Endothermic reactions, Enthalpy, Thermal capacity.

School Science Lessons
2024-06-11

Thermochemistry
(topic14)
Contents
14.3.0 Chemiluminescence, bioluminescence
14.1.0 Exothermic and endothermic reactions
14.01 Energy of reactions
14.04 Enthalpy
14.03 Fluorophores
14.06 Heat of reaction
14.07 Hess's law
14.05 Thermal capacity

Experiments
14.2.6 Combustion of potassium nitrate
14.1.2 Endothermic reactions take in heat energy, (List), (Experiments)
14.3.2 Enthalpy of reaction, heat of reaction
14.1.1 Exothermic reactions give out heat energy, (List), (Experiments)
14.1.5 Heat of neutralization with a calorimeter
14.3.3 Heat of reaction, anhydrous copper (II) sulfate with water
14.1.3 Heat of reaction, acids with water
14.1.4 Heat of rusting, steel wool
14.3.1 Luminol tests for blood, Cu, Fe, Cn^-
14.2.8 Methylene chloride, Drinking bird, (Experiments)
14.2.4 Reaction of ammonium nitrate with water
14.2.1 Reaction of ammonium salts and potassium salts with water
14.3.4 Reaction of copper (II) sulfate solution with magnesium
14.2.5 Reaction of potassium nitrate with water
14.2.2 Reaction of urea with water
14.2.3 Reaction of urea ammonium carbonate with ethanoic acid

14.1.0 Exothermic and endothermic reactions
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.
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.

14.1.1 Exothermic reactions give out heat energy
14.3.2 Enthalpy of reaction, heat of reaction
Experiments
14.2.7 Exothermic reactions, (Experiments)
14.2.6 Combustion of potassium nitrate, fire line on paper, string fuse
14.3.4 Heat of reaction, copper (II) sulfate solution with magnesium
14.1.6 Heat of displacement reaction
14.1.5 Heat of neutralization with a calorimeter, (Experiments)
14.1.9 Heat of reaction, chromium (VI) oxide with ethanol
14.1.8 Heat of reaction, potassium permanganate with ethanol
14.1.7 Heat of reaction, potassium permanganate with glycerol
14.1.10 Heat of reaction, potassium with diethyl ether
14.1.4 Heat of rusting
14.3.3 Heat of reaction, anhydrous copper (II) sulfate with water

14.1.12 Heat of reaction, iron with oxygen gas, iron powder heat pack, hand warmers
14.1.11 Heat of reaction, crystallization of sodium acetate, sodium acetate heat pack
14.1.13 Heat of reaction, crystallization of sodium thiosulfate, sodium thiosulfate heat pack
14.1.3 Heat of reaction, acids with water
14.8.0 Reaction of iron, Fe.

14.1.2 Endothermic reactions take in heat energy
14.2.4 Reaction of ammonium nitrate with water, Ammonium nitrate cold pack
Experiments
14.2.3 Reaction of ammonium carbonate with ethanoic acid
14.2.8 Methylene chloride
14.2.5 Reaction of potassium nitrate with water
14.2.1 Reaction of ammonium salts and potassium salts with water
14.2.2 Reaction of urea with water

14.01 Energy of reactions
Energy of reactions, enthalpy, thermal capacity, heat of reaction, Hess's law
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.

14.04 Enthalpy
Enthalpy, heat content, refers to the energy stored in a substance.
Enthalpy, H = U + PV, where U = internal energy, P = pressure and V = volume, of a system.
The SI unit is the joule.

14.05 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 C^o).

14.06 Heat of reaction
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.

14.07 Hess's law
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.

14.3.0 Chemiluminescence, bioluminescence
14.3.01 Chemiluminescence, bioluminescence
14.03 Fluorophores
14.3.1 Luminol tests for blood, Cu, Fe, Cn^-

14.8.0 Reaction of iron, Fe
14.8.21 Burn steel wool
14.8.27 Black iron oxide is a mixed base, Fe3O4
14.8.15 Detect iron in fruit juice using black tea
14.8.30 Heat hydrated iron chlorides
14.8.25 Heat iron (II) sulfide
14.8.23 Heat iron filings with powdered sulfur
14.8.33 Heat iron with sulfur
14.8.29 Iron displaces hydrogen from hydrochloric acid
14.8.28 Iron displaces hydrogen from sulfuric acid
14.8.16 Iron (II) sulfate (ferrous sulfate) with sodium carbonate
14.8.17 Iron (II) sulfate (ferrous sulfate) with ammonia
14.8.18 Iron (II) sulfate (ferrous sulfate) oxidation to iron (III) sulfate (ferric sulfate)
14.8.19 Iron (III) sulfate (ferric sulfate) reduction to iron (II) sulfate (ferrous sulfate)
14.8.20 Oxidation of iron (II) salts
14.8.14 Prepare iron (II) ammonium sulfate
14.8.24 Prepare iron (II) oxide
14.8.31 Prepare iron (II) sulfate (ferrous sulfate) crystals with iron filings
14.8.32 Prepare iron (III) hydroxide and iron (III) oxide
14.8.26 Prepare iron (III) oxide
14.8.22 Reduce iron (III) salts
14.8.1 Reaction of iron (II) salts & iron (III) salts

14.2.1 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.

14.2.2 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.

3.81 Endothermic reactions
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.
Endothermic dissolving processes of a solid in a solvent are associated with low solubility.
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 90^oC.
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.

14.2.3 Reaction of ammonium carbonate with ethanoic acid
Ammonium ethanoate, carbon dioxide and water form.
The reaction is very cold.

14.2.4 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.
Experiment
Investigate which substances would make the best cold pack: potassium nitrate, sodium chloride, calcium chloride, ammonium nitrate.

14.2.5 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.

14.2.7 Exothermic reactions
Be careful! The reactions may be vigorous.
The reactants of exothermic reactions form products with rise in temperature.
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. Citric acid with sodium hydrogen carbonate solution
5. Thermite reaction
6. Dilute hydrochloric acid with dilute sodium hydroxide solution.
(Neutralization reactions)
7. Burning substances and combustion of fuels.
8. Setting of cement and concrete
9. Corrosive oxidation of metals.

14.2.6 Combustion of potassium nitrate
1. Draw a line on newspaper or duplicating paper or paper towel with a glass rod or cotton bud dipped in potassium nitrate 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 potassium nitrate solution to write their own 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 a potassium nitrate solution.
The end of the dried string can be ignited to make a fuse.

14.2.8 Methylene chloride
Pass a current of air through methylene chloride, CH2Cl2, dichloromethane, organic solvent, in paint strippers.
It has a low boiling point of 39.6^oC and is used in the drinking bird heat engine.

14.3.01 Chemiluminescence, bioluminescence
Luminescent substances emit light not, because of a rise in temperature of the substance.
Chemical reactions that produce energy not as heat, but as light are called chemiluminescent reactions.
So chemiluminescence is luminescence resulting from a chemical change
Such chemical reactions in living organisms are called bioluminescent reactions, e.g. the "cold light" from the abdomen of the firefly, glow worms, luminescent fish.
Chemiluminescence is used in a "light stick", "light necklace", "light bracelet" or "glow stