal was going on in the desert. One of the suspects used a shotgun which contained steel shot and shot a vehicle leaving traces of iron metal. The “drugs” were actually a fake. Instead of being cocaine, the drug dealer substituted sugar which is a white powder that is soluble in cold water.

3. A drug deal was going on in the desert. One of the suspects used a shotgun which contained steel shot and shot at a vehicle leaving traces of iron metal. The “drugs” were actually a fake. Instead of being cocaine, the drug dealer substituted benzoic acid which is a white powder that is soluble in hot water but not soluble in cold water.

4. A drug deal was going on in the desert. One of the suspects used a shotgun which contained steel shot and shot at a vehicle leaving traces of steel metal. The “drugs” were actually a fake. Instead of being cocaine, the drug dealer substituted sugar which is a white powder that is soluble in cold water.

References and Additional Reading1. Suits, Jerry Chemistry 111 Laboratory Manual, University of Northern Colorado

20092. Max M. Houck, Jay A. Siegel, Fundamentals of Forensic Science Academic Press,

20063. Picture credit: http://www.canstockphoto.com/police-line-hazmat-0469933.html


Experimental Data

Collect the following data during your experiment:

a) Mass of starting mixture: __________________________________

Component Separation Technique Mass of Dry Componentb)

c)

d)

e)

f) Total mass of components: _________________________________g) Mass lost in separation(a-f): _________________________________h) Percent Recovery ((f/a) x100%) _________________________________


Experiment 2 Postlab Questions

1. Describe your process for separating the mixture. Which techniques did you use and which order did you use them?

2. Is there anything you would have done differently in your separation?

3. Which possible crime scene scenario seems to best fit your data? Why?


Experiment 3: Acids and Bases

IntroductionYou have probably heard about acid rain on

the news. This is a very serious environmental phenomenon that has devastating effects. But what exactly is an acid?

You may have even seen a movie or two where some fight scene results in the bad guy being tossed into a vat of boiling acid. So are all acids bad?

The purpose of this week’s lab is to learn about acids and bases. An acid is a molecule which is capable of producing protons or H+ in water. These protons can then react with all kinds of substances that they may come into contact with. A base is a molecule which produces hydroxide or OH-

in water. Bases will actually take protons from things that they come into contact with.

Acids are actually more common than you think. Many of the foods that you enjoy every day contain acids. Things like soda pop, sour candy, lemons, vinegar, even grape juice. Often when you eat things that have an acid in them it gives the food a sour taste.

Your body also contains a number of acids. You have acid in your stomach to help protect your body and digest food, and tears are acidic too (that is why they make your eyes red when you cry). In the case of the body, usually these acids are there as a protective mechanism to prevent infection. So not all acids are bad! However, as you will be learning about this week, the environment can be severely impacted by acid present in rain.

From a chemistry standpoint, here is what happens when you add a normal acid to water:

HX + H2O → H3O+ + X-

The X here could be any number of things (your text has a list of different acids where everything else but the proton that is being donated could be written as X, see page 488). The species H3O+ is called hydronium, and this is what form an acid exists in when it is formed in water. Some of the strongest acids are things like H2SO4, sulfuric acid, which is the acid in the battery of your car. This is a very powerful acid that can dissolve metal and burn your skin very badly. Another acid we will talk about is HNO3, nitric acid, which is used in metal working and is also used to manufacture explosives.

We have talked about a number of common acids, but at this point you may be wondering how acid rain forms. Are the things produced in pollution acidic? The figure on the next page should help give you a picture of how acid rain forms.


Burning sulfur in the presence of oxygen is just an oxidation reaction. When this happens you form sulfur dioxide and sulfur trioxide. Both of these are gases.

S(solid) + O2(gas) → SOx(gas) (where x can be 2 or 3)

Watch what happens when sulfur trioxide (SO3) is reacted with water in the following chemical equation:

SO3 + H2O → H2SO4 (sulfuric acid!)

This is the same strong acid that we just said was in car batteries! So when we burn sulfur in the presence of oxygen and those gases escape into the atmosphere they form sulfuric acid in clouds. This sulfuric acid then precipitates as acid rain. The same process with NO2 produces HNO3 (nitric acid) as well. Nitrogen dioxide is the reddish brown gas that gives smog its characteristic nasty color.

These reactions can actually be done in the laboratory, but why would we want to do that? The environment already has to deal with enough acid rain without us adding more!

Instead we are going to study acid rain where we have prepared a solution of acid from environmentally safe chemicals. We will use this to study the effects of acid rain on


lakes, plant matter, and also marble statues. Remember though, that since we are doing this in that laboratory, it may not perfectly represent what happens in the environment. Take for example the amount of materials we use. Don’t just assume that because one simulated “rainstorm” causes the lake to become acidic that this is what happens on a larger scale in the real world. The damage from acid rain takes time!

Before we get to the experiment itself, there is another thing we need to talk about. You have already learned about concentrations in your lecture course. How much of something we have is important to know. This is true for acids too. Often we want to know how much acid we have in a given solution. We could talk about the concentration in terms of how many moles of acid we have per liter of solution like you are used to talking about with other chemicals. But acids present a problem with this. The range of concentration with acids is really huge and can include some really small numbers. For example, you might have 1 x 10-7 moles of H+. 10-7 is a really small number isn’t it! To simplify things, chemists developed a shorthand method to talk about the concentration of acid in a solution. We like to talk about the negative log of the H+ concentration. We call this the pH. This is what happens when we take the negative log:

-log (1x10-7) = 7 (so this would have a pH of 7)

Isn’t that a lot nicer to work with? Now we just have a single digit to work with to tell us how much acid is in the solution. But, sometimes this can get a little tricky. For example, we might have a solution that has 1 x 10-4 moles of H+ per liter. 10-4 is actually a bigger number than 10-7 if you remember back to what you have learned about exponents (if you don’t believe me, I will write them out without using scientific notation: .0001 (which is 1 x 10-4) vs. .0000001(which is 1 x 10-7)). If we take the negative log of 1 x 10-4 we get 4. This means that: the lower the pH, the higher the amount of acid in the solution. I know this can be tricky but it is an important thing to remember.

When we talk about the pH scale and acids and bases it is important to talk about where acids and bases fall on the pH scale. A pH of 7 is actually neutral. This is neither acidic nor basic. Any pH below 7 is considered acidic. Any pH above 7 is considered basic. Now that we have talked about this, we can use pH to help us study acids and acid rain.

SafetyYou will be working with a vinegar solution in this laboratory. You should wear

safety glasses and avoid getting the vinegar solution in your eyes. Do not eat or drink in the laboratory. All solutions can go down the drain once you are finished with the experiment.

Experimental ProcedureIn this experiment you are going to be doing three different activities. Part A

involves exposing a “lake” of red cabbage juice to “acid rain” while monitoring the pH using a computer and an electronic pH probe. In Part B you will be exposing apple peels, pieces of chalk, and marble chunks to the same acidic solution and observing the effects.


Part AFor this part of the experiment you will be working with a partner. There are only

four computers in the lab so partnerships will have to take turns using the computers while other students go on ahead to part B while they are waiting.

1. Acquire a 100 mL beaker and fill with 25 mL of red cabbage solution.2. Open the Logger-Pro software package on your computer. Once you open the

program, the computer should recognize that a pH meter is attached to the computer. Check the plot on the screen to see if pH is on the y-axis and time is on the x-axis.

3. Adjust the time to 600 seconds to give you enough time to run the experiment before the computer stops collecting data.

4. Rinse the pH probe with distilled water (do this every time it changes solutions!!!). Place the pH probe into the solution of red cabbage juice. Record the initial pH of the solution

5. Acquire a 50 mL beaker and add about 25 mL of vinegar to the beaker.6. Click collect with the computer7. Take an eye-dropper and begin to add drops of vinegar to the red cabbage

solution.8. Record information about the color of the solution and what the pH is doing on

the computer.

Part BFor this part of the experiment you will work either in pairs or individually.

There are three different experiments in part B. 1. Take two pieces of fresh apple peel and place them in test tubes. Label one of

these “acid” and the other “control”.2. Add 5 mL of vinegar to the test tube labeled acid. Add 5 mL of water to the test

tube labeled control. 3. Record your observations of what happens to the apple peel.4. Select two marble chunks and get the mass of each one. Place one of them in a

tube labeled “control” and the second one in the tube labeled “acid”. Make sure you keep track of which piece (mass) is going into which tube!

5. Add 5 mL of vinegar to the acid sample and 5 mL of H2O to the control.6. Now take two pieces of chalk and repeat the same process. Be sure to get the

masses of each piece of chalk before adding acid and water. Prepare an acid sample and a control each with 5 mL of liquid.

7. Record your observations of what happens to the marble chunks and the chalk.8. After 30 minutes, pull the marble chunks out of the tubes and let them dry.

Measure the mass of the chunks to see if there is any change.9. Thought questions: Why does the acid cause a change in the apple peel? Why

does the acid cause the observed changes in the marble chunk? Why is the reaction with chalk more aggressive than the reaction with the marble chunk?

References and Additional Reading1. http://www.haverford.edu/educ/knight-booklet/acidrain.htm2. Picture Credits: http://www.maltaweather.info/pollution.html, http://www.maine.gov/dep/air/acidrain/3. http://www.epa.gov/acidrain/


Experimental Data

Part A

Initial Color of Solution

Initial pH of Solution

Volume of Acid added (mL)

Final Color of Solution

Final pH of Solution

*Include a printout of your pH plot in your lab notebook.

Part B1. Observations of apple peel in acid:

2. Mass of Marble chunks: Control__________ (g) Acid Sample _________(g)

3. Observations of Marble chunk/acid reaction:

4. Observations of Marble chunk/water reaction:

5. Mass of Marble chunks after experiment: Control________(g) Sample________(g)

6. Mass of Chalk pieces: Control__________ (g) Acid Sample _________(g)

7. Observations of Chalk /acid reaction:

8. Observations of Chalk/water reaction


Experiment 3 Post Lab Questions

1. Was your hypothesis supported by the experimental data? Explain.

2. Based on your observations, how long do you think it takes for acid rain to affect the pH of a lake? (Think about the volume of a lake vs. the volume of rain that will be added in a storm) Explain.

3. Based on your observations, how long do you think it takes for acid rain to affect plant life? Explain.

4. Based on your observations, how long do you think it takes for acid rain to affect marble statues? Explain.

5. You probably noticed some bubbles forming with both the chalk and the marble chunk experiments. What do you think those bubbles are? (Hint: the formula for both chalk and marble is the same, CaCO3. What common gas do you think you can form from this compound in acid?)


Experiment 4: Titrations

IntroductionHave you ever wondered how

someone can test water for the presence and concentrations of acid or different types of metals? This is very important for the study of natural waters and their health. Sometimes scientists use very advanced equipment for these types of tests. But there is a method that many of them still use, and that is the method of titration.

In general a titration is a technique where you add a reacting species to a sample and use the stoichiometry of the reaction to determine the concentration of something of interest in the sample. In the case of determining the concentration of acid, a base is added to a sample of acid in the presence of a pH indicator. Hopefully the indicator (if well chosen) will change color at the same time that the concentrations of acid and base are equal. The pH where the indicator changes color is called the endpoint. It is very important that an indicator is carefully selected so that the endpoint occurs when the concentration of acid and base are equal. Titrations can also be used to determine how “hard” a water sample is (how much calcium and magnesium are present) as well as concentrations of heavy metal ions.

Typically, when titrations are done in a general chemistry laboratory, a strong acid and a strong base are used. However, since we are trying to teach chemistry principles with “green” chemicals, we are going to have to be a little creative with our titration. For example, instead of using a synthetic indicator we can use a natural indicator. You worked with one of these in the acid/base lab when you used red cabbage juice in one of your experiments. This week you are going to explore another natural indicator: the anthocyanins in grape juice. You will be using sodium carbonate (used as washing soda, and also used to raise the pH in pools) as a base, and Sierra Mist as the acid. All of these chemicals are relatively safe to work with, and are not harmful to the environment. Also of interest, is the fact that sodium carbonate is very similar to calcium carbonate which is an ingredient commonly found in antacids. So this experiment is in a sense looking at how much antacid is required to neutralize the acid found in soda pop. Why didn’t we use calcium carbonate? It is because this compound is not very soluble in water so it would be hard to use for a titration!

SafetyAcids and bases can be very damaging to your eyes. It is very important that you

were your safety glasses when working with these chemicals. All reagents can go down the sink upon completion of the experiment. Do not eat or drink in the laboratory.


Experimental ProcedureThe first part of this experiment will involve determining how much base it takes

to neutralize the grape juice indicator. Grape juice itself is acidic and so this will require some of the base in order to be neutralized. After you have determined how much base is required to neutralize your indicator, you will then titrate a soda pop solution and determine the actual concentration of acid present.

Part I. Titration of Grape Juice.1. The first step in this experiment is to prepare your titration setup. This involves

the use of a buret, a buret stand, and Erlenmeyer flasks. Carefully rinse out the buret several times before using to make sure that it is clean. See Figure 2 below:

2. Now that you have setup your titration apparatus, you need to prepare the base solution that you are going to be using in the titration. You are provided with 0.500 M sodium carbonate (Na2CO3). You need to do a dilution to prepare 100 mL of 0.250 M sodium carbonate. The formula for dilution is M1V1 = M2V2. Measure out the volume of 0.500 M Na2CO3 needed using a graduated cylinder. Pour this into a 100 mL volumetric flask and fill to the line with distilled water. Stopper the flask and mix well.

3. Using a glass funnel, pour just a couple of mLs of sodium carbonate solution into the buret, rotating the funnel inside the buret as you are pouring so that you rinse the inside of the buret with the sodium carbonate solution. Allow this “rinse”


carbonate to drain out of the bottom of the buret into a beaker. This is a great opportunity to play with the stopcock and figure out how it works. Make sure that you can control the flow well, especially to get it to run about 1 drop per second. The sodium carbonate that you collect in the beaker can go down the sink.

4. Now fill the buret up to the 0 mL line with sodium carbonate. You should read this just like an upside-down graduated cylinder. The bottom of the meniscus should be even with the 0 mL line while looking at eye-level. If the liquid level is not exactly at the 0.0 mL line that is ok. It just needs to be close.

5. Measure out 5.00 mL of grape juice using a pipete and add it to a 125 mL Erlenmeyer flask. Measure out 45 mL of distilled water in a 100 mL graduated cylinder and then add the 45 mL of distilled water to the Erlenmeyer flask.

6. Place the Erlenmeyer flask directly under the buret as shown in Figure 2.7. Now carefully observe your starting volume. If it is not exactly 0.0 mL that is

okay, but you need to know exactly what it is. 8. Once you have recorded your starting volume, you can begin the titration. Slowly

open the stopcock on the buret and allow the base to flow into the flask with grape juice at a rate of about 1 drop per second.

9. At the first sight of a color change in the Erlenmeyer flask (it should turn green), shut off the stopcock and swirl the flask. You may find it helpful to place a scrap of white paper below the flask so that it is easier to see the color change.

10. At this point, add the titrant (base) very slowly and swirl with each drop added to the flask.

11. When it seems like the green color is not going away, stop the titration.12. The goal is to add just enough base to retain the green color. You don’t want to

add more than you have to.13. Once it seems as though the color is no longer changing and you have stopped

your titration, take note of where the volume is in the buret.14. Add just one more drop to the Erlenmeyer flask and see if the color changes at all.

If it does not change, you are done and the volume you noted before that last drop is your final volume. If it does change, continue to note the volume and then add an additional drop. Repeat this process until there are no longer any changes in the color of the solution.

15. Now determine your delivered volume. This is going to be the difference between your starting volume and your final volume. Record this data on your experimental data sheet.

16. Repeat this process again for an additional trial. Depending on the volume of base used, you may not need to refill the buret. Just make sure that you record your initial volume before you begin the titration.

17. Once you have completed the second trial, compare the volumes of sodium carbonate needed to neutralize the grape juice. Discuss any concerns you have with your instructor if the two numbers are not very close to the same.

Part II.1. Refill your buret and prepare your titration setup as at the beginning of part 1.

Make sure you carefully record your initial volume of sodium carbonate in the buret.


2. This time you will use 5.0 mL of grape juice, 15.0 mL of water, and 30.0 mL of Sierra Mist in your titration. Add these amounts of each liquid to a 125 mL Erlenmeyer flask.

3. Begin the titration of the soda pop with the grape juice indicator. Once again, you can begin at the pace of 1 drop per second. Once the solution starts to turn green you should slow down as you did when you titrated grape juice alone.

4. This will take much more volume than it did previously.5. Repeat the same process as before once the solution turns green. Once the

solution does not change color when you add a drop of base, you have completed the titration.

6. Repeat this process for two additional trials with soda pop for a total of three different soda pop trials.

7. Hint: Once you have completed your first trial you have a good idea about how much base is needed to neutralize the acid. On your second and third trial you can quickly add base until you are just a few mLs short of where the first titration was complete. Then go slowly during the last couple of mLs to make sure you do not miss the end point. This will save you a lot of time on your additional trials!

8. All materials can go down the sink once the experiment is complete.

References and Additional Reading1. Introductory Chemistry Version 3.1 Escience laboratory manual2. http://www.education.com/activity/article/Explore_Color_Science_with_Grape/3. Picture Credit:

http://www.townofstratford.com/content/1302/402/625/1100/5066/5511.aspx


Experimental Data

Part I.

Trial # Initial Volume Na2CO3 (mL)

Final Volume Na2CO3 (mL)

Volume of Na2CO3 Delivered (mL)

1

2

Average Volume of Na2CO3 Delivered:________________________ (mL)

Part II.

