School Science Lessons
(topic14)
2025-02-19a
Thermochemistry, Heat of reaction
Contents
Experiments
14.0 Exothermic and endothermic reactions
14.1.0 Exothermic reactions, Experiments
14.2.0 Endothermic reactions, Experiments
14.15.0 Reaction of iron, Fe
Exothermic reactions
14.3 Heat of reaction, Hess's law
14.4 Heat of rusting, steel wool
14.5 Heat of neutralization with a calorimeter
14.6 Heat of displacement reaction, Experiments
14.7 Heat of reaction, potassium permanganate with glycerol
14.8 Heat of reaction, potassium permanganate with ethanol
14.9 Heat of reaction, chromium (VI) oxide with ethanol
14.10 Heat of reaction, potassium with diethyl ether
14.11 Heat of reaction, crystallization of sodium acetate, sodium acetate heat pack
14.12 Heat of reaction, iron with oxygen gas, iron powder heat pack, hand warmers
14.13 Heat of reaction, crystallization of sodium thiosulfate, sodium thiosulfate heat pack
14.14 Reaction of ammonium salts and potassium salts with water
14.21 Combustion of potassium nitrate and sugar mixture
14.22 Dehydration of sugar with sulfuric acid
17.3.12 Burning sugar cube, combustible cube
17.7.20 Elephant's toothpaste, Hydrogen peroxide with yeast
12.3.15 Thermite reaction, Al and iron oxide
12.6.6 Citric acid with sodium hydrogen carbonate
8.2.1 Sugar with potassium chlorate, spontaneous combustion
Endothermic reactions
14.16 Reaction of ammonium carbonate with ethanoic acid
14.17 Reaction of ammonium nitrate with water, Ammonium nitrate cold pack
14.18 Reaction of potassium nitrate with water
14.19 Reaction of copper (II) sulfate solution with magnesium
14.20 Reaction of urea with water
Drinking bird heat engine, drinking duck, dippy bird, (Experiment)
14.15.0 Reaction of iron, Fe
14.15.1 Burn steel wool
14.15.2 Black iron oxide is a mixed base, Fe3O4
14.15.3 Detect iron in fruit juice using black tea
14.15.4 Heat hydrated iron chlorides
14.15.5 Heat iron (II) sulfide
14.15.6 Heat iron filings with powdered sulfur
14.15.7 Heat iron with sulfur
14.15.8 Iron displaces hydrogen from hydrochloric acid
14.15.9 Iron displaces hydrogen from sulfuric acid
14.15.10 Iron (II) sulfate (ferrous sulfate) with sodium carbonate
14.15.11 Iron (II) sulfate (ferrous sulfate) with ammonia
14.15.12 Iron (II) sulfate (ferrous sulfate) oxidation to iron (III) sulfate (ferric sulfate)
14.15.13 Iron (III) sulfate (ferric sulfate) reduction to iron (II) sulfate (ferrous sulfate)
14.15.14 Oxidation of iron (II) salts
14.15.15 Reaction of iron (II) salts & iron (III) salts
14.1.0 Exothermic reactions
The reactants of exothermic reactions form products with rise in temperature.
Exothermic reactions give out heat energy
In an exothermic reaction, the amount of energy absorbed when chemical bonds brake is less than the amount of energy released when chemical bondsform.
The reaction is exothermic if the energy absorbed in bond breaking < energy released when bonds form.
In an exothermic reaction, water containing the reacting ions become hotter, because of the heat energy released by the ions.
Example of exothermic reactions include; Combustion of gases, mixing strong acids and bases, adding water to strong acids and any anhydrous salt. dissolving laundry detergent in water, respiration, freezing water, burning candles, mixing cement and water.
Uncontrolled exothermic reactions cause damaging fires and explosions.
14.2.0 Endothermic reactions
If less energy is needed to break the bonds in a reaction than the energy released by making the bonds, the reaction is endothermic.
The reaction is endothermic if the energy absorbed in bond breaking > energy released when bonds form.
In an endothermic reaction, water containing the ions becomes colder, because the ions absorb heat energy.
Energy is measured in joules (J) or kilojoules (kJ).
In an endothermic reaction, the energy absorbed when chemical bonds break is greater than the energy released when chemical bonds form
C + y kJ --> D, i.e. C -->D
DH = +y kJ / mol C.
In an endothermic reaction, the amount of energy absorbed when chemical bonds are broken is greater than the amount of energy released when chemical bonds are formed.
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.
If less energy is needed to break the bonds in a reaction than the energy released by making the bonds, the reaction is endothermic.
Endothermic dissolving processes of a solid in a solvent are associated with low solubility.
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. Pass dry ammonium chloride over barium hydroxide octahydrate crystals.
7. Add sodium carbonate to ethanoic acid.
8. Add thionyl chloride (SOCl2) to cobalt (II) sulfate heptahydrate.
14.3 Heat of reaction, Hess's law
Heat of reaction (enthalpy of reaction) DH, can be expressed as heat of combustion, heat of crystallization, heat of formation, heat of neutralization, heat of solution.
