Tie dye

by M. J. Simpson

Note: Plan on dressing in old clothes for this lab. These dyes are powerful and may permanently stain your clothes if you tiedyesplash or drip them on yourself.

Tie dye is a classic color chemistry lab. There are many different “recipes” for tie dye, but we will be following the procedure provided by Colorado Wholesale Dye Corp to ensure the most vibrant colors possible.

The wash and dyes will be prepared for you, so these are the instructions you need to know:

Before coming to lab: Prewash your 100% cotton, white garment before coming to lab.

In lab:

  1. Use a permanent marker to write your name on the tag of your garment.
  2. Soak your garment in the Sodium Carbonate wash for at least 10 minutes. This solution is very basic, so use gloves when in contact with the wash. Wring out the garment to dry it as much as you can.
  3. Fold and tie with waxed sinew your garment however you like.  Here are some designs you might like to try. This video shows you some techniques:


  4. Working on a surface covered in newspaper, apply the dye. Be sure to saturate the fabric with dye, but don’t put so much on that the dye pools underneath the fabric. Again, wear gloves when working with dye or it will dye your skin!
  5. Wrap your garment in newspaper and put it in a sealed plastic bag. Write your name and section on the bag, then give it to the instructor.

Molecular Modeling – Digital and Analog

adapted by M. J. Simpson from a lab originally by S.F. Sontum, S. Walstrum, and K. Jewett


Theories of chemical bonding allow us to understand the electronic structure and geometrical arrangement of atoms in a molecule or ion. Models provide a useful way of visualizing the arrangement of electrons in a molecule. As you learned in class, there are several types of structure representations used by the chemist at different times to explain chemical phenomena. These include Lewis structures (including formal charges) based on a simple counting rule; valence bond models (including hybridization, and resonance) based on orbital overlap; valence shell electron pair repulsion theory (VESPR theory) based on electron repulsion, to predict overall shapes of molecules/ions; and molecular orbits to predict certain electrical and magnetic properties.Model

The figure above is a ball and stick model for an early precursor to one of the anti-HIV drugs known as Protease inhibitors developed at Abbott Laboratories, Inc. We need models like this because actual molecules are too small to see. Modeling of molecular structures allows us to make predictions about the behavior of these invisible molecules so that we can design new chemicals with beneficial properties such as these inhibitors. The ball and stick model kits you use in this chemistry laboratory are very useful for visualizing chemical structures but they do not give quantitative information needed to design and build new molecules with specific properties.  ChemDraw is a computer “modeling kit” which we use to augment our visualization of the molecule with quantitative information.


To illustrate how molecular geometry can be obtained from Lewis structures and valence shell electron pair repulsion, we will use molecular models. With models, it is relatively easy to see both geometry and polarity, as well as to deduce Lewis structures. You may want to initially generate your Lewis structures before you come to the laboratory. When you come to the laboratory use the molecular models to check and refine your Lewis structures. In this exercise you will assemble models for a number of common chemicals and interpret them in the ways we have discussed.

The models consist of plastic bonding centers and bonding tubes. The bonding centers represent the hybrid valence electron distribution about the atom. The tubes represent the bonding electron pairs or valence bonds. As you build the models also draw the three dimensional structures on paper, so that you can develop skills at representing these three dimensional structures on paper.

Most of the structures will obey the octet rule. Often the central atom will have four electron pairs and will have a tetrahedral bonding center. In your model kits, multiple bonds are modeled by the long, flexible bonds.  We will see from our molecular orbital studies that the bent “banana” p bonds of our model kit are not good representations of the electron distribution but never the less correctly represent the Lewis structure.

Methane: Construct a model of methane, CH4.methane

Place the model on the desk top and note the symmetry of this tetrahedral molecule. The molecule looks the same regardless of which three hydrogens are resting on the desk. All four hydrogens are said to be equivalent. Is this molecule polar? Draw a three dimensional picture of this molecule.

Dichloromethane: Replace two of the hydrogen atoms in methane with a chlorine to make dichloromethane CH2Cl2. Is dichloromethane polar? Which Lewis structure best represents dichloromethane or are they the same?


Ethyl Alcohol and Dimethyl Ether: Construct ethyl alcohol, CH3CH2OH, and dimethyl ether, CH3OCH3.  The following is a chart of the physical properties of these two isomers.


Boiling Point Melting Point Solubility in water
Isomer 1 78.5 oC -117 oC Infinite
Isomer 2 -24.0 oC -139 oC Slight


Which isomer is more soluble in water? Which would have a higher boiling point? Why?

Constructing Lewis Structures

Now that you have some experience building molecular models from a Lewis structure you should be ready to fill in a geometrical structures chart starting only from the molecular formula.  In assembling a molecular model of the kind we are considering, it is possible, indeed desirable, to proceed in a systematic manner. We will illustrate the recommended procedure by developing a model for the first molecule on the chart CH2O. Please work along with your model as we describe the procedure.


  1. Choose a skeleton with the least electronegative atom for the central atom.
    For CH2O, this is carbon. Remember, ring structures may be possible with six member rings being the most stable and 3 member rings the least stable.
  2. Determine the total number of valence electrons.
    For CH2O this total would be 12 valence electrons. Now collect 6 bonding links representing the 6 pairs of bonding electrons.
  3. Assemble a single bonded skeleton structure.
    In CH2O the model can be assembled by using two links to form s bonds between C and H, and then using a third link to form a s bond between C and O.
  4. Add lone pairs.
    Use the remaining tubes to fill all the unfilled sites on the bonding center and to satisfy the octet rule. In the model, three links may be used to fill the O site leaving one unfilled sites on C.
  5. Form multiple bonds as needed to satisfy the octet rule.
    Move one of the lone pair electrons off of the oxygen to generate a double bond with carbon and satisfy the octet rule on carbon.
  6. Interpretation of the Lewis Structure.
    The tubes and spatial arrangement of the bonding centers will closely correspond to the electronic and atomic arrangement in the molecule. Be sure to check that you have a valid Lewis structure (octet rule) and the right total number of electrons. The Lewis structure of the molecule is given below:nnn

Formal Charge: Since all of the octet atoms satisfy the 8n rule there is no formal charge.

Hybridization: The angle between VSEPR groups is 120° so the hybridization is sp2 on both the carbon and the oxygen.

Resonance structures: The formation of other resonance structures would entail the moving of p bonds or the formation of new p bonds from lone pairs. For this molecule, there are no other ways to rearrange the electrons over the molecular skeleton and still satisfy the octet rule. This is the only possible full octet resonance structure for this molecular skeleton.

Atom Geometry: Given our model, we would describe the CH2O molecule as being planar with single bonds between carbon and hydrogen atoms and a double bond between the C and O atoms.

Polarity: Since all bonds are polar and the molecular symmetry does not cancel the polarity in CH2O, the molecule is polar.  (The compound having molecules with the formulas CH2O is well known and is called formaldehyde. The bonding and structure of CH2O are correctly given by this model.)

Isomerism: A possible alternate skeleton and Lewis structure would be H–C=O–H but this structure results in a formally positive oxygen next to a negative carbon. Because of the unfavorable charge separation this molecule does not exist.

