topic17

School Science Lessons
Topic 17 Catalysis, chemical equilibrium, enzymes and biological catalysts, gravimetric analysis, rates of reaction
2012-01-25 SP
Please send comments to: J.Elfick@uq.edu.au
See: History of this document

Table of contents
17.3.0 Catalysis
17.5.0 Chemical equilibrium
17.4.0 Enzymes and biological catalysts
17.2.0 Factors affecting rates of reaction
17.6.0 Gravimetric analysis
17.1.0 Rates of reaction
17.1.1 Rates of reaction, clock reactions

17.3.0 Catalysis
17.3.0 Catalysis, catalysts
13.6.6.1 Catalytic oxidation of ammonia forms nitrogen monoxide, with red-hot platinum wire
13.6.6.2 Catalytic oxidation of ammonia, with chromium (III) oxide catalyst
17.3.13 Bromine catalyses the oxidation of sulfur to sulfuric acid
17.3.3 Ethanedioic acid-2-water (oxalic acid) with potassium manganate (VII), autocatalysis
17.3.15 Ethyl acetate with sodium hydroxide, autocatalytic hydrolysis
17.3.11 Heat potassium chlorate, manganese dioxide catalyst
17.3.1 Hydrogen peroxide decomposition, with different catalysts
17.3.2 Hydrogen peroxide decomposition, with manganese (IV) oxide catalyst
17.3.8 Hydrogen peroxide with catalase enzyme in raw beef liver
17.3.6 Hydrogen peroxide with potassium sodium tartrate, cobalt (II) chloride catalyst, visible activated complex
17.3.9 Hydrogen peroxide with sodium thiosulfate, ammonium molybdate catalyst
17.3.12 Methanol with platinum wire catalyst, catalytic oxidation
17.3.10 Oxidation of acetone vapour, copper oxides catalyst
17.3.7 Potassium bromate with propanedioic acid, double autocatalytic reaction, oscillating reaction
17.3.5 Potassium iodide with hydrogen peroxide, reverse colour change
17.3.14 Sodium hypochlorite decomposition, cobalt sulfate catalyst
17.3.4 Sugar with potassium chlorate, spontaneous combustion

17.5.0 Chemical equilibrium
17.5.0 Chemical equilibrium, the law of chemical equilibrium (law of mass action)
17.5.5.3 Common ion effect to precipitate sodium chloride from solution
17.5.5.4 Common ion effect to precipitate barium chloride from solution
17.5.5.5 Effect of temperature on chemical equilibrium, thermal dissociation of ammonium chloride
17.5.7.0 Explanation of group analysis
17.5.6.2 Heat nitrogen tetroxide (dinitrogen tetroxide, N₂O₄)
17.5.5.2 Hydrolysis of antimony chloride
17.5.5.1 Hydrolysis of bismuth chloride
17.5.5.0 Law of mass action and reversible reactions, effect of alteration of concentration

17.4.0 Enzymes and biological catalysts
17.4.1 Breakdown starch to sugar
17.4.3 Bromelain enzyme from pineapples
17.4.01 Commercially available enzymes
17.4.2 Fermentation using yeast

17.6.0 Gravimetric analysis
17.6.2 Weight of aluminium in aluminium sulfate
17.6.3 Weight of calcium in marble, calcium carbonate
17.6.1 Weight of iron in iron (II) ammonium sulfate
17.6.4 Weight of magnesium in magnesium sulfate
17.6.5 Weight of sulfate radical in sodium sulfate
17.6.6 Weight of tin in solder

17.1.0 Rates of reaction
17.1.0 Rates of reaction
3.7.2 Rates of reaction
17.1.4 Balloons, dilute hydrochloric acid with marble chips
3.94 Catalysts and rate of reaction
3.7.1 Concentration and rate of reaction
3.92 Concentration and rate of reaction, sodium thiosulfate with hydrochloric acid
17.1.1 Count bubbles, dilute hydrochloric acid with granulated zinc
17.2.0 Factors affecting rate of reaction
17.1.3 Gas burette, dilute hydrochloric acid with marble chips (calcium carbonate)
17.1.5 Hydrogen peroxide with manganese (IV) oxide, rate of reaction, height of suds
17.5.4.0 Rate of reaction depends on concentration
3.91 Size of particles and rate of reaction
3.93 Temperature and rate of a reaction
17.1.2 Volume of gas, dilute hydrochloric acid with zinc

17.1.1 Rates of reaction, clock reactions
17.1.8 Hydrogen peroxide clock reaction, Briggs-Rauscher oscillating reaction, hydrogen peroxide, potassium iodate
17.1.7 Old Nassau flag clock reaction, (orange and black) sodium metabisulfite, mercury (II) iodide
17.1.6 Iodine clock reaction, hydrogen peroxide, potassium iodide
17.1.9 Iodate clock reaction, potassium iodate, sodium metabisulfite (sodium bisulfite)
17.1.10 Persulfate-iodide clock reaction

17.5.4.0 Rate of reaction depends on concentration
17.5.4.1 Magnesium with hydrochloric acid reaction
17.5.4.2 Potassium iodide with potassium iodate

17.2.0 Factors affecting rates of reaction
17.2.2.1 Concentration of reactants, hydrochloric acid with magnesium
17.2.2 Concentration of reactants, hydrochloric acid with sodium thiosulfate (hypo)
17.2.1 Particle size, dilute hydrochloric acid with marble chips
17.2.4 Rates of reaction of aspirin
17.2.3 Temperature and rate of reaction, sulfuric acid with iron

17.1.0 Measure rates of reaction
Rates of reactions for laboratory experiments should not be so fast that an explosion occurs. Also, they should not be so slow that you cannot observe or measure any change in reasonable time. To measure the speed of a reaction you must measure some change, e.g. time taken for colour change, pH change, precipitate appears, reagents disappear, temperature change, volume of gas that forms weight of reactants used. Chemical reactions occur only if the particles of reactants can collide with sufficient energy, the activation energy, E_a. Reactions will become faster if the number of collisions increases and if the energy of collision increases. You can increase the rate of reaction by increasing concentration of reactants, light energy, particle size, pressure, temperature, and with catalysts.

17.1.1 Count bubbles, dilute hydrochloric acid with zinc
See diagram 17.1.1: Count bubbles, dilute hydrochloric acid with zinc
Put a piece of granulated zinc in dilute hydrochloric acid in a test-tube. Count the number of bubbles of hydrogen gas that reach the surface of the solution every 30 seconds. Draw a graph by plotting the number of bubbles along the vertical axis and time along the horizontal axis. The reaction starts quickly then becomes slower until it stops.

17.1.2 Volume of gas, dilute hydrochloric acid with zinc
See diagram 17.1.2: Volume of gas, dilute hydrochloric acid with zinc
Repeat the experiment with a flask connected to a gas syringe to measure how much hydrogen gas forms in the reaction.

17.1.3 Gas burette, dilute hydrochloric acid with marble chips
See diagram 17.3.3: Gas burette | See: Saturation vapour pressure over water
Add zinc or marble chips to dilute hydrochloric acid. Collect the gas in a burette inverted over water. Compare how much gas forms in unit time for each size of marble chips or zinc Weigh the container and note the loss in mass every half minute while the reaction goes on.

17.1.4 Balloons, dilute hydrochloric acid with marble chips
See diagram 3.2.91: Balloons, dilute hydrochloric acid with marble chips
Add marble chips to dilute hydrochloric acid. Show the relative production of hydrogen gas by attaching a previously stretched balloon to each test-tube.

17.1.5 Hydrogen peroxide with manganese (IV) oxide (manganese dioxide), height of suds
Manganese (IV) oxide does not take part in the chemical reaction. Its function is to provide an increased surface area for the reactants. Pour 10 mL of hydrogen peroxide, 5 mL of detergent and 5 mL of glycerine into two identical measuring cylinders, Measuring cylinder 1. and Measuring cylinder 2. Stir both solutions.
Measuring cylinder 1. Add crystals of manganese (IV) oxide (MnO₂, manganese dioxide) and stir again. A mass of bubbles arises to form suds. The glycerine increases the surface tension of the liquid to delay the collapsing of the bubbles. The hydrogen peroxide decomposes into water and the oxygen that forms the bubbles. The suds are higher in the measuring cylinder containing the manganese (IV) oxide because it acts as a catalyst.
Measuring cylinder 2. Control without manganese (IV) oxide.
Repeat the experiment with dust or dirt or ashes instead of manganese (IV) oxide. The heights of the detergent suds are similar to the heights using manganese (IV) oxide. This observation shows that manganese (IV) oxide does not take part in the chemical reaction. Its function is to provide an increased surface area for the reactants.

