Iodine Clock Reaction
In this demonstration an almost clear liquid will suddenly turn dark blue. You will need:
 1000-mg Vitamin C tablet
 Hydrogen peroxide (3%)
 Tincture of iodine (2%)
 Liquid laundry starch (or corn starch)
1. Create “Solution A” – Pulverize the Vitamin C tablet by placing it into a plastic bag and crushing it into a fine
powder. Put all the powder in the first cup and add (60 ml) of warm water. Stir for at least 30 seconds.
2. Create “Solution B” – Put 5 ml of Solution A into a new cup and add to it: 60 ml of warm water and 1
teaspoon (5 ml) of the iodine. Notice the brown iodine turned clear.
3. Create “Solution C” – In another cup, mix 60 mL warm water, 15 ml hydrogen peroxide and 2.5 ml starch.
4. Perform the Reaction -- Pour all of LIQUID B into LIQUID C. Then pour them back and forth between the 2 cups
a few times. Place the cup down and observe….be patient....somewhere between a few seconds and a few
minutes, the liquid will suddenly turn dark blue!
Class Discussion Questions:
 What measurements could you take to “quantify” the kinetics of this reaction?
 What factors could you alter to turn this demonstration into an experiment? What questions could you ask?
The Chemistry of the Reaction:
Tincture of iodine contains 0.08M iodine and 0.16 M sodium iodide in solution of ethanol and water. Hydrogen
peroxide slowly oxidizes the iodide in acidic solution to iodine (equation 1 below). Once the iodine is formed
vitamin C rapidly reduces it to regenerate I- (equation 2).
H2O2(aq) + 2 I-(aq) + 2 H+(aq) → I2(aq) + 4 H2O(l)
(1)
C6H8O6(aq) + I 2(aq) → C6H6O6(aq) + 2 H+ (aq)+ 2 I-(aq)
(2)
When the vitamin C has been consumed I 2 accumulates, reacts with I- to form the I3- ion, which reacts with the
starch to form an iodine/starch complex with a purple color.
Foaming Hydrogen Peroxide – Reaction Rate vs. Concentration.
In this experiment students will investigate and quantify the relationship between concentration of a reactant and
the time it takes for a reaction to proceed. You will need, per student group:
 4 clean small plastic soda bottles
 1 package dry yeast
 Medicine measuring cup
 A cup of warm distilled water
 100 mL 20-volume hydrogen peroxide
 Liquid dish washing soap
(6% solution)
1. Create four different concentrations of hydrogen peroxide, each 30 mL in volume. For 5%, mix 25 mL 6%
hydrogen peroxide and 5 mL water. For 4% mix 20 mL and 10 mL, and for 3% mix 15 mL and 15 mL.
2. Pour your peroxide solutions into the four water bottles.
3. Add about 1 mL of liquid dish soap into each bottle and swish the bottle around to mix it.
4. In a separate small cup, combine 60 mL of warm water and the yeast and mix for about 30 seconds. Separate
into four equal portions.
5. While one student notes the time, have four students simultaneously pour one yeast solution into each bottle.
6. Record the time it takes for the foam to reach the opening of each bottle.
Data Table and Graphing
 Have students prepare a table to record their data – time and concentration.
 Each group should create a graph of time vs. concentration.
Class or Lab-Group Discussion Questions
1. Consider just one of the reactions you observed. If you could have measured the change in the moles of
hydrogen peroxide with time, what would the graph look like? Sketch this graph, without specific values on the
axes.
2. On the same sketched graph, draw curves for the changes in moles of oxygen and moles of water that would
also have occurred with time. Why does the stoichiometry of the chemical reaction help here?
The Chemistry of the Reaction
The decomposition of hydrogen peroxide into water and oxygen gas is a very slow process. Here it is speeded up
by the presence of a catalyst, the yeast. Each tiny foam bubble is filled with oxygen. Also note that the bottle gets
warm – this is an exothermic reaction. You can test the foam for the presence of oxygen by lighting a match,
letting it burn for a moment, blowing it out, and quickly placing the glowing match into the pile of foam. It should
relight if there is not too much water in the foam.
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Exploding Baggies – Temperature vs. Time
In this demonstration, a gas is created when vinegar and baking soda are mixed, causing the baggie to explode.
The temperature of the reaction will be related to the time it takes for this to happen. You will need:
 Three small zip-lock “freezer” bags
 Vinegar
 Baking soda
 Three tissues
NOTE: you should do this demonstration outside.
1. Bring three 150-mL vinegar samples to different temperatures: room temperature, warmed gently, and quite
warm (but not too hot to touch).
2. Create three baking soda “packets”: For each packet, put two large spoonfuls of baking soda into the
middle of a tissue and wrap by folding the tissue around the powder.
3. Partially zip each of the bags closed, leaving enough space to add the baking soda packet.
4. While one student marks the time, have three students simultaneously add the tissue with the baking soda into
the bag and quickly zip the bag completely closed.
5. Put the bag down on the ground (outside) and step back. Note the time it takes for the bag to go… POP!
Class Discussion Questions:
What is the reaction that is occurring in the baggie?
At the molecular level, how is a warmer solution different than a cooler solution? Why would this influence how
fast the reaction occurs? What is the connection to collision theory and activation energy?
The Chemistry of the Reaction
Vinegar is acetic acid (CH3COOH) and baking soda is sodium hydrogen carbonate (NaHCO3). The reaction
produces water, sodium acetate, and carbon dioxide. The production of carbon dioxide over time causes
pressure to build within the baggie.
Exploring Le Chatelier’s Principle using a Copper Chloride Solution
In this experiment, students prepare a reversible (equilibrium) reaction and explore how the equilibrium shifts when
changes are made to the reaction solution by observing color changes. You will need, per lab group:

100 mL of sodium chloride (or 100 grams)

Boiling water

200 mL of copper(II) sulfate (or 200 grams)

Ice water

Distilled water

4 medicine cups

Concentrated hydrochloric acid

Optional: 0.10 M silver nitrate solution
Prepare the equilibrium solution to be distributed to the students.
1. Place the solid sodium chloride and solid copper(II) sulfate into a large glass jar. Add 500 mL of distilled water.
2. Place the jar into a shallow pan filled with boiling water and stir while the solution heats. If the jar has a lid,
periodically shake the solution. The solids should dissolve and the solution should turn a bright green color.
NOTE: If solids remain in the solution, allow to sit for a few minutes and carefully pour off the clear solution.
Pre-Lab Discussion Questions:
Show the students the balanced equilibrium equation for the reaction, indicating the colors of the complex ions,
and have them answer the following questions either in groups or as a class.
 Does the constant color mean the reaction has stopped?
 What does the green color mean in terms of the relative concentrations of the copper-water ion of the
products and the copper-chloride ion of the reactants? (i.e. which one is more prevalent?)
Investigating the effect of reactant and product concentration on equilibrium
Each group should be provided with about 100 mL of the equilibrium solution in a small bottle.
1. Half-fill four small medicine cups with the equilibrium solution.
2. To the first cup slowly add five drops of water, one drop at a time and record your observations.
3. To the second cup add enough water so that its color matches that now in the first cup. Then slowly add
drops of concentrated (12 M) hydrochloric acid, HCl, and stir. Record your observations.
4. To the third cup, slowly add about 10 drops of 0.10 M silver nitrate solution, if available. (You can also do this
as a thought exercise.) Record your observations.
5. To the fourth cup, add enough isopropyl alcohol (C3H7OH) to double the volume of the solution.
Investigating the effect of temperature on equilibrium
1. Place the original bottle of solution in an ice bath and wait for three minutes. Record your observations.
2. Now place the bottle in a pan of boiling water and wait for three minutes. Record your observations.
Class Discussion Questions:
 What does each color change indicate about the direction of the "shift" in equilibrium position?
 How does LeChatelier's principle explain why your system responded as it did?
The Chemistry of the Reaction
The equilibrium present in the solution is best described as: CuCl42-(aq) + 4 H2O ⇄ Cu(H2O)42+(aq) + 4 ClThe copper-chloride ion is green and the copper-water ion is light blue. The reaction is exothermic in the forward
direction.
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Milk of Magnesia Color Changes
In this demonstration, vinegar is added in small amounts to milk of magnesia (a suspension of Mg(OH) 2) that
contains cabbage juice indicator. Color changes continue to occur as the limited amount of acid is slowly
consumed, and then when an excess of acid is finally attained. You will need:

Milk of magnesia

Spoon for stirring

Vinegar

Cabbage juice indicator
1. Place 100 mL of Milk of Magnesia in a large glass jar and dilute with about 150 mL of tap water.
2. Add about 20 mL of cabbage juice indicator: red in acidic solutions and blue in alkaline solutions.
3. Have a student continually stir the solution, which is light blue, indicating that it is slightly basic.
4. While stirring add 10-20 mL of vinegar. The color changes immediately to red and then slowly back to blue.
5. Continue to add small amounts of vinegar until the color change back to blue does not occur.
The chemistry of the reaction
Initially, the acid neutralizes the small amount of hydroxide ion from the Mg(0H)2 that has dissolved, then turns the
solution acidic. However, as more of the Mg(0H)2 from the suspensions gradually dissolves into solution, the acid is
neutralized and eventually the solution becomes basic again. When all of the hydroxide has been neutralized,
the acid is in excess and the indicator will remain red.
Calculating the Molar Mass of Oxalic Acid via Syringe Titration
In this experiment, students prepare a standard solution of sodium hydroxide, and use this solution in a titration to
experimentally determine the molar mass of oxalic acid. The limiting factor here is the necessity of a balance that
can be used to measure masses to the 10th of a gram. You will need:

0.1-g balance

35-mL syringe

Solid sodium hydroxide (lye)

Medicine cups

Solid oxalic acid

Distilled water

Curry powder indicator
1. Create 30 mL of a 1.0-M NaOH solution. [Have students perform the calculation and discuss the proper
procedure: Place 1.2 grams NaOH into a medicine cup, dissolve in a small amount of distilled water, then fill to
the 30-mL mark.]
2. In a clear cup, place 0.5 grams of oxalic acid. Add just enough water to dissolve it.
3. Add a good pinch of curry powder indicator to the acid solution and stir. The indicator is yellow in acid, red in
base.
4. Fill the syringe with 25 mL of the NaOH solution, being careful to release any air bubbles.
5. With continual stirring, slowly add the NaOH solution, drop-by-drop, into the acid solution. Continue until the
solutions remains a red color after stirring. Record the volume of NaOH added to the acid.
Analysis and Discussion
The students have collected enough data to determine, experimentally, the molar mass of oxalic acid. They need
to be told that oxalic acid is a diprotic acid (if they do not already know the formula). They can evaluate the
accuracy of their titration by looking up the formula and using a periodic table to determine the true value. They
can then evaluate the accuracy of their work by calculating a percent yield.
The chemistry of the reaction
Oxalic acic, H2C2O2, is neutralized by the sodium hydroxide solution. When the solution “just” turns pink (and
remains pink) the end point of the titration has been reached. It is assumed that this is the point of equivalence…
when the moles of hydroxide have been matched by the moles of hydrogen ions released by the acid.
Create a Cola Can Battery
In this activity, each student group prepares one “cola” galvanic cell battery. When the batteries are connected
in series, they should create enough current to light an LED bulb. You will need, per lab group

A can of cola

Two pieces of copper wire (preferably with clips)

A copper coin

Steel wool

Can opener
1. Open your can of cola and pour the contents into a cup and set it aside.
2. Using a can opener cut off the top of the can. Clean and dry the inside of the can.
3. To remove the protective plastic coating on the inside of the can, scrub with steel wool.
4. Pour the contents from the cup back into the can.
5. Take one of your wires and wrap it around (or clip it to) the copper coin, then dip it into the drink. Attach
another wire to the top of the cola can.
6. When the two wires are connected, they produce a small current.
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The chemistry of the reaction
Solid aluminum is oxidized to Al3+ (Eºcell = -1.66) at the Al electrode, and water is reduced to H2(g) and OH- (Eºcell = 0.83) at the Cu electrode. The wire allows electrons to flow between these two processes. The phosphoric acid in
cola allows for charge neutrality
Analysis and Discussion:
Sketch a schematic diagram of the cell, showing the half reactions and the direction of electron flow through the
wire. Identify which metal is the anode and which is the cathode. Discuss what will happen over time as the
battery continues to supply current. Why will it go “dead?”
ALTERNATIVE: A Salt-Water Battery
Create a battery in a plastic cup by filling the cup with water, adding plenty of salt, and hanging a zinc-coated
nail from one side and a strip of aluminum (perhaps from a soda can) on the other side. Connect the metals with
a wire, and the battery is complete. As for the Cola Can battery, you would need to connect batteries in series to
light an LED bulb.
Electrolysis of Water
In this demonstration, electricity from a battery is used to force a non-spontaneous redox reaction. If at all
possible, don't make this a demonstration. Allow student groups to make their own. For each apparatus you will
need:

9 volt battery

Clear tape

Small rubber band

Foam meat tray

2 test tubes (optional)

2 short pencils,

Distilled water

Match to test for hydrogen gas
sharpened on both

Eyedropper

Diluted blue food coloring
ends

1 teaspoon salt

One clear plastic cup
1. Hal-fill the cup with distilled water and add a spoonful of salt. Stir till dissolved.
2. Cut the foam to create rectangular piece with a curved end. Poke the pencils through the curved end of
the foam holder so they are the same distance apart as the terminals of the 9 volt battery.
3. Set foam holder/pencil assembly in the plastic cup. Balance the 9 volt battery on the pencil leads and put a
small rubber band around the battery to hold it to the foam holder (rubber band should not be tight, so that
gravity will hold battery terminals against pencil leads.) IMPORTANT: The battery must be above water level.
4. If you do not see bubbles rising make sure the battery terminals are resting on the pencil leads.
5. Test the solution: Add blue coloring to the solution and observe. Add an acid/base indicator and observe.
6. Test the gases. This is tricky. Fill a test tube with water and, while holding your finger over the opening, turn it
upside down in the cup, releasing your finger when the test tube top is under water. Position the test tubes
over the pencil tips to catch the gas bubbles. When they are half or more full you can test the solutions.
Analysis and Discussion:
 What do the tests indicate about the identity of the gases?
 What is the net reaction occurring in this process?
The chemistry of the reaction
Electrolysis is the process of using electricity to force a non-spontaneous chemical reaction. Graphite in pencil
“leads” is a good conductor of electricity from the battery. Since water is a relatively poor conductor of
electricity, sodium chloride is added to the water. However, this causes an unexpected reaction to occur at the
anode. Instead of oxidizing water to O2 gas, the chloride ion in the salt solution is oxidized to chlorine gas. At the
cathode, the expected reduction of water to hydrogen gas and hydroxide ions is formed. The hydrogen gas is
flammable and a small amount will “pop” if collected in a small test tube. The chlorine gas dissolved in water will
“bleach” the blue food coloring. The production of hydroxide ions at the cathode will cause an indicator to
change color.
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Calculations Involving Equilibrium Expressions
1.
2.
3.
4.
The following reaction has the equilibrium pressures listed. What is the value of K at this temperature?
 CS2(g) + 4 H2(g) ⇄ CH4(g) + 2 H2S(g)
 PCH4 = 0.131 atm, PH2S = 0.084, PCS2 = 0.428 atm, PH2 = 0.921 atm
 Answer: K = 0.00300
The deadly gas phosgene is formed from CO and Cl2 according to the reaction below. Use the information
provided to determine K for this reaction at at 600oC.
 CO(g) + Cl2(g) ⇄ COCl2(g)
 The initial pressures of of the gases are PCO = 0.60 atm, PCl2 = 1.10 atm and PCOCl2 = 0 atm.
 The final pressure of COCl2 is 0.10 atm at 600oC.
 Answer: K = 0.20
The following reaction has a K value of 0.64 at 900oC. Use the information below to calculate the equilibrium
partial pressures of the all of the gases in the reaction.

CO2(g) + H2(g) ⇄ CO(g) + H2O(g)

The initial partial pressures are PCO2 = 1.00 atm, PH2 = 1.00 atm, PCO = 0.00 atm and PH2O = 0.00 atm
 Answer: PCO2 = 0.56 atm; PH2 = 0.56 atm; PCO = 0.44 atm; PH2O = 0.44 atm
The reaction below has a K value of 1.8 at room temperature. Which direction will the reaction proceed (will it
produce more reactants or more products) under the given situation below?
 2 NO(g) + O2(g) ⇄ 2 NO2(g)
 The initial pressures of all species are 0.50 atm.
 Answer: the reaction will move to the left
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