School Science Lessons
Chemistry experiments
Please send comments to: J.Elfick@uq.edu.au
2012-05-14 SPwp
Table of contents
3.95 Breakdown large molecules to small molecules
3.70 Chemical reactions
3.84 Electrical energy from chemical reactions
3.80 Energy from chemical reactions
3.31 Hygroscopic, deliquescent and efflorescent chemicals
3.91 Rate of reaction
3.95 Breakdown large molecules to small molecules
3.95 Breakdown starch to sugars
3.96 Breakdown ethanol to ethene (ethylene)
3.97 Breakdown polymers using heat
3.98 Elements in foods
3.99 Gases from wood
3.70 Chemical reactions
12.2.0 Chemical reactions, types of chemical reactions
3.71.5 Blackboard chalk, (school chalk)
3.72 Magnesium displaces copper from copper ions
3.74 Metals displace hydrogen from acids
3.79 Prepare soap with fats and oils
3.77 Reactions of magnesium with carbon dioxide, sparkler experiment
3.75 Reactions of salts with water
3.73 Reactions of sodium with water
3.71.1 Solubility table and solubility rules
3.71.2 Sodium chloride solution with copper (II) sulfate solution
3.71.3 Test if precipitate forms when solutions of salts are mixed
3.71.4 Test if precipitate forms when solutions are added to lead (II) nitrate
3.78 Titration of acids with bases
3.84 Electrical energy from chemical reactions
3.85 Daniell cell, porous pot, Zn in ZnSO4 soln. / Cu in CuSO4 soln.
3.84 Electrical energy from a simple cell, displacement of copper by zinc
3.84.1 Electrochemical cell, voltaic cell, galvanic cell
3.86 Electrode potentials of metals
3.87 Lead accumulator cell, lead-acid rechargeable battery
3.88 Leclanché cell, dry cell
3.89 Movement of copper and chromate ions
3.90 Movement of ions between microscope slides, Cu2+ ions, CO2+ ions
3.84.7 Pencil sharpener galvanic cell, Mg, Fe, NaCl, cola
3.84.2 Simple cell with different metals
3.84.3 Simple electric cell, copper with zinc in dilute sulfuric acid
3.84.4 Simple galvanic cell, Zn and Cu in dilute HCl
3.84.6 Standard electrode potential
3.84.5 Voltaic cell with a salt bridge, Zn in ZnSO4 soln. / Cu in CuSO4 soln.
3.80 Energy from chemical reactions
3.80 Exothermic reactions give out heat energy
3.81 Endothermic reactions take in heat energy
3.82 Heat of neutralization reactions
3.83 Heat of reaction when metals displace copper
3.31 Hygroscopic, deliquescent and efflorescent chemicals
3.31.01 Hygroscopic substances
3.31.02 Deliquescent substances
3.31.03 Efflorescent substances
3.31.04 Desiccants, drying agents
3.31.1 Expose different salts to the air
3.31.2 Expose sodium carbonate decahydrate, washing soda, to the air
5.3.9 Tests for presence of water, cobalt (II) chloride paper
3.91 Rate of reaction
3.7.1 Concentration and rate of reaction
3.91 Size of particles and rate of reaction
3.92 Concentration and rate of reaction, sodium thiosulfate with hydrochloric acid
3.93 Temperature and rate of a reaction
3.94 Catalysts and rate of reaction
3.31.01 Hygroscopic substances
The hygroscopic solids absorb water without liquefaction. Hygroscopic substances absorb water from the air. Hygroscopic chemicals include the following: (anhydrous) calcium chloride (CaCl2), glycerol (CH2OH.CHOH.CH2OH), iron (II) sulfate-7-water (FeSO4.7H2O), (concentrated) nitric acid (HNO3), potassium carbonate (K2CO3), potassium chloride (KCl) (slight), potassium iodide (KI) (slight), silica gel desiccant (SiO2), (anhydrous) sodium carbonate (Na2CO3), (concentrated) sodium hydroxide (NaOH), sodium nitrite (NaNO2), (anhydrous) sodium sulfate Na2SO4, (concentrated) sulfuric acid (H2SO4). Hygroscopic substances should be treated as potentially highly corrosive.
3.31.02 Deliquescent substances are hygroscopic substances that absorb water to such an extent that they form a concentrated solution of the substance. Deliquescent chemicals absorb water from the air and dissolve in it to form a concentrated solution, e.g. citric acid (slight), cobalt (II) nitrate Co(NO3)2.6H2O, magnesium chloride MgCl2, potassium hydroxide KOH, potassium iodate KIO3 (slight), potassium iodide KI (slight), sodium nitrate (if NaNO3, in moist air) sodium thiosulfate Na2S2O3.5H2O (in moist air). Store deliquescent chemicals in an airtight container or in a desiccator. When exposed to the air, sodium chloride neither gains nor loses water. Pure NaCl is not hygroscopic. However, sodium chloride as table salt in a salt shaker may become sticky and hard to shake out because it contains deliquescent magnesium chloride as an impurity. Add calcium carbonate or rice grains to table salt to stop it deliquescing. Deliquescent substances should also be treated as potentially highly corrosive.
3.31.03 Efflorescent substances have crystals containing water of crystallization but may lose some of that water by evaporation and begin to powder. Also, in a cellar, efflorescence refers to when dissolved salts rise to the surface and crystallize out of the ground or on a wall. Efflorescent chemicals include the following: copper (II) sulfate (CuSO4.5H2O) (in dry air), iron ammonium sulfate (Fe(NH4)2(SO4).6H2O), lead acetate [(CH3COO)2Pb.3H2O] (slow), magnesium sulfate (MgSO4.7H2O), di-sodium orthophosphate (Na2HPO4.12H2O), sodium sulfate decahydrate (Na2SO4.10H2O), sodium tetraborate decahydrate (borax) (Na2B4O7.10H2O) (in dry air), zinc sulfate (ZnSO4.7H2O).
3.31.04 Desiccants, drying agents
See diagram 3.31.04: Drying agents
Desiccants, drying agents, are used where drying by heating may decompose the substance to be dried. Drying agents, desiccants, may just absorb water into their surface structure, e.g. silica gel. Dry silica gel is hygroscopic. It absorbs water from the air, but does not dissolve in the water. Little white bags containing silica gel desiccant are often included in electrical equipment and other products just before the close of the manufacturing process.
Powerful desiccants may remove water by a chemical reaction, e.g. concentrated sulfuric acid.
Concentrated sulfuric removes the water of crystallization from blue copper sulfate crystals to form white anhydrous copper sulfate
CuSO4.5H2O --> CuSO4 + 5H2O
Concentrated sulfuric acid removes water from the white to colourless sucrose molecules to leave black carbon.
C12H22O11 --> 12C + 11H2O
However, concentrated sulfuric acid is a powerful acid and oxidizing agent so it must be used with great care.
Laboratory desiccators are usually heavy glass containers sealed with a ground glass lid lubricated with Vaseline to make the container air tight. The substance to be dried is on a watch glass on a shelf above the desiccant. The base of the desiccator may contain various desiccants, e.g. phosphorus pentoxide, anhydrous magnesium sulfate, melted caustic potash (potassium hydroxide), concentrated sulfuric acid (it turns black with time), silica gel, anhydrous calcium oxide, anhydrous calcium chloride, anhydrous zinc chloride, anhydrous copper sulfate. Phosphorus pentoxide is a much more powerful drying agent then anhydrous copper sulfate. The desiccator may be used as a suitable cooling vessel so that the substance can cool yet not at the same tome absorb water from the atmospheres. Some liquids, e.g. organic compounds, may be dried by dropping in calcium chloride that is removed hours later by a glass wool filter. Gases are dried by passing the damp gas through an U-tube or glass tower containing anhydrous calcium chloride between glass wool plugs with airtight taps at each end. Gases can be dried by bubbling through concentrated sulfuric acid but some acid droplets get carried into the dried gas. To avoid acid droplets, the sulphuric acid can be absorbed into lumps of pumice that are then used in a glass drying tower. Drying gases is usually a very slow process.