DH is the heat change for a reaction in kJ per mol of reactant or product.
If DH is negative the reaction is exothermic.
If DH is positive the reaction is endothermic.
Hess's law, law of additivity, law of constant heat summation, states that the overall energy change from reactants to products is the same by direct reaction or any other route.
So if equations can be added to give a final equation the heats of reaction of each equation can be added to give the heat of reaction of the final equation.
This law is an example of the principle of conservation of energy.
For example:
Reaction 1 H2SO4 (10 M) + NaOH (1 M) --> NaHSO4 (0.5 M) + H2O, dH1
Reaction 2 H2SO4 (10 M) + solvent --> H2SO4 (1 M) dH2
Reaction 3 H2SO4 (1M) + NaOH (1 M) --> NaHSO4 (0.5 M) + H2O, dH3,
dH1 = dH2 + dH3.
Thermal capacity
Thermal capacity (heat capacity) is the ratio of how much heat is supplied to the resulting rise in temperature.
Specific heat capacity refers to the mass of the substance and is measured in J K-1 kg-1.
Molar heat capacity refers to amount of substance and is measured in J K-1 mol-1 (K = kelvin, oK = - 273.15 Co).
Energy of reactions
1. Pairs of atoms may be bound together by the sharing of electrons between them in a covalent bond.
Two or more atoms bound together by one or more covalent bonds form a molecule, with definite size, shape and arrangement of bonds.
An atom or group of atoms covalently bound together may gain or lose one or more electrons to form ions.
2. Forces weaker than covalent bonding exist between molecules.
The structure of a metal involves positive ions embedded in a sea of electrons.
3. All chemicals contain two kinds of energy, kinetic energy of the particles and the energy stored in their chemical bonds.
Energy is absorbed when chemical bonds are broken and energy is released when chemical bonds form.
In a chemical reaction, some bonds are broken in the reactants and some bonds form in the products.
"Be Careful! When diluting strong acids always slowly add ACID to WATER.
Never add water to acid.
1. Add concentrated sulfuric acid very slowly to water.
Stir the mixture thoroughly each time a small amount of acid is added.
Note any change in temperature.
100 mL H2SO4 + 100 mL H2O (20o C) → H3O++ HSO4– (130o C, after about one minute!)
2. Pass hydrogen chloride gas into water.
Note any change in temperature.
3. Add acetic acid to water.
Acetic acid, a weak acid, produces less heat than the strong acids sulfuric acid and hydrochloric acid.
4. 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.
Note any changes in the thermometer reading.
5. Pour 2 cm of water into a test-tube.
Add concentrated sulfuric acid drop by drop down the side of the tube.
BE CAREFUL!
Stir gently with a thermometer after the addition of each drop.
Record the changes in temperature of the solution.
6. Pour dilute sodium hydroxide solution into the test-tube.
Test with litmus paper.
Red litmus paper turns blue.
Add dilute hydrochloric acid until the litmus paper turns to a colour between blue and red.
Record the changes in temperature of the solution.
Energy from chemical reactions
The heat of reaction, DH (δH) is the heat change for a reaction in kJ per mol of reactant or product.
This is also called the enthalpy of reaction.
For endothermic reactions, DH (δH) is positive.
For exothermic reactions, DH (δH) is negative.
A --> B + xkJ
i.e. A --> B, DH is -xkJ / mol A
Be careful! The reactions may be vigorous!.
Experiments
1. Use a test-tube containing a thermometer.
Record the initial temperature.
Put anhydrous copper (II) sulfate powder in the test-tube.
Add water drop by drop.
Record the changes in temperature of the solution.
2. Put white anhydrous copper (II) sulfate powder to a depth of 1 cm in a test-tube.
Hold a thermometer with the bulb in the powder.
Add water drop by drop.
Record the changes of the thermometer reading.
CuSO4 (s) (white) + 5H20 (l) --> CuSO4.5H2O (s) (blue)
14.4 Heat of rusting
1. Moisten some steel wool with iron chloride solution to accelerate rusting.
Wrap the bulb of a thermometer in the steel wool.
Hang in a draught free place.
Note the temperature changes as rust forms.
2. Roll some steel wool into a ball and weigh it.
Use tongs to hold the ball of steel wool over a sheet of paper.
Heat the steel wool over a burner until red-hot.
Remove the burner and blow gently on the red hot steel wool until it stops burning.
Weigh the burned steel wool and any fragments that have fallen on to the sheet of paper.
The weight is greater, because the iron oxide that forms is heavier than the steel wool.
14.5 Heat of neutralization with a calorimeter
See diagram 3.5: Heat of neutralization
Neutralization heat is the formation heat of one mole of water molecules from H+ and OH- ions.
The measured value of neutralization heat is thus Q / 005.
If no insulated cup is available in the laboratory, the following simple apparatus can be used instead.
The heat of neutralization reaction of strong acids with bases is -58 kJ / mol.
The heat of neutralization = the heat of formation of one mole of water molecules from the ions.