17.1.6 Iodine clock reaction, hydrogen peroxide, potassium iodide
Order online: Iodine Clock Reaction, sodium salts, citric acid, starch, MSDS, eco-foam
1. Solution A: Add drops of water to 0.2 g soluble starch, pour this paste into a beaker containing 1 cc boiling water and stir. Pour this solution into a 1 litre beaker and dilute to 800 mL. Add 30 mL glacial ethanoic acid (glacial acetic acid, CH₃CO₂H) 4.1 g sodium ethanoate (sodium acetate, CH₃CO₂Na) 50 g potassium iodide (KI) and 9.4 g sodium thiosulfate-5-water (Na₂S₂O₃.5H₂O). Dissolve all solutes by stirring and when cooled to room temperature make up solution to 1 litre. The solution is slightly cloudy.
Solution B: Dilute 500 mL 20 vols hydrogen peroxide to 1 litre.
Put 100 mL of each solution in separate beakers. Pour the contents of one beaker into another and stir quickly or use a magnetic stirrer. After about 20 seconds the colourless mixture suddenly turns dark blue. Repeat the experiment at room temperature but note the time taken for the colour change after mixing.
2. Repeat the experiment at 10^oC above room temperature. Note the time taken for the colour change after mixing, about double the time compared to 1.
3. Repeat the experiment at 10^oC below room temperature. Note the time taken for the colour change after mixing, about half the time compared to 1.
4. Repeat the experiment at room temperature using half the concentration of solution B, i.e. 50 mL solution B + 50 mL water. Note the time taken for the colour change after mixing, about double the time compared to (1.) because the reaction rate is halved. The amount of hydrogen peroxide has been halved.
5. Repeat the experiment at room temperature using half the concentration of solution A, i.e. 50 mL solution A + 50 mL water. Note the time taken for the colour change after mixing, the same as for (1.). The reaction rate has been halved but the amount of thiosulfate also has been halved so you only need to make half the amount of iodine to combine with all the thiosulfate. If the sodium thiosulfate was not in solution A but in a solution C, the result of 5. would be the same as in 4.
H₂O₂ (aq) + 2I^- (aq) + 2H⁺ (aq) --> I₂ (aq) + 2H₂O (l)
Hydrogen peroxide reacts with iodide ions to form iodine. The ethanoic acrid with sodium ethanoate buffer the reaction
I₂ (aq) + 2S₂O₃^2- (aq) --> 2I^- (aq) + S₄O₆^2- (aq) +
The iodine reacts with thiosulfate to form tetrathionate ions and iodide ions return to the solution. As soon as all the thiosulfate is converted to tetrathionate ions the remaining iodine reacts with the starch solution to form a blue black colour.

Hydrogen peroxide with potassium iodide, iodine clock reaction
Second version
Solution A Prepare a solution of potassium iodide, sodium thiosulfate and starch
Solution B Add sulfuric acid to hydrogen peroxide solution
Add solution A to solution B
H₂O₂ (aq) + 3I^- (aq) + 2H⁺ (aq) --> I₃^- (aq) + 2H₂O (aq) slow reaction
hydrogen peroxide + iodide ion + hydrogen ion --> triiodide ion + water
I₃^- (aq) + 2S₂O₃^2- (aq) --> 3I^- (aq) + S₄O₆^2- (aq)

17.1.7 Old Nassau flag clock reaction
This experiment should NOT be done in schools because it uses mercuric (II) chloride, HgCl₂!
1. Sodium metabisulfite reacts with water to form sodium hydrogen sulfite, colourless reaction
Na₂S₂O₅ (aq) + H₂O (l) --> 2NaHSO₃ (aq)
2. Hydrogen sulfite ions reduce iodate (V) ions to iodide ions, colourless reaction
IO₃^- (aq) + 3HSO₃^- (aq) --> I^- (aq) + 3SO₄^2- (aq) + 3H⁺ (aq)
3. With excess iodide ions, mercury (II) iodide, HgI₂, forms as an orange precipitate
Hg²⁺ (aq) + 2I^- (aq) --> HgI₂ (s)
4. When all the mercury is used up reacting with iodine, the remaining iodate and iodide ions react to form iodine and a blue black iodine starch complex forms.
IO₃^- (aq) + 5I^- (aq) + 6H⁺ (aq) --> 3I₂ (aq) + 3H₂O (l)
The orange precipitate and the black iodine complex are the colours of Princeton University, sometimes referred to as "Old Nassau".

17.1.8 Hydrogen peroxide clock reaction, Briggs-Rauscher oscillating reaction
Order online: Oscillation Reaction, Briggs-Rauscher Reaction, catalysts, equilibrium
1. Solution A. Add 43 g potassium iodate, (KIO₃) + 5 mL sulfuric acid (H₂SO₄), to 800 mL deionized water, stir to dissolve then dilute to 1 litre.
2. Solution B. Add 15.6 g malonic acid (propanedioic acid, CH₂(COOH)₂) + 3.4 g manganese sulfate monohydrate, (MnSO₄.H₂O), to 800 mL deionized water + 4 g of freshly made soluble starch starch, stir to dissolve then dilute to 1 litre.
3. Solution C. Dilute 400 mL 30% hydrogen peroxide, (H₂O₂) to 1 litre.
4. Mix thoroughly 300 mL of solutions A. and B., then add 300 mL of solution C. The 3 mixed colourless solutions oscillate through colourless --> amber --> blue --> colourless, for about 5 minutes then stays blue-black
Hydrogen peroxide + iodate --> iodine + oxygen (auto-catalytic) (blue)
IO₃^- + 2H₂O₂ + CH₂(COOH)₂ + H⁺ --> ICH(COOH)₂ + 2O₂ + 3H₂O

17.1.9 Iodate clock reaction
Concentration and temperature affects rate of reaction, potassium iodate, sodium metabisulfite (sodium bisulfite)
Solution A. Dissolve 4.3 g potassium iodate, KIO₃, in on litre of water
Solution B. Prepare a starch solution by dissolving 5 g of starch in 800 mL hot water, boil, then leave to cool. Add 0.2 g sodium metabisulfite. Add 5 cc of M sulfuric acid. Dilute to 1 litre.
1. Note the time for colour change, colourless to dark blue, by adding 50 mL of solution A in beaker A to so 50 mL of solution B in beaker B. Pour the solution back and forth, B --> A --> B --> A. The reaction may take 5 to 6 minutes. The reaction is designed to be a slow colour change reaction, a clock reaction.
2. Repeat the experiment with half concentration of solution A
3. Repeat the experiment with solution A warmed in a water bath to 35^oC.
Increase of concentration increase the rate of reaction because of increased chance of particles reacting
Increased temperature increases th rate of reaction because of increase of speed on particles increases chance of particles reacting.
IO₃^-(aq) + 3HSO₃^- (aq) --> I^- (aq) + 3HSO₄^-(aq)
iodate ion + bisulfite ion --> iodide ion + bisulfate ion
IO^3- (aq) + 5I^- (aq) + 6H⁺ (aq) --> 3I₂ + 3H₂O (l) iodate ion oxidizes the iodide ion to iodine
iodate ion + iodide ion + hydrogen ion --> iodine + water
However, the iodine is reduced immediately back to iodide by the bisulfite:
I₂ (aq) + HSO₃^- (aq) + H₂O (l) --> 2I^- (aq) + HSO₄^-(aq) + 2H⁺ (aq) bisulfite reduces iodine back to iodide ion
iodine + bisulfite ion + water --> iodide ion + bisulfate ion + hydrogen ion
When no more bisulfate ion exists because it is all used up reducing the iodine, the remaining excess iodine forms blue-black colour with starch.