Trial #

Initial Volume Na2CO3 (mL)

Final Volume Na2CO3 (mL)

Total volume-added (mL)

Volume of Na2CO3 added to neutralize soda (mL)

1

2

3

Trial # Volume of Na2CO3

added (mL)Moles of Na2CO3

added (mol)**Moles of acid present (mol)

1

2

3

Average

**Sodium carbonate is dibasic (which means it can accept two protons) and citric acid (the main acid in clear soda) has three different protons it can donate. Thus, it takes 3

moles of sodium carbonate to neutralize 2 moles of acid in Sierra Mist. Once you get the moles of sodium carbonate, divide it by 1.5 and that will give you the moles of acid

present.



1. Was your hypothesis supported by your experimental data?

2. Explain the process of titration in your own words. What are the important details to keep in mind when doing a titration?

3. How much acid was in the soda pop (on average)? Does this surprise you?

4. What effect do you think this amount of acid has on your body? Your teeth?

5. How close where your three trials? What sources of error do you think may have contributed to these differences? (Hint: There are always some errors in your experiment, this cannot be avoided no matter how good a scientist you are)


Experiment 5: Chemical Reactions Part I: An Iodine “Clock”

IntroductionWhen you hear the phrase

“chemical reactions” you may still be thinking about something going on in a big vat in a chemical plant or in a fancy glassware setup in a chemistry laboratory.

But in reality, chemical reactions are going on around and inside of us all the time. This is much more common than you might think. Some chemical reactions go unnoticed, like many of the metabolic reactions that go one inside of your body 24 hours a day, while other chemical reactions attract a great amount of our attention such as a wildfire.

So what is a chemical reaction? Basically when two or more compounds or elements interact with each other in a way in which one or both is changed into something new. These kinds of reactions also often involve heat. They either produce heat like a forest fire (this one produces a lot of heat!), or they absorb (or take in) heat like the instant cold pack that feels so good on an injury.

One of the ways that chemists often explore chemical reactions is through some kind of physical change in the reaction mixture. When you have a lot of specialized equipment available you can explore what is going on with a chemical reaction more carefully. But sometimes you can monitor the progress of a chemical reaction just by watching a color change.

Today we are going to be exploring a very interesting reaction using household chemicals.

We will be studying the reaction between iodine and vitamin C which has the formula C6H6O6. The reaction would be written as follows:

I2 + 2(C6H6O6)(aq) → 2I- + 2(C6H6O6)(aq)

Tincture of iodine is a solution of iodine, water, and alcohol that is about 2 % elemental iodine (I2). This iodine solution is a brown color. When elemental iodine is reacted with starch it forms a very deep blue color. So starch can be used as an indicator for the presence of elemental iodine.

However, iodide (I-) is colorless (clear) in solution. When it interacts with starch it does not form a blue color. We will perform a couple of experiments to look at this reaction in more detail. There is also another reaction that we are going to explore:

2I-(aq) + H2O2(aq) + 2H+

(aq) → I2(aq) + H2O(l)

In this second reaction we are doing the exact opposite of what we did with vitamin C. Iodide is reacting with hydrogen peroxide to form water and elemental iodine.


SafetyIodine is toxic and should not be ingested. Iodine solution will stain skin and

clothing so contact with skin or clothing should be avoided. No food or drink should be consumed in the laboratory. Iodine solutions should be “neutralized” with vitamin C solution (should not have a blue color) before pouring down the sink.

Experimental ProcedureIn the first part of the experiment you are going to explore the effect of vitamin C

on iodine (I2). You will be doing this in the presence of starch so that you get a feel for how starch indicates the presence of iodine. In the second part of the experiment you will be combining both reactions.Part A

1. Measure out 3 mL of iodine solution using your graduated cylinder and add it to a 4” test tube in a test tube rack. All of these items will be provided for you on the bench.

2. Add 2 mL of starch solution to the 3 mL of iodine.3. Mix the tube thoroughly by flicking the bottom of the tube with one hand while

holding the top of the tube loosely in the other hand (make sure you don’t flick it out of your hand!)

4. The resulting mixture should be a deep blue/black color. Record your observations. If the solution appears to be too dark to tell if it is blue or not add 2 mL of distilled water to the solution.

5. Now add 5 mL of vitamin C solution and mix the tube thoroughly6. Record your results on the experimental sheet7. Save your solution for the next part of the experiment8. Thought questions: What happened to the color of the tube? Why did the color

change? Which species, iodine (I2) or iodide (I-) is present in the solution now?Part B

1. Measure out 2 mL of iodine solution and pour it into a 4” test tube2. Measure out 2 mL of vitamin C solution and pour it into the same 4” test tube.

The resulting mixture should be clear.3. Take a second 4” test tube and to this add 2 mL of hydrogen peroxide and 2 mL

of starch solution.4. Pour the contents of test tube A into test tube B. Start the stop watch. Mix well by

pouring back and forth between the two tubes a couple of times.5. Watch the solution carefully. It may take a couple of minutes for anything to

happen but you are going to want to be watching when it does.6. Stop the stopwatch when the solution begins to change color.7. Record your observations.8. Thought questions: Which species (iodine (I2) or iodide (I-) ) is present at the

beginning after you have mixed the two solutions? What is going on during the time you are waiting for the color to change? What species (iodine (I2) or iodide (I-)) is present after the color change? Why is this the major species now?

References and Additional Reading1. http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/MAIN/CLOCKRX/PAGE1.HTM1. Picture Credit: http://www.erh.noaa.gov/okx/okxfirewx.html


Experimental Data

Part A

1. Describe your observations when mixing the iodine solution and the starch solution.

2. Describe your observations when you added the vitamin C solution to the iodine and starch solution.

3. Write the reactions happening in both steps (adding iodine to starch and adding vitamin C to the previous mixture)

Part B

1. How many times did you mix the two tubes?

2. How long did it take for the mixture to change color?

3. Describe your observations of the reaction process.



1. What was going on in the reaction during the time you were waiting for the color change? Think about both reactions that were going on at the same time.

2. Why did the reaction eventually change color?

3. What would happen if we added more vitamin C?

4. What would happen if we added more hydrogen peroxide?


Experiment 6: Chemical Reactions II Metals and Oxidation

IntroductionHave you ever wondered what causes

steel to rust? Or what about tarnish on that brass door plate that used to be nice and shiny? Both of these changes in metal are caused by a process called oxidation. Just like it sounds, oxidation can be the addition of oxygen to a compound. But in general, oxidation means that something is losing electrons (or donating electrons to something else). There are a couple of different ways to remember whether oxidation is a loss of electrons or gaining them. OIL RIG stands for “Oxidation Is Loss of electrons, Reduction Is Gain of electrons” and serves as a very helpful way to remember which is which.

You have probably noticed that not all metals act the same way when they are exposed to “the elements”(a phrase used to describe the outside, things like rain and snow). That is what we are going to be exploring this week, how different metals react with the same oxidizing conditions. As a result of this experiment, you should be able to develop an activity series for the metals you are testing.

Oxidation is a process which generally takes a long time. We are working hard in this lab to use very mild and environmentally friendly materials, this makes studying oxidation a little bit tricky. We have a few tricks up our sleeve to get over this problem though. And these tricks will make this lab more colorful too!

Any time you have an oxidation going on, you have a reduction going on too (the electrons have to go somewhere). In this experiment you will be oxidizing metal and you will also include a dye in your solution which will be reduced (the dye serves as the electron acceptor). When the dye is reduced, it “loses” its color. So say for example you are using a red dye, when it is reduced it will become clear. Different colored dyes have different chemical compositions, so they will not all react to being reduced in the same way (some may be faster than others, some may not change color at all).

SafetyYou will be working with several different metals and vinegar. You need to wear

safety glasses and avoid getting vinegar in your eyes. Food coloring will stain your clothing. Liquid solutions can go down the sink. Metal wool should be placed in the container for each type of metal wool (example: “used zinc wool”).

Experimental ProcedureIn the first part of the experiment you are going to be testing four different dye

combinations to figure out which one is the best indicator of an oxidation process occurring. In the second part of the experiment you will take your indicator of choice and use it to study the oxidation of three different metals.


Part A

1. First you are going to acquire 4 different 6” test tubes. Make sure that the test tubes are very clean before you begin (but they don’t have to be dry).

2. Place the test tubes in the test tube rack and add 5 mL of distilled white vinegar to each tube. Put on a pair of gloves. Do not handle metal wool without gloves!

3. Place a drop of food coloring in each tube. You should have four different colors available: red, yellow, green, and blue.

4. Prepare four clumps of steel wool that all have a mass of 2 grams. They don’t have to be exactly two grams but they should be really close.

5. Spread the clumps of steel wool out to maximize the surface of steel wool that is available to react.

6. Place a piece of steel wool in each test tube and press them in using a glass stir rod (try and keep the steel wool spread out, don’t compact it all in one place in the solution). Try to add the four pieces to the four tubes quickly.