Gas to be dried
 Drying agent
Ammonia
 Calcium oxide
Carbon dioxide
 Any in this column
Carbon monoxide
 Phosphorus pentoxide
Chlorine
 Concentrated sulfuric acid
Hydrogen
 Calcium chloride
Hydrogen chloride
 Concentrated sulfuric acid
Hydrogen sulfide
 Calcium chloride
Nitrogen
 Calcium chloride
Nitrous oxide
 Concentrated sulfuric acid
Oxygen
 Phosphorus pentoxide
Sulfur dioxide
 Any in this column
3.31.1 Expose different salts to the air
Put equal volumes of different salts on separate watch glasses, e.g. dry calcium chloride, copper (II) sulfate-5-water, iron (II) sulfate-7-water, potassium carbonate, dry silica gel, pure sodium chloride, and sodium sulfate-10-water. Examine the salts after one hour and after one day.
3.31.2 Expose sodium carbonate decahydrate, washing soda, to the air
The relative molecular mass of sodium carbonate-10-water = 286.15.
The relative molecular mass of anhydrous sodium carbonate = 105.99.
The relative molecular mass of sodium carbonate in sodium carbonate crystals = (105.99/286.15) × 100 = 37%
Open an unopened packet of washing soda. Put 10 g of washing soda in an evaporating basin. Record the time for the crystals to change into a white powder. Leave the basin for two days. Record the time for all the crystals to change into a white powder. Weigh the powder. Calculate the weight of the powder expressed as the percentage of the original weight of the crystals.
3.31.3 Tests for water with cobalt (II) chloride
Tests for the presence of water with blue cobalt (II) chloride paper. Soak paper in anhydrous cobalt (II) chloride and store in a desiccator. Heat cobalt (II) chloride-6-water crystals. The reaction forms the dark blue anhydrous cobalt (II) chloride with the loss of water. Add water to anhydrous cobalt chloride. The solution becomes pink. Evaporate the pink solution to form purple crystals.
[In this direction, heat enters the reaction. -->]
CoCl2.6H2O (s) [pink] <--> CoCl2 (s) [blue] + 6H2O (l)
[<-- In this direction, heat leaves the reaction.]
3.71.1 Solubility table and solubility rules
1. All ethanoates (acetates) are soluble, but the Ag+ salt is slightly soluble.
2. All carbonates are insoluble, except the Na+, K+ and NH4+ salt.
3. All chlorides are soluble, except the Ag+ and Hg+ salt. The Pb2+ salt is slightly soluble, but more soluble in hot water.
4. All hydroxides are insoluble, except the Na+, K+ and NH4+ salt. The Mg2+ and Ca2+ salts are slightly soluble.
5. All nitrates are soluble.
6. All phosphates are insoluble, except the Na+, K+, NH4+ salts and some acid phosphates.
7. All common sodium, potassium and ammonium salts are soluble.
8. All sulfides are insoluble, except the Na+, K+, NH4+, Mg2+, Ca2+ and Ba2+ salts.
9. All sulfates are soluble, except the Ba2+, Pb2+, Ca2+ and Hg2+ salts. The Ag2+ salt is slightly soluble.
10. All salts of silver are insoluble, except silver nitrate and silver chlorate. Silver ethanoate (silver acetate) and silver sulfate are slightly soluble.
Test if a salt is soluble in water. Select salts from the laboratory, e.g. ammonium chloride, barium chloride, barium sulfate, calcium sulfate, copper nitrate, copper (II) carbonate, copper (II) sulfate, lead (II) nitrate, potassium nitrate, potassium chloride, potassium sulfate, sodium chloride, sodium ethanoate (acetate) sodium sulfate, sodium carbonate. Put 5 g of each salt in a test-tube. Note the room temperature. Add 10 mL of water and stir or shake vigorously. Note whether the temperature of the mixture changes. Classify each salt as soluble or slightly soluble or insoluble. Check whether the results agree with the solubility rules.
3.71.2 Sodium chloride solution with copper (II) sulfate solution
If neither of the possible products of a reaction is insoluble, then the reaction does not occur and go to completion. The reaction reaches equilibrium with both forward and reverse processes continuing at equal rates.
1. Add sodium chloride solution to copper (II) sulfate solution. No reaction occurs to completion because the solubility rules show that both possible products, Na2SO4 and CuCl2, are soluble. The equilibrium only involves ions interacting with the water. There is no "Na2SO4" as such.
Na2SO4 + CuCl2 <--> 2Na+ + SO42- + Cu2+ + 2Cl-
2. Repeat the experiment with a mixture of sodium nitrate solution and potassium chloride solution. Nothing happens. Check potassium nitrate and sodium chloride in the solubility table.
3.71.3 Test if precipitate forms when solutions of salts are mixed
Mix 5 mL of each of the different solutions of salts available in your laboratory. Record observations for each pair of solutions.
3.71.4 Test if precipitate forms when solutions are added to lead (II) nitrate
Add lead (II) nitrate to the following:
1. dilute hydrochloric acid,
2. dilute sulfuric acid,
3. sodium hydroxide solution.
A precipitate forms in each test-tube.
3.71.5 Blackboard chalk, (school chalk)
The chemical in blackboard chalks are harmless if ingested but chalk dust may cause respiratory problems in children with asthma and in teachers after exposure for many years. The new product called "dustless chalk" is chemically the same but the chalk dust it produces consists of much heavier particles that quickly drop down instead of being suspended in the air. Chalk dust should be regularly cleaned out of the classroom . Chalk dust may damage electronic equipment, e.g. computer motherboard and playback heads of VCRs. Yellow chalk on a green "blackboard" may be more visible to most students but may cause problems for colour blind students.
1. Solubility of blackboard chalk in water. Shake powdered blackboard chalk (school chalk), mainly CaSO4.2H2O, with water in a test-tube. Filter the mixture and collect the filtrate in an evaporating basin. Evaporate the water by heating the evaporating basin over a beaker of boiling water. Examine the inside surface of the basin. If any residue is found, then some chalk is soluble in water. School chalk may also be calcium sulfate.
2. Blackboard chalk in weak acids. Put three sticks of blackboard chalk in beakers partly filled with water, lemon juice, vinegar, so that about half of each stick is still dry. Observe the chalk sticks over the next days. The chalk sticks dissolve in the weak acids.
3.72 Magnesium displaces copper from solution of copper ions
A metal higher in the activity order can displace copper metal from a solution of copper ions.
1. Put 10 mL of an M copper (II) sulfate solution in a small beaker. Clean magnesium ribbon and cut into 0.5 cm pieces. Add these pieces to the copper (II) sulfate solution one at a time. Be careful! The reaction can be vigorous. Copper metal deposits and the blue colour gradually disappears as the magnesium displaces the copper ion. Note any heat given out by the reaction. When the solution is colourless, decant the solution from the red copper powder at the bottom of the beaker. Collect the copper and dry it.
Mg (s) + Cu2+ (aq) ---> Mg2+ (aq) + Cu (s)
2. Repeat the experiment by attempting to displace copper metal using powdered zinc and iron metal. Note the comparative activity of the metals.
3.73 Reactions of sodium with water
See diagram 3.2.73: Sodium in water and under kerosene | See diagram 3.73.1: Sodium with water
2Na (s) + 2H2O (l) --> 2NaOH (aq) + H2 (g)
Be careful! Check the safety rules in your school system before doing this dangerous experiment. Use safety glasses and nitrile chemical-resistant gloves
1. The following experiments use a test-tube to observe the reaction of sodium with water but some teachers use a beaker full of water instead to avoid problems of confining the reaction in a test-tube. Cover the beaker with wire gauze immediately after adding the sodium. Do not try to hold the sodium in place on the surface of the water to collect the hydrogen produced because the gas may ignite or explode. Some teachers have done the experiment using the school swimming pool.