Introduction to ChemDraw

ChemDraw is another “modeling kit” which we use to augment our information about molecules, visualize the molecule, and give us quantitative information about the structure of molecules.

The following is a brief introduction on how to use ChemDraw.


  1. Open Chem3DPro
  2. Click on the white panel on the right labeled ChemBioDraw – LiveLink
  3. Make sure the Text tool is selected when the Tools menu pops up
  4. Type the formula. For example, SO32- is entered as SO3-2
  5. If you want a different isomer, you can enter the formula more specifically. For example, instead of C3H4, you might enter H2CCCH2
  6. Hit enter. The molecule will appear. If it is outlined in red, that means something is wrong with your molecule. Consult the lab instructor or TA if you cannot figure out the problem
  7. Hit control-M to minimize energy. The total energy will be displayed at the bottom under Output
  8. To measure bond lengths, hover the mouse over the bond
  9. To explore the structure in different modes, go to View → Model Display → Display Mode, and then pick another mode. Try a few different ones
  10. To show or hide lone pairs, go to Structure → Lone Pairs, and make your selection
  11. To save a structure as an image, go to File → Save As, and then change the file format to whatever you like (such as .png or .jpg)
  12. Hit control-N to start over
  13. Another way to build a molecule is to build a backbone with the solid bond building tool and then change the atom types manually. After building the backbone, just click on the intersection you want to change and type the symbol of the atom you want to change. Double click the bond to make a double bond. This is especially useful for building the P4 isomers


Report:  Fill out the worksheet for the report.

CO2: Enthalpy of Sublimation, Reaction and Metabolism

by S. F. Sontum and K. Jewett (edited by R. Sandwick)


The oxides of carbon could not be more chemically different. Carbon monoxide (CO) is a polar basic gas that binds strongly to metals while carbon dioxide (CO2) is a non-polar gas that reacts with water. Carbon monoxide is produced when hydrocarbons are burned in a limited amount of oxygen while carbon dioxide is produced when hydrocarbons are burned in excess oxygen. Carbon monoxide is a deadly poison while carbon dioxide is essential to the metabolism of plants and animals.

Carbon dioxide is the fourth most common gas in our atmosphere and its concentration is increasing. Atmospheric carbon dioxide plays an important role in maintaining surface temperatures. It is one of the “greenhouse” gases responsible for global warming. The biological processes of metabolism, which produces carbon dioxide, and photosynthesis, which consumes carbon dioxide have functioned to maintain the levels of carbon dioxide in our atmosphere. The industrial revolution has shifted this balance. The concentration of carbon dioxide has increased more than 30% to the present day value of 400 ppm. Global temperatures are rising.

This experiment focuses on a physical property of CO2 – its energy content, and a chemical property of CO2 –  its ability to dissolve in water to make the solution acidic. Like all molecules CO2 stores potential energy in the molecular bonds that hold the oxygen to the carbon and in the intermolecular forces between molecules. We will measure the heat of sublimation of dry ice to investigate the intermolecular forces between CO2 molecules.

The phase diagrams below indicate the stable forms of compounds at different pressures P and temperatures T. The lines describe when two phases are in equilibrium and the triple point is the place where all three phases can coexist in equilibrium. For example, the triple point of water occurs at 0.01 oC and 0.006 atm. In order to convert ice directly into steam the partial pressures of water would have to be lower than 0.006 atm. Because we live at higher pressures (i.e., at 1 atm), we always see ice melt before it vaporizes to steam. The process of converting a solid to its liquid is called melting (aka, fusion). The triple point of carbon dioxide on the other hand occurs at a much higher pressure and lower temperature, -57 oC and 5.11 atm.  We call solid carbon dioxide “dry ice” because at normal atmospheric pressures it converts directly from the solid into the gas. The process of converting a solid into a gas is called sublimation. Both melting and sublimation require the input of heat.

Phase diagrams

Phase diagrams


When a solid substance is heated, it absorbs thermal energy and its temperature increases.  When the melting point of the substance is reached, addition of further thermal energy breaks up the forces holding the solid together, and a liquid begins to form. A balance is established between solid and liquid, and as more thermal energy is added the temperature remains constant while the amount of solid decreases and the amount of liquid increases. The heat absorbed when one mole of solid is melted at constant pressure is called the molar enthalpy of fusion (Δ Hfus). When heat is absorbed, the sign of Δ Hfus is, by convention, positive. Reactions that absorb heat are said to be endothermic reactions. Those which produce heat are said to be exothermic.

On the molecular scale, many processes occur as the ice absorbs heat: the ordered array of molecules in the crystal lattice of the solid is broken down into a collection of mobile, disordered liquid-phase molecules. The water molecules in ice are also losing potential energy associated with intermolecular forces between molecules and gaining kinetic energy of motion.  Meanwhile, the temperature of the water, associated with the average kinetic energy of motion, rises as the molecules move about and vibrate more rapidly.

The change in internal energy for a reaction Δ Erxn can be measured by running the reaction in a constant volume bomb calorimeter. By designing the calorimeter so that no heat leaks out to the surroundings, the heat absorbed by the reaction should equal the heat lost by the calorimeter. The heat absorbed by the calorimeter can be measured by the heat capacity of the calorimeter (Ccalorimeter) times the change in temperature.

Δ Erxn  = – Ccalorimeter Δ T

The melting of ice or the sublimation of carbon dioxide is not a constant volume process but rather a constant pressure process. To investigate the heat flows in a constant pressure process we will have to define a new form of internal energy called enthalpy, H. At constant pressure the change in enthalpy is related to the change in internal energy minus the work done on the system due to volume changes (Δ V).

ΔH  = ΔE  + Δ(PV) = ΔE  + PΔV

At constant pressure, enthalpy changes are a direct measure of the heat absorbed by the system. Enthalpy changes are easier to measure than internal energy changes because it is easier to maintain the constant pressure than constant volume.

Diagram of a constant pressure calorimeter

Diagram of a constant pressure calorimeter

Δ H is measured with a calorimeter where the amount of heat flowing is reflected in a temperature change of a known mass of water. In fact, the unit of energy, the calorie, has been defined to be that amount of heat that will raise one gram of pure water one degree Celsius in temperature. Although the calorie is convenient energy unit, the scientific community now uses the Joule as its standard unit of energy. (1 cal =  4.184 J) Assuming no heat loss during the transfer, the heat of a reaction will be equivalent to the heat absorbed by g grams of water when the temperature raises Δ T degrees Celsius, or

ΔHrxn = – g H2O (4.184 J/g H2O oC ) ΔT

This is the basic equation describing a solution calorimeter that is intended to measure the change in enthalpy during a constant pressure process.

Enthalpy of Reaction (see Hess’s law in the textbook)

Sometimes it is difficult or even impossible to measure an enthalpy of formation directly.  In this experiment we will be determining the enthalpy of formation for CO32-(aq).