7. Start the stopwatch as soon as you have added all of the pieces of steel wool.8. Make observations about the process. Do you see bubbles forming in the test

tubes? Are there any color changes taking place?9. Every 5 minutes or so, check the test tubes to see what changes are taking place.10. It may take 15-20 minutes for the reactions to proceed. 11. Based on the time it takes for the color to change (if it does) as well as how

drastic the color change is, decide which color you would like to use as your indicator. Everyone does not have to use the same indicator.

Part B

1. The setup here is going to be very similar to the setup you used in part A. You want to set up three different 6” test tubes.

2. Place 5 mL of distilled white vinegar in each test tube. 3. Add 1 drop of the food coloring that you chose to use as your indicator to each

test tube.4. Label the three tubes with a sharpie (if they are not already labeled). You should

have one tube for iron (steel is iron + carbon), one for copper, and one for zinc.5. Weigh out 0.5 grams of each type of metal wool. You should have one clump of

steel, one clump of copper, and one clump of zinc.6. Once you have clumps that are 0.5 g, spread them out (fluff them) so that you can

maximize surface area.7. Add the clumps of metal the appropriately labeled tube.8. Start the stopwatch and begin your observations of the different metals.9. You are looking for bubble formation, how aggressive the bubble formation is,

and any color change in the solution. 10. Once you have completed the experiment, pour the liquid down the sink and put

the pieces of metal wool in the appropriate metal wool container.References and Additional Reading

1. http://pubs.acs.org/doi/abs/10.1021/ed083p1792A1. Picture credit: http://www.flickr.com/photos/mach86/galleries/72157622790404297


Experimental Data

Part A

Color Mass of Steel Wool Time to Color Change

Color Change

Red

Yellow

Green

Blue

Other observations:

Your Chosen Indicator:__________________________________

Why did you choose this indicator for your experiment?

Part B

Metal Mass of Wool Bubbles Color Change Reactivity RankIron

Copper

Zinc

Other Observations:



1. Write out the activity series for the metals that you used in your experiment based on your results.

2. Look up the activity series of metals in your textbook for lecture. How does the activity series that you determined in the laboratory compare to the one in the textbook?

3. Would this experiment have worked without the indicator? Why or why not?

4. What would happen if we used Pb or Ag in this experiment? Where would they fit in to your activity series? Why do you think we didn’t use them?

5. What would happen if we used a much stronger acid (like concentrated HCl) in this experiment? Why do you think we didn’t do this?


Experiment 7: Stoichiometry and the Determination of Molar Mass

IntroductionWhen you explore the structure of

molecules you discover that elements form compounds in definite proportions. This property is true of all chemical compounds.

In a chemical reaction, the amount of product that can be formed will be determined by the material present in the lowest quantity, which we refer to as the limiting reagent. Let’s look at the picture to the right of birds and tree limbs. These birds are large enough that you can only get one bird per limb. So the interaction between birds and limbs is a one to one ratio. If I wanted to know how many bird occupied limbs could be produced from a combination of limbs and birds, the answer would depend on which of these two things I have less of. For example, if I have a hundred birds but only 30 limbs, the maximum number of limbs that could have birds would be 30. In this case the limbs act as the “limiting reagent” because they limit the amount of final “product” (birds on limbs) that can be produced. It could certainly go the other way too. If I had 30 limbs but only 17 birds, it will be only possible to have 17 limbs occupied with birds. In this case, birds would be the limiting reagent.

When we talk about chemical reactions, the process in understanding the amounts in a chemical reaction is basically the same that we just used with birds and limbs. This process has a fancy name; stoichiometry. Stoichiometry is simply a tool for understanding the amounts of various reactants required to complete a reaction. It also allows us to determine how much product we should make. Sometimes reactions are a one to one ratio (this is nice!) but other times there are other ratios, maybe one to two or two to three. Here is an example of a chemical reaction:

2 H2 + O2 → 2 H2O

In this reaction we have reacted hydrogen and oxygen together to form water. Let’s talk about the two different kinds of numbers shown in this reaction. The coefficients (the 2 in front of H2 and H2O) tell us the ratio of each type of molecule (there are twice as many hydrogen molecules needed in the reaction as there are oxygen molecules). The subscripts tell us the relationship between atoms in a molecule. So O2 tells us that there are two oxygen atoms in that molecule. Be careful not to get these two confused.

Notice that if you count the total number of each type of atom on both sides of the equation it is equal (we are not losing or gaining any atoms). This is a very important part of stoichiometry, that you balance the number of each atom type on both sides of the equation.


I know we have already talked about quite a few new terms, but there is one other important concept to discuss and that is the use of moles. If you take a look at a periodic table (there is one in the front of your lecture textbook) you will notice the mass number for each element listed below the element symbol (e.g. Oxygen is “O” and has an atomic mass of 16.00). Hydrogen has an atomic mass of 1.01. Since these two different elements have very different masses, how do we study relationships between them? Mass is an easy way to measure the amount of something, but it is not easy for us to study relationships between two different elements with such different masses.

This is why chemists developed the mole. The whole purpose of this unit is for us to “standardize” the mass of different elements so that we can understand their relationships. To make things easy (well kind of…) one mole of each atom is equal to their atomic mass in grams (for example, 1 mole of Oxygen weighs 16.00 grams). Here is an example of how moles can be used to study relationships between atoms:

One mole of Hydrogen is 1.01 grams, but hydrogen exists in the diatomic (two atom) form of H2. As a result, H2 gas has a molar mass of 2.02 g/mol (grams/mole). Oxygen gas is also diatomic, and has a molar mass of 32.00 g/mol (2 x 16 = 32). Let’s look at the same equation for the formation of water:

2 H2 + O2 → 2 H2O

These coefficients that we talked about are actually referring to a mole ratio not a mass ratio. So for every one mole of oxygen gas present, we need to have two moles of hydrogen gas present in order to make two moles of water.

Now let’s work through a limiting reagent problem using this same reaction.

Example Problem: You have 5.0 grams of Hydrogen gas and 27.0 grams of oxygen gas. How many grams of water can you produce from this reaction?

To start with, don’t forget that you can’t just relate the masses together; you have to convert to moles first to put both reactants on the “same plane”. Begin by converting the masses of reactants into moles:

= 2.5 moles of H2

= 0.844 moles of O2

Now that we have converted these masses into moles, we need to figure out which reactant is the limiting reagent. In order to do this we must consider the ratio of the two reactants as we saw in the chemical reaction.

We need twice as much hydrogen as we do oxygen. So, to use up all of the oxygen we would need 2 x (0.844 moles) = 1.69 moles of hydrogen. Looking above we


see that we have more than enough hydrogen to use up all of the oxygen (along with some left over!). So oxygen is our limiting reagent. This is an important lesson to learn: just because you need twice as much of something in the chemical reaction does not mean it will automatically be the limiting reagent! You have to do the calculation to find out.

Now we want to determine how many moles of water we can produce with the given amount of oxygen and hydrogen that we have. We have determined that oxygen is the limiting reagent, so the number of moles of oxygen we have will determine how many moles of water we can produce.

We have 0.844 moles of O2 and we produce two moles of water for every one mole of oxygen we have. So we would produce 1.69 moles of water.

We are not quite home yet! Remember that the question asked us for a mass of water that we would produce; right now we are still talking in moles.

The molar mass of water can be calculated using information from the periodic table. We know that the atomic mass of H is 1.01 grams/mole, and oxygen is 16.0 grams/mole. The formula for water is H2O:

2(1.01 g/mol) + 1(16.0 g/mol) = 2.02 g/mol + 16.00 g/mol = 18.02 g/mol

This same approach works for calculating the molar mass of any compound. Add up all of the atomic mass of atoms. Don’t forget to keep track of the ratio of atoms in the molecule. Now that we have calculated the molar mass of water we have one calculation left. Now we multiply the moles of water produced in the reaction by the molar mass of water:(1.69 moles of H2O) x (18.02 g/mol of H2O) = 30.5 grams of water produced. That is the final answer for the problem we were given. This number that we just calculated represents something called the theoretical yield. The theoretical yield is the maximum amount of product that could be produced if everything goes perfectly. Yet we live in an imperfect world. So something else that we would want to know after the experiment is how close our experiment got to the theoretically ideal experiment. This leads us to the calculation of the percent yield, where we can determine how close we are to the theoretical yield. The equation is as follows:

% yield = x 100%

For example, running this experiment that we have just done the calculations for may yield 28.6 grams of water. The % yield would then be (28.6 g / 30.5 g) x 100 % = 93.8 % yield. Pretty good!

For this week’s experiment you will be using a precipitation reaction to study stoichiometry. A precipitation reaction is a reaction where a solid forms as a result of the reaction between two solutions. The reaction that you are going to be studying is as follows:


CaCl2(aq) + Na2CO3(aq) → CaCO3(s) + 2 NaCl(aq)

NaCl is just table salt and is very soluble in water, but CaCO3 is the chemical formula for the substance found in marble and chalk. This compound is not soluble in water.