Since the reacting particles release energy by giving this to the solution, the energy change can be written:
H (change of heat) = x -kJ / mol, the heat released when one mole of hydrogen ions (H+) reacts with one mole of hydroxide ions (OH-).
Both 1 mol / L hydrochloric acid and 1 mol / L sodium hydroxide have a density of 1 g / mL.
The mass of 50 mL of 1 mol / L HCl =o 005 kg, i.e. m1 = 005 kg, so is the mass of 50 mL of 1 mol /L NaCl solution, i.e. m2 = 005 kg.
Consequently, the heat released by this neutralization reaction can be calculated as follows:
Quantity of heat = mass × specific heat × change in temperature.
Q = (005 + 005) × 42 × (t1 - t2).
H+ (aq) + OH- (aq) ---> H2O (l).
Experiments
1. Make a simple calorimeter by using a plastic cup inside an insulated box.
Pour 50 mL of 2 M hydrochloric acid into the plastic cup and record the initial temperature.
Pour 50 mL of 2 M sodium hydroxide solution into a beaker and note the original temperature.
When the initial temperatures are the same, add sodium hydroxide solution to the plastic cup while stirring constantly with the thermometer.
Record the highest temperature.
Assume that the specific heat capacity of this weak solution is the same as water = 42 kJ / kg / oC.
Also, assume that all the heat from the reaction heats the water, raising the temperature from t1 to t2.
Calculate how much heat that would be produced if 1 M of sodium hydroxide is neutralized by 1 M of hydrochloric acid.
2. 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 stir with a thermometer.
Note the maximum temperature reached.
The increase of temperature should be 13 oC.
You have doubled the volume of water 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.
3. Place some strips of paper in the bottom of a large beaker, and then stand a small beaker on the paper strips.
Stuff the space between the two beakers with a lot more strips of paper.
Cover the mouth of the large beaker using a piece of cardboard to reduce heat loss.
Repeat the experiment with 1 mol / L ethanoic acid (acetic acid) replacing hydrochloric acid.
The determined value of neutralization heat will be lower, because ethanoic acid is a weak acid, mainly in molecular form in aqueous solution.
So some energy released by the neutralization must be used to ionize the ethanoic acid molecules.
4. Repeat the experiment with equal concentrations of other strong acids and bases.
The heat of neutralization, J, is the same, because the same chemical reaction above occurs.
Na+ + OH- + H+ + Cl- --> Na+ + Cl- + H2O + J joules
OH- + H+ --> H2O + J joules.
5. Repeat the experiment using 2 M ethanoic acid, acetic acid, HAc.
The heat of neutralization, J1, is lower.
Ethanoic acid is a weak acid mainly in molecular form.
Some energy from the heat of neutralization is used to ionize the ethanoic acid molecules.
HAc + J1 --> H+ + Ac-
So the heat evolved = J - J1
CH3COOH (aq) + NaOH (aq) --> CH3COONa (aq) +H2O (l).
14.6 Heat of displacement reaction
See diagram 3.2.83: Heat from a displacement reaction
1. Use 02 M copper (II) sulfate solution and zinc or iron.
Use a plastic container, insulated with a polystyrene jacket for insulation, with a one-hole stopper fitted with a thermometer.
Put 25 mL of 02 M copper (II) sulfate in the container.
Replace the stopper invert and shake gently.
Record the initial temperature of this solution.
Add 05 g of zinc dust.
The amount is more than needed to ensure that all the copper (II) sulfate is used up in the reaction.
Replace the stopper, invert the bottle and shake gently.
Record the highest temperature reached.
The temperature difference should be about 10 oC.
Zn (s) + Cu2+ (aq) --> Zn2+ (aq) + Cu (s).
2. Put 25 mL 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.
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 9 oC and 10 oC.
Zn (s) + Cu2+ (aq) --> Zn2+ (aq) + Cu (s).
3. 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 6 oC and 7 oC.
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.
14.7 Heat of reaction, potassium permanganate with glycerol
BE CAREFUL! This is a dangerous experiment.
Use very small quantities and follow your safety rules.
Remember that strong oxidants should be stored separately from flammable organic chemicals.
Put a few drops of glycerol on a few fine crystals of potassium permanganate in an evaporating basin.
Observe the effect of heat of reaction.
14.8 Heat of reaction, potassium permanganate with ethanol
BE CAREFUL! This is a dangerous experiment so use very small quantities and follow your safety rules.
Remember that strong oxidants should be stored separately from flammable organic chemicals.
Experiment
Add alcohol to cotton wool in an evaporating basin.
Dip a glass rod into concentrated sulfuric acid, then touch crystals of potassium permanganate.
Touch the cotton wool with the glass rod.
BE CAREFUL! The heat from the formation of manganese (VII) oxide on the glass rod ignites the alcohol.
14.9 Heat of reaction, chromium (VI) oxide with ethanol
This experiment is too dangerous for schools.
BE CAREFUL!
Use very small quantities and follow your safety rules.
Remember that strong oxidants should be stored separately from flammable organic chemicals.