17.1.10 Persulfate-iodide clock reaction
Use sodium persulfate (sodium peroxodisulfate) Na₂S₂O₈, or potassium persulfate (potassium peroxydisulfate or KPS) K₂S₂O₈, or ammonium persulfate (ammonium peroxydisulfate), (NH₄)₂S₂O₈.
The reaction mixture remains colourless for several minutes after the reactants are mixed.
1. Solution A. Prepare 10 mL of 0.1 M potassium iodide solution + 5 mL of 0.01 M sodium thiosulfate solution
Solution B. Prepare 10 mL of 0.1 M ammonium persulfate solution + 1 drop of starch solution
Pour solution B into solution A. Pour the solution back and forth, A --> B --> A --> B
Record the time when the starch-I2 complex forms. Record the temperature of the solution
2. Solution A. Prepare 10 mL of 0.1 M potassium iodide solution + 5 mL of 0.01 M sodium thiosulfate solution
Solution B. Prepare 5 mL of 0.1 M ammonium persulfate solution + 1 drop of starch solution. Add 5 mL of water to keep the volume constant
Pour solution B into solution A. Pour the solution back and forth, A --> B --> A --> B
Record the time when the starch-I2 complex forms. Record the temperature of the solution
S₂O₈^2- (aq) + 2I^- (aq) --> 2SO₄^2- (aq) + I₂ (aq) Slow reaction, but the iodine formed is soon changed to iodine in the next reaction
I₂ (aq) + 2S₂O₃^2- (aq) --> 2I^- + S₄O₆^- (aq) Fast reaction
I^- (aq) + starch --> blue-black complex, When all the thiosulfate ion is used up, the concentration of I₂ increases and forms a blue complex with starch,
The rate of reaction is the rate of consumption of the S₂O₈^2- ion.

17.2.1 Particle size, dilute hydrochloric acid with marble chips (calcium carbonate)
The smallest particles show the most vigorous reaction because if you grind two pieces of substance with the same weight into coarse and fine sizes, the fine size pieces would have total surface area greater than the coarse pieces. BE CAREFUL! Hydrogen gas forms.
Break marble chips or granulated zinc into four sizes with a hammer Put 2 g of each into four test-tubes: Test-tube 1, original size as control, Test-tube 2, rice grain size, Test-tube 3, half rice grain size,
Test-tube 4, coarse powder. Add the same volume of 2 M hydrochloric acid to each test-tube. Compare the reactions. Rate of reaction of Test-tube 4 > Test-tube 3 > Test-tube 2 > Test-tube 1.

17.2.2 Concentration of reactants, hydrochloric acid with sodium thiosulfate
See diagram 17.4.1: Measure the cloudiness in the solution
1. The reaction slowly forms sulfur that makes the solution cloudy. Measure the rate of reaction by measuring the cloudiness in the solution. Note the time until you can no longer observe a black cross on a piece of paper below the beaker. Change the concentration of the sodium thiosulfate and keep the concentration of hydrochloric acid constant. Dissolve 20 g sodium thiosulfate in 500 mL of water. Pour 50 mL of this solution into a 100 mL beaker. Put the beaker on a sheet of paper marked with a black cross. Add 5 mL of 2 M acid and stir the acid into the solution. Record the time when the cross is no longer visible through the precipitated sulfur in the solution.

2. Repeat the experiment with a smaller concentration of thiosulfate as follows:

.	Sodium thiosulfate Solution	deionized water
Beaker 1	30 mL	20 mL
Beaker 2	20 mL	30 mL
Beaker 3	10 mL	40 mL

Stir the solution, then add 5 mL of 2 M acid as before. The time for the cross to become invisible is greater. The reaction takes longer as the concentration decreases. Draw a graph to plot concentration of the thiosulfate solution against time taken for the reaction. Express concentration values as the volume of the original thiosulfate solution used, e.g. 50 mL, 30 mL, 20 mL and 10 mL.
Na₂S₂O₃ (aq) + 2HCl (aq) --> H₂O (l) + SO₂ (g) + S (s)

17.2.2.1 Concentration of reactants, hydrochloric acid with magnesium
Pour 50 mL 6 M hydrochloric acid solution into a 100 mL standard flask. Make up the solution to the mark and mix it well. This is Solution B, 3 M. Pipette 50 mL Solution B into another 100 mL standard flask. Make up this solution to the mark and mix it well. This is Solution C, 1.5 M. Pour the remaining solution into a 100 mL beaker. Pipette 50 mL solution C into another 100 mL standard flask and make this solution up to the mark and mix it well. Label this solution D, 0.75 M. Pour the remaining solution into a 100 mL beaker. Pipette 50 mL of solution D into a 100 mL beaker. Cut off four 0.5 cm pieces of magnesium ribbon. Put one strip into solution A and record the time of the reaction until you have dissolved all of the magnesium. Repeat with a new magnesium strip for each of solutions B, C, and D. 9. Graph the results with time as the vertical axis and concentration as the horizontal axis. Note whether a linear relationship exists between acid concentration and the rate at which magnesium dissolves.
Mg + 2HCl --> MgCl₂ + H₂

17.2.3 Temperature and rate of reaction, sulfuric acid with iron
Increase in temperature of 10^oC usually doubles the rate of reaction. The rate increases because collisions between particles are more frequent and with more energy. Add enough potassium manganate (VII) solution to dilute sulfuric acid to make it pink. Put in a nail and record the time for the solution to lose the pink colour at room temperature. Repeat the experiment at 10^oC above room temperature. Repeat the experiment at 20^oC above room temperature. Draw a graph to show time taken to decolorize against temperature. The rate of reaction may double for each 10^oC rise in temperature.

17.2.4 Rates of reaction of aspirin
Be careful! Aspirin is a drug so do not let the students take them!
Soluble aspirins dissolve in water with a well defined end point and constants interaction time. Use these tablets to investigate the effects of temperature, stirring, crushing and varying the water volume on reaction time. Investigate the temperature change required to halve the reaction time between one tablet and known quantity of water at room temperature. Investigate the number of tablets or fractions of tablets required to exactly double the reaction time in a known quantity of water at room temperature.

17.3.0 Catalysis
Catalysts increase the rate of reactions without themselves being chemically changed. A catalyst can change the rate of a chemical reaction without itself being permanently changed. They provide an alternative pathway for the reactions and so decrease the activation energy needed. Substances that slow the rate of reactions are called inhibitors.

17.3.1 Hydrogen peroxide decomposition, with different catalysts
1. Some substances can increase the rate of reaction for the decomposition of hydrogen peroxide. Put two small equal amounts of hydrogen peroxide solution gently into four test-tubes. Note the bubbles. Tests for oxygen with a glowing splint. Add the following substances: Test-tube 1, Nothing added (control) Test-tube 2, Manganese (IV) oxide, Test-tube 3, Iron (III) chloride. Test-tube 4, Copper (II) sulfate. Note the bubbles. Tests for oxygen with a glowing splint.
2. Decomposition of hydrogen peroxide by heat. Heat drops of hydrogen peroxide in a test-tube. When effervescence begins, put a glowing wood spill into the test-tube. The spill relights, showing that oxygen is being produced. Observe the rate at which hydrogen peroxide decomposes by the addition of acid or alkali. Put 2 cm of hydrogen peroxide into each of two test-tubes. To test-tube 1, add drops of a solution of sodium hydroxide or strong ammonia solution. To test-tube 2, add sulfuric acid or citric acid. Put both test-tubes into a beaker of hot water for one minute. Test both test-tubes for oxygen. Oxygen forms from test-tube 1 containing the alkali but very little from the other test-tube. Alkalis speed up the formation of oxygen, while acids slow it down. For this reason acid is often added to commercial hydrogen peroxide to help it to keep its oxygen. A catalyst is a substance which helps a chemical action without being changed itself. A substance which hinders instead of helping a chemical action is called a “negative” catalyst. Thus we say that sulfuric acid is a negative catalyst for the decomposition of hydrogen peroxide. Sodium hydroxide and ammonia are “positive” catalysts.

17.3.2 Hydrogen peroxide decomposition, with manganese (IV) oxide catalyst
1. Fill a test-tube with hydrogen peroxide to a depth of about 1 cm and add a little of manganese dioxide. Tests for oxygen with a glowing splint. A rapid evolution of oxygen occurs and the manganese dioxide is not lost.
2H₂O₂ (l) –> 2H₂O (l) + O₂ (g)
2. Put 2 cm of 20 vols hydrogen peroxide (20 vol. solution) in two test-tubes, Test-tube 1 and Test-tube 2.
Test-tube 1. Add 5 drops of sodium hydroxide solution.
Test-tube 2. Add 5 drops of dilute sulfuric acid solution.
Immerse both test-tubes in a beaker half full of hot water. Use a glowing splint to tests for oxygen in the mouth of the test-tubes. Oxygen is found in Test-tube 1 but not in Test-tube 2.
3. Set up a conical flask fitted with a one hole stopper and delivery tube that leads into a beaker of water. Invert a closed burette full of water over the end of the delivery tube. Pour 50 mL of water in the flask and add 2 mL of 20 vols (6%) hydrogen peroxide solution. Add 1 g manganese (IV) oxide and immediately insert the stopper with the delivery tube into the flask. Note how much oxygen forms every 15 seconds. Plot on a graph how much oxygen forms every 15 seconds against the time of the reaction.
4. Repeat the experiment by adding more hydrogen peroxide solution to the same test-tube. The manganese (IV) oxide is not "used up" because more oxygen forms. Repeat the experiment with 1 g copper (II) oxide. Add 2 mL of 20 vols (6%) hydrogen peroxide solution.
5. Repeat the experiment with 1 g zinc oxide. Add 2 mL of 20 vols (6%) hydrogen peroxide solution. Plot a graph for each experiment. Manganese (IV) oxide is the better catalyst in these reactions. Warm some hydrogen peroxide solution gently in a test-tube and hold a glowing splinter of wood in the mouth of the test-tube.
6. Place two small equal amounts of hydrogen peroxide in separate test-tubes. Add some iron chloride to one test-tube and a little manganese dioxide to the other. Apply the glowing splinter test to any gas given off. Hydrogen peroxide decomposes to give off oxygen when heated but will decompose without heating when iron chloride and manganese dioxide are added to it. The black colour of manganese dioxide and the brown colour of iron chloride remains after the reactions, so these chemicals may not have been altered during the reactions. Iron chloride and manganese dioxide are catalysts or "chemical accelerators".