C(s) + 3/2 O2(g) + 2 e →  CO32-(aq)

Direct measurement by burning C(s) in oxygen would not give us carbonate, so we will be using an indirect method. According to Hess’s Law (the conservation of enthalpy), if two or more reactions can be added to give a net reaction, Δ H for the net reaction is simply the sum of the Δ H’s for the reactions which are added (energy is additive). Consider the following four reactions:


(1) 2H+(aq) + 2e- → H2(g)                             Δ H1

(2) CO32-(aq) → C(s) +3/2 O2(g) + 2e-       Δ H2

(3) H2(g) + 1/2 O2(g) → H2O(l)                    Δ H3

(4) C(s) + O2(g) → CO2(aq)                           Δ H4


(5) 2H+(aq) + CO32-(aq)  H2O(l) + CO2(aq)

Δ H5 = Δ H1 + Δ H2 + Δ H3 + Δ H4

You will measure Δ H5 directly. We will combine this with the literature values for the heats of formation of H+(aq), H2O(l), and CO2(aq).

H2(g) + 1/2 O2(g)(g)  → H2O(l)

Δ H3 =   Δ Hfo (H2O(l))= -285.83 kJ/mole

C(s) + O2(g)(g)  → CO2(aq)

Δ H4 = Δ Hfo (CO2(aq))= -413.80 kJ/mole

H2(g) → 2 H+(aq) + 2 e-

-Δ H1 = 2 x(Δ Hfo (H+(aq))=  2 x 0 kJ/mole

You will thus be able to calculate Δ H2, which is the enthalpy of formation of CO32-(aq) reaction written backwards.

Bomb Calorimetry

Not like this bomb

Not like this bomb

As we said above, the daily caloric intake of an adult is about 3000 kcal/day. To maintain an energy balance an average adult consumes six tons of solid food over a 40 year period, which amounts to about 12.5 g/hr. Typically we obtain one third of our energy from each primary type of food substance — carbohydrates, proteins and fat. The average enthalpy content of carbohydrates and proteins are about the same (4 kcal/g) while fats average twice as much (9 kcal/g). We will measure the caloric content of food directly by using a bomb calorimeter. We will run a one-gram sample of food in the bomb calorimeter, determine its heat of combustion, and compare the value with what the package says is its caloric content.

Experimental Procedures

Part 1   Enthalpy of Sublimation

Clean, dry, and weigh the calorimeter (2 styrofoam cups, one inside the other – the inner cup will have a series of holes, plus the cover with one hole taped closed).

Fill a 100 mL volumetric flask to the mark with distilled water adjusted to room temperature and weigh it. Place a split-holed rubber stopper carefully over the top of the high precision thermometer.  Carefully insert the thermometer into the volumetric flask and stabilize the thermometer by attaching the split-holed stopper to a ring stand via a Bunsen burner clamp. Record the stable temperature; this is the starting temperature of the water.

Once you have weighed the calorimeter and measured the stable temperature reading, take a piece of CO2 (between 2 and 10 grams) with tweezers and put it between the two cups and weigh the entire assembly. The mass will keep rolling down as it sublimes, therefore after a few seconds record the mass, not bothering with the last digit, immediately take the assembly back to your station where your partner has the thermometer ready to put through the cover. Pour the 100 mL of water into the cups, replace the cover, and gently push down, trying not to spill too much water as it seeps into the lower cup. Reweigh volumetric flask to determine the mass of water transferred.

Even though you see your system “smoking”, the “smoke” shouldn’t be cold to the touch as the CO2 gas has to absorb the heat of the water as it travels through it. Swirl the cups gently, again making sure to hold the two cups together as much as possible and to keep the cover on. Watch the temperature drop. When it stabilizes and you don’t feel the bubbles in solution, swirl again and look in the calorimeter. If it’s cloudy, return to swirling. If it’s clear, record the stabilized temperature.

Empty your calorimeter and thoroughly dry the cups, cover, and thermometer.  Repeat the experiment.  Calculate enthalpy.

Calculations:  The heat gained by the CO2Hsub) is equal to the heat released by the water.  Knowing that the specific heat of water is 4.184 J g-1 Co, calculate the average ΔHsub/g CO2 and the average Δ Hsub/mol CO2.  Compare versus the literature value.

Part 2   Enthalpy of Reaction for formation of H2O and CO2 from CO32- and H+

The calorimeter and thermometer setup will be two small cups, the same size, with no holes.  Rinse the inner cup of your calorimeter with distilled water and dry the cup, cover, thermometer and stirrer thoroughly.

  1. Obtain 100 mL each of the 1.00 M solution of K2CO3 and the 3 M HCl in clean, dry, labeled beakers.
  2. Add exactly 50.0 mL of your solutions to each of two 50 mL volumetric flasks. Weigh both. Load the calorimeter with the 50.0 mL of 3 M HCl , then reweigh the volumetric flask. Insert the thermometer into the 50 mL of K2CO3 solution and determine the initial temperature. Rinse the thermometer and assemble the calorimeter (thermometer and cover).
  3. Begin measuring the temperature of the HCl solution and record the temperature every 15 seconds. After the temperature stabilizes (three consistent measurements), add the K2CO3 solution at a moderate pace so that it does not bubble out (but within 10 seconds), then re-cover the calorimeter. Swirl the solution to release all CO2 and continue recording the temperature until it is approximately constant for several readings (at the end you may see a slight decrease due to heat leakage). Reweigh the 50.0 mL K2CO3 volumetric flask to determine how much was transferred.

Carefully empty the calorimeter and clean the cup, cover, and thermometer. Rinse the cup with distilled water. Dry all pieces thoroughly. Repeat the experiment once a second time. Record your data in a table.

Calculations:  Assume the mixture of the two solutions has the same specific heat as water (4.184 J g-1 Co) and perform a calculation similar to Part 1 to determine Δ Hrxn and then Δ Hrxn/ mol CO32-.  (This is Δ H5.) Use Hess’s law to determine Δ Hf for CO32- (Δ H2 backwards). Compare ΔH2 to literature values.

Part 3   Bomb Calorimetry

Perform bomb calorimetry on a sample at some point during the lab period. Each pair of students should bring and be prepared to analyze a uniform dry (or chocolate) food sample using the bomb calorimeter. Weigh out one gram +/- .0005 g. Crush (if it’s dry) or break the sample into small pieces (if it’s chocolate) and fill the calorimeter pan. Procedures for using the Parr bomb calorimeter will be demonstrated in class.

Calculations:  The read-out from the calorimeter will give you cal/grams. Remembering that a food calorie (Cal) is actually a kcal, compare your results to the value listed on the package.


Complete the worksheet for your lab report. You should have all calculations clearly shown and your final results compared to literature values.