2. Pour a 2 cm layer of kerosene on to the surface of water in a test-tube. Drop a 3 mm diameter side cube of sodium into the kerosene. Be careful! Sodium sinks in the kerosene and float in the water. Adjust the layer of kerosene to be shallow enough to allow the top of the sodium to protrude above the surface. This reaction of sodium with the water is much slower than if the sodium had been dropped directly on to the water. You can watch the reaction through a magnifying glass held at the side, but never look down into a test-tube. Sodium metal is lighter than water but heavier than kerosene. A small area of the sodium suddenly reacts causing a stream of hydrogen bubbles to appear. The stream of bubbles at one side causes movement. The irregular shape of the sodium changes to a sphere. The sodium melts because the reaction gives off heat. Note any variations in light refraction and reflection below the sodium that suggests something dissolving in the water. Slight smoke where the hot sodium is above the kerosene level suggests a slight reaction with air. Test the gas bubbles for oxygen gas or hydrogen gas.
3. Be careful! Use a piece of sodium the size of a wheat grain or rice grain. Do not point the open end of a test-tube containing sodium at anybody.
The experiment can be done on an overhead projector or drop the sodium in a swimming pool.
Put water into a Petri dish and add a few drops of phenolphthalein. Put the Petri dish on an overhead projector. Use forceps to take a small lump of metallic sodium stored under kerosene and wipe with a piece of filter paper. Cut a piece of sodium the size of a red bean from the lump and then put it in the Petri dish. Observe from the screen that the sodium grain floats on the water to melt into a small sphere, which moves very rapidly in all directions, becoming smaller and smaller. At last, the sodium sphere wholly disappears and the solution in the dish also changes its colour from colourless to red.
4. Put a piece of sodium in 2 cm of water in a test-tube.
Be careful! Observe the reaction through the side of the test-tube. Do not look down the test-tube!
Do not point the open end of the test-tube at anybody! Tests for hydrogen gas with a lighted splint. Test the contents with litmus paper or phenolphthalein. The reaction forms sodium hydroxide.
5. Pour a 3 mm layer of kerosene (paraffin oil) on to the surface of water in a test-tube.
Be careful! Observe the reaction through the side of the test-tube. Do not look down the test-tube!
Drop a small piece of sodium into the test-tube. The sodium sinks in the kerosene and floats in the water. The layer of kerosene should be shallow enough to let the top of the sodium protrude above the surface. The reaction of sodium with water is much slower than if the sodium had been dropped directly on to the water. A small area of the sodium suddenly reacts which causes a stream of bubbles to appear. The irregular shape of the sodium changes to that of a sphere. The sodium melts because the reaction forms heat. Something dissolving in the water below the sodium can be seen. Some smoke where the hot sodium is above the kerosene level suggests a reaction with air. Test the gas bubbles for hydrogen gas.
6. Support a short length of glass tubing vertically with one end 2 cm below the surface of water. Test the water with moist litmus paper. Add litmus paper to the water. The upper part of the tube must be quite dry. Drop a piece of sodium down the tube. Note the reaction. Tests for hydrogen gas with a lighted splint. Test the water with litmus paper. The water becomes alkaline.
2Na (s) + 2H2O (l) -->H2 (g) + 2NaOH (aq)
6. Remove a small lump of sodium from storage under oil. Observe the layer of oxide / hydroxide on the surface of the sodium. Cut a pea size piece from the surface of the sodium. Observe the shiny sodium metal surface that rapidly darkens because of reaction with air and moisture.
7. Demonstrate the reaction of sodium with water in a fume cupboard or outside if students are protected by a safety shield or wearing safety glasses. Use a large beaker filled with water to within 1 cm of the top, so that there is no space for air / hydrogen mixture to accumulate. Add a piece of sodium no greater in size than a rice grain, 3 mm diameter. Use a wire gauze on top of the beaker to prevent the ejection of sodium. The sodium fizzes around the surface of the water for a few seconds, reacting violently. Add an acid / base indicator to the water, e.g. universal indicator or phenolphthalein, to demonstrate the alkalinity of the reaction products. Demonstrate the disposal of waste sodium by reaction with ethanol or methylated spirit in a beaker in the fume cupboard as described above. The sodium and may catch fire with the following reaction:
Na + H2O –> Na+ + OH- + 1/2H2 (g)
7. Place sodium on wet absorbent paper and observe the ignition of the hydrogen gas, This experiment can be used to test for the presence of alcohols.
3.74 Metals displace hydrogen from acids
1. Pour 5 cm of the acids in the table below into test-tubes. Place a piece of metal foil in each test-tube. Note the formation of hydrogen and compare the different rates at which the bubbles are formed.
Rate of formation of hydrogen gas with 3M hydrochloric acid and 3M sulfuric acid
Magnesium: 3M hydrochloric acid - Very rapid reaction, 3M sulfuric acid - Rapid reaction
Aluminium: 3M hydrochloric acid - Slight reaction, 3M sulfuric acid - No reaction
Zinc: 3M hydrochloric acid - Moderate reaction, 3M sulfuric acid - Slight reaction
Iron: 3M hydrochloric acid - Very slight reaction, 3M sulfuric acid - Very slight reaction
Tin: 3M hydrochloric acid - No reaction, 3M sulfuric acid - No reaction
Lead: 3M hydrochloric acid - No reaction, 3M sulfuric acid - No reaction
Copper: 3M hydrochloric acid - No reaction, 3M sulfuric acid - No reaction
2. Recover the zinc after the reaction has stopped. Evaporate the solution to leave zinc sulfate crystals. Dissolve the colourless zinc sulfate crystals in water and put two carbon electrodes (central poles of dry cell batteries) in the solution. Connect the electrodes to a 6 V or 12 V d.c. supply. Zinc forms rapidly on the cathode.
3.75 Reactions of salts with water
Water and salts do not usually react but sometimes hydrolysis occurs and the solution becomes either acidic or alkaline.
Dissolve a small amount of the following salts in demineralized water and test each solution with red and with blue litmus paper: sodium chloride, sodium carbonate, copper (II) sulfate, sodium acetate, iron chloride. copper (II) sulfate and iron chloride give acidic solutions. Sodium carbonate and sodium acetate give alkaline solutions. Sodium chloride solution is neither acidic nor alkaline.
3.77 Reactions of magnesium with carbon dioxide, sparkler experiment
1. Fill a gas jar with carbon dioxide. Hold a piece of clean magnesium ribbon in a pair of tongs, ignite the magnesium with a Bunsen burner flame and plunge it into the carbon dioxide gas. The magnesium continues to burn. If the magnesium is taking oxygen from the carbon dioxide for burning then you would find carbon in the gas jar. Look for carbon specks in the gas jar. To make the carbon more visible, you can add drops of sulfuric acid to remove the magnesium oxide and any unburned magnesium.
2. The sparkler experiment
Use Plasticine to stick a small birthday candle to the bottom of a cut-off plastic drink bottle and light the candle. Mix one teaspoon of bicarbonate of soda in half a cup of water. Mix one teaspoon of cream of tartar (tartaric acid) in another half a cup of water. Pour both solutions into the drink bottle (but not enough to cover the candle!). Bubbles of carbon dioxide appear and then the candle goes out. You cannot relight the candle with a lighted match because of the carbon dioxide around the candle and carbon dioxide is heavier than air. So the match goes out before you can light the candle. Some people can blow a soap bubble with a bubble pipe and sit the bubble on top of the layer of carbon dioxide - but doing this is not easy. Light the sparkler and hold the sparkling end in the drink bottle. The sparkler does not go out because the sparkles come from burning magnesium powder and magnesium reacts with carbon dioxide. You can now relight the candle with a match because all the carbon dioxide has reacted with the magnesium in the sparkler and oxygen has returned to the cut-off drink bottle. Have a dish of sand nearby to take the hot end of the sparkler. Teachers refuse to do this experiment with some classes because undisciplined children may burn themselves or children, or leave the hot sparkler on the desk.
You may see some black bits of carbon form on the side of the bottle. They come from the carbon dioxide.
2Mg + CO2 --> 2MgO + C
3.78 Titration of acids with bases
Put 20 drops of a dilute acid, e.g. vinegar in a test-tube. Add one drop of indicator, e.g. methyl orange or phenolphthalein. Add a dilute base drop by drop, and count the drops. Within experimental error, it will always take the same number of drops to neutralize the 20 drops of acid if the same dropper, teat pipette, is used. If the concentration of the acid is known, the concentration of the base can be estimated by comparing the numbers of drops of acid and drops of base that just react.
3.79 Prepare soap with fats and oils
Soap is made by using a double displacement reaction of a strong base, e.g. sodium hydroxide and many fats and oils. Castile soap, seafarer's soap, is made from vegetable oil, usually olive oil, but also coconut oil and other vegetable oils.
1. Dip the tips of two fingers in < 0.1 M dilute potassium hydroxide solution to feel the soap made. Rinse hands thoroughly after this experiment.
2. For this experiment, wear safety glasses and the area must be well-ventilated. Dissolve 2 g of sodium hydroxide or potassium hydroxide in 25 mL of methylated spirit in a small beaker. Heat the solution on an electric hot plate or in a water bath. Pour 10 mL of olive oil into an evaporating basin, add the sodium hydroxide solution then stir with a glass rod. Put the evaporating basin on a large beaker filled with hot water. Lave the methylated spirit to evaporate until a soft semi-solid mixture of soap and glycerol remains. Dissolve the mixture in 25 mL of hot water and pour it into a 50 mL beaker. Add 5 g of sodium chloride and stir. Add 2 mL ethanol and warm the mixture, not to boiling. Leave the mixture to cool, then cool remove the top layer of soap and test it by shaking with water.
3. Obtain animal kidney fat from a butcher. Boil this fat in water and the oil will separate on the surface. When cold, the fat will solidify and it can be separated from the water. Melt the fat again and strain through several layers of cloth. Weigh this fat and then weigh out one third as much sodium hydroxide pellets. Be careful! Do not touch the solid sodium hydroxide or the solution because both are caustic. Use safety glasses and nitrile chemical-resistant gloves. Heat the fat in an iron dish. When it is molten, slowly add the sodium hydroxide solution with continuous stirring. Heat with a small flame to avoid boiling over. Allow the fat and the sodium hydroxide to boil for 30 minutes. Stir the mixture frequently. Weigh sodium chloride using twice the weight of the sodium hydroxide pellets. After the 30 minutes boiling, stir the sodium chloride into the mixture and leave to cool. The soap separates as a layer at the top. Separate this soap from the liquid below, melt and pour into matchboxes where it will solidify again as small bars of soap.
3.80 Exothermic reactions give out heat energy
Be careful! The reactions may be vigorous. If less energy is needed to break the bonds in a reaction than the energy released by making the bonds, the reaction is endothermic.
1. Put 1 cm of white anhydrous copper (II) sulfate powder in a test-tube. Hold a thermometer with the bulb in the powder. Add water drop by drop. Record any change in the thermometer reading. (Water with anhydrous compounds reactions.)
CuSO4 + 5H2O --> CuSO4.5H2O
2. Put 10 mL of 0.4 M copper (II) sulfate solution into a wide test-tube. Support a thermometer with the bulb in the solution. Add magnesium powder, or magnesium ribbon, a little at a time, until the blue colour disappears. Record any change in the thermometer reading. Magnesium is higher in the reactivity series than copper so it displaces copper from its sulfate.
3. To a little water in a wide test-tube, add concentrated sulfuric acid, drop by drop, down the side of the test-tube. Stir gently with a thermometer after the addition of each drop. Record any change in the thermometer reading. (Concentrated acid with water reactions)
4. Dilute hydrochloric acid with dilute sodium hydroxide solution. (Neutralization reactions)
5. Citric acid with sodium carbonate solution.
C6H8O7+ 3NaHCO3 --> C6H5Na3O7 + 3CO2 + 3H2O
citric acid + sodium hydrogen carbonate --> sodium citrate + carbon dioxide + water
6. Burning substances and combustion of fuels.
7.Thermite reaction, Reactions of aluminium, Al, aluminate, AlO2: 12.01.1 (See 3.)
8. Setting of cement and concrete
9. Corrosive oxidation of metals.
3.81 Endothermic reactions take in heat energy
See diagram 3.81: Temperature of potassium nitrate solution
An endothermic reaction drops in temperature as it absorbs heat energy. If more energy is needed to break the bonds in a reaction than the energy released by making the bonds, the reaction is endothermic.
1. Put 10 mL of water in a test-tube. Read the temperature of the water. Dissolve 2 g of potassium nitrate in the water. The temperature should fall through 90oC. So while dissolving, the particles are absorbing heat energy. This energy is taken from the surrounding water. Repeat the experiment with potassium chloride.
Mg +CuSO4 --> MgSO4 + Cu displacement
2. Add citric acid to sodium hydrogen carbonate solution in styrofoam cup. Stir the solution and note the temperature change. When the reaction stops the temperature should return to room temperature.
H3C6H5O7 (aq) + 3NaHCO3(s) --> 3CO2 (g) + 3H2O (l) + Na3C6H5O7 (aq)
3. Dissolve ammonium chloride in water.
4. Dissolve ammonium nitrate in water.
5. Dissolve potassium chloride in water.
6. Dissolve urea in water.
7. Pass dry ammonium chloride over barium hydroxide octahydrate crystals.
8. Add sodium carbonate to ethanoic acid..
9. Add thionyl chloride (SOCl2) to cobalt (II) sulfate heptahydrate.
3.82 Heat of neutralization reactions
1. Dissolve 40 g of sodium hydroxide pellets in water and make up to 500 mL, a 2M solution. Prepare 500 mL of a 2M hydrochloric acid solution and leave to cool. Note the temperature of the solutions when cool. Quickly add the acid to the base and carefully stir with a thermometer. Note the maximum temperature reached. The increase of temperature should be 13oC. The volume of water has been doubled by adding one solution to the other, so the final solution contains 1 mole of OH- (aq) ions that reacted with 1 mole of H+ (aq) ions to form 1 mole of water molecules. Assume that the specific heat of this weak solution is the same as the specific heat of water.
2. Put 25 mL of 2.0 M sodium hydroxide solution into a polystyrene cup and measure the temperature of the sodium hydroxide solution. Stand the polystyrene cup in a beaker. Put 25 mL of 2.0 M hydrochloric acid in a measuring cylinder and measure the temperature of the acid. The sodium hydroxide and the hydrochloric acid should have the same room temperature. Pour the acid into the sodium hydroxide, stir the solution with the thermometer, and note the highest temperature of the mixture. Record the temperature difference reached in the mixture. After each temperature measurement rinse the thermometer and dry it with absorbent paper.
Repeat the experiment with 50 mL of 2.0 M sodium hydroxide and hydrochloric acid.
Repeat the experiment with 50 mL of 1.0 M sodium hydroxide and hydrochloric acid.
Compare the temperature differences reached in the 3 experiments.
3.83 Heat of reaction when metals displace copper
See diagram 3.2.83: Temperature rise of the reacting solution
1. Put 25 mL of 0.2 M copper (II) sulfate solution in a 100 mL plastic bottle fitted with a one-hole stopper and thermometer. Replace the stopper, invert the bottle and shake it gently. Record the temperature of this solution. Turn the bottle the right way up, remove the stopper and add 0.5 g of zinc dust. The quantity of zinc powder is in excess to ensure that all the copper (II) sulfate is used up in the reaction, so some zinc will remain when the reaction stops. Replace the stopper, invert the bottle, and shake gently. Record the highest temperature reached. Calculate the rise of temperature. This rise of temperature in not affected by the volume of 0.2 M copper (II) sulfate used for the experiment. For a 1 M solution, multiply the rise in temperature by 5 (5 × 0.2M = 1.0 M). The reactants lost energy to the solution. The temperature change is usually between 9oC and 10oC.
Zn (s) + Cu2+ (aq) ---> Zn2+ (aq) + Cu (s)
2. Repeat the experiment with 0.5 g of iron powder or iron filings. This amount is again in excess so that all the copper (II) sulfate will be used up in the reaction. The temperature change is usually between 6oC and 7oC. The zinc metal became zinc ions and copper ions became copper metal due to transfer of electrons from zinc metal to the copper ion. To get electrical energy, these electrons must flow in an external conductor, e.g. a wire, from the zinc to the copper. The potential or voltage will reflect the greater activity of zinc over copper. The current flowing will depend on the extent and rate of the reaction.