SafetySodium carbonate is basic. You must wear eye protection during this experiment.

If you get sodium carbonate on your skin, wash it off with copious amounts of water and let your instructor know. Calcium carbonate should be placed in the trash after you have your final mass. Everything else can go down the sink. No eating or drinking in the laboratory.

Experimental Procedure1. Obtain the mass of a 100 mL beaker.2. Use the beaker to weigh out 2.0 grams of calcium chloride (CaCl2) (it does not

have to be exactly 2.0 but it should be close. Record your actual mass on the data sheet.

3. Add 50 mL of water to the beaker using a graduated cylinder (don’t use the volume graduation on the side of the beaker to do this!)

4. Now use a 50 mL beaker to weigh out 2.5 grams of sodium carbonate (Na2CO3). Record your actual mass in the data table.

5. Measure out 25 mL of distilled water using a graduated cylinder and then add the water to the sodium carbonate.

6. Stir both solutions with a glass stir rod (rinsing in between!) until all of the solids have dissolved.

7. Pour the sodium carbonate solution from the 100 mL beaker into the 100 mL beaker containing the calcium chloride solution. Use a couple of mLs of distilled water to rinse out the sodium carbonate beaker.

8. Stir the mixture and record your observations of what happens in the beaker. 9. Continue to stir the mixture until it appears that no additional precipitation is

taking place.10. Take a piece of filter paper and write “1” on it with a pencil.11. Measure the mass of the piece of filter paper.12. Set up your funneling apparatus. This involves a glass funnel, a piece of filter

paper (the one you just got the mass of), and a 125 mL Erlenmeyer flask. 13. Allow the solution to sit, wait to filter until after you have prepared your

apparatus. It helps to add a little water to the filter paper to help it stick in the funnel.

14. Carefully pour the solution with the precipitate through the filter paper in the funnel. Allow the solution to drain through the paper.

15. After all of the liquid has filtered through the paper, add about 5 mL of distilled water carefully to ensure that no sodium chloride remains in the filter paper.


16. Now rinse the filter paper with about 5-10 mL of ethanol to help the filter paper and the solid dry.

17. Pull the filter paper out of the funnel and spread it out on a paper towel.18. Now repeat this process in the exact same way for two more trials. Use the

same amounts of materials and follow the exact same process as closely as you can. You can reuse all of the same glassware (rinse everything in between trials). Make sure you keep the three trials separate! You will want to write a number on each piece of filter paper with a pencil before weighing it to keep track of which trial is which.

19. Once you have completed your three trials, check the filter paper for the level of dryness. If needed the filter paper can be gently heated using a hot plate with the help of the instructor. Check with your instructor for help with this part.

20. Obtain the mass for the three different trials and record all observations and data on the experimental data sheet.

21. Perform the calculations to determine the limiting reagent, theoretical yield, and percent yield for this reaction as well as the averages for all three trials.

References and Additional Reading1. Introductory Chemistry Version 3.1 Escience laboratory manual2. http://www.learnchem.net/tutorials/stoich.shtml3. http://www.shodor.org/unchem/basic/stoic/4. Picture Credit: http://photos.ibibo.com/photo/1529333/nature-beautiful-photo-

birds-tree


Experimental Data

Trial #

Mass of CaCl2 (g)

Mass of Na2CO3(g)

Filter Paper Initial Mass (g)

Filter Paper Final Mass (g)

Mass of CaCO3 (g)

1

2

3

Observations of Experiment:

Calculations:

1. Determine the molar mass of CaCl2: _______________________g/mol

2. Determine the molar mass of Na2CO3: _____________________g/mol

3. Determine the molar mass of CaCO3: ______________________g/mol

Trial #

Moles of CaCl2

Moles of Na2CO3

Theoretical Moles of CaCO3

Theoretical Yield of CaCO3 (g)

1

2

3

Trial # Percent Yield (%)1

2

3

Average



1. Was your hypothesis supported by the experimental data?

2. What was the limiting reagent in this reaction? Was there a lot of excess of the reagent that was not limiting?

3. What mass of the limiting reagent would be needed in order to use all of the other reagent? (in other words, what mass of the limiting reagent would be an equal number of moles to the number of moles of the excess reagent in this reaction)

4. In what parts of the experiment do you think you may have lost product? How might this be improved?

5. What principle of Green Chemistry would you be employing if you made the number of moles of both reactants equal? (Hint: it would not be so wasteful)

Experiment 8: Reaction Rates and Catalysis


IntroductionAs we have been learning over the course of

the semester, there are chemical reactions going on all around us and inside us. Many of these reactions have something called an “activation energy” which is how much energy input is required for the reaction to take place. There are a couple of ways to deal with this energy need. You can either introduce energy from outside of the system (you could start adding heat energy by increasing the temperature), or you can add a catalyst. Catalysts allow a reaction to take a different energy path, so in a sense they lower the activation energy of the reaction allowing the reaction to occur much more easily. There are many examples of catalysts in the natural world. Fireflies like the one in the picture to the right use an enzyme called luciferase to harvest energy from ATP (the main energy currency of your body) and produce light. This is just one of thousands of reactions involving enzymes in nature. You will be exploring several biological catalysts and their action as part of this week’s lab.

One area where catalysts are getting a lot of attention these days is in the area of alternative energy. Based on what we have talked about with catalysts so far, this makes sense right? If there is a way that we can lower the amount of energy we have to put in to producing fuel we could be more efficient in our fuel production process.

Here is an example. One of the most popular ideas for alternative energy under consideration today is the use of hydrogen gas as a fuel. The ideal source of hydrogen is from water, splitting water into hydrogen and oxygen. The equation for this reaction is as follows:

2 H2O(liquid) → 2 H2(gas) + O2(gas)

For this reaction to happen it takes : 282.1 kJ/mol of energy. Don’t let the units here worry you. Kilojoules (kJ) are units that we use to describe a quantity of energy.

If I burned hydrogen gas in the presence of oxygen the reaction would go in the exact reverse direction of the reaction shown above. And how much energy to I get out of burning it? 282.1 kJ/mol! (and that is under ideal conditions). Energy can neither be created or destroyed. So in this reaction you get out of it exactly what you put into it. And in fact, you will lose (it is not destroyed it is just converted to another form) some of the energy in the process and so you will not be able to harvest all of that energy from the reaction.

This is where catalysts come in. By using catalysts we can use less energy to separate water into hydrogen and oxygen because the catalyst allows us to use a different energy pathway to get to the same place. For additional helpful reading about how a catalyst lowers the activation energy of a reaction, see your textbook for CHEM 130 lecture (Introductory Chemistry by Tro 3rd addition, pages 562-563).


You will get the chance to make your own hydrogen gas using solar power in this laboratory. You will then fuel up and run a hydrogen powered remote controlled car.

SafetyCarefully follow the instructions for the use of the hydrogen powered car to

prevent the uncontrolled release of hydrogen gas into the room. Hydrogen gas is extremely flammable and can be explosive. Do not eat or drink in the laboratory. All waste materials in this laboratory can go down the sink when you are finished. Place light sticks in the appropriate container. The materials from today’s lab cannot be re-used.

Experimental Procedure:You are going to perform four different mini-experiments with a partner. It seems

like a lot, but it is actually going to be fun. You do not need to perform the experiments in any particular order. Just find available materials for one of the experiments at the back bench and get started.

Experiment A Hydrogen Powered CarAt this station you are going to be catalytically producing hydrogen gas from

water and then using it to run a remote controlled car. 1. Acquire a box containing a hydrogen powered remote controlled car and accessories.2. Slowly pour distilled water into the tank opening on the hydrogen gas generating station. 3. Allow the water to remain in the tank for at least 5 minutes to allow the electrolyzer time to absorb the water.4. In the next step you get to create hydrogen gas from water. Depending on the conditions, you will either be using solar energy (on a sunny day in a daytime lab) or DC power (from batteries) to create hydrogen gas. Check with your TA/instructor for which energy source is recommended. 5. If you are using the solar option, connect the solar panel to the cable matching the colors of the cables to the colors of the ports on the solar panel (red : red, black : black). Place the solar panel in direct sunlight.6. Switch the hydrogen generator station from “off” to either “DC” or the icon of a sun depending on which energy source you are using.7. The little blue lights should now be flashing and you are now producing hydrogen gas. You should be able to observe little bubbles of oxygen gas on the right side of the tank.8. While you are generating hydrogen you can be purging the balloon system in the car. Start by closing the plunger on the syringe to remove as much air from the syringe as possible.9. Attach the syringe to the side of the car by inserting the syringe’s connector valve gently into the car’s input valve and turning clockwise.10. Pull gently on the plunger of the syringe to remove air from the balloon in the car. Stop pulling on the syringe when the balloon no longer contains any air.11. Disconnect the syringe connector valve by turning it counterclockwise. It should just pop out due to the spring tension in the connector. Your tank