Net Ionic Reactions

adapted by M. J Simpson from a lab originally by S. F. Sontum, K. Jewett and R. Sandwick

Learning goals: work collaboratively with a lab partner, follow instructions to complete a chemistry experiment, collect experimental data, formulate logical conclusions based on experimental results



Much as a cook uses recipes to guide her/him in preparing dishes, chemists use chemical equations as ways of expressing the results of chemical reactions.  Sometimes it is more descriptive to use the ions involved in a reaction rather than the ionic compounds. For example, to express that hydrochloric acid reacts with sodium hydroxide to give water and sodium chloride, we could write:

Molecular equation: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

The above equation is fine in some ways, but in actuality NaOH and HCl solutions do not contain NaOH and HCl molecules, and there are no NaCl molecules in the solution after the reaction occurs. The better description for this reaction is:

Total ionic equation: H+(aq) + Cl(aq) + Na+(aq) + OH(aq)  →  Cl(aq) + Na+(aq) + H2O(l)

Net ionic equation: H+(aq) + OH-(aq) → H2O(l)

The first equation is called the total ionic equation, while the second is termed the net ionic equation.  A net ionic equation summarizes the changes that have taken place as a result of a chemical reaction.


The solubility of a substance in a solvent is the maximum amount of the substance that can be dissolved in a given amount of solvent. While there is no exact definition for the boundary between soluble and not soluble, a general guide might be:

SOLUBILITY                                      TERM

> 0.1 M                                                 Soluble

0.01 to 0.1 M                                       Moderately soluble

< 0.01 M                                              Slightly soluble

Some guidelines have been prepared to help estimate a compound’s solubility. These are written in a manner where statements above other statements have priority. For example, an alkali metal carbonate is soluble since “all alkali metals (#2) are soluble” have priority over “All carbonates are insoluble.” [There are exceptions that are not included here. (See textbook)]    

Solubility Rules:

  1. All nitrates (NO3-) and acetates (C2H3O2-) are soluble.
  2. All salts of alkali metals and NH4+ are soluble.
  3. All common halides (Cl-, Br, I) are soluble, except those of Ag+, Pb2+, Cu+, Hg22+
  4. Hydroxides (OH-) are insoluble except for Na+, Ca2+, and Ba2+.
  5. All sulfates (SO42-) are soluble, except BaSO4 which is slightly soluble, and CaSO4 and Ag2SO4, which are moderately soluble.
  6. All carbonates (CO32-) and phosphates (PO43-) are insoluble.


Note: perform this experiment with a partner, but be sure to observe the tests together.
Make up the following standard solutions:

50.0 mL of 0.10 M KI(aq) from KI(s)

25.0 mL of 0.10 M CuCl2 (aq) from CuCl2 ● 2H2O (s)

50.0 mL of 0.10  M H2SO4 (aq) from 1.00 M H2SO4(aq)

25.0 mL of 0.10 M NaOH (aq) from 1.00 M  NaOH(aq)

50.0 mL of 0.10 M Na2CO3 (aq) from Na2CO3(s)

Three test solutions will be prepared for you:  0.10 M AgNO3, 0.10 M BaCl2, and 0.10 M Co(C2H3O2) 2.

To determine how to prepare a solution from a solution of higher concentration, use the formula: Ci Vi = Cf Vf


Ci = initial concentration of higher concentration solution

Vi = initial volume of the higher concentration solution

Cf = final concentration (of the diluted solution)

Vf = final volume of the diluted solution (volume of the initial solution plus water)

Tests for chemical reactions:

For each of the solutions you prepared, test each solution with the following test solutions:

0.10 M AgNO3

0.10 M BaCl2

0.10 M Co(C2H3O2) 2

0.10 M Na2CO3 (the solution you prepared)

A test involves taking 5 mL of your solution and adding 3 – 5 drops of the test solution. Make observations in the table on the worksheet.

Things to note:

(a)  a change of color;

(b)  a separate phase is formed;

(c)  any other change.


Complete the worksheet for your lab report. Optional survey.

The following procedure may help you when writing a net reaction for any chemical reaction.

  1. Determine the principal forms present of each reactant in solution.
  2. Determine the principal forms present after the reaction has occurred.
  3. Write a balanced equation for the reaction.
  4. Cross out reactants and products that do not change (the spectator ions).

Example:  Mix HNO3 solution + KOH solution

Principal substances before reaction:  H+, NO3-, K+, OH-

Principal substances after reaction: NO3-, K+

Product formed:  H2O

Net reaction:  H+(aq) + OH-(aq) → H2O (l)

Glucose in Natural Products


Structure of D-glucose

Structure of D-glucose

As the world’s most common biomolecule, glucose (C6H12O6) has numerous cellular functions including serving as a metabolic fuel, as a precursor to energy storage molecules such as starch and glycogen, and as a building block molecule to important structural components such as cellulose. Glucose is a simple sugar or monosaccharide, meaning it contains a carbonyl group (C=O) along with several hydroxyl (-OH) groups. It is just one of many natural molecules that fall in the simple sugar class (fructose, galactose, ribose are others); these vary by the number of carbons and hydroxyls and by the spatial arrangement of the –OH groups around the carbons. Due to the multiple hydroxyl groups, monosaccharides are highly soluble in water.


In this laboratory, you and your partner will design a research plan to investigate the glucose levels in a system of your choosing. The study should not be a simple comparison of the glucose levels in two different, unrelated products; you should instead design the experiment to ultimately “tell a story” relating the glucose levels to something about the 18430036-Vector-illustration-of-Cartoon-Scientist-Stock-Vector-scienceproducts. For example, an appropriate title of your project may be “An Investigation of the Comparative Glucose Concentrations of Two ____(fill in the blank)_________.” You can pick different products to compare or you might want to pick one product which somehow gets manipulated (temperature, storage, etc.). The plan obviously needs some discussion and some preliminary research of the literature (i.e., Does the product you intend to use actually have glucose in it?). Creativity is admired. The experiment cannot include the use of animals nor the use of human body fluids. The lab instructor will work with you to determine the feasibility of the project.

Analysis of Glucose

A spectroscopic procedure will be made available to you to determine the glucose levels in your samples. Like all simple sugars, glucose does not absorb light in the visible range. We will therefore need to use a combination of reagents to generate colors that are proportional to the amount of glucose in solution. Here, the mixture of reagents includes phenol, aminoantipyrine, glucose oxidase (GO), and horseradish peroxidase (HRP); these oxidize glucose while simultaneously producing a rose-colored product that absorbs light at 505 nm.

The important aspect of this procedure is that the absorbance produced at 505 nm in this reaction is proportional to the amount of glucose you add to the reagent mixture. The instructions on how to perform the analysis is at the end of this write-up. In olden days, chemistry was mostly done using sight and smell as the analytical tools. For example, if you worked as a chemist for Gatorade in the mid-1800s, you would have to use your eyeballs to adjust the color of Gatorade XTREMO Tropical Intensio. Late in the 1800s, the colorimeter was invented by …. uh, uh, … John Colorimeter…? and it proved useful in comparing the colors of two solutions (one being a standard) and was a blessing for the colorblind chemist. The unit used sunlight as the light source and colored filters to select a particular light type. In the 1940s, Arnold Beckman and his associates in what was later to become the Beckman Instrument Co. developed and highly commercialized the first ultraviolet/visible spectrophotometer, a unit that used bulbs as a light source, a prism to select a desired wavelength, and a phototube to measure the light that made it through the sample. You are already familiar with the spectrophotometer, the concept of absorbance, and Beer’s Law.