3.84 Electrical energy from a simple cell, displacement of copper by zinc
See diagram 3.2.84: Copper and zinc foil in a voltmeter
1. Put concentrated copper (II) sulfate solution in a beaker. Connect copper foil to the positive terminal, red wire, of a voltmeter and a zinc foil to the negative terminal, black wire. Simultaneously dip the two metals briefly into the copper sulfate solution. Record the readings on the voltmeter. The voltage falls to zero after a short time because black copper deposited on the zinc and caused the reaction to stop. When copper deposits on the zinc electrode, it prevents more zinc from entering the solution. This causes the voltage to fall to zero after a short time and the cell becomes "dead". You can separate the electrolytes to prevent the voltage fall by using a Daniell Cell that has a porous pot or 2. a salt bridge.
2. Pour concentrated copper (II) sulfate solution into a beaker. Connect a copper rod to the positive terminal of a voltmeter and a zinc rod to the negative terminal. Dip the two metals briefly into the copper (II) sulfate solution. Zinc dissolves and hydrogen bubbles form on the surface of the copper. The voltmeter reads 1.1 V, so electrons are moving from the zinc to the copper.
3.84.1 Electrochemical cell, voltaic cell, galvanic cell
Order online: Volta Cell Kit
Electrochemical cells form electricity from chemical reactions. The cell is made up of two half cells. Each half cell consists of an electrode in contact with an electrolyte. It is usually a metal in contact with a solution of one of its salts. The half cells are separated by a porous membrane or salt bridge which allows movement of ions between the half cells.
Put a piece of zinc metal in a zinc sulfate solution to form a Zn / Zn2+ half cell. The zinc metal atoms dissolve as zinc ions, leaving negative charges on the electrode until the increased charge stops the process. Connect the zinc foil to the negative terminal of a 5 V voltmeter.
Zn (s) ---> Zn2+ (aq) + 2e-
2. Put a piece of copper foil in a concentrated copper (II) sulfate solution to form a Cu / Cu2+ half cell. The copper metal ions in solution take electrons from the electrode and deposit on the copper electrode as copper atoms.
Cu2+ (aq) + 2e- ---> Cu (s)
3. Connect copper foil to the positive terminal of the 5 V voltmeter. Record any changes in the voltmeter reading. Note the maximum reading. Note any changes at the copper foil and the zinc foil. The voltage falls to zero after a short time because copper deposits on the zinc and causes the reaction to stop. Zinc metal becomes zinc ions and copper ions become copper metal. Electrons transfer from the zinc metal to the copper ions by moving from the zinc to copper along a wire. The potential difference or voltage reflects the greater activity of zinc over copper. The current flowing depends on the size and rate of the reaction.
Zn (s) + Cu2+ (aq) ---> Zn2+ (aq) + Cu (s)
3.84.2 Simple cell with different metals
Use alligator clips to connect a zinc strip to the negative terminal of a voltmeter and a copper strip to the positive terminal. Dip the metals in dilute sulfuric acid separated from one another. Note whether any reaction occurs. Note the reading on the voltmeter. The zinc dissolves in the acid. Hydrogen comes from the copper. From these reactions electrons flow through the external circuit producing a voltage of 1.1 V. If a 1.5 V light bulb is connected in series the light bulb glows for a short time then fades because the hydrogen bubbles collecting on the copper strip reduces the flow of electric current. Make the filament continue to glow by mechanically removing the bubbles or by adding the oxidizing agent potassium dichromate to the acid in the cell.
Repeat the experiment using magnesium, iron and lead in place of zinc. The greater the difference in activity of the two metals the greater the voltage. Test this by substituting magnesium, iron and lead for zinc, then record the voltages.
3.84.3 Simple electric cell, copper with zinc in dilute sulfuric acid
See diagram 3.84.3: Voltaic cell
1. Put a clean piece of zinc and a clean piece of copper in separate test-tubes of dilute sulfuric acid. Bubbles of hydrogen gas may come from the surface of the zinc. No bubbles appear on the copper.
2. Put the zinc and copper in the same beaker containing dilute sulfuric acid. Arrange the metals so that they touch, or connect them through an ammeter. Many bubbles of hydrogen come from the copper and few or no bubbles come from the zinc. The zinc is going into solution. The electrons move towards the copper. The zinc is the negative electrode. The chemical energy of the zinc causes the electrons to flow.
Zn (s) ---> Zn2+ (aq) + 2e-
At the surface of the copper, the electrons are transferred to the hydrogen ions of the sulfuric acid to form hydrogen gas. The copper acts as a sort of catalyst here. The copper is the positive electrode. The copper is positively charged with respect to the zinc.
2H+ + 2e- ---> H2
or
2H3O+ (aq) + 2e----> H2 (g) + 2H2O (l)
3.84.4 Simple galvanic cell, Zn and Cu in dilute HCl
Add zinc to dilute hydrochloric acid. Hydrogen forms and the test-tube becomes hot. Use this reaction as a source of electrical energy. Connect a strip of zinc and a strip of copper with wire to a 1.1 V light bulb. Pour the dilute acid into the beaker. Note the hydrogen bubbles liberated at the copper strip and note the glow of the light bulb. The chemical energy liberated as heat in the test-tube reaction is liberated as electrical energy in this simple galvanic cell.
Zn (s) + 2HCl (aq) ---> ZnCl2 (aq) + H26.
The half reaction equations are as follows:
Oxidation: Zno (s) - 2e- ---> Zn2+ (aq)
Reduction: 2H+ (aq) + 2e- ---> H2 6.
Zno (s) + 2H+ (aq) ---> Zn2+ (aq) + H2 (g) (Net reaction of electron transfer)
3.84.5 Voltaic cell with a salt bridge, Zn in ZnSO4 soln. / Cu in CuSO4 soln.
See diagram 3.2.85.2: Voltaic cell with a salt bridge
Put a zinc rod in a beaker containing zinc sulfate solution and put a copper rod in a beaker containing copper (II) sulfate solution. Connect the copper to the positive terminal of a voltmeter and the zinc to the negative terminal of the voltmeter. The reading of the voltmeter is zero. Make a simple salt bridge by soaking filter paper in a concentrated solution of an electrolyte, e.g. sodium chloride or potassium nitrate. Fix the filter paper to dip into the zinc sulfate and copper (II) sulfate solutions. The voltmeter shows that current is flowing. Read the voltmeter. Disconnect the voltmeter and substitute: 1.5 V light bulb, ammeter, conducting wire. Record the observations. Examine the electrodes after 2 minutes. The zinc corrodes and new copper has deposited on the copper electrode. The copper (II) sulfate solution loses some of its blue colour. A more permanent salt bridge is made from a glass U-tube filled with a 1 M potassium nitrate solution. The solution may be mixed with agar gel to keep it in the U-tube. Put cotton wool plugs at each end of the U-tube.
3.84.6l Standard electrode potential
See diagram 3.84.1.1: Standard electrode potential apparatus
Standard electrode potential of half reactions 1.0 molar solutions of some metals at 25oC is compared to the hydrogen half reaction with electrode potential assumed to be zero. A negative vale for E0 shows a poorer electron attracting ability than in the hydrogen half cell.
Half reaction and E0 in volts, V
Zn2+ (aq) + 2e- ---> Zn (s) -0.76 V (Zn2+ / Zn)
2H+ (aq) + 2e- ---> H2 0.00 V <----- Hydrogen (H+ / H2 Standard Hydrogen Electrode, SHE)
Cu2+ (aq) + 2e- ---> Cu (s) +0.34 V (Cu2+ / Cu)
The E0 value = higher value - lower value (+0.34) - (-0.76) = 1.1 V.
The saturated calomel electrode, SCE, (Hg2Cl2 / Hg) uses saturated KCl electrolyte
Hg2Cl2 (s) + 2e- --> 2Hg (l) + 2Cl-
E = 0.24 V
3.84.7 Pencil sharpener galvanic cell, Mg, Fe, NaCl, cola
Use a screwdriver to remove the steel blade from a magnesium pencil sharpener. Use crocodile clips to connect the positive wire from a voltmeter to the steel blade and the negative wire to the magnesium. Dip the steel and magnesium electrodes into a dilute sodium chloride solution in a Petri dish. Record the voltage measured by the voltmeter (about 1.3 V). Repeat the experiment by substituting cola for the sodium chloride solution. Record the voltage measured by the voltmeter (about 1.9 V). Observe tiny bubbles on the steel blade electrode.