17.3.3 Ethanedioic acid-2-water (oxalic acid) with potassium manganate (VII), autocatalysis
Solution A: 6 g of ethanedioic acid-2-water (oxalic acid) in 300 mL of water.
Solution B: 100 mL of 0.001 M potassium manganate (VII) solution, (potassium permanganate).
Pour 150 mL of Solution A into two beakers, then add 5 mL of concentrated sulfuric acid to each beaker. Add 50 mL of the potassium manganate (VII) solution to each beaker. Add a small crystal of manganese (II) chloride, MnCl_2, to one beaker then stir the solutions in both beakers The solution containing the crystal of manganese (II) chloride starts to lose colour and becomes colourless in about a minute. The other solution does not change in colour for two or three minutes but, when sufficient Mn²⁺ ions are present, it starts to become colourless. The Mn²⁺ autocatalyses the solution as follows.
2MnO₄^- (aq) + 5C₂O₄²⁺ (aq) + 16H₃O⁺ (aq) --> 2Mn²⁺ (aq) + 10CO₂ (g) + 24H₂O (l)

17.3.4 Sugar with potassium chlorate, spontaneous combustion
BE CAREFUL! Do this experiment in a fume cupboard or in the open, behind a glass screen.
1. The reaction of the concentrated sulfuric acid with the sugar released heat. The heat then activated the release of oxygen from the potassium chlorate. The oxygen released by the potassium chlorate further oxidized the sugar. This further oxidation released so much heat that the sugar bursts into flames.
Mix sugar or powdered sugar (castor sugar) with an equal amount of potassium chlorate crystals in an evaporating dish. Push a dent in the top of the heap of powder. Add one drop of concentrated sulfuric acid.
A spontaneous combustion occurs. BE CAREFUL!
2KClO₃ --> 2KCl + 3O₂
2. Repeat the experiment with potassium nitrate. The reaction is slower because while potassium chlorate loses all its oxygen, potassium nitrate loses only one third of it oxygen.
2KNO₃ --> 2KNO₂ + O₂

17.3.5 Potassium iodide with hydrogen peroxide, reverse colour change
Dissolve potassium iodide crystals in water. Add drops of starch solution and dilute hydrochloric acid. The solution is colourless. Add drops of dilute hydrogen peroxide solution. The solution turns blue black. Iodide ions are oxidized to iodine that gives starch a blue black colour.
2KI (aq) + H₂O₂ (l) --> 2KOH (aq) + I₂ (g)
Add drops of dilute sodium thiosulfate solution. The solution turns colourless. The sodium thiosulfate reduces the iodine back to iodide ions that are colourless.
I₂ (g) + Na₂S₂O₃ (aq) --> NaI (aq) + Na₂S₂O₅ (aq)
Wait until the blue black colour returns. Add drops of sodium thiosulfate solution and it disappears again. The first reaction is still going slowly. The second reaction is much slower.

17.3.6 Hydrogen peroxide with potassium sodium tartrate, cobalt (II) chloride catalyst, visible activated complex
Use a 250 mL beaker. Add 20 mL 6% hydrogen peroxide solution to a solution of 5 g potassium sodium tartrate-4-water (Rochelle salt) in 60 mL water. Heat to 75^oC. Stir and observe that gases are formed and a pink colour forms. Add 5 mL of solution containing 0.2 g cobalt (II) chloride-6-water. Observe frothing and pink colour. Be careful!
Test gases with limewater. Carbon dioxide is in gases produced by the reaction. Test gases with a glowing splint, not extinguished because some oxygen from hydrogen peroxide. Later, observe frothing stops and pink colour returns. Equation in 2 parts:
C₄H₄O₆^2- (aq) + 3H₂O₂ (aq) --> 2HCOO^- (aq) + 2CO₂ (g) + 4H₂O (l)
Co²⁺ --> Co³⁺
Pink cobalt ions oxidized to a green activated complex with tartrate
2HCOO^- (aq) + 2H₂O (aq) --> 2OH^- (aq) + 2CO₂ + 2H₂O
Co³⁺ --> Co²⁺
Basic equation: Green activated complex with tartrate reduced to pink cobalt ions
C₄H₄O₆^2- (aq) + 5H₂O₂ (aq) --> 4CO₂ (g) + 2OH^- (aq) + 6H₂O (l)

17.3.7 Potassium bromate with propanedioic acid, double autocatalytic reaction, oscillating reaction
In this reaction, bromine ions form to give a red colour but some intermediate product also forms to react with bromine ions to give a colourless solution.
Use a clean beaker washed in deionized water. Add 75 mL concentrated sulfuric acid to 750 mL deionized water (NOT tap water!). Be careful!
Leave the hot acid solution long enough to cool to room temperature slowly. Stir the cooled solution fast enough to form a vortex. Add 9 g propanedioic acid HOOC.CH_2.COOH, malonic acid. Add 8 g potassium bromate (V), KBrO₃. Add 1.8 g manganese (II) sulfate, MnSO₄.H₂O. Observe a red colour that oscillates from red to colourless, with increasing time between oscillations.
3HOOC.CH_2.COOH (aq) + 4BrO₃^- --> 4Br^- (aq) + 9CO₂ (g) + 6H₂O (l)

17.3.8 Hydrogen peroxide with catalase enzyme in raw beef liver
Hydrogen peroxide, acting as an oxidizing agent, it is toxic to cells, so it is a useful disinfecting agent that disrupts the metabolism of bacteria. Our body cells contain an enzyme called catalase that accelerates the conversion of toxic hydrogen peroxide to water and oxygen gas. It is a reactive oxygen metabolic by-product that regulates some oxidative stress-related states related to asthma, inflammatory arthritis, atherosclerosis, diabetic vasculopathy, osteoporosis, and some neurodegenerative diseases.
BE CAREFUL! This reaction can be violent and the steam formed may be hot. Use safety glasses and nitrile chemical-resistant gloves.
Put 10% hydrogen peroxide in a clear plastic container and record its temperature. Put a 6 mm piece of chopped liver in the hydrogen peroxide. The mixture starts to bubble and foam. Record the temperature each minute for 5 minutes. The temperature immediately rises, levels, then decreases. The hydrogen peroxide is decomposed, oxygen gas is given off, bubbles form and heat energy is given off.
Put a glowing splint in the test-tube of liver and hydrogen peroxide. The splint flames in the oxygen. If you boil the liver before the experiment, no bubbling occurs when hydrogen peroxide is added because you have destroyed the enzyme catalase in the beef liver, i.e. it is denatured.
Put 6 g manganese dioxide into 100 mL 10% hydrogen peroxide. Observe bubbling and relights a glowing wooden splint. Repeat the experiment using freshly cut potato instead of raw beef liver.
2H₂O₂ (l) --> H₂O (g) + O₂ (g)

17.3.9 Hydrogen peroxide with sodium thiosulfate, ammonium molybdate catalyst
Dissolve together in deionized water: 8.7 g sodium thiosulfate-5-water, 3.8 g sodium ethanoate-3-water (sodium acetate tri-hydrate) or 2.3 g anhydrous sodium ethanoate, and 0.5 g sodium hydroxide. The sodium ethanoate buffers the sodium hydroxide. Make up the solution to 1 litre and add universal indicator until each solution is blue. Pour 225 mL of the solution into 3 test-tubes labelled 1. "With catalyst" 2. "No catalyst" 3. "Control". Dissolve 14 mL of 20 vols hydrogen peroxide in deionized water. Make up to 40 mL and divide into two 20 mL portions. Add 0.08 g ammonium molybdate to test-tube 1. "With catalyst", and shake to dissolve. Add the 20 mL portions to test-tubes 1. "With catalyst" and 2. "No catalyst". Observe colour changes after 5 minutes: 1. "With catalyst": changes from blue to green to yellow to orange to orange red 2. "No catalyst": Same colour changes but slower 3. "Control": No colour change.
Na₂S₂O₃ (aq) + 4H₂O₂ (aq) --> Na₂SO₄ (aq) + H₂SO₄ (aq) + 3H₂O (l)