Experiment
Add ethanol to a piece of mineral wool in an evaporating dish.
Drop a very small amount of chromium (VI) oxide on the mineral wool.
BE CAREFUL! The heat of reaction ignites the alcohol.
Red chromium (VI) oxide (chromium trioxide), is reduced to green chromium (III) oxide.
2CrO3 (s) + C2H5OH (l) + 11/2O2 (g) ---> Cr2O3 (s) + 2CO2 (g) + 3H2O (l).
14.10 Heat of reaction, potassium with diethyl ether
This experiment is too dangerous for schools.
BE CAREFUL! THIS IS A DANGEROUS EXPERIMENT!
Use very small quantities and follow your safety rules.
Remember that strong oxidants should be stored separately from flammable organic chemicals.
Diethyl ether has an ignition temperature about 80 oC. A mixture of diethyl ether vapour and air is explosive!
Experiment
Put a very small piece of potassium metal and a few drops of diethyl ether in a beaker covered with a watch glass.
Pour water into the large beaker.
BE CAREFUL! The potassium metal reacts violently with the water producing heat.
The heat ignites the hydrogen gas produced in the reaction.
Then the heat ignites the diethyl ether.
K (s) + 2H2O (l) ---> 2KOH (s) + H2 (g).
14.11 Heat of reaction, crystallization of sodium acetate, sodium acetate heat pack
Reusable hand warmers have a supersaturated sodium acetate solution inside the packet.
As the solution crystalizes, it releases energy as heat.
To reuse these hand warmers, boil the closed packet in water to return the crystals to a liquid.
1. Prepare a supersaturated solution of sodium ethanoate-3-water to make a "heat pack".
Dissolve 125 g sodium ethanoate-3-water, CH3CO2Na.3H2O, in 12.5 mL water.
Heat to form a clear solution, cover with a watch glass and leave to cool.
Hold a watch glass in the palm of your hand, pour in some solution then add a few crystals of sodium ethanoate-3-water.
The supersaturated solution immediately crystallizes.
Feel the heat given out.
The exothermic property of the crystallization of saturated solutions is used in "heat packs" sold by chemical suppliers.
2. Heat packs provide instant, portable and reusable heat and generate heat for two to three hours.
To reactivate after use, boil the heat pack in water until it is clear and then remove and let cool.
Heat packs may contain sodium acetate, which will freeze at 54 oC in an open container.
However, when this solution is in a sealed container, the solution can be cooled below this temperature, as low as -10 oC.
Flexing a metal "trigger" within the sealed container causes a few molecules of liquid to crystallize, which starts a chain reaction causing the supercooled solution to change from a liquid to a solid as crystals form.
This phase change causes the pack to give out heat.
When the heat pack contents crystallize, its temperature returns to its freezing point.
This supercooled solution can be stored for extended periods and still crystallizes on demand,
Once the unit has given off all of its heat, it is then recycled by heating it in boiling water.
The crystals dissolve in their own water of crystallization, the heat pack returns back to a liquid state and then cools below its freezing temperature.
It is then ready to be activated again.
14.12 Heat of reaction, iron with oxygen gas, iron powder heat pack, disposable hand warmers
1. Disposable heat packs are used for transporting small animals that need heat to survive the journey, e.g. sugar gliders.
Open the outer wrapper and remove the inner pad.
Shake the contents in open air and heat will begin to be generated in 4-5 minutes.
Place the heat source in shipping containers.
After use, dispose of an outer wrapper and expired heat pack.
The contents are high grade iron powder that undergoes rapid rusting with heat as a by-product, activated charcoal powder, cellulose, zeolite and water.
2. Use an "instant hot pack".
Remove inner pack.
Squeeze or shake several times.
Allow a few minutes to warm up.
Keep covered in pocket, glove or clothing for maximum warmth.
Caution: Store in cool dry place.
The hot pack has an outer plastic bag.
The inside bag is made from cloth or a paper with many tiny holes and contains a mixture of iron powder, salt, charcoal and sawdust, all dampened with water.
When the paper bag is removed from the plastic bag and shaken vigorously it gets hot.
Iron is reacting with oxygen gas in the air to make iron oxide or rust.
3. Hand warmers
To make an air-activated hand warmer, put 25 g of iron powder or very fine iron filings and 1 g of sodium chloride in a small plastic bag.
Shake the bag to mix.
Add about a tablespoon of vermiculite or sawdust or sand to the bag and shake well.
Add 5 mL of water and seal the bag without squashing out all the air.
Shake the bag vigorously.
A reaction should start after about a minute.
4Fe (s) + 3O2 (g) --> 2Fe2O3 (s)
"Hand warmers" as sold by chemical suppliers contain iron powder, water, sodium chloride, activated charcoal, and vermiculite, in a polypropylene package.
4. Put iron powder in a plastic bag, e.g. a "Ziploc" bag.
Add sodium chloride and mix contents by shaking the closed bag.
Add 1 tablespoon of small vermiculite pieces and mix again.
Add 5 mL water to the bag and seal with a twist tie.
Squeeze and shake the bag.