Procedure for Determination of Glucose

The following procedure is appropriate for glucose concentrations in the 0 – 2.0 g glucose/L range. If your sample has a glucose concentration higher than this value, the solution will need to be diluted and this dilution will need to be factored back in when figuring out the original solution’s concentration.
Since the intensity of color that a given amount of glucose will generate by this procedure is unknown, you will need to prepare some glucose standards and run them in the procedure to generate a [glucose] vs Absorbance standard relationship. (It should be linear.) Prepare a 2.0 g/L glucose solution using crystalline glucose. Using this solution and test tubes, prepare a group of standards (volume = 4.0 mL) in the range 0 – 2.0 g glucose/L. These samples should be spread out fairly evenly throughout the range and should include a 0.0 g/L sample and a 2.0 g/L sample. A typical standard relationship is typically generated with four or five samples throughout this range.

To perform the procedure, set the spectrophotometer to 505 nm. In the cuvette (sample holder) add 1.0 mL of the color reagent and 0.10 mL of your standard or sample. Using a piece of Parafilm, flip the cuvette quickly to mix, put it immediately into the spectrophotometer, and start monitoring the rise in absorbance. Measure the absorbance change over a two-minute period. It is the rate of absorbance change at 505 nm that should be proportional to glucose concentration. You will need to prepare an Excel graph that shows the relationship of absorbance increase to the glucose concentration. Let Excel determine the slope of the line and the correlation coefficient (R2) of the fit. Before leaving the laboratory, look at your data closely. If a point on your relationship seems off, you may want to re-run that standard.

You should run your samples in the same way as you ran the standards, i.e., 1.0 mL of the color reagent to 0.10 mL of sample. For accuracy the rate you obtain for your sample needs to be in the range of the rates you measured for the standards. If it the rate is too high, perform a dilution on the sample and re-run the procedure on the diluted sample. (Again, you will then need to back-calculate to the concentration of the sample by knowing this dilution.)

Final Report

Upon completing the analysis, you and your partner should combine to put together a report on your findings using a typical journal format of:

  • Title: specific, descriptive, concise;
  • Abstract: short summary of what you did, why, and what you found;
  • Introduction:
    • include appropriate background (such as how a spectrophotometer works, how the spectrophotometry data will help you answer your question, and the reaction you are using to detect glucose, the Trinder glucose activity test);
    • introduction of the question (such as why you think it is interesting, and what you think will be the outcome and why);
  • Experimental Procedures: in narrative form because this is a formal report, and specific enough for someone to repeat it exactly how you did it;
  • Results: well-formatted tables and/or graphs with captions. Show all of your raw data in whatever form is clearest;
  • Discussion: show and explain how you used your raw data to find the answer (such as plotting your calibration curve and sample data together). Give a detailed, logical argument leading to your final answer of how much glucose is in your sample (either in terms of %wt or g/mL concentration of your full strength solutions). Compare your findings to your hypothesis and attempt to explain any discrepancies. If possible, compare your answer to a cited literature value;
  • Conclusion: the answer you came up with and a summary of how you solved the problem. Discuss possible sources of error that lead to uncertainties about the accuracy of your conclusion. Suggest a future experiment that would continue this work.

Note: this is a formal scientific writing assignment. Grammar and spelling count. Please proof-read your work before submitting it.

Format: Section headers are recommended. Citations must be formatted appropriately. You may want to take a look at a biochemistry or chemistry journal paper to see a properly formatted paper and/or consult the ACS style guide. I will gladly proof-read your papers, but you must show me your draft at least a week before it is due because it may take me a few days to read through it.

Turn in your report electronically with Canvas.

Titration of Citric Acid

Learning Goals

  1. Perform a titration and analyze the uncertainties associated with titrations;
  2. Employ a color pH indicator;
  3. Apply stoichiometry to analyze titration data;
  4. Observe an acid-base reaction.

by S. Choi, R. Gleason, and K. Jewett (with edits by R. Sandwick and M. J. Simpson)


Citric acid is a polyprotic acid (can release three H+s) that is a bit on the weak side (i.e., tends not to ionize completely).  In solution in fruit juices, it lets a small portion of the H+ go, however this small amount of acid is enough to create a pH = ~3 solution and a sharp taste on the palate.  If strong base is added to citric acid it will sequentially lose its three protons in the following manner:

Citric acid deprotonation in 3 steps. Credit: CrystEngComm, 2014,16, 3387-3394

Citric acid deprotonation in 3 steps. Credit: CrystEngComm, 2014,16, 3387-3394


In today’s experiment we are going to determine the amount of citric acid in a fruit juice by using a base-acid reaction. If we know exactly how much base we add to completely remove all the H+ ions (called “deprotonating”) from the citric acid in the juice, we can calculate how much citric acid is in the solution. This process of employing one reagent of known concentration to determine a compound of unknown concentration in solution is termed titration. It usually involves slowly adding small amounts of the titrant to the analyte until a reaction is just barely complete. The apparatus typically used is a buret. The equivalence point (or end point) is the exact point where all the analyte in solution has reacted. Since the equivalence point of many titrations do not result in observable changes, end point indicators are added to (are you ready for this?) indicate the end point.

In this experiment you will use a solution of NaOH to titrate the acid in a fruit juice.  To be accurate, the exact concentration of the NaOH solution you prepare must be known. Powdered NaOH (from which you will make your solution) is known to slowly decompose upon reaction with CO2 in the air to generate NaHCO3. Thus you will need to standardize, or precisely determine the concentration, the NaOH solution using a stable primary standard. Only after you have satisfactorily determined the exact concentration of your NaOH solution can you accurately determine the citric acid concentration in the fruit juice.

There are several good primary standards for standardizing base solutions, but one of the best and cheapest is the compound oxalic acid dihydrate, H2C2O4 ● 2 H2O. (Oxalic acid is a natural acid found in rhubarb leaves; it is toxic so don’t eat it.)  It is important to note that the chemical equation (shown below) shows a stoichiometry of one moles of oxalic acid to every two mole of NaOH in this reaction.

H22O4(aq) + 2 NaOH(aq)  → C2O42-(aq) +  2 Na+(aq) + 2 H2O(l)

The indicator we will use in both is phenolphthalein, a common indicator of acid-base titration. Phenolphthalein was the active ingredient in Ex-lax until recently when it was phased out due to its carcinogenicity.


Note: do this lab with a partner.