17.3.10 Oxidation of acetone vapour, copper oxides catalyst
1. Heat a copper wire coil or copper coil to red heat in a Bunsen burner flame then hang it just above a very thin layer of acetone in a beaker. Note the shimmering colours of the copper surface caused by the heat being maintained in the copper coin from the heat of the chemical reactions on its surface, and the colours of black copper (II) oxide, red copper (I) oxide (red) and pink copper metal. You can warm the beaker in hot water to produce sufficient acetone vapour.
Be careful! Do not smell the oxidation products because they contain ketones that may be a health hazard.
The red heat is maintained as long as some acetone remains because it keeps warm through the heat of the exothermic reaction. The reaction is safe except that at the top of the beaker where air dilutes the vapour a flame may occur. If this happens move the hot copper quickly out and cover the beaker to extinguish the flame.
2. Drop the red-hot copper into a beaker with a 2 mm thin layer of acetone at the bottom. A sizzling sound caused by the cooling of the copper ends in a loud crescendo when liquid acetone contacts the copper to increase the rate of cooling. Be careful!

17.3.11 Heat potassium chlorate, manganese dioxide catalyst
Be careful! In some school systems this experiment is not allowed because potassium chlorate may explode.
Mix 0.5 g of manganese dioxide with 2 g of potassium chlorate and put the mixture in Ignition tube 1. Put 0.5 g of manganese dioxide in Ignition tube 2. Put 2g of potassium chlorate in Ignition tube 3. Insert the ignition tubes vertically and close together in a sand tray and place a safety screen between you and the sand tray. Slowly heat the sand tray. Use a glowing splint to tests for oxygen at the openings of the ignition tubes. Oxygen appears first from Ignition tube 1 and later from the other Ignition tubes. When the reaction is complete, wash the contents into a beaker. Stir the contents to dissolve all the potassium chloride and any remaining potassium chlorate. Filter the mixture, dry the residue on the filter paper and weigh the manganese dioxide residue to show that there is no loss in weight.

17.3.12 Methanol with platinum wire catalyst, catalytic oxidation
Be careful! Have ready a piece of cardboard or glass to put over the beaker if the methyl alcohol ignites.
See diagram 17.3.12
1. Make a platinum spiral by winding platinum wire around a glass rod and leave a length of wire above the spiral. Put 1 cm of methanol (methyl alcohol) in a small beaker and warm it gently with an electric heater. Do NOT use a Bunsen burner. Heat the spiral strongly with the electric heater and transfer the glowing spiral to the beaker, holding the length of wire above the spiral so that the spiral is just above the methyl alcohol. The spiral continues to glow and you can notice the smell of formaldehyde given off from the reaction
2CH₃OH + O₂ --> 2HCHO + 2H₂O
2. Repeat the experiment with 880 ammonia solution instead of methanol and pass oxygen gas from an oxygen cylinder through the mixture. Observe brown fumes of nitrogen dioxide or white fumes of ammonium nitrate and ammonium nitrite.
4NH₃ + 7O₂ --> 4NO₂ + 6H₂O

17.3.13 Bromine catalyses the oxidation of sulfur to sulfuric acid
Do the experiment in a fume cupboard.
Put 1 cc of flowers of sulfur into two evaporating basins. Add 5 mL of concentrated nitric acid to each evaporating basin. Be careful! Add one drop of bromine to Evaporating basin 1. Be careful!. Warm each evaporating basin for 2 minutes. Pour the solutions into Test-tube 1 and Test-tube 2. Add hydrochloric acid then barium chloride solution to each test-tube. A precipitate of barium sulfate occurs only in Test-tube 1. In this catalysis intermediate compounds form which are more readily decomposed.
2S + Br₂ --> S₂Br₂ sulfur monobromide
2S₂Br₂ + 2H₂O –> SO₂ + 4HBr + 3S
SO₂ + 2HNO₃ –> H₂SO₄ + 2NO₂
2HBr + 2HNO₃ –> 2H₂O + 2NO₂ + Br₂

17.3.14 Sodium hypochlorite decomposition, cobalt sulfate catalyst
Warm a test-tube half full of sodium hypochlorite solution and observe that no decomposition occurs. Add a two drops of cobalt sulfate solution and tests for oxygen with a glowing splint.
2NaOCl –> 2NaCl + O₂ (g)

17.3.15 Ethyl acetate with sodium hydroxide, autocatalytic hydrolysis
1. Make the following solutions: Flask A 100 mL of 0.5 M sulfuric acid, Flask B 0.5 M hydrochloric acid, Flask C 0.5 M acetic acid. Leave the 3 flasks in a thermostat at 25^oC so that all the contents are at the same temperature. Titrate 2 mL of each acid in Flask A, Flask B and Flask C separately against 0.1 M sodium hydroxide, with phenolphthalein indicator and record the results.
2. Add 5 mL of ethyl acetate solution to Flask A, Flask B and Flask C then leave them in a thermostat for 15 minutes. Titrate 2 mL of the Flask A, Flask B and Flask C solutions (+ ethyl acetate) against 0.1 M sodium hydroxide solution, with phenolphthalein indicator. Repeat the titrations every 15 minutes for 2 hours and tabulate the results. Increased titration with time is because of the formation of acetic acid by hydrolysis of ethyl acetate. The catalytic action of the hydrogen ions in the original acid is increased by the hydrogen ions from the acetic acid produced by the hydrolysis. The rate of hydrolysis in the presence of a mineral acids in Flask A and Flask B is higher than in Flask C where the hydrogen ion concentration is lower. This reaction is an example of autocatalysis.
CH₃COOC₂H₅ + H₂O --> CH₃COOH + C₂H₅OH

17.4.01 Commercially available enzymes:
Enzymes are organic catalysts, but are usually used up in the reaction. An enzyme activity experiment uses a Thunberg tube to estimate dehydrogenase activity in plant material with DCPIP as indicator. Enzyme activity is affected by time, temperature, substrate concentration, enzyme concentration, pH, substrate specificity of enzyme, heat denaturation, and inhibition of enzyme activity by chemicals.
Commercially available enzymes: Enzyme biotechnology, e.g. enzymes in washing powders; amylase, starch --> maltose; cellulase, leaf materials --> various breakdown products; diastase, starch --> maltose; invertase, sucrose --> (+) glucose + fructose; lactase, milk --> (+) glucose + galactose; pancreatin (amylases, proteases, lipases); food --> various breakdown products; pectinase, fruit pulp or leaves --> various breakdown products; pepsin, skimmed milk --> peptides, protease, skimmed milk --> peptides + amino acids (in meat tenderizers); rennin (rennet powder, rennet tablets, rennilase, junket tablets) milk --> caseinogen --> casein; trypsin, proteins --> peptides + amino acids; urease, urea --> ammonia and carbon dioxide.

17.4.1 Breakdown starch to sugar
See 9.130: Hydrolysis of starch by salivary amylase (ptyalin)
See 9.142: Tests for starch, Fehling's tests for starch
Salivary amylase enzyme breaks down starch into the reducing sugars (+) glucose and maltose. Reducing sugars do not react with iodine solution and starch does not react with Fehling's solution. The sugars reduce copper (II) in Fehling's solution to brick-red copper (I) oxide.
Pour 10 mL of dilute starch solution into a test-tube. Add 1 mL of saliva and stir. After 2 minutes use a dropper put 2 drops of the solution on a white tile. To tests for starch, add iodine solution and note the intensity of the blue black colour. To tests for reducing sugars, add Fehling's No. 1 and No. 2 solutions and heat. Repeat the experiment every 2 minutes with clean droppers. Note the decreasing intensity of the blue colour that shows that starch is being used up. Repeat the experiment. Put 3 drops of the reaction mixture in a test-tube. Add 3 mL of Fehling's solution. Heat the mixture. Note the intensity of the brick-red colour increasing with time.

17.4.2 Fermentation using yeast
See 3.38: Carbon dioxide and fermentation for brewing
During fermentation, enzymes breakdown carbohydrates and other organic molecules in the absence of oxygen.