After 2 minutes feel the bag and observe the heat produced.
The iron powder and the oxygen in the bag react to form iron oxide.
Salt speeds this reaction and is therefore a catalyst.
The vermiculite insulation ensures that the heat stays in the bag.
The iron oxide formed is a compound.
2Fe + 3O2 ---> Fe2O3 + heat.
hand warmers make use of the oxidation of iron to achieve an exothermic reaction
4Fe + 3O2 → 2Fe2O3 ΔH⚬ = - 1648 kJ/mol
14.13 Heat of reaction, crystallization of sodium thiosulfate
Sodium thiosulfate heat pack
Fill a test-tube 3 / 4 full of sodium thiosulfate crystals.
Heat the crystals over a Bunsen burner until all of the crystals have melted.
Let the clear colourless liquid cool to room temperature.
It contains supercooled sodium thiosulfate and it should not recrystallize.
Place one seed crystal of sodium thiosulfate into the solution of sodium thiosulfate.
If nothing happens after a minute, add another crystal.
Put your hand around the test-tube.
When a seed crystal is added, it starts the change from supercooled liquid to solid.
As the sodium thiosulfate becomes solid, it releases heat energy.
14.14 Reaction of ammonium salts and potassium salts with water
Use test-tubes containing 10 mL of water.
Put a thermometer in each test-tube and record the initial temperature.
Put the same mass of ammonium chloride, ammonium nitrate, potassium nitrate, and potassium chloride in each test-tube.
Record the changes in temperature of the solution.
14.15.1 Burn steel wool
Wear safety glasses and safety apron.
Handle steel wool with tongs.
Small pieces of steel, e.g. pins, needles, nails, will not ignite when heated with a lighter or Bunsen burner, because the surface area / volume ratio is too small.
However, a grinding wheel can be used to break steel into tiny pieces and heat them by friction to form incandescent pieces of iron with large surface / volume ratio, that react with oxygen in the air to form sparks.
Pull out strands of steel wool from a steel wool pad and use them to connect the terminals of a 6 volt battery.
The strands become hot caused by the high resistance of the iron and the surface starts to oxidize until all the strands are converted to iron oxides.
The strands burn brighter and faster if you blow on them to increase the oxygen supply.
4Fe + 3O2 --> 2Fe2O3 + energy
Burn steel wool in air with a Bunsen burner over a heat resistant mat to form black magnetite, FeOFe2O3, that is weakly magnetic.
14.15.2 Black iron oxide is a mixed base, Fe3O4
Cover the bottom of a test-tube with black iron oxide and add 3 cm of concentrated acid.
Heat the solution slowly then filter it.
Divide the filtrate into two parts.
Test one part for iron (III) ions.
Test the other part for iron (II) ions.
Both ions are present.
Fe3O4 + 8HCl -->- 2FeCl3 + FeCl2 + 4H2O.
14.15.3 Detect iron in fruit juice using black tea
Add strong black tea to samples of fruit juice, e.g. apple, pineapple, cranberry.
Note the time for a cloudy precipitate of iron compounds to form.
The precipitate may not appear for hours or days and the time for precipitation may depend on the temperature and concentrations of the tea and fruit juice.
Pineapple juice should give the shortest time for precipitation.
The precipitate is formed by a reaction between the ferric, Fe3+, non-haem iron from the fruit juice with the tannins in the black tea.
The non-haem iron is an important component in your diet, but black tea may make this iron indigestible so that we cannot absorb it.
Perhaps we should drink black tea only between meals and not with meals.
The ferrous, Fe2+, haem iron comes mainly from haemoglobin and myoglobin in red meat.
14.15.4 Heat hydrated iron chlorides
1. Prepare iron (II) chloride solution by dissolving iron filings in concentrated hydrochloric acid.
Evaporate in a test-tube until crystals appear.
Heat strongly and test the vapour for hydrogen chloride with silver nitrate solution on a glass rod.
Note the residue of iron (III) oxide formed when the iron (II) oxide is oxidized in the air.
FeCl2 + H2O --> FeO + 2HCl
2FeO + O (air) --> Fe2O3.
2. Heat iron (III) chloride in a test-tube.
Test the gas for hydrogen chloride and note the residue of iron (III) oxide.
Hydrolysis has occurred.
2FeCl3 + 3H2O --> Fe2O3 + 6HCl
14.15.5 Heat iron (II) sulfide
FeS2 (pyrite) fool's gold
Iron (III) oxide and sulfur dioxide forms.
(FeS2 is not iron (IV) sulfide.)
4FeS2 (s) + 11O2 --> 2Fe2O3 (s) + 8SO2 (g).
14.15.6 Heat iron filings with powdered sulfur
Grey iron (II) sulfide forms, FeS.
It is ferrimagnetic.
8Fe + S8 --> 8FeS (direct union of elements to form compounds).
14.15.7 Heat iron with sulfur
Heat iron filings with sulfur powder (synthesis reaction):
See diagram 12.2.1: Iron (II) sulfide, FeS
S8 (s) + 8Fe (s) --> 8FeS (s).
Be careful! The following reactions are vigorous.
Do not use large quantities of the chemicals.
Be careful! The reaction of iron (II) sulfide with hydrochloric acid will form the poisonous gas, hydrogen sulfide, with an odour of rotten eggs.
Experiments
1. Mix half a metal bottle top of powdered sulfur with the same volume of iron filings.
Heat a small portion of the mixture on the metal bottle top with the cork removed or in a hard glass test-tube.
When the reaction begins, i.e. the mixture starts to glow, stop heating by moving the Bunsen burner to the side.
If the glow stops, heat the test-tube again.
The reaction of a mixture of iron with sulfur gives out so much heat that the mixture becomes red hot.
Note the following different properties of powdered sulfur, iron filings and the iron (II) sulfide: | appearance | colour | hardness | magnetism|.
Iron is magnetic so is easily removed from a mixture of iron and sulfur, but iron (II) sulfide is not magnetic.
Fe + S --> FeS (s).
2. Use < 10 g total material iron with sulfur in a fume cupboard.
Heat the mixture to start the reaction.
However, be aware that unreacted sulfur may catch fire and produce sulfur dioxide gas to irritate the lungs.
3. Mix uniformly reduced iron powder and powdered sulfur in a weight ratio of seven to four.
Carve the word "FeS" on a red coloured brick with a knife.
Spread the iron sulfur mixture throughout the word groove and press the powdered mixture solid.
Heat one tip of a glass rod until red hot with an alcohol burner and then immediately dig the hot tip into the mixture at one end of the word groove.
A chemical reaction is starts immediately.
The reaction continues violently to release a large amount of heat and meanwhile to develop rapidly a red glow, which looks like a small "fiery dragon".
The heat lost by the reaction is more than the heat needed to start the reaction.
The reaction produces a new black solid substance, iron (II) sulfide, that has different properties from the two reactants, iron and sulfur.
Compare iron powder, powdered sulfur and iron (II) sulfide.
Note their appearance.
Test them respectively with a magnet.
Add in drops hydrochloric acid solution to them respectively.
4. Mix equal amounts of iron filings and powdered sulfur.
Heat the mixture in a crucible or a small tin with sand in the bottom.
The sand prevents the bottom of the tin from melting by spreading the heat.
Heat the mixture strongly until you see a red glow spreading through the mass.
The heat lost by the chemical reaction is more than the heat needed to start the reaction.
The reaction forms a new substance iron (II) sulfide that has different properties from the two elements used to make it.
Compare iron filings, powdered sulfur, and iron (II) sulfide.
Note their appearance.
Test with a magnet.
Add drops of hydrochloric acid.
5. Make a mixture of 7 parts of iron filings with 4 parts of sulfur powder in a sealed plastic bag.
Hold a magnet over the plastic bag to show that the iron filings can be easily separated from the mixture.
Quarter fill an ignition tube with the mixture.
Near an open window or in a fume cupboard heat the end of the ignition tube with a Bunsen burner.
When the mixture glows move the Bunsen burner away, but when the glow stops move the Bunsen burn back again until all the mixture reacts.
Leave the ignition tube to cool, then move the magnet near it.
The magnet can no longer attract iron filings or the iron sulfide in the ignition tube.
14.15.8 Iron displaces hydrogen from hydrochloric acid
Iron displaces hydrogen from hydrochloric acid to form iron (II) chloride
Fe (s) + 2 HCl (aq) --> FeCl2 (aq) + H2 (g)
Evaporate the solution to form crystals of pale green FeCl24H2O.
In the air, the iron (II) is oxidized to FeCl3 and Fe2O3.
14.15.9 Iron displaces hydrogen from sulfuric acid
Iron displace hydrogen from sulfuric acid to form iron (II) sulfate
Fe (s) + H2SO4 (aq) --> FeSO4 (aq) + H2 (g)
Evaporate the solution to form blue-green crystals of FeSO47H2O, green vitriol.
In air, iron (II) salts are oxidized to iron (III) salts, so brown iron(III) hydroxide and iron (III) sulfate forms on the blue-green crystals.
14.15.10 Iron (II) sulfate (ferrous sulfate) with sodium carbonate
Add sodium carbonate (washing soda), solution to iron (II) sulfate solution.
A green precipitate of iron (II) carbonate forms iron (II) sulfate + sodium carbonate --> iron (II) carbonate + sodium sulfate
14.15.11 Iron (II) sulfate (ferrous sulfate) with ammonia
Add drops of dilute ammonia solution to 2 cm of iron (II) sulfate solution in a test-tube and shake the test-tube.
Observe a green-grey precipitate of iron (II) hydroxide.
The precipitate left on the side of the test-tube quickly turns brown, because oxygen from the air turns it into iron (III) hydroxide, ferric hydroxide.
14.15.12 Iron (II) sulfate oxidation
Iron (II) sulfate (ferrous sulfate) oxidation to iron (III) sulfate (ferric sulfate)
Heat iron (II) sulfate solution with a substance rich in oxygen, e.g. hydrogen peroxide.
Boil 2 cm of iron (II) sulfate solution in a test-tube with drops of hydrogen peroxide.
The green colour changes to yellow or brown.
Cool the test-tube under the tap and test the liquid, iron (III) sulfate solution (ferric sulfate) by adding dilute ammonia solution.
A brown precipitate of iron (III) hydroxide (ferric hydroxide) forms.
14.15.13 Iron (III) sulfate reduction
Iron (III) sulfate (ferric sulfate) reduction to iron (II) sulfate (ferrous sulfate)
Put 2 cm of ammonium iron (III) sulfate solution in a test-tube with an equal amount of dilute sulfuric acid or sodium hydrogen sulfate solution.
Add 2 mL of iron filings or coiled iron wire or steel wool or small pieces of zinc or zinc powder
Heat the test-tube until effervescence starts, because of the formation of hydrogen.
The hydrogen is the reducing agent.
Leave the test-tube to stand.
The brown / yellow colour of the solution vanishes, and when it is pure the solution is light green.
2Fe3+ + Zn --> 2Fe2+ + Zn2+
Test 1. Add dilute ammonia solution.
A dirty green precipitate forms to indicate iron (II) sulfate.
The precipitate does not dissolve in excess ammonia solution.
Test 2. Add dilute sodium hydroxide solution.
A dirty green precipitate forms that does not dissolve in excess sodium hydroxide.
14.15.14 Oxidation of iron (II) salts
1. Use 2 cm of iron (II) sulfate solution in a test-tube.
Add just more than an equal volume of dilute sulfuric acid and three
drops of concentrated nitric acid.
Heat until the solution boils.
Leave to cool and add sodium hydroxide solution until a red precipitate
of iron (III) hydroxide forms.
6FeSO4 + 3H2SO4 + 2HNO3 -->
Fe2(SO4)3 + 4H2O + 2NO
Iron (II) ions are oxidized to iron (III) ions by electron loss.
Fe2+ - e- --> Fe3+.
2. Pass chlorine gas through an iron salt solution.
Brown iron (III) chloride solution forms.
2Fe2+ (aq) --> Fe3+ (aq) + e- (oxidation of iron when it loses an electron)
Cl2 (aq) + 2e- --> 2Cl- (aq) (reduction of chlorine when it gains electrons)
2Fe2+ (aq) + Cl2 (g) --> 2Fe3+ (aq) + 2Cl- (aq)
Chlorine is the oxidizing agent.
Its oxidizing number drops from 0 to -1 when it gains an electron.
Iron is oxidized as it increases from Fe (II) to Fe (III) when it loses an electron.
3. Repeat the experiment by substituting other oxidizing materials, e.g. bromine, KMNO4 or hydrogen peroxide, for nitric acid in the above experiment.
14.15.15 Reaction of iron (II) salts & iron (III) salts
1. Add sodium hydroxide or ammonia solution, NH3 (aq) ("ammonium hydroxide")
Iron (II) salt: Green precipitate of iron (II) hydroxide Fe(OH)2.
Iron (III) salt: Red precipitate of iron (III) hydroxide, hydroxide Fe(OH)3.
2. Add acidified potassium permanganate.
Iron (II) salt: Permanganate manganate loses its colour.
Iron (II) salts are reducers.
Iron (III) salt: Does not reduce.
3. Add potassium ferrocyanide, K4[Fe(CN)6].
Iron (II) salt: Light blue precipitate.
Iron (III) salt: Deep blue precipitate, Prussian blue.
4. Add potassium ferricyanide K3[FeCN)6].
Iron (II) salt: Deep blue precipitate "Turnbull's blue".
Iron (III) salt: Brown colour.
5. Add potassium or ammonium thiocyanate solution
to a freshly made iron (II) sulfate solution.
Iron (II) ammonium sulfate should give a negative result.
Iron (II) salt: No action.
Iron (III) salt: Blood red coloration of iron (III) thiocyanate Fe(CNS)3
6. Add iron (III) ions to thiocyanate ion solutions to form bright red complexes, e.g. (Fe(SCN)3, Fe(SCN)63-.
So a thiocyanate solution can be used as a test for iron (III) ions, because iron (II) ions do not cause a colour change.
7. Iron (II) thiocyanate oxidizes pale green Fe(SCN)23H2O crystals to red iron (III) thiocyanate and so can be used as a test for the presence of oxygen gas and peroxides.
8. Iron (II) ions and iron (III) ions react with ferrocyanide ion (Fe(CN)64+), and ferricyanide ion, (Fe(CN)63+), to form the coloured pigment, Prussian blue.
K4Fe(CN)6 (aq) + Fe3+ (aq) --> KFe[Fe(CN)6] (s) + 3 K+ (aq).
9. Iron (II) react with ferricyanide ions to form the same coloured pigment.
K3Fe(CN)6 (aq) + Fe2+ (aq) --> KFe[Fe(CN)6] (s) + 2 K+ (aq)
10. In blueprinting, the undeveloped paper is covered with iron (III) ferricyanide ion, and citrate.
In the light, the citrate reduces the iron (III) to iron (II).
With the addition of water, the deep blue pigment forms.
14.16 Reaction of ammonium carbonate with ethanoic acid
Add an excess of ethanoic acid to ammonium carbonate.
Keep heating until no more carbon dioxide leaves the mixture.
The endothermic reaction is very cold.
2CH3COOH + (NH4)2CO3 --> 2CH3COONH4 + H2O + CO2
ethanoic acid + ammonium carbonate --> ammonium ethanoate + water
Heat the ammonium ethanoate to form ethanamide, C2H5NO, CH3CONH2, acetamide, acetic acid amide.
CH3COONH4 --> CH3CONH2 + 2H2O
ammonium ethanoate --> ethanamide + water
14.17 Reaction of ammonium nitrate with water
Ammonium nitrate cold pack
Cold packs contain chemicals that mix when the cold pack is squashed.
The cold pack has two sealed bags, one inside the other.
The outer bag is made of thick strong plastic.
It contains a white powder ammonium nitrate and a second plastic bag.
The inner bag is made of a thin weak plastic and contains water.
When the cold pack is punched, the inner bag breaks.
The water mixes with the powder, dissolves it and the solution becomes very cold.
When ammonium nitrate dissolves in water, it absorbs heat, i.e. "it gets cold".
This type of cold pack is not reusable.
NH4NO3 + H2O --> NH4+ (aq) +NO3- (aq)
Experiment
Investigate which substances would make the best cold pack: potassium nitrate, sodium chloride, calcium chloride, ammonium nitrate.
14.18 Reaction of potassium nitrate with water
See diagram 3.81: Potassium nitrate with water
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 90 oC.
This means that while dissolving, the particles have absorbed energy.
This energy has been taken from the surrounding water in the form of heat.
Repeat the experiment with potassium chloride.
14.19 Reaction of copper (II) sulfate solution with magnesium
1. Pour concentrated copper (II) sulfate solution into the test-tube.
Add very small pieces of magnesium ribbon until the blue colour disappears.
Record the change in temperature of the solution.
BE CAREFUL!
The reaction is vigorous!
Mg (s) + CuSO4 (aq) --> Cu (s) + MgSO4 (aq)
magnesium + copper sulfate --> copper + magnesium sulfate.
Magnesium is more reactive than copper.
When a piece of magnesium is dipped into blue copper sulfate solution, a displacement reaction occurs.
The magnesium displaces the copper from the copper sulfate solution to give copper and a solution of magnesium sulfate.
2. Put 10 mL of strong aqueous copper (II) sulfate solution into a wide test-tube or small container.
Support a thermometer with the bulb in the solution.
Add magnesium powder, or ribbon, a little at a time until the blue colour disappears.
Note any changes in the thermometer reading.
Mg(s) + CuSO4(aq) → MgSO4(aq) → Cu(s)
Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s)
3. Magnesium displaces copper that is lower in the activity series from its salt copper (II) sulfate.
Add magnesium ribbon to a test-tube of copper (II) sulfate solution.
The reaction can be vigorous with the magnesium.
Copper metal deposits and the blue colour gradually disappears as the copper ion is displaced by the more reactive metal that is higher in the activity series.
The reaction loses heat.
When the solution is colourless, decant the solution leaving red copper powder at the bottom of the test-tube.
Mg (s) + CuSO4 (aq) --> MgSO4 (aq) + Cu (s)
Mg loses electrons: Mg --> Mg2+ + 2e- (oxidation)
Cu gains electrons: Cu2+ + 2e- --> Cu (reduction).
14.20 Reaction of urea with water
Urea (carbamide, H2NCONH2) is a crystalline solid that is very soluble in water.
Use a test-tube containing 10 mL of water.
Put a thermometer in the test-tube and record the initial temperature.
Put 5 g of urea in the test-tube.
Record the changes in temperature of the solution.
14.21 Combustion of potassium nitrate and sugar mixture
Dissolve 3 parts of potassium nitrate and 2 parts of white household sugar in just enough water to dissolve the components.
1. Draw a line on newspaper or duplicating paper or paper towel with a glass rod or cotton bud dipped in the solution, then leave the paper to dry.
Put the paper in a safe place on a fire resistant surface.
Light a match, blow out the flame and touch one end of the potassium nitrate line with the glowing end of the match.
A flame races along the line as the potassium nitrate and paper near it burns.
Use the solution to write a student's name or the name of their school on paper, then see the name burst into flame.
2. Wash clean string in soapy water to dissolve away any preservative.
Rinse the string in running water then leave the wet string in the solution.
The end of the dried string can be ignited to make a fuse.
14.22 Dehydration of sugar with sulfuric acid
Use a fume hood.
Put household white sugar in a tall beaker.
Add a little water to just dampen the sugar.
Add concentrated sulfuric acid.
Be Careful!
A highly exothermic raeaction occurs, with a caramel sulfurous odour.
C12H22O11 + H2SO4 --> 12 C + 11H2O