(balloon) is now purged and ready for you to fill with hydrogen from the refueling station. 12. Make sure the switch underneath the car is in the off position. 13. Attach the connector from the hydrogen generating station to the side of the car in the same fashion as you attached the purging syringe. 14. It will take about 10 minutes in strong sunlight to fill the tank of the car with hydrogen. If using DC current from batteries, this will take about a minute. 15. DO NOT allow the hydrogen generating station to run longer than 20 consecutive minutes. Once you have filled the tank on the car, turn the hydrogen generation station to the “off” position. Allow the generating station to rest for at least 10 minutes in between running if allowed to run for close to 20 minutes. If using DC current and the station is not run for as long, it does not need to rest as long.16. Keep the hydrogen generation station running until the balloon inside the car is full. Once the balloon is full, turn the hydrogen generation station to the “off” position. 17. Now that the fueling station is turned off, the first tank of hydrogen is going to be used to purge the balloon of any gases other than hydrogen. Take the car to the hood in the back of the classroom. Press the purging valve and release all of the hydrogen gas from the car.18. Now reattach the fueling station and fill the balloon again (once the hydrogen generating station has rested for an appropriate amount of time). Run the hydrogen station until the balloon is full. Turn the hydrogen station off.19. Move the switch on the bottom of the car to “warm up” and allow it to work through the balloon of hydrogen that you have just filled. Allow it to warm up until the balloon is empty.20. Now fill the balloon with the hydrogen generation station one last time. This third tank of hydrogen you can now use to run the car!21. Thought questions: How does a fuel cell work? Does the car run off of hydrogen only, or does it use oxygen too?

Experiment B Light Sticks and TemperatureIn this station you are going to explore the effect of temperature on the intensity

of light produced by a light stick. 1. Acquire a light stick2. Prepare two baths in 400 mL beakers. One of them should be a hot water bath

(just using hot tap water) and the other should be an ice-water bath.3. Place the glowstick in the hot water bath and make observations about how

intense the light is. Compare the intensity of light to the light produced when the stick was at room temperature.

4. Pull the glowstick out of the hot water and let it sit at room temperature for 5 minutes. What happens to the intensity of the light as it is cooling down?

5. Now place the glowstick into the ice-water bath. Record your observations about the intensity of light in cold water.

Experiment C Hydrogen Peroxide and yeast


In this experiment you will be exploring the catalytic breakdown activity of the enzyme catalase. Catalase is found in normal baker’s yeast. You will be exploring the effects of this enzyme on hydrogen peroxide and water. Hydrogen peroxide has an interesting structure, compare the structure to water:

H-O-O-H (Hydrogen peroxide) H-O-H (water)

This experiment will allow you to determine if the catalase enzyme breaks down H-O bonds or O-O bonds. Water contains only H-O bonds. Hydrogen peroxide contains both H-O bonds and O-O bonds.

1. Acquire two 4” test tubes and label them sample and control.2. Add 5 mL of water to the control tube.3. Measure the mass of the control tube + water4. Weigh out two samples of 0.5 grams of baker’s yeast and pour one of the piles of

granulated yeast pellets into the control tube.5. Add 5 mL of hydrogen peroxide to the sample tube (without yeast!)6. Measure the mass of the sample tube + hydrogen peroxide.7. Add 0.5 grams of baker’s yeast to the sample tube (containing peroxide).8. Record your observations of the sample and control tubes.9. Once things are no longer changing, measure the mass of the sample and control

tubes again.10. Thought question: What gas is being formed when bubbles are produced? How

might you test your theory?

Experiment D PhET Simulation of Reactions vs. Temperature.For this experiment you will explore a computer simulation of how molecules

react to temperature changes.1. Click on the Reactions vs. Temperature experiment simulation icon on the

computer. This will open up the simulation.2. Click on the “Many collisions” tab.

The reaction that you are looking at is shown in the bottom left hand corner. (A +BC AB + C).

3. You can add the molecule of your choice by using the little red pump handle. You can select the molecule(s) that you want to add below the pump. Simply click and drag the pump handle with the mouse and molecules will enter the reaction window.

4. Add some of both reactants and then observe how the molecules interact.5. Pay attention to the plot to the right of the reaction window. This plot is an

energy diagram that gives you a visual representation of the energies of the reactants and the products. Notice that it also includes the activation energy, which as we talked about in the introduction is the energy that must be added to the system for the reaction to take place.

6. Experiment with the temperature. The temperature control is directly below the reaction window. Try lowering the temperature and notice what happens with the


molecules. Also notice what happens with the energy of the system in the energy diagram on the right

7. Now try heating the reaction up. Notice what is happening with the molecules inside. Pay attention to the energy diagram as well. Are you getting closer to the activation energy/

8. Now increase the temperature enough to bring the energy in the system above the activation energy. What is happening to the molecules now?

9. You can pause the simulation at any time by clicking the pause button down at the bottom of the simulation window.

10. Play around with the simulation for a few minutes and study the effects of temperature, adding different molecules (and different amounts of them).

References and Additional Reading1. http://onlinelibrary.wiley.com/doi/10.1111/j.1471-5740.2005.00106.x/pdf2. http://www.usc.edu/CSSF/History/2003/Projects/J0411.pdf3. http://www.stevespanglerscience.com/experiment/light-sticks-the-science-of-

liquid-light4. http://auto.howstuffworks.com/fuel-efficiency/alternative-fuels/fuel-cell.htm5. Picture Credits:

http://animal.discovery.com/tv/a-list/creature-countdowns/cheats/cheats-05.html


Experimental Data

Experiment AObservations of the formation of hydrogen and oxygen gas and the running of a small car with that fuel.

Experiment BObservations of the effect of temperature on the light intensity from a light stick.

Experiment CSample Mass of Yeast

(g)Mass of Liquid (g)

Starting Mass (g)

Final Mass (g)

Control (H2O)

Sample (H2O2)

Observations of the effect of the enzyme catalase in yeast on hydrogen peroxide vs. control.

Experiment D

Observations of the effect of temperature on the rate of a reaction.



1. Was your hypothesis supported by your experimental data? Explain.

2. What is the relationship between experiment 2 (light sticks and temperature) and experiment 6 (PhET simulation)?

3. Which bonds does the enzyme catalase (in yeast) break in hydrogen peroxide? Defend your ideas with the experimental data you collected.


Experiment 9: Liquids, Solids, and Intermolecular Forces

IntroductionWater is one of the most amazing

chemicals in the universe! We have already learned that the density of liquid water is greater than the density of solid water. This results in ice floating on top of water and water freezing from the surface down allowing fish to survive the winter in deep water. But there are a lot of other cool properties of water too. Have you ever seen a water strider? These little insects like the one shown in the picture to the right are able to suspend themselves on top of water. How is this possible?

Water (and some other liquids too) has an interesting property called surface tension. Water demonstrates surface tension because of hydrogen bonding. Hydrogen bonding is a different kind of bonding than those that we have talked about so far. It is not like ionic or covalent bonding. There are three different elements which are willing to accept a hydrogen bond, and these are nitrogen, oxygen, and fluorine. In order to have hydrogen bonding you also have to have a hydrogen bond donator which will also be one of these three atoms. So water participates in hydrogen bonding with itself because it has an oxygen which can act as a hydrogen bond acceptor, and it has two hydrogens that it can use to donate hydrogen bonds. Take a look at the figure on the right. The solid lines represent covalent (polar covalent if you remember) bonds and the dashed lines represent hydrogen bonds. So the water molecule in the middle is being shown hydrogen bonding with four other water molecules.

The hydrogen bonding in water results in surface tension. This surface tension is what is responsible for the meniscus that we talk about when measuring out volumes of water in a graduated cylinder. This same surface tension is what is responsible for water striders being able to suspend themselves on the surface of water.

Another thing that we are going to be looking at this week is miscibility. It’s a big word, and just means how well two liquids mix together. Solubility describes how well a solid dissolves in a liquid, and miscibility describes how well two liquids mix together.

One of the main things that you are going to learn about this week is intermolecular forces. Just like an interstate goes between different states, these types of bonds and interactions are between different molecules. Hydrogen bonding, which we have just learned about, is an intermolecular force. But there are a couple of other intermolecular forces as well:Dipole-Dipole interactions: These interactions occur in polar molecules where the dipoles (separation of charge) in different molecules are attracted to each other. This type of interaction does not occur in nonpolar liquids.


Dispersion forces: This is a very weak force that occurs when electrons are attracted to nuclei. This type of force occurs with all molecules and atoms. Since it is found in all molecules, even if a molecule does not have hydrogen bonding or dipole-dipole interactions it will still have dispersion forces.

Intermolecular forces have a big effect on the miscibility of liquids as you will explore in this weeks lab. You will also be looking at how intermolecular forces are involved in an oil spill and the cleanup process. Oil is a nonpolar material that does not have any hydrogen bonding.