17.4.3 Bromelain enzyme from pineapples
Add pineapple juice to milk. The milk protein begins to coagulate and degrade as it reacts with the bromelain. Also, pineapple juice will also remove the gelatine emulsion surface on black and white photographic film!

17.5.0 Chemical Equilibrium
1. Equilibrium exists only in a closed system. No reactants are put in and no reactants are taken out. A closed system cannot exchange matter with its surroundings but it may exchange energy with its surroundings.
2. Limestone (calcium carbonate) is heated in a furnace to form quicklime (calcium oxide) and carbon dioxide. The reverse reaction cannot occur because the carbon dioxide is sucked out of the furnace. Also, the calcium oxide is steadily and replaced by limestone. This is a steady state system. If the furnace were closed, at high temperature both decomposition and formation of calcium carbonate would occur. When these process occur at the same rate then equilibrium exists.
CaCO₃ (s) --> CaO (s) + CO₂ (g)
3. If different concentrations of iron (II) nitrate solution are added to potassium thiocyanate solution, at equilibrium the concentration of the FeSCN²⁺solution has a certain blue colour. It is a property of this equilibrium reaction that does not change. If sodium thiocyanate NaSCN or iron (III) nitrate Fe(NO₃)₃ is added to the equilibrium mixture it changes colour.
Fe³⁺ (aq) + SCN^- (aq) <--> FeSCN²⁺ (aq)
(d) The following equation shows that both forward and reverse reactions are going on. It does not show the position of equilibrium. The position at equilibrium shows whether there are more reactants or more products at equilibrium. It can shift if reactants or products are added or removed.
N₂O₄ (g) <--> 2NO₂ (g)
In the above reaction, if, at equilibrium, N₂O₄is added, the system moves to the right, i.e. some N₂O₄changes to NO₂ until equilibrium is reached.
In the above reaction, if, at equilibrium, N₂O₄is removed, the system moves to the left, i.e. some NO₂ changes to N₂O₄ until equilibrium is reached.
Position at equilibrium can shift if the temperature changes but it will not shift because solid is added or removed from a system.
Position at equilibrium does not change if a catalyst is added to the reaction but if the system is not at equilibrium it will reach equilibrium quicker because of the influence of the catalyst.
(d) Le Chatelier's principle states that if the conditions of a system at equilibrium are altered, changes will occur in the system towards counteracting the change in conditions, when a system in equilibrium is subjected to a change in conditions, it adjusts itself so as to oppose that change, Henri Le Chatelier, (1850 - 1936).

The law of chemical equilibrium (law of mass action)
For reaction aA + bB <--> eE + fF (lower case = number of species, upper case = different types of species, e.g. 3 species of oxygen gas = 3O₂) at equilibrium: [E]^e X [F]^f / [A]^a X [B]^b= K, where K is the equilibrium constant of that reaction at that temperature. If K is large, at equilibrium, the concentration of products is much greater than the concentration of the reactants. If K is small, at equilibrium, the concentration of products is much smaller than the concentration of the reactants.

17.5.1 Concentration and temperature, cobalt (II) chloride-6-water
Test the effect of concentration and temperature on the equilibrium position. Dissolve 4 g cobalt (II) chloride-6-water in 40 mL water. The solution contains Co(H₂O)₆²⁺. It is pink. A dd concentrated hydrochloric acid until total volume is 100 mL. The solution is violet (between pink and blue). Add more concentrated hydrochloric acid to the violet solution. The solution contains CoCl₄^2-. It is blue. Add water to the violet solution. It turns pink. Heat the solution. It turns blue. Cool the solution with ice water. It turns pink. Add sodium chloride to the pink solution. It turns blue.
Co(H₂O)₆²⁺ (aq) + 4Cl^- (aq) <--> CoCl₄^2- (aq) + 6H₂O (∆H +ve)

17.5.2 Common ion effect, sodium ethanoate and ethanoic acid
A common ion occurs in solution and in another substance added to the solution. The common ion effect refers to the change in concentration of solute ions on addition of a substance to a solution of another substance.
Hydrochloric acid is a completely dissociated strong acid. Ethanoic acid is a less dissociated weak acid.
Make three solutions and add universal indicator to each:
Solution A: 100 mL 2 M hydrochloric acid, then add universal indicator until red.
Solution B: 100 mL 2 M ethanoic acid, then add universal indicator until orange.
Solution C: Add 13.5 g sodium ethanoate-3-water to 80 mL 2 M ethanoic acid. Add more ethanoic acid to 100 mL, then add universal indicator until yellow.
Put 1 g calcium carbonate powder + 1 mL liquid detergent into 3 measuring cylinders, (A1, B1, and C1), containing 100 mL water.
Simultaneously, put solution A into measuring cylinder A1, put solution B into measuring cylinder B1, and put solution C into measuring cylinder C1.
Observe the measuring cylinders: A1 has the fastest froth produced, height h. B1 has the second fastest froth produced, height h. C1 has the slowest froth produced, height h/2. In this mixture of sodium ethanoate and ethanoic acid, the equilibrium moves left, decreasing the concentration of hydrogen ions.
CH₃CO₂H (aq) <--> CH₃CO₂^- (aq) + H⁺ (aq)

17.5.3 Equilibrium between ICl and ICl₃
See diagram 17.5.3: Equilibrium reaction
Put 0.1 g iodine in the U-tube. With the 3-way tap in the CB position, turn on the filter pump to pass air bubbles through the sodium hydroxide solution. With the 3-way tap in the AB position, turn on the filter pump and let drops of concentrated hydrochloric to fall on 10 g potassium permanganate. As the chlorine passes over the iodine, first iodine monochloride, ICI, forms as a brown liquid. then iodine trichloride, ICI₃, forms as a yellow solid. Turn the 3-way tap in the CB position, remove stopper 1 from the U-tube to draw in air. The yellow solid turns into a brown liquid. Replace stopper 1 in the U-tube and turn the 3-way tap to the AB position. The brown liquid turns into a yellow solid. Increasing concentration of chlorine moves equilibrium to the right, obeying Le Chatelier's principle. With the 3-way tap in the AC position to get rid of the chlorine, tighten the screw clips. When you dip the U-tube in water just below boiling point, the yellow solid turns into a brown liquid. When you dip the U-tube in ice water, the brown liquid turns into a yellow solid. Increase of temperature moves the equilibrium to the left. Decrease of temperature moves the equilibrium to the right. Dispose of the chemicals by using different sinks. Pour the sodium hydroxide solution into a laboratory sink and wash it down with plenty of water to follow. Wash the contents of the flask generating the chlorine down the sink in the fume cupboard with plenty of water to follow.
ICl (l) + Cl₂ (g) --> ICl₃ (s)

17.5.4.1 Magnesium with hydrochloric acid reaction
A reversible reaction can proceed in either direction by altering the conditions of the reaction, e.g. 1. altering the relative concentrations, active mass 2. altering the temperature 3. altering the pressure.
Measure 10 mL. of concentrated hydrochloric acid into 4 beakers: beaker A add nothing, beaker B add 10 mL water, beaker C add 30 mL water, beaker D add 70 mL water. Simultaneously add 6 cm of magnesium ribbon to each beaker. Note how the rate of reaction in each beaker is proportional to the concentrations of the acid.

17.5.4.2 Potassium iodide with potassium iodate
Put 30 mL of solution A containing 5 g of potassium iodate per litre in a measuring cylinder and dilute to 200 mL with water, then transfer to a beaker and add drops of starch solution. Put 30 mL of solution B containing 10 g of hydrated sodium sulfite per litre and 2.5 mL of 2M sulfuric acid in a measuring cylinder and dilute to 200 mL with water. Pour this solution B into the beaker containing the diluted solution A and note the starting time in seconds. Record the time when a dark blue coloration appears. Repeat the experiment with 5 mL, 7.5 mL, 10 mL, 12.5 mL and 15 mL of 2M sulfuric acid. Plot a graph of volume of acid to reciprocal of time and note the straight line result. Reaction 1. is a slow and reaction. Reaction 2. is rapid but does not occur until reaction 1. is complete. The presence of iodine as a dark blue coloration indicates the completion of reaction 1.
1. IO₃^- + 3HSO₃^- --> I^- + 3HSO₄^-
2. 5I^- + IO₃^- + 6H⁺ –> 3I₂ (s) + 3H₂O