This experiment shows that cola would function better as a salty bridge then sodium chloride solution, perhaps because acids in the cola contribute hydrogen ions that were reduced to hydrogen gas.
At the anode, oxidation occurs: Mg --> Mg2+ + 2e-
At the cathode, reduction occurs: 2H+ + 2e- --> H2
Repeat the experiment with more tan one galvanic cell (with cola) in series to try to light a LED, (light emitting diode).
3.85 Daniell cell, porous pot, Zn in ZnSO4 soln. / Cu in CuSO4 soln.
See diagram 3.2.85.1: Daniell cell
The Daniell cell uses a porous pot to prevent copper depositing on the zinc.
1. Put a porous pot in a beaker. Pour 0.5 M zinc sulfate solution into the porous pot. Pour concentrated copper (II) sulfate solution into the beaker and fill to the same level as the zinc sulfate solution. Make a cylinder shape with copper foil and place it in the beaker to surround the porous pot. Connect the copper foil to the positive terminal of a 1 to 5 V voltmeter. Connect a zinc rod to the negative terminal of the voltmeter and lower the zinc rod into the zinc sulfate solution. Note the reading of the voltmeter. Insert a 1.5 V light bulb in place of the voltmeter and note whether it lights. Insert an ammeter into the circuit to find the current flowing. Note whether you can vary the current by moving the copper electrode nearer to the zinc electrode, or by altering the surface area of the copper foil by raising and lowering it in the solution.
Oxidation occurs at the zinc anode, losing electrons to the electric circuit. Reduction occurs at the copper cathode, gaining electrons from the electric circuit.
At the anode, zinc atoms lose two electrons to become zinc ions, Zn2+. At the cathode, copper ions, Cu2+ receive two electrons from the electric circuit to become copper atoms. The zinc ions in solution gather around the zinc anode. The copper ions are removed from solution when they are reduced to copper metal and join the copper cathode. In the solution, sulfate ions move towards the zinc anode and copper (II) ions move towards the copper cathode. Positive ions, cations, move towards the cathode and negative ions, anions, move towards the anode. So charge is carried by ions in solution and carried by electrons in the electric circuit. The reactions continue until all the zinc in the anode joins the solution or all the copper ions are plated onto the cathode
2. Instead of a porous pot, repeat the experiment with a salt bridge made from a U-tube containing M potassium nitrate solution and agar gel.
3.86 Electrode potentials of metals
See diagram 3.2.86: Electrode potentials
1. Electrode potentials of metals are calculated from comparisons with the hydrogen cell under standardized conditions. However, you can use a copper and copper (II) sulfate solution as a standard. Lay filter paper soaked with copper (II) sulfate solution on clean copper foil. Use a short length of wire and crocodile clips to connect the copper foil to the positive terminal of the 1 to 5 V voltmeter. Similarly connect the specimen metal to the negative terminal of the voltmeter. Clean the surface of the specimen metal and press it firmly on absorbent paper. Record the voltage for this metal.
2. Test the following metals: magnesium, tin, lead, iron, zinc aluminium, and silver. After testing a metal, clean the copper again with a fine emery cloth and replace the absorbent paper, then test another metal. The metal surfaces must be clean and the absorbent paper must contain enough copper (II) sulfate solution for a steady reading on the voltmeter. If the voltage starts at a high value and then falls as a deposit forms on the metal, record the highest value.
3. Test aluminium after dipping it in concentrated hydrochloric acid then press it on absorbent paper to remove the layer of aluminium oxide. The voltage reading will start at a low value then increase as remaining aluminium oxide dissolves. Record the maximum value.
3.87 Lead accumulator cell, lead-acid rechargeable battery
See diagram 3.2.87: Lead cell accumulator
See 8.6: Prepare electrolyte for a lead accumulator cell
The most common motor car battery, the 12 volt battery, contains six cells connected in series, each of which produces 2 volts.
Charged positive plate: lead (IV) oxide, Discharged positive plate: lead (II) sulfate
Charged negative plate: lead, discharged negative plate: lead (II) sulfate
Electrolyte: sulfuric acid
1. Make two sheets of lead foil, 40 cm × 10 cm, and two lead strips 2 cm × 14 cm for terminals. Clean the lead thoroughly with steel wool and degrease them by soaking in 0.5 M sodium hydroxide solution for 15 minutes, then wash in deionized water. Fold the long sheets of lead tightly around the shorter strips to make good electrical contact. The projecting ends serve as terminals. Make a sandwich of alternating strips of lead foil and absorbent paper. Roll it tightly and put elastic bands around it. Mark one terminal positive and the other negative. Cover the roll of lead with a sodium sulfate solution made by dissolving 40 g of anhydrous sodium sulfate crystals in 200 mL of water. Use a 6 V battery charger or with any low voltage direct current supply giving up to 10 amps. Connect the positive terminal on the charger to the positive terminal on the cell. After charging for some minutes the cell lights a 1.5 V light bulb. The more times the accumulator is charged and discharged, the more efficient it becomes. When discharging at the negative terminal electrode, the electrons move through the circuit. When charging at the negative terminal electrode, the concentration of sulfuric acid increases.
Pb (s) + SO42- (aq) ---> PbSO4 (aq) + 2e-
When discharging at the positive terminal electrode, the concentration of sulfuric acid decreases. When charging at the positive terminal electrode, the electrons move through the circuit.
PbO2 (s) + 4H3O+ (aq) + SO42- (aq) + 2e- ---> PbSO4 (aq) + 6H2O (l)
When a lead cell accumulator is fully charged, the concentration of sulfuric acid is at maximum. When the accumulator is fully discharged, "flat battery", the concentration of sulfuric acid is at minimum. Use a battery hydrometer to read the relative density (specific gravity) of sulfuric acid in the electrolyte and check how charged the battery is. The density varies from about 1.28 in a fully charged battery to 1.15 in a discharged battery. The density of sulfuric acid purchased for use in accumulators is about 1.25 at 20oC.
Charging cell
Pb (s) + 2H2O (l) --> PbO2 (s) + 4H+ (aq) + 4e- (anode)
4H+ (aq) + 4e- --> 2H2 (g) cathode
Discharging cell
PbO2 (s) + 4H+ (aq) + 2e- --> Pb2+ (aq) + 2H2O (l) (anode)
Pb (s) --> Pb2+ (aq) + 2e- (cathode)
2. Charge and discharge a lead accumulator cell. Prepare a simple lead accumulator cell consisting of two pieces of lead dipping into sulfuric acid. To charge the cell connect three 1.5 volt batteries in series for two minutes. To discharge the cell connect a torch cell lamp, light globe. Note the time that the lamp remains lit. The cell is now discharged. Repeat the experiment by increasing the time of charge and noting the time until discharge. Draw a graph of the results with time charged (x axis) against time discharged (y axis)..
3. Fix two lead foil strips in a beaker and add 200 mL of 1 mol per litre sulfuric acid. Connect the lead electrodes to a power pack set at 2 V and switch it on for two minutes. The lead strip connected to the positive terminal becomes covered with brown lead dioxide. Disconnect the power pack and connect the lead strips to a torch battery. The battery glows but the brown leads dioxide on the positive terminal does not disappear.
Repeat the experiment with increasing charging times. The time the battery glows increases with charging time up to 30 seconds then hardly changes.
Repeat the experiment with different charging voltages. Different charging voltage makes hardly any difference in the time the battery glows. However, at high charging voltages hydrogen is produced at the negative electrode and oxygen gas at the positive electrode.
Charging
At the positive electrode: Pb (s) + 2H2O (l) ---> PbO2 (s) + 4H+ (aq) + 4e-
At the negative electrode: 2H+ (aq) + 2e- ---> H2 (g)
Also lead reacts with the sulfuric acid to produce lead sulfate
At the positive electrode: PbSO4 (s) + 2H2O (l) ---> PbO2 (s) + 4H+ (aq) + SO42- + 2e-
At the negative electrode: PbSO4 (s) + 2e- ---> Pb (s) + SO42- (aq)
So sulfuric acid is produced during charging and is consumed during discharging. As sulfuric acid has about twice the density of water, the density of the electrolyte shows the state of charge of the battery.