SafetyOil should not go down the sink. A mixture of oil, soap, and water can go down

the sink. Do not eat or drink in the laboratory. Sand, soil and feathers should be placed in their respective containers.

Experimental ProcedurePart A

For this part of the experiment you will be exploring how food coloring, oil, water, and soap interact.

1. Take a 6” test tube and add 5 mL of oil to the test tube.2. Add 1 drop of food coloring (use red, blue or green; your choice)3. Record your observations of what happens when the oil and food coloring mix (or

not). Include things like: does the food coloring mix with the oil? What do the droplets of food coloring look like in the oil?

4. Now take an eye dropper and begin to add drops of water to the top of the oil in the tube.

5. Record your observations of how the water and the oil interact. (Do they mix? What do the droplets of water look like in the oil? What does the area where the intersect look like? Etc.) Also identify how the food coloring and the water interact (you may have to do some gentle mixing by tapping the side of the test tube).

6. Continue to add water until you have added 5 mL of water.7. Mix gently until you have only two layers.8. Now add a drop of soap to the top of the oil.9. Carefully watch what the drops of soap look like (do they have the same shape in

the oil as the water did?)10. Once you have made your observations of the soap droplets, go ahead and mix the

tube vigorously.11. Record your observations of what the mixing process looks like now.12. Thought questions: What was it that caused the different droplets to adopt the

shape that they did? Why was the soap different than the water and the food coloring? What did the soap do when it was mixed with the other layers? Is a soap molecule polar, nonpolar or both?

13. The mixture (as long as it is well mixed and looks the same throughout the tube) can go down the sink with water.


Part BIn this part of the experiment you will be exploring the effect of soap on the

surface tension of water.1. Take a petri dish and fill it about halfway with water. 2. Sprinkle some pepper on the surface of the water.3. Record your observations of how the pepper and the water interact.4. Now add a drop of soap to the petri dish.5. Record your observations 6. Thought questions: What is it that causes the pepper to interact with water the

way it does? Why did this change when I added the soap? How does soap affect hydrogen bonding?

7. Water, pepper and soap can go down the sink.Part C

In this part of the experiment you will be simulating an oil spill and learning about the cleanup process. **Important note: Cooking oil is much more mild and easy to work with than crude oil. Keep this in mind when you are learning about the cleaning process. Working with material contaminated with crude oil is actually much worse than what you are doing! And it is not clean unless it is really clean!!! Make sure you understand this before you come to conclusions about this experiment.

1. Set up three 50 mL beakers2. Place about 10 mL of soil in one beaker (very rarely do we measure solids by

volume! But it is OK in this qualitative lab) 10 mL of sand in another beaker, and place a feather in the third beaker.

3. Add 5 mL of cooking oil to the beakers with sand and oil. 4. Use a glass stir rod to mix the oil and sand as well as the oil and soil.5. Drizzle a little cooking oil over the feather in the third beaker.6. Record your observations about what happened with the three different items

when they encountered the oil.7. Now add about 10 mL of water to all three beakers.8. Record your observations. Does the water help or does it make more of a mess?9. Now add a drop of soap to each beaker. 10. Mix the soap in well with the soil and sand using the stir rod. Also try and scrub

the feather with the soap.11. Now take a piece of fabric and pour the sand mixture into the fabric over the sink.

Try and wash the sand in the fabric at the sink and see how clean you can get it.12. Repeat this process with the soil and see how clean you can get it.13. Wash the feather in the sink and see if you are able to remove all of the oil.14. Now place all three items on a dish on your bench top. Spread them out so they

will dry.15. Record your observations of the washing process in trying to get these three

materials free of oil.16. Once the materials have dried, make observations about how well you were able

to remove the oil from these things with dish soap.17. Please place feathers, sand, and soil in their appropriate “waste” containers so that

they can be cleaned and reused in another class.


References and Additional Reading1. http://www.chem.ufl.edu/~itl/2045/lectures/lec_g.html2. http://cost.georgiasouthern.edu/chemistry/general/molecule/forces.htm3. Picture Credits:

http://www.treknature.com/gallery/Asia/Thailand/photo154910.htm, http://vinstan.wikispaces.com/Chapter+2+++++Chemical+Bonding


Experimental DataPart A

1. Observations of interactions between oil and food coloring:

2. Observations of interactions between oil and water:

3. Observations of interactions between food coloring and water:

4. Observations of interactions between soap and oil:

5. Observations of mixing once soap was added:

Part B1. Observations of interactions between pepper and water:

2. Observations of interactions between pepper and water once soap had been added:

Part C1. Observations between oil and the three substances (sand, soil and feather):

2. Observations once water had been added to the mixtures:

3. Observations once soap had been added to the mixtures:

4. Observations of how clean the three substances were after processing with soap:




2. What do you think a soap molecule looks like? Is it polar, nonpolar, or both? (you can draw a picture or write out an answer)

3. How hard was it to clean up sand, soil, and a feather that had oil on them? How much harder do you think it would be if it were crude oil contaminating these things instead of vegetable oil?

4. How well do you think a water strider would be able to suspend itself in soapy water? Explain your answer and use data from your experiment.

5. What affect does soap have on hydrogen bonding? Use data from your experiment to answer this question.


Experiment 10: The Ideal Gas Law

IntroductionHave you ever been in an accident and had

an airbag help keep you safe? Maybe you have only seen them on television, but it really is an amazing invention. Nowadays you can’t buy a car that does not come with one (or several!). So how do these things work? How do they produce just the right amount of gas? If you think about it, forming just the right amount of gas is extremely important. If they form too much, the bag will burst and only make the accident worse. If the airbag forms too little gas it may not stop the occupant from being injured.

Inside the airbag a chemical reaction is taking place:

2NaN3(s) → 2Na(s) + 3N2(gas)

The compound shown here is called sodium azide. It is actually a powerful explosive and is very toxic. But, when it explodes the gas it produces is nitrogen which makes up 78% of our air anyways! The sodium that results is eventually converted to glass from subsequent reactions that take place (we won’t worry about those just now). The cool thing is that we can predict how much gas is going to be produced by the reaction using stoichiometry (which we worked on in a previous lab) and the ideal gas law. The ideal gas law is as follows:

PV=nRT

Where P is the pressure, V is the volume of the “container”, n is the number of moles of gas, R is the gas constant (always remains the same) and T is the temperature. So as long we know four of these five variables, we can determine the 5th one.

In this week’s lab you are going to determine the actual concentration of hydrogen peroxide in a solution by reacting hydrogen peroxide with the catalase enzyme in yeast. This is what the reaction looks like:

2 H2O2 + Catalase → 2 H2O + O2 + Catalase

You will measure the volume of gas produced and you will be given the pressure and temperature information. This will allow you to calculate the amount of hydrogen peroxide in the solution you used. We want to set up the reaction so that hydrogen peroxide is the limiting reagent. We do this for two reasons: 1st, it is hard to calculate the actual concentration of catalase enzyme when we add yeast, 2nd we want to determine the concentration of hydrogen peroxide in the solution. We are able to determine how much product (O2 gas) we made. The amount of product made is dependent on the hydrogen


peroxide concentration (because it is the liming reagent). This way we can use the amount of oxygen produced to determine the concentration of hydrogen peroxide we started with.

SafetyHydrogen peroxide is a strong oxidizer. Avoid contact with your clothing as it

may bleach some of the color out. Wear safety glasses when performing the experiment. Do not eat or drink in the lab. Once the experiment is complete, the reaction mixture can be poured down the sink.

Experimental Procedure

In this experiment you will be producing oxygen gas with the reaction between hydrogen peroxide and yeast.

1. Prepare the reaction setup. This uses a 125 mL Erlenmeyer flask, a rubber stopper, clear plastic tubing, a buret, a 250 mL beaker, and an eye dropper bulb. See Figure 2 below.

Figure 2. Gas Collection Apparatus


2. Fill the 250 mL beaker about 1/3 of the way full of water. With the eye dropper bulb in place on the end of the buret, fill the buret up to the top with water.

3. Now insert the plastic tubing into the end of the buret underwater. It is ok if some bubbles of air enter the buret.

4. Place your finger on the open end of the buret and the tubing. Quickly invert the buret and place it in the beaker.

5. Release your finger from the bottom once the end of the buret is below the water level in the beaker.

6. Now use an eye-dropper to add bubbles of air to the buret to get the water level down to near the 25 mL mark.

7. If you introduced more than about 10 mL of air, go ahead and refill the buret with water and try again. A volume anywhere between 20 and 25 mL for the starting volume is fine. Once you have an acceptable volume of air in the buret, attach it to the ring stand with a clamp.

8. Take the temperature of the water in the beaker with a thermometer. The barometric pressure on the day you perform the lab should be written on the board.

9. Use a graduate cylinder to measure out 1.0 mL of hydrogen peroxide. Add this hydrogen peroxide to the 125 Erlenmeyer flask.