17.5.5.0 Law of mass action and reversible reactions, effect of alteration of concentration
A chemical reaction may stop although some of the reacting substances remain. If A and B are the reacting substances and C and D are the resulting substances, an equilibrium occurs with some A and B remaining unchanged and some C and D formed. Initially A and B react at a rate which depends on their concentrations. A change in the concentration of either A or B produces a change in the rate of the reaction. The rate of the forward action is proportional to the product of the concentrations of A and B, so rate of reaction of A and B is proportional to (Concentration of A) X (Concentration of B) = k1 (Concentration of A) X (Concentration of B). However, as soon as A and B react, their concentrations decrease, so the rate of reaction continuously decreases. The reaction between A and B will have formed some C and D, and the concentrations of A B will increase and in turn react to form A and B with the rate of reaction proportional to the product of their concentrations. Rate of reaction of C and D is proportional to (Concentration of C) X (Concentration of D) = k2 (Concentration of C) X (Concentration of D). When there is apparently no further action, an equilibrium is reached with the rate of reaction of A and B forming C and D, equal to the rate of reaction of C and D forming A and B. At equilibrium, k1 (Concentration of A) X (Concentration of B) = k2 (Concentration of C) X (Concentration of D). k1 / k2 = the equilibrium constant, k. So at equilibrium, the product of the concentrations of C and D, divided by the product of the concentrations of A and B has a definite value. Hence, if at equilibrium the concentration of, say, A is increased by an of more of it, the concentrations of B, C, and D will assume new values that the value of the expression (K) will remain unchanged. This will involve the combination of some A and B to form more C and D, i.e., the previous equilibrium concentration of B will be decreased and those of C and D will be increased.

17.5.5.1 Hydrolysis of bismuth chloride
Put some bismuth chloride in a test-tube, add 1 mL of water and observe white bismuth oxychloride forming because equilibrium was reached when certain concentrations of bismuth oxychloride and hydrochloric acid form.
BiCl₃ + H₂O <–> BiOCl (s) + 2HCl
[BiOCl] [HCl]² / [BiCl₃] [H₂O] = k (Equilibrium constant)
Add more drops of water and observe the white bismuth oxychloride reappear.
Add drops of concentrated hydrochloric acid until the white precipitate disappears because some bismuth oxychloride and some hydrochloric acid reacted to form more bismuth chloride and water. By these changes the value of the expression assumed the original mathematical value of K.

17.5.5.2 Hydrolysis of antimony chloride
Repeat the above experiment using antimony chloride in the place of bismuth chloride. The reactions and the explanations are similar to those for bismuth.
SbCl₃ + H₂O <–> SbOCl (s) (antimony oxychloride) + 2HCl

17.5.5.3 Common ion effect to precipitate sodium chloride from solution
Add a 3 drops of concentrated hydrochloric acid to a saturated solution of table salt. Sodium chloride precipitates as white crystals. Increase in the concentration of the chloride ion favours the backward reaction with subsequent precipitation of common salt. Increase in the concentration of the chloride ion favours the backward reaction with subsequent precipitation of common salt.
NaCl <–> Na⁺ + Cl^-
HCl --> H⁺ + Cl^-

17.5.5.4 Common ion effect to precipitate barium chloride from solution.
Add a 3 drops of concentrated hydrochloric acid to a saturated solution of barium chloride. Barium chloride precipitates as white crystals. Increase in the concentration of the chloride ion favoured the backward reaction with subsequent precipitation of barium chloride. Increase in the concentration of the chloride ion favoured the backward reaction with subsequent precipitation of barium chloride. The increased concentration of ammonium ion from ammonium chloride reduces the (OH^-) concentration in ammonia solution, NH₃ (aq) ("ammonium hydroxide") solution. Add a drop of phenolphthalein to a dilute solution of ammonium hydroxide. The solution goes pink. Add solid ammonium chloride drop by drop until the colour disappears. NH₄Cl produces NH₄⁺ ions which increases the speed of the back reaction.
NH₄OH <–> NH₄⁺ + OH^-

17.5.5.5 Effect of temperature on chemical equilibrium, thermal dissociation of ammonium chloride
Heat 1 mL of ammonium chloride in a dry test-tube damp red litmus paper fixed inside the mouth. The ammonium chloride sublimes and condenses on the side of the test-tube, leaving a clear space where the tube is hot. The damp litmus paper turns blue. The clear space contains the colourless gases ammonia and hydrogen chloride formed by the decomposition of the ammonium chloride. Ammonia is lighter than hydrogen chloride so it diffuses faster and reaches the red litmus paper first. Recombination to form ammonium chloride occurs in the cooler part of the test-tube
NH₄C1 <–> HCl + NH₃ (hot –>, cold <–)

17.5.6.2 Heat nitrogen tetroxide (dinitrogen tetroxide, N₂O₄)
See diagram 17.5.6.2: Heat nitrogen tetroxide
Heat a test-tube containing the lead nitrate and also heat the centre of the long horizontal delivery tube. The colour of the gas in the hot part of the delivery tube is a darker brown than in the cooler part of the tube because more brown NO₂ molecules are there. The gas phase dissociates to form nitrogen dioxide.
N₂O₄ (pale yellow) <–> 2NO₂ (brown) (pale yellow <– cooling | heating –> brown)

17.5.7.0 Explanation of group analysis
A saturated solution can remain in equilibrium with undissolved molecules of the solute. Two equilibria exist: 1. an equilibrium between the undissolved solute and dissolved molecules, and 2. an equilibrium between dissolved molecules and ions formed by dissociation
(XY) <--> XY <--> X⁺ + Y^-
Equilibrium 1.: (XY) <--> XY
Equilibrium 2.: XY <--> X⁺ + Y^-
(XY) = undissolved molecules
XY = dissolved but unionized molecules
X⁺ + Y^- = ions
The tendency of the solid to pass into solution depends on its active mass, solution pressure. If the temperature remains constant, the active mass remains constant because, by the law of mass action, concentration dissolved molecules / concentration undissolved molecules = the constant, K. The concentration of dissolved molecules is also a constant. Also, by the law of mass action, if [concentration X⁺] [concentration Y^-] / [concentration dissolved molecules] = a constant, in a saturated solution, the product of the concentrations of the ions is a constant. So if large concentrations of ions are brought together into the same solution, ions of X⁺ and Y^- will precipitate out of solution as solid molecules until the concentrations of the remaining ions in solution have such values that the product of their concentrations equals the specific constant, the solubility product.

Group I
Lead, silver and mercury (I) are precipitated as chlorides by chloride ions from hydrochloric acid. The concentrations of silver and chloride ions which can remain in solution are small. So when a solution of a silver salt containing silver ions is mixed with hydrochloric acid, most of the silver and chloride ions form molecular silver chloride and leave the solution as a solid phase until the remaining ions attain equilibrium.
(concentration silver ions) x (concentration chloride ions) = 1 x 10^-10, the solubility product.
The solubility product has a constant value, so adding excess chloride ions reduces the concentration of the silver ions to a negligible quantity. Only silver chloride, lead chloride and mercury (I) chloride have low solubility products. So the ions of other metals remain in solution in the presence of high concentrations of chloride ions.

Group II
Assume hydrogen sulfide is ionized
H₂S <–> 2H⁺ + S^2-
By the law of mass action (concentration H⁺)² x (concentration S^2-) / (concentration unionized H₂S) = the constant, 1.1 x 10^-22. In a neutral solution, the concentration of sulfide ions is low because hydrogen sulfide is a weak electrolyte. The concentration of hydrogen ions is also low. In the acid solution used for Group II, the concentration of hydrogen ions is increased by the presence of the strong acid. So to maintain the value of the solubility product constant, the concentration of the sulfide ion is reduced below its already small value in neutral solution. However, the amount of sulfide ions is enough to allow the solubility products of the sulfides of mercury (II) lead, copper and bismuth to be exceeded. Also, cadmium sulfide may precipitate if the acid is not too concentrated. So in Group II all the sulfides of mercury (II) lead, copper, bismuth and cadmium precipitate.
Solubility products: lead sulfide 4 X 10^-28, copper sulfide 8 X 10^-45, mercury (II) sulfide 4 x 10^-54, cadmium sulfide 3.6 X 10^-29
manganese sulfide 1.4 x 10^-15, zinc sulfide 1.2 X 10^-24
The concentration of sulfide ion in acid solution is not enough to allow the higher solubility products of the sulfides of manganese, zinc, cobalt or nickel to be reached with any possible concentration of the metal ion, so these sulfides do not precipitate. They precipitate later in Group IV, where the precipitating agent is the highly ionized salt, ammonium sulfide, and the concentration of the sulfide ion from it is high.