4. When the battery is fully charged, the specific gravity (relative density) = 1.280, electrode A is lead and electrode B is lead dioxide. When the battery is discharging, electrode A changes from lead to lead sulfate, electrode B changes from lead dioxide to lead sulfate, and the concentration of sulfuric acid decreases. When the battery is being charged, these processes are reversed. The concentration of sulfuric acid suggests the state of charge of the battery so this concentration can be measured with a battery hydrometer.
Electrode A: Pb + SO42- ---> PbSO4 + 2e-
Electrode B: PbO2 + 4H3O+ + SO42- + 2e- ---> PbSO4 + 6H2O
In a motor car battery, the electrodes have a coat of lead (II) oxide (PbO) and lead powder (Pb). In the electrolyte, electric current converts the PbO to Pb on the negative plate, and the PbO to lead (IV) oxide (lead peroxide) PbO2 on the positive plate.
Discharging --->
PbO2 + 2H2SO4 + Pb < = > 2PbSO4 + 2H2O
<--- Charging
If you pass electricity through the battery after it is fully charged, "gassing" occurs, i.e. water is decomposed into hydrogen and oxygen gas. Never smoke or allow a naked flame near a charging battery because it may produce the inflammable gas hydrogen.
Also, the lead-acid battery is dangerous because it can produce very high currents and contains sulfuric acid that may be spilt. It is the only common wet battery.
3.88 Leclanché cell, dry cell
See diagram 3.2.88: Leclanché cell
A Leclanche cell (Georges Leclanché 1839-1882) is a primary voltaic cell with a carbon rod anode, zinc cathode, dilute ammonium chloride solution electrolyte and e.m.f. approximately 1.5 volts.
Zn + H2SO4 --> (discharge) ZnSO4 + H2O + H2 (g)
A torch "battery" is the dry cell version of the Leclanché cell. It has manganese dioxide [manganese (IV) oxide] around the carbon rod to oxidize hydrogen and so depolarize the anode.
2MnO2 + H2 --> Mn2O3 + H2O
The electrolyte is in the form of a water paste so the dry cell is not really "dry".
1. Add 1 mL of phenolphthalein indicator to 2 cm of ammonium chloride solution in a shallow dish. Attach a crocodile clip and conducting wire to a carbon rod from a dry cell battery and a piece of zinc foil. Join the conducting wires. Dip the carbon rod and zinc foil into the solution to act as electrodes. Note any changes around the electrodes.
2. The voltage from a single commercial "battery" is usually about 1.5 V, however maximum current varies with the type of battery. Batteries for radios produce a small current for long periods, e.g. 4 amps. Batteries for torches or flashlights produce a large current for short periods, e.g. 6 amps. Connect an ammeter directly across the terminals of different types of batteries.
3. Use a saw to cut through a used dry cell lengthways. The central carbon rod is the positive electrode. The zinc container is the negative electrode. The moist paste contains the electrolyte ammonium chloride that acts like an acid in dissolving zinc and the black manganese (IV) oxide that acts as a depolarizer by oxidizing hydrogen. It slowly oxidizes hydrogen to water so that any hydrogen formed cannot block the flow of electric charge. If the power from a dry cell fades after continual use, the power may be restored after some hours because of the depolarizing action of the manganese (IV) oxide. In a used battery the zinc appears eaten away because some of it has gone into solution.
4. Cut down one side of a dry cell battery with a hacksaw. Scoop out the contents and put them in a beaker half full of hot water, then stir. Put cotton wool in the neck of a funnel to act as a filter. Pour the contents of the beaker into the funnel. The black substance remaining in the cotton wool is mainly manganese dioxide. Evaporate the clear solution that had passed through the filter. White ammonium chloride, NH4Cl, and some zinc chloride, ZnCl2, remain after evaporation..
5. Mix 4 g of carbon black, powdered carbon, with 10 g manganese (IV) oxide, the oxidant. Stir in ammonium chloride solution to make a thick paste. Make a zinc can made from zinc foil rolled into a cylinder or a cleaned zinc can from an old dry cell. Cut absorbent paper to make a cylinder to line the zinc can. Place the mixture on the absorbent paper, compress it into a cylinder and wrap the absorbent paper around the cylinder so that it just fits inside the zinc can. Pour ammonium chloride solution between the paper and the zinc to ensure good contact. Press the mixture into the zinc can firmly and tightly. Attach a crocodile clip and conducting wire to a carbon rod from an old dry cell. Press the carbon rod down the centre of the mixture so that it almost touches the bottom of the zinc can. Connect another crocodile clip and conducting wire to the zinc container. This cell should light a 1.5 V light bulb and run a small 1.5 V electric motor. The carbon lowers the internal resistance of the cell. Test the voltage and the current from the cell.
Zn (s) ---> Zn2+ (aq) + 2e- (oxidation)
2NH4+ (aq) + 2MnO2 (s) + 2e- ---> Mn2O3 (s) + 2NH3 (aq) + H2O (aq) (reduction)
3.89 Movement of copper and chromate ions
See diagram 3.2.89: Movement of ions 1
Copper chromate forms two coloured ions, the blue-green positive copper ion and the orange negative chromate ion. Prepare copper chromate by adding 100 mL of 1 M copper (II) sulfate solution to 100 mL of 1 M potassium chromate solution. Use a Buchner funnel fitted with a filter pump to separate the copper chromate precipitate. Wash the precipitate with demineralized water then transfer it to a beaker. Dissolve the precipitate in the minimum volume of dilute hydrochloric acid. Dissolve excess urea in the copper chromate solution. Add dilute hydrochloric acid to a U-tube until it is one third full. Position the jet of a pipette full of copper chromate solution at the bottom of the U-tube. Slowly deliver the copper chromate solution so that it pushes the hydrochloric acid up and forms a separate layer below. Take out the pipette carefully to avoid mixing. The carbon electrodes must be in contact with the hydrochloric acid and also connected to a 20 V d.c. supply. After some minutes, note the blue-green colour of the copper on the negative side, and the orange chromate colour on the positive side. The boundaries of these coloured ions will move very slowly towards the electrodes.
3.90 Movement of ions between microscope slides, Cu2+ ions, CO2+ ions
See diagram 3.2.90: Movement of ions 2
Show the movement of positive coloured ions towards a negative electrode. The electrolyte is held in a strip of filter paper sandwiched between two microscope slides. Use carbon rod electrodes to lead the current through the filter paper. With a 10 to 20 V d.c. supply, use the width of the slide, not as in the diagram, because greater voltage is needed if the length of the slide is used. Cut a strip of dry filter paper 1 cm wide. Make a pencil mark across the centre of the paper. Moisten the paper with tap water so that it is damp but not wet. Use a fine capillary tube to apply the coloured ion solution, e.g. Cu2+ or CO2+ ions, along the pencil mark. Fix the strip of filter paper between the two slides and fold the ends around the carbon rods. Use a paper-clip to the slides together. Connect the carbon rods to the 20 V d.c. supply. Wait for some minutes and observe the coloured ions moving towards the negative electrode. By contrast, repeat the experiment with potassium manganate (VII) and observe the coloured permanganate ion moving towards the positive electrode. If you want to make a dispenser for the solution of coloured ions, fold a strip of filter paper 1 cm wide around a thin piece of plastic material to form a firm wick. Wedge the wick between a split cork and put it in a test-tube containing the solution of coloured ions. First touch the wick on absorbent to remove excess solution and then lightly touch it on the pencil mark.
3.91 Size of particles and rate of reaction
See diagram 3.2.91: Size of particles and rate of reaction
1. Use a hammer to break marble chips into four sizes: 1.1 coarse powder, 1.2 half a rice grain, 1.3 rice grain, 1.4. original marble chips. Put 2 g of each size separately into four test-tubes. Blow up four balloons several times to stretch them. Put 5 mL of dilute hydrochloric acid into each of the four balloons. Slip the mouths of the balloons over the tops of the test-tubes but do not let any acid enter the test-tubes. Tip the acid from each balloon into the attached test-tube. Note which balloon is the fastest and the slowest to expand because of the production of carbon dioxide. The coarse powder produces carbon dioxide in the shortest time.