Standardization of a Solution of Sodium Hydroxide
  1. Thoroughly clean and rinse with distilled water your supplies: a burette, a 25-mL graduated cylinder, a 500 mL boiling flask, and three 250 or 300-mL Erlenmeyer flasks.
  2. Put your NaOH pellets (mass calculated in the pre-lab) in the 500 mL boiling flask. Cover them with a small amount of water and swirl until they dissolve. The heat of solution produced by the NaOH helps to speed the dissolving process. Once the pellets are dissolved, fill the flask to the base of the neck. Using Parafilm, mix the solution thoroughly by inverting several times (the flask not you!). It is best to keep the flask covered with Parafilm when not in use as CO2 from the air can slowly neutralize the NaOH.
  3. Using an analytical balance, weigh out the oxalic acid dihydrate in a weigh boat (mass calculated in the pre-lab). Place it in a clean, clearly labeled Erlenmeyer flask. Note: It is not necessary to weigh out exactly the amount calculated (although you should be close), however it is imperative that the mass of each sample of oxalic acid dihydrate be known precisely for each flask. Label each clearly.
  4. Dissolve the oxalic acid dihydrate in 25 mL of distilled water and add three drops of phenolphthalein indicator solution.
  5. Repeat steps 3 and 4 for 2 more samples of oxalic acid dihydrate so you have a total of 3 flasks containing oxalic acid solutions.
  6. Rinse your buret once with water and then twice with 5-mL portions of the solution of sodium hydroxide which you have prepared, draining the solution off through the burette tip into a beaker for waste reagents. Fill the buret nearly to the top of the graduated portion with the solution of sodium hydroxide you have prepared, making sure that the buret tip is completely filled with the solution. Touch the inner wall of a beaker for waste reagents to the buret tip to remove any hanging drop of solution.
  7. Make a preliminary titration using one of your solutions of oxalic acid to learn approximately how the neutralization proceeds. Place a sheet of white paper under the Erlenmeyer flask so that the color of the solution is more easily observed. Make sure to either adjust the meniscus to the 0.00 mL line or record the exact volume to the nearest 0.05 mL. Swirl the sample in the flask throughout the titration. Add sodium hydroxide rather rapidly from the buret until the color of the solution where the sodium hydroxide is entering the solution begins to linger, then add the base dropwise until, finally, one drop of the alkaline solution of base changes the colorless solution to a permanent pink, not red. A drop should not be left hanging on the buret tip. You probably will overrun the endpoint in this first titration, but it will provide a useful rough measure of the volume of the sodium hydroxide solution needed to neutralize the acid and provide you with valuable experience. Read and record the level of the meniscus in the buret (to the nearest 0.05 mL), and compute the volume of basic solution used in the titration.
  8. Now titrate the remaining two samples of standard acid, being certain each time to refill the burette nearly to the top graduation with your sodium hydroxide solution and to record the burette reading. In these runs, add the sodium hydroxide from the buret very rapidly, again with swirling, until you are ~ 2 mL short of the volume that you estimate will be needed on the basis of your first titration. Then carefully add base drop by drop so that you can determine the equivalence point accurately. Record the data for your titrations in a table and perform the calculations necessary to yield the exact molarity of the NaOH solution. Show your results to your instructor. The concentrations of your titrations should agree within 5 % of each other.
  9. The solution of sodium hydroxide that you have just standardized will be used in Part II, so do not waste it.
Total acidity of a citrus fruit

Several different types of samples will be available for you to use. When using fruit, squeeze the juice into a 250-mL beaker by cutting the end from the fruit and applying pressure. Use a Büchner funnel and filtering flask to vacuum filter the juice to remove the pulp. If you are using one of the juice samples, simply use the juice as is. Deliver 5.00 mL of juice (or 2.00 mL if its highly acidic) into a tared 250 or 300-mL flask (you’ll need to reuse one from the previous part of the experiment). Weigh the tared flask and its contents to the nearest 0.0001 g and then dilute the juice to approximately 50 mL with distilled water. Add three drops of phenolphthalein indicator.

Titrate the prepared juice solution to the phenolphthalein end­point. Record the final burette reading to the nearest 0.05 mL. Repeat the experiment until you are satisfied with the precision of successive runs. Put all data into an appropriate table.

All waste can go down the sink drain.


Calculate the moles of oxalic acid reacted, the moles of NaOH titrated, and the molarity of the NaOH.

Having discovered the exact molarity of the NaOH you used, calculate the number of moles of citric acid in each sample. Convert this to grams and then to grams citric acid per grams sample. Convert this to a percent citric acid by weight.

The reaction of sodium hydroxide with citric acid is:

3 NaOH   +  H3C6H5O7 → 3 Na+  +  3 H2O  +   C6H5O73-


Fill out the worksheet for the report. Look up values for citric acid concentration in juices and compare your value vs. those literature values. Give a proper citation for the source.


Spectrophotometric Determination of Iron

by E. L. Pool and D. Copeland (with edits by R. Sandwick and M. J. Simpson)

Learning Goals

  1. Practice handling hazardous chemicals safely;
  2. Use a bunsen burner to heat a solution;
  3. Vacuum filter a solution;
  4. Practice using a transfer pipet and a mortar and pestle;
  5. Manipulate experimental conditions to illustrate the concept of limiting and excess reagents;
  6. Use Beer’s Law to create and apply a calibration curve for quantitative analysis;
  7. Consider the consequences of the limitations of the spectrophotometer and the analytical balance.


You will use spectrophotometry to determine the amount of iron in a multivitamin to see if the manufacturer’s claim is correct. Iron itself is not a huge absorber of light, but when it (in solution in the Fe2+ form) binds to 1,10 phenanthroline (C12H8N2), it forms a highly stable red/orange-colored species. By quantifying the color with spectrophotometry, we can deduce the concentration of iron in the solution and back-calculate the amount of iron in the original pill. This can be compared to the manufacturer’s claim.

phenanthroline1,10 Phenanthroline


The formation of the red/orange-colored iron-phenanthroline complex requires the iron to be in the Fe2+ form, and the procedure thus includes the reagent hydroxylamine hydrochloride (NH2OH ● HCl) that will reduce all iron in the sample to the ferrous form. Sodium acetate (NaC2H3O2) is used to control the pH of the solution since this also affects the absorbance.

The red/orange iron-phenanthroline complex absorbs light at 508 nm, which is green light. This should make sense if you remember the Kool-Aid lab. To quantify the intensity of the color, you can use the spectrophotometer to measure the “absorbance.”

A key assumption in this experiment is that all of the iron will be converted to the colored complex, meaning there must be excess phenanthroline present. In preparation for the experiment, you will analyze the effects of varying the amounts of the two reagents that are directly involved in the production of the colored complex: the iron and the phenanthroline. This will show you the conditions under which this assumption is valid (or not!).

Once you measure the absorbance of the pill solution, how will you know how to correlate the absorbance with a concentration?  One way is to know the extinction coefficient and to plug the appropriate numbers into Beer’s law. Another way is to prepare iron standards and to generate a standard curve of absorbance versus iron concentration. This latter technique is preferred as it accounts for any small changes one might have in her/his particular laboratory situation. We will use this method in our iron determination.

An online tutorial of this lab is available here.


Note: do this lab with a partner.

Sample (Unknown) Preparation

The sample preparation requires crushing a pill, boiling the powder in acid to release the iron, and filtering the solution to remove particulates that could interfere in the spectrophotometry measurements.

  1. Each pair should obtain one vitamin pill. Determine the mass of the pill. Using a mortar and pestle, crush the pill into a fine powder. (You can take your angst out on the pill, but be careful none of it escapes the mortar.) Transfer the pulverized pill powder quantitatively to a 400 mL beaker. Determine the mass of the powder. In your final calculations, you will need to scale up your final iron concentration in the pill to account for any powder that you did not use.