Group III
The precipitating agent is ammonia solution.
NH₄OH <–> NH4⁺ + OH^-
By the law of mass action (concentration NH₄⁺) (concentration OH^-) / (concentration unionized NH₄OH) = a constant.
Ammonia solution is a weak base, so does not ionize much and the value of the constant is only about 2 X 10^-5. Most of the ammonia solution will be dissolved but not ionized. The small hydroxyl ion concentration in a solution of ammonia solution which is also fairly concentrated with respect to ammonium chloride is still large enough to cause a precipitation of the hydroxides of ferric iron, chromium and aluminium, but not great enough to precipitate those of zinc, manganese, cobalt and nickel. Manganese hydroxide may precipitate slightly if the concentration of ammonium chloride is not sufficiently great.

Group IV
In Group II the presence of hydrogen ions from the added acid reduced the concentration of sulfide ions but this reduced value was enough to allow the solubility products of the metallic sulfides in the group to be exceeded. In Group IV, hydrogen sulfide is added to a solution made alkaline with ammonia solution and so contains excess of hydroxyl ions.
(concentration H⁺)² (concentration S^2-) / (concentration unionized H₂S) = a constant. The ionic product of water [H⁺] x [OH-¹] = (10^-14). The hydroxyl ions from the ammonia solution lower the concentration of hydrogen ions causing an increased concentration of S^2- ion. In Group IV, the metal sulfides not already precipitated in Group II because of their high solubility products, are here precipitated. So the ionic concentration of the sulfide ion is controlled by variation of the concentration of the ions it is associated with, i.e. hydrogen ions.

Group V
The metals still remaining in solution include barium, strontium, calcium, magnesium, sodium and potassium. Barium, strontium and calcium are precipitated as carbonates by the addition of CO₃^2- ions in alkaline solution because of the low values of the solubility products [X²⁺] [CO₃^2-] = K, between 10^-8 and 10^-9. The solubility product of magnesium carbonate is low, 10^-5, and in neutral solution would be precipitated, but its precipitation is prevented in this group by the ammonium ions, mainly from the ammonium chloride added before Group III. With the large NH₄⁺ ion concentration from this source, the concentration of CO₃^2- ions is reduced below the concentration to precipitate magnesium as a carbonate.
[NH₄⁺]² [CO₃^2-] / [NH₄)₂CO₃] = K

17.6.1 Weight of iron in iron (II) ammonium sulfate.
Dissolve 2 g of iron (II) ammonium sulfate in 50 mL of demineralized water. Heat to boiling then add 2 mL of concentrated nitric acid and continue boiling for two minutes. Leave to cool. Add ammonium hydroxide and stir until precipitation is complete. Test with litmus paper. Heat again to boiling and filter. Use a wash bottle to wash any iron (III) hydroxide from the beaker and stirring rod. Leave the filter paper and filtrate to dry. Weigh a crucible. Transfer the dry filter paper and precipitate to the crucible. Ignite the filter paper, leave to cool and weigh. Heat the crucible until it becomes red, leave to cool and weigh until again until the weight remains constant.
2Fe₂O₃ –> 4Fe + 3O₂
Molecular weight of Fe₂O₃ = 160
Molecular weight of Fe = 56
So 2 Fe, 2 x 56 is equivalent to 1 Fe₂O_3, 160
Let weight of iron (II) ammonium sulfate = 1.96 g
Let weight of iron (III) oxide = 0.40 g
So weight of iron in iron (III) oxide = 112 x 0.40 / 160 = 0.28 g
Percentage weight of iron in iron (II) ammonium sulfate = (0.28 / 1.96) X 100 = 14.29%.

17.6.2 Weight of aluminium in aluminium sulfate.
Dissolve 5 g of aluminium sulfate in 50 mL of demineralized water. Add drops of dilute sulfuric acid and heat to boiling. Add excess of ammonium hydroxide and heat to boiling. Filter, leave to dry, ignite the filter paper in a crucible, leave to cool, then weigh.
Al₂SO₄ + NH₄OH –> Al(OH)₃ + NH₄SO₄
AL(OH)₃ –> Al₂O₃ + H₂O
Molecular weight of Al₂O₃ = 102.
Let weight of aluminium sulfate = wa g,
Let weight of aluminium oxide = xb g.
So weight of aluminium in oxide = (54 / 102) x wb g.
Percentage weight of aluminium in aluminium sulfate = [(54 x wb) / 102] / wa x 100%.

17.6.3 Weight of calcium in marble, calcium carbonate.
Stir 1.5 g of marble powder into 10 mL of demineralized water then add 5 mL of concentrated hydrochloric acid. Heat the solution until the marble dissolves, then add water to make the volume 50 mL. Add ammonium solution until the solution is alkaline by litmus test, then heat to boiling. Add 3 g of crushed ammonium oxalate, stir, and again heat to boiling. Leave the solution to form a precipitate. Wash the liquid and precipitate into a filter paper. Wash the precipitate with water until the filtrate contains no more chloride ions as tested with silver nitrate solution. Dry the precipitate of calcium oxalate, ignite the filter paper in a furnace, leave to cool, then weigh the final residue of calcium oxide.
CaCO₃ + 2HCl –> CaCl₂ + CO₂ + H₂O
CaCl₂ + (NH₄)₂C₂O₄ —> CaC₂O₄ + NH₄Cl
CaC₂O₄ + "O" –> CaO + 2CO₂
(The “O” is from air)
Molecular weight of CaO = 56
Let weight of calcium carbonate = wa g.
Let weight of calcium oxide = wb g.
So weight of calcium in calcium oxide = (40 / 56) x wb g.
Percentage weight of calcium in marble = (40wb / 56wa) X 100%.

17.6.4 Weight of magnesium in magnesium sulfate
Dissolve 2 g of magnesium sulfate in 50 mL of demineralized water. Add 10 mL of ammonium chloride solution and then ammonium hydroxide until alkaline after stirring. If magnesium hydroxide precipitates add more ammonium chloride. Heat and add sodium phosphate solution in excess, stir, and leave to settle. Filter then wash the precipitate with ammonium hydroxide to remove ammonium chloride. Test presence of chloride with silver nitrate solution acidified with nitric acid. Heat the solid gradually then strongly, leave to cool, then weigh.
MgSO₄ + NH₄OH + Na₂HPO₄ --> Mg(NH₄)PO₄ + Na₂SO₄ + H₂O
magnesium ammonium phosphate
2Mg(NH₄)PO₄ --> Mg₂P₂O₇ + 2NH₃ + H₂O
Molecular weight of Mg₂P₂O7, magnesium pyrophosphate = 222
Let weight of magnesium sulfate = aw g.
Let weight of magnesium pyrophosphate = bw g.
So weight of magnesium in the pyrophosphate = (48 /222) x bw g.
Percentage magnesium in magnesium sulfate [(48 bw /222)] / aw X 100%.

17.6.5 Weight of sulfate radical in sodium sulfate
Dissolve 3 g of sodium sulfate crystals in 50 mL of demineralized water. Add 5 mL dilute hydrochloric acid and 5 mL of ammonium chloride solution, then heat to boiling. Add excess barium chloride solution to assist precipitation of barium sulfate, then and heat to boiling. Leave to settle and then decant the liquid into filter paper. Wash the precipitate onto a filter paper then dry in an oven. Add a drops of concentrated nitric acid then heat the crucible until it becomes red. Leave to cool then weigh.
Molecular weight of BaSO₄ = 233
Let weight of sodium sulfate crystals = aw g.
Let weight of barium sulfate = bw g.
So weight of sulfate in barium sulfate = (96 / 233) x bw g.
Percentage weight of sulfate in sodium sulfate = [(96 bw / 233) / aw] x 100%.

17.6.6 Weight of tin in solder.
Make finely divided solder with a rough file. Add 10 mL of concentrated nitric acid to 1 g of finely divided solder in an evaporating basin. Heat gently in a fume cupboard. When the reaction stops, add two drops of nitric acid and heat gently again, until no more nitrogen dioxide forms. Dilute the contents to 50 mL and filter. Wash the precipitate with dilute nitric acid. Dry the precipitate, ignite the filter paper, leave to cool, then weigh the final product. The tin was first oxidized to metastannic acid, hydrated tin (IV) oxide, formula H₂Sn₅O₁₁ or SnO₂.xH₂O. It is used as an opacifying colour in ceramics and an abrasive in the glass industry.
5Sn + 20HNO₃ –> H₂Sn₅O₁₁ + 20NO₂ + 9H₂O
H₂Sn₅O₁₁ –> 5SnO₂ + H₂O
Molecular weight of SnO₂ = 151
Let weight of solder = aw g
Let weight of tin (IV) oxide = bw g
So weight of tin in tin (IV) oxide = 119 bw / 151 g
Percentage of tin in solder = (119 bw /151 aw) x 100.