2. Repeat the experiment without balloons but with four conical flasks on a sensitive top balance. Add 5 mL of dilute hydrochloric acid to 2 g of coarse powder and note the loss in weight every 30 seconds. Then continue the experiment with the other sizes of marble chips.
3.92 Concentration and rate of reaction, sodium thiosulfate (hypo) with dilute hydrochloric acid
See diagram 3.2.92: Black cross no longer visible
1. Dissolve 5 g sodium thiosulfate crystals in 500 mL water. Add 5 mL hydrochloric acid to 50 mL the sodium thiosulfate solution.
2. Dissolve 20 g of sodium thiosulfate in 500 mL of water and put 50 mL of the solution in a container. Place the container on a black cross marked on a sheet of paper. Add 5 mL of dilute hydrochloric acid to the container and note the time. Look down through the solution and note when the black cross is no longer visible. Sulfur is produced during the reaction making the solution cloudy. Repeat the experiment with 40 mL of sodium thiosulfate solution and 10 mL of water. Add 5 mL of dilute hydrochloric acid. The time when the black cross is no longer visible is greater. Repeat the experiment with 30 mL of sodium thiosulfate solution and 20 mL of water. The time when the black cross is no longer visible is still greater. Repeat the experiment with 20 mL of sodium thiosulfate solution and 30 mL of water. Use graph paper to plot the volume of the thiosulfate solution (concentration) against time taken for the reaction.
Na2S2O3 (aq) + 2HCl (aq) ---> H2O (l) + SO2 (g) + S (s) [The S (s) causes the solution to become cloudy.]
(S2O3)2- (aq) + 2 H+ (aq) --> H2O (l) + SO2 (g) + S (colloidal)
3.93 Temperature and rate of a reaction
Increasing the temperature by 10oC usually doubles the rate of reaction. The increased rate of reaction occurs because the average collision between particles occurs with greater energy. Also, more collisions per second may occur.
Dissolve 20 g of sodium thiosulfate in 500 mL of water and put 10 mL of the solution in a container. Add 40 mL of water to the container. Note the room temperature and add 5 mL of dilute hydrochloric acid. Look down through the solution and note when the black cross is no longer visible. Repeat the experiment at 30oC, 40oC, 50oC and 60oC. Use graph paper to plot the temperature of the sodium thiosulfate solution against time taken for the reaction. The time when the black cross is no longer visible becomes less as the temperature rises.
3.94 Catalysts and rate of reaction
See diagram 3.2.94: Catalysts and rate of reaction
1. Put a cube of sugar in a glass or aerated water, e.g. fizzy lemonade. The drink fizzes more because the sugar cube provides more sites for the formation of carbon dioxide bubbles on the sharp corners of the sugar crystals. The sugar is said to act as a physical catalyst.
2. Fill a burette with water and invert it in a container of water to measure the volume of a gas in the burette by downward displacement of water. Use a flask or test-tube fitted with a one-hole stopper and bent delivery tube. Add 2 mL of 20 volumes to 50 L of water in the flask. Note the time then immediately add 1 g of manganese (IV) oxide to the flask, close the stopper and adjust the end of the delivery tube inside the burette. Note the volume of hydrogen gas in the burette at intervals of 15 seconds. Repeat the experiment with copper (II) oxide, nickel oxide and zinc oxide. Use graph paper to plot the volume of oxygen gas produced every 15 seconds against the time of the reaction. Manganese (IV) oxide is the best catalyst for this reaction.
3.95.01 Prepare starch solution
Prepare a clear solution of starch by adding a mixture of 1g starch in 10 ml of water to 500 mL of boiling water, then leave the solution to cool to room temperature.
3.95 Breakdown starch to sugars
9.130 Hydrolysis of starch by salivary amylase (ptyalin)
1. Prepare a clear solution of starch by adding a mixture of 1g starch in 10 ml of water to 500 mL of boiling water, then leave the solution to cool to room temperature.
2. Put 10 mL of dilute starch solution into a test-tube. Add to this 1 mL of saliva and stir this into the starch solution. Record the time of adding the saliva. At 5 minute intervals remove three drops with a dropper and put them on a clean white tile taking care to keep them from running into other. The dropper must be washed between each test. Put some iodine solution on each drop. The decreasing intensity of the blue colour shows the decreasing amount of starch. To tests for increasing amounts of sugar, put three drops of the reaction mixture into a small test-tube. Add 3 mL of Fehling's solution and heat this mixture almost to boiling point. The test should show that there is more sugar after boiling.
3.96 Breakdown ethanol to ethene (ethylene)
See diagram 3.2.96: Breakdown of ethanol
Push cotton wool soaked in methylated spirit to the bottom of a hard glass test-tube. Pack small pieces of porous pot, unglazed porcelain, in the middle of the test-tube. Fit a delivery tube to collect ethene gas over water in a receiving test-tube. With the hard glass test-tube in a horizontal position, heat the porous pot strongly, then gently heat the cotton wool to produce ethanol vapour. The ethanol vapour breaks down over the hot porous pot to produce ethene gas and water vapour. Ethene is insoluble in water and collects in the receiving test-tube. Collect three receiving test-tubes full of ethene then immediately disconnect the delivery tube when you stop heating to avoid a suck back of water on the hot porous pot. In test-tube 1, burn ethene with a lighted taper. Shake test-tube 2 with drops of dilute potassium manganate (VII) solution and sodium carbonate solution. The colour disappears. Shake test-tube 3 with bromine water. The colour disappears.
C2H5OH (l) ---> C2H4 (g) + H2O
3.97 Breakdown polymers using heat
See diagram 3.2.97: Breakdown of polymers | See diagram 1.13a: Simple fume hood
Do this experiment in a fume cupboard, fume hood. Put pieces of polymer, e.g. Perspex or polystyrene, in a hard glass test-tube. Fit a one-hole stopper with a delivery tube. Be ready to cool the receiving test-tube with cold water because the fumes produced by the reaction are harmful. Slowly heat the hard glass test-tube. The polymer melts then produces vapours to be collected in the receiving test-tube. Keep heating until all the fumes from the reaction are condensed to the liquid in the receiving test-tube. The polymer has been broken down by heat to smaller molecules.
3.98 Elements in foods
See diagram 3.2.98: Find nitrogen in foods
1. Collect small pieces of different foods together, such as cheese, bread, flour, sugar, leaves, maize. Heat a piece of each, about the size of a rice grain, on a tin lid or metal bottle top. Hold the lid with tongs. Black carbon is always left on the lid
2. Heat small amounts of food with copper oxide in a small test-tube. Copper oxide releases oxygen to the food. Test the gas in the test-tube with limewater by withdrawing a little gas in a teat pipette and bubbling the gas through the limewater. The limewater turns milky indicating the presence of carbon dioxide. Also, water is condensed on the cooler parts of the tube.
3. Put a small amount of crushed food in a test-tube and add three times that volume of soda lime. Mix the substances thoroughly then heat the test-tube. Use your hand to fan gases from the mouth of the test-tube towards you to smell ammonia at the mouth of the tube. Test the gases with wet blue and red litmus paper. The red litmus paper turns blue. If the food gives off ammonia gas, the nitrogen in the ammonia must have come from the food.
4. Mix separately cane sugar, vegetable oil and egg white with soda lime, then heat the mixtures. Note any smell of ammonia at the mouth of the test-tube containing the egg white. The nitrogen in the ammonia came from the protein in the egg white.
5. Heat a mixture of 0.5 cm of sucrose and 1.0 cm of concentrated sulfuric acid gently for 2 seconds and then leave to stand. Note the vigorous reaction and the colour change from white sugar to black carbon.
C12H22O11 + (H2SO4 catalyst) --> 12C + 11H2O
3.99 Gases from wood
See diagram 3.2.99: Gases from wood
Heat sawdust in the hard glass test-tube, gently then strongly until almost red-hot. Ignite the wood gas coming from the jet. The black residue is charcoal, carbon.