10. Rinse the graduated cylinder out several times to remove all of the hydrogen peroxide. This rinse water can go down the sink.

11. Use a balance to weigh out 0.25 grams of yeast pellets. But don’t add it to the Erlenmeyer flask yet.

12. Place the rubber stopper tightly in the flask. Read the volume of water that is in the buret and write it down on your data sheet. Make sure that the stopper fits very tightly in the flask as you don’t want any of the gas to leak out.

13. With the stopper removed from the Erlenmeyer flask but ready to be quickly re-inserted, pour the yeast pellets into the flask and quickly add the stopper.

14. Make sure the stopper is in the flask tightly so that none of the oxygen gas leaks out. You should see bubbles forming where the pellets are dissolving in the hydrogen peroxide.

15. Swirl the Erlenmeyer flask so that the solution mixes well.16. Watch the level of the liquid in the buret, it should be going down (you should

have bubbles coming out of the tube inside the buret)17. Continue to swirl the flask until it is no longer forming bubbles and the level of

liquid in the buret is no longer going down. This should take anywhere from 5 to 15 minutes.

18. Once bubbles have stopped forming, carefully read the new volume in the buret and write it down

19. Clean up your glassware well before placing them back in the center of the bench.Calculations


The main equation that we will be working with in this lab is the ideal gas law. This equation can be rearranged to solve for n, or the number of moles of gas:

The pressure that you need to plug into this equation will be written on the board. It should be converted into atmospheres (atm) for the calculation. The volume will be the volume of gas collected in your experiment. The temperature will be the temperature of the water (because of the glass sides on the buret and the direct interface between the gas and the water in the beaker the temperature of the gas should be almost exactly the same as the water). This temperature needs to be converted into units of K. And finally, R is the gas constant, and its value is 0.08206 L*atm/mol*K. Using this equation will allow you to calculate the moles of oxygen gas produced in the reaction.

To figure out the actual percent hydrogen peroxide in the solution, you need to calculate how many moles of oxygen would be produced if the solution you added was 100 % hydrogen peroxide.

This is done by first converting the volume of hydrogen peroxide you used into the mass of hydrogen peroxide you would have if it were 100%. To do this, multiply the volume of hydrogen peroxide by the density:

Volume of H2O2 (mL) x Density of H2O2 (1.02 g/mL) = mass of H2O2 (@100%)

Now this needs to be converted to moles. The molar mass of H2O2 is 34.0 g/mole. We can get the moles of hydrogen peroxide if the solution were 100 % by multiplying the mass from the above equation by the inverse of the molar mass of hydrogen peroxide (so that grams cancel out):

Mass of H2O2 (grams) x = moles of H2O2 (@ 100 %)

If you remember from the reaction equation, it takes two moles of hydrogen peroxide to make one mole of oxygen:

Moles of H2O2 x = moles of O2 (@ 100%)

Now all that is left is to compare the moles of O2 your reaction produced the moles of O2 that the same volume of 100 % hydrogen peroxide would produce:

% hydrogen peroxide = x 100%


References and Additional Reading1. http://faculty.plattsburgh.edu/tom.moffett/che111/gaslab111.pdf2. http://www.cerlabs.com/experiments/10534977685.pdf3. Picture Credit: http://www.carbuyersnotebook.com/missing-airbag-

responsible-for-1-in-5-fatal-accidents/


Experimental Data

Barometric Pressure (mm Hg)

Temperature (oC)

Volume of H2O2 added (mL)

Starting buret volume (mL) [A]

Final buret volume (mL) [B]

Volume of gas collected (mL) [A-B]

Calculations

Conversion of Pressure:

1. ________ mm Hg x = ___________________ atm

Conversion of Temperature:

2. ________ oC + 273.15 K = ________________________ K

Conversion of Volume:

3. ________ mL x = ____________________ L

Ideal Gas Law:

4. n = = ______________ moles O2(gas)

Theoretical moles of O2:

5. ________ mL H2O2 x 1.02 x x = ________ moles

of O2 (@ 100%)

Percent hydrogen peroxide:


6. x 100% = ___________ %

7. Congratulations!




2. Based on your observations of the rate of this reaction, how well do you think hydrogen peroxide and yeast would work as a green substitute in an airbag? Explain.

3. You used standard household hydrogen peroxide for this experiment. The label on the bottle advertises that the bottle contains 3% hydrogen peroxide. How well did your experimental data agree with this?

4. Explain possible reasons for the difference between your experimentally derived percentage and the one on the bottle. Can you think of areas in your experiment where error might have been introduced? (Hint: There are always some errors in your experimental data no matter how good it is!)


Appendix A.

CHEM 130 Laboratory Technique Guide

I. Measurement of Mass

There are many different types of balances or scales available to measure the mass of an object. The selection of the balance depends on the mass of the object or sample and the precision needed for the measurement. In this course the measurement of mass will be done using top loading electronic balances. Our balances accurately measure mass to ±0.001 g or ±1 mg, so always record masses determined on these balances to this precision (three digits to the right of the decimal place in grams) even if the last digit is zero. All mass measurements will be done in grams. Please note that balances will change in the last digit ±0.002 g; this is to be expected.

Figure 1. Top loading electronic balance

There are many different types of electronic top balances but they all use the same two simple procedures. To simply weigh an object tare the balance to zero and then place the object on the balance to measure its mass. Weighing-by-difference is used to measure the mass of a sample being transferred from one container to another and will be used in experiment 6. A few rules need to be followed when using a balance.

Balance Rules and Instructions


Figure 1 illustrates one type of top loading electronic balance, which is one type of balance used in CHM 130L. Refer to this figure when following the steps and precautions for using the balance listed below:

1. Never pour or transfer chemicals over the balance. Spilled chemicals can damage the balances, which are very expensive to repair or replace. Never weigh warm or hot objects; if you can feel any heat, the weighing will not be accurate. Always use a container such as a vial, beaker, flask, or watch glass to weigh a solid or liquid chemical on the balance to protect the balance pan.

2. Make sure your hands are clean and dry before you handle containers or objects that are to be weighed. The outside of these containers or objects must also be clean and dry. Clean up any spills on the balance pan or lab bench around the balance immediately with a clean, damp sponge.

3. First open or remove the draft lid or cover (if there is one) and check to make sure that the balance pan is clean. If the pan is dirty, have your TA show you how to clean it and gently place it back on the balance.

4. Close or put the cover back on the balance and zero the balance by pressing the "T" or "on/tare" button. Wait 5-10 seconds for the weight display to stabilize. (If the object to be weighed is so large that the draft lid can't be used, do this step without the draft lid in place.)

5. Open or remove the draft lid and place the object to be weighed on the balance pan. Then close or place the draft lid back on the balance. (As long as it does not touch the object to be weighed, leave the lid off if it does touch the object.) After 5-10 seconds the weight display will stabilize and then record mass to ±0.001 g.

6. Never unplug the balance but be sure to turn it off at the end of the day.

Weighing Solids and Liquids

Since using the top loading electronic balance is so much easier than using the old triple beam balances or 1 mg mechanical analytical balances, very few errors are made measuring the mass of an object. Most errors are made when trying to measure the mass of solid or liquid transferred from one container to another (weighing by difference). The following are some helpful hints to keep in mind when weighing by difference.

1. Be very careful to avoid spilling material outside the target container.

2. If you are weighing the container that the material is being transferred from, do not use a spatula to transfer the material, but gently tap the container to slowly transfer the material into a new container.


3. Make sure the outside of the container is clean and dry before you weigh it for the first time and then touch it as little as possible until after the final weighing.

4. Set containers to be weighed on clean surfaces only.

5. Always cool containers or samples to room temperature before you weigh them.

6. It is sometimes helpful to preweigh the sample before it is transferred.

Some of these hints will be more important in future experiments.

Balance Calibration Check

All of the balances are regularly checked for correct calibration by the stockroom. To insure that you are using the balance correctly and that it is properly calibrated, the mass of a calibration weight will be measured. Calibration weights are provided in small wooden boxes. Make sure your hands are clean and dry before you touch the weights. (Normally calibration weights are not touched with your hands.)

Select a calibration weight and record the "Known Mass" for the weight (stamped on the weight) on the report sheet found at the end of this experiment. Measure the mass on a balance and record it on the report sheet. If the measured and known mass differ by more than 0.01 g reread the instructions for using the balance and measure the mass of the calibration weight again. If you get the same results again, see your instructor or TA.

II. Measurement of Volume

The best piece of equipment for measuring standard volumes in the laboratory is the graduated cylinder. Beakers and Erlenmeyer flasks are not accurate measures of the volume of liquids. To use a graduated cylinder, pour the liquid inside and then use the lines to identify the volume of the liquid. When using liquids like water, the surface tension will produce a meniscus (in other words the water on the top will be curved and not flat). The reading for the volume should be taken at the bottom of the meniscus (see figure at right). The correct reading for the volume in the figure shown is 36.5 mL.


Documents

CHEM 130 Lab Manual Fall 2011