    Apparatus to heat a solution

  2. Add approximately 150 mL of 1.0 M hydrochloric acid (HCl) [CAUTION!]. Add a few boiling chips. Use an apparatus similar to the illustration to heat the solution to near boiling (simmer) and keep heating at that temperature for 10 minutes. Do not leave your solution unattended. Avoid loss by splattering. Record observations about the appearance of your solution.
  3. Set up a vacuum filter apparatus with a Buchner filter, a vacuum flask, and a piece of filter paper. Hold the flask firmly or clamp it to a ringstand so it does not tip over. Pour 50 mL of DI water through the filter first, then turn on the vacuum and pour the hot solution through the filter. Your final solution should be clear and yellow. Record observations about the clarity of your solution in Table 3.
  4. Pour the filtered solution into a volumetric flask. Dilute to exactly 1.000 L with distilled water and mix thoroughly.

Solution Preparation

  1. Label 7 clean test tubes with the concentration of iron that they will contain (in units of mg/L): 0, 5, 10, 15, 20, 25, and the unknown vitamin solution.
  2. Label 4 small beakers with the reagents you will acquire from the stock solutions: sodium acetate, hydroxylamine HCl, 1,10-phenanthroline, and DI water. Fill these beakers with about 10 mL of their respective reagent. Put a plastic pipet in each. 
  3. To each test tube, add 1 mL of each: sodium acetate, hydroxylamine HCl, and 1,10-phenanthroline. Note: you are using a concentrated phenanthroline solution that ensures it is in GREAT excess.
  4. To each test tube, add appropriate amounts of water and/or 25 mg/L iron stock solution to make a total of 5 mL of the concentration of iron specified on the label. Record the amounts of water and iron stock solution you used in Table 4. In the vitamin solution test tube, put 5 mL of the vitamin solution. 
  5. Mix test tubes with gentle agitation and let the reaction complete for 10 minutes. Although you should use this time wisely, I advise against cleaning up, as you may need to remake a solution or two before the lab ends.
  6. Record observations about the color change and any trends you notice in Table 4. Note in your observations what concentration iron solution appears most similar to your vitamin sample. Pro-tip: these qualitative measurements should agree with your quantitative spectrophotometric measurements that you make in the next step. Your eyes are good spectrophotometers, and something is probably wrong if the quantitative measurements disagree with your visual inspection.

Spectrophotometric Analysis

  1. Get 7 clean cuvettes in a cuvette tray. Label them with the concentration of iron that they will contain (in units of mg/L): 0, 5, 10, 15, 20, 25, and the unknown vitamin solution. Transfer approximately 1 mL of each solution from its test tube to its corresponding cuvette.
  2. Using a Spectronic 401 at 508 nm, read and record the Absorbance (A508) of your solutions. Record this data into Table 4.

Clean up

Check that your data in Table 4 increases linearly from low concentration of iron to high concentration of iron before cleaning up any part of your experiment. Ask your TA for help neutralizing the pill solution, then pour it down the drain. Pour the contents of the test tubes and cuvettes into the waste container. Throw away the cuvettes.

Data Analysis

The absorbance measurements of the standard solutions are used to determine the concentration of iron in the pill solution, which is used to determine the amount of iron present in the original pill. You can plot the concentration versus absorbance, then perform linear regression to find the formula that relates concentration and absorbance.

With this formula, you can simply plug in the pill solution absorbance and calculate the unknown concentration of iron in the pill solution. Using the concentration of iron in the pill solution, calculate the amount of iron present in the original pill. (Note: you will have to scale up your value to account for mass lost between the original pill and the pill powder.) You can then compare this value to the amount listed on the bottle in terms of percent error. The percent error is the absolute difference between the experimental and theoretical value divided by the theoretical value, then multiplied by 100 to get a percent.


Fill out this worksheet. Turn in either a paper or digital copy. Optional survey.

Chemistry of Colors

by M. J. Simpson


Colored chemicals absorb and sometimes emit light in the visible portion of the electromagnetic spectrum, which is 400 – 700 nm. Absorption occurs when an electron absorbs the energy from the light, temporarily promoting the electron to a higher energy orbital. Light emission can occur when the electron relaxes back to the ground state and produces light, but emission is less common than absorption because there are non-radiative ways for the electron to relax.



Unless a chemical is emitting light, the color that a chemical absorbs is the opposite of the color that it appears. The color wheel shows which colors are opposite. For example, -carotene, a pigment found in many fruits and vegetables including carrots, absorbs purple and blue light (400 – 500 nm), so it appears yellow/orange.

In today’s lab, you will investigate a variety of colored dyes. In the first part of the lab, you will use paper chromatography to determine the number of dyes present in your Kool-Aid packet. Then, you will measure the absorption spectrum of several food dyes and your Kool-Aid packet. This will allow you to identify the dye(s) present in your Kool-Aid packet. In the second part of the lab, you will use prepared dyes to alter the absorption properties of a cellulose t-shirt, which does not absorb visible light in its natural state.


Paper chromatography experiment

chromotograpguPrepare a dilute salt water solution using approximately 0.1 g of NaCl per 100 mL of DI water. Take a small amount of your Kool-Aid packet. Mix with a few drops of water to make a dark-colored slurry. A good concentration is about 15% wt. Kool aid powder in water. Using a wooden stick, make a dot of the mixture on a rectangular piece of filter paper about an inch from the bottom. Wait for it to dry. If it is not clearly visible, make another dot in the same spot. You can accelerate the drying process by blowing air on it. Suspend the paper in a beaker containing a small amount of salt water using the same wooden dowel. Allow the filter paper to just touch the top of the water (you don’t want the water to touch the spot of Kool aid). Leave the paper in the water until the dye(s) sufficiently separate or until you are sure there is only one dye present (wait at least 10 minutes).

Measuring the absorption spectrum of dyes and Kool-Aid

Microplate reader

Microplate reader

Mix up the Kool-Aid solution with DI water in a beaker according to the packet’s instructions. You will only need a small sample of the Kool-Aid, so you may want to scale the instructions on the packet to mix up a small portion of the Kool-Aid. Use a micropipette to put a 200 μL sample of each dye and of your Kool-Aid mixture in a microplate and measure the absorption spectra with the microplate reader. Be sure to also run a blank sample of pure water. Record the major peak(s) present in each spectrum in a table. If you have never used a micropipette before, read these instructions for use of a micropipette. If your Kool-Aid absorbance measures greater than 1 at any wavelength, dilute the solution and try again. Be sure to record in your notes how you diluted it. Note: “tropical punch” flavor must be diluted extra.



Use this worksheet for collecting your data. For the lab report, you and your partner will be making a poster using either PowerPoint or Illustrator. Use photographs of your experiment and the graphs and tables you generated during the experiment to make a visually appealing and informative poster. It should include the following:

  1. A concise, descriptive title;
  2. A brief abstract summarizing the experiment and the results;
  3. A short introduction to the question you are answering;
  4. A summary of the procedure you used (including photographs);
  5. The results you found in the form of tables and graphs;
  6. A discussion of the results, which must answer the following questions:
    • How many dyes are present in your Kool-Aid? How did you know?
    • What dye(s) are present in your Kool-Aid? How do you know?
    • How do your results compare to the ingredients listed on the packet? If your results disagree with the ingredients list, what may have caused the disagreement?
    • What are some sources of uncertainty? How might you resolve them?
    • What other experiments could you do to confirm your results?
  7. A conclusion summarizing your findings.

Use this website for help creating an effective poster. Turn in an electronic version of your poster on Canvas and also post a .pdf version on the shared folder. We will present them on a high resolution screen, but if you would like to print a full size version, there is a plotter in the library you can use.

Penny Statistics

adapted by M. J. Simpson and R. Sandwick from a lab originally written by by K. Jewett and S. Sontum

Learning goals: analyze experimental data, formulate a logical conclusion, explain likely sources of experimental error


A basic understanding of statistics is required in most scientific disciplines in order to analyze experimental data. For example, statistical analysis tells a pharmaceutical company whether or not a new drug is effective. One of the most common statistical tests is Student’s t-test, which helps to determine whether two populations are significantly different. The critical value that comes out of a t-test represents the probability that the populations are NOT significantly different. A common definition of “statistical significance” is p < 0.05, although this has been a topic of recent debate. This laboratory is designed to give you an introduction to routine techniques such as weighing and performing basic calculations and statistics.


In 1795 at the age of eighteen a German mathematician Karl Friedrich Gauss devised the Gaussian curve or normal distribution (symmetrical about the average) in which the scatter of a data set is summarized by the standard deviation, or spread, of the curve. Gauss found that the best estimate of the area of uncertainty was the standard deviation, which for N measurements is given by:

The Gaussian curve has turned out to be a nearly universal description of the distribution of random data sets. Student’s grades measured by an exam follow this bell shaped curve, as well as the distribution of the weight of each grain of wheat in a field or the repeated measurements of the weight of a single wheat grain. In the bell shaped Gaussian curves depicted above, 67 % of the points fall within one standard deviation on each side of the average while 95 % of the points are with ± 2 standard deviations around the average.


When two averages are two standard deviations or more apart we can say they are significantly different at a 95% confidence level, which corresponds to p = 0.05. The smaller the p value, the more confident you can be that the two averages are significantly different.

In today’s lab, you will determine whether or not there are more than one kind of U.S. minted penny based on the mass of the pennies. You will record the masses of pennies, perform statistical analysis on the raw data, and draw conclusions based on those analyses.


Each group will pick 20 pennies out of the pool. Weigh each to four decimal places and record the mass and the year in a table. If you have never used an analytical balance before, consider reviewing this information.

Statistical Analysis

Try to group the pennies into two groups with significantly different average masses. Calculate the averages and standard deviation for the two groups. Then, calculate the p-value for the hypothesis that there are two kinds of U.S. minted pennies. This is the equation for the t-value, which you can use to look up the p-value on the t-table. Alternatively, you may use Excel or some other software to calculate the p-value, but you must understand the math in order to answer the post-lab questions. Here is an example to show you how to do the math on Google Sheets.

Note: the number of significant digits for the average depends on the standard deviation. The last significant digit of the average is the first decimal place in the standard deviation. For example, if your average is 3.025622 g and your standard deviation is 0.01845 g, then this is the correct number of significant figures for the average: 3.03 g, because the first digit of the standard deviation is in the hundredths place, so the last significant digit of the average is in the hundredths place. For the p-value, just report the range. For more information about significant figures, refer to this tutorial.

Report: Fill out the worksheet for the report. Turn in either a paper or digital copy. Optional survey.

Lab documentation

During the semester you will be asked to develop various aspects of your scientific record keeping and report writing. These talents will hopefully be expanded and built upon during subsequent chemistry and other science lab courses. The exact documentation requested for each laboratory exercise is explained at the end of each experimental write-up.

Preparation for an Experiment

Before you arrive in the laboratory to perform a given experiment it is essential that the experiment be studied carefully, with special emphasis on the overall experimental design and on the general flow of the procedure. The laboratory will run smoother and more safely if everyone arrives prepared.  Occasionally, it will be necessary to make changes in a given procedure. These changes will be brought to your attention early in the lab period; please note them immediately in the appropriate place in this manual.

Prior to starting an experiment, a researcher will often clearly record her/his objectives in an attempt to clearly define the purpose of the experiment in order to best design procedures that can achieve these purposes. We will ask you to prepare for each laboratory experiment by reading the lab handout and then by writing a short objectives section prior to arrival at lab.   A sample objective section for an experiment might be:

Experiment 1

Volume Calibration of a Kimax Buret and Density of a Metal

In this laboratory, I am going to establish a calibration curve for a 50.0 mL buret.  Each buret is slightly different and the volumes they deliver may or may not agree with the volume markings.  I will determine this by dispensing volumes from the buret and then correlating these volumes to their actual volumes by weighing the dispensed amounts and converting the determined mass to volume by knowing the density of water at room temperature. (This value is 0.998204 g/mL.)  Linear regression will be used to statistically treat the data. In the second portion of the experiment I will determine the density of a metal by measuring its volume and mass.  Again, the data will be treated statistically to determine the accuracy of my measurements.

Laboratory Notebook Report

For at least one laboratory, you will be asked to prepare and submit a document in the form of a laboratory notebook. In official laboratories (for example, industrial or research labs) where documentation is critical, it is customary to record experimental events directly into a bound notebook with dated and numbered pages.  The flow of information for a particular experiment should tell a story – it should have an opening, a plot, and a conclusion.  When complete, the documentation should provide a second reader a clear understanding of what happened in the experiment and should include all the necessary information to permit the reader an opportunity to repeat and validate the experiment.  Through the years, the scientific community has established a standard sequence to laboratory reports; it is:

  1. Title: A concise, specific, and descriptive title for the experiment.
  1. Objectives: A brief statement about the experiment and what you plan to accomplish (what you wish to accomplish) in the experiment. Your objective is not “to learn how to….” but to do something. (See example.)
  1. Procedure: A list or narrative of the steps you did in performing the experiment. Your experimental record should be complete enough to permit someone else to perform the same experiment in the same way.
  1. Observations: Your record of measured data and observations.
  1. Results: A section where results are calculated and presented in tabular form.
  1. Discussion:  A section to show how you analyzed your results and to present a discussion that eventually leads to a conclusion.
  1. Conclusions: A final paragraph should summarize the conclusions of the study. This section is also where you can mention future work that you would like to do to follow-up this experiment and where you can discuss errors and uncertainties associated with your results and conclusions.

Instead of a notebook we will provide worksheets. You should treat these write-up sheets as your laboratory notebook for these experiments. All primary data should be recorded directly onto your write-up as you do the experiment.

For one laboratory report (the glucose lab), you will be asked to prepare a report in scientific journal form.  You should review chemistry journal articles in preparation for preparing this report to see what needs to be included and how data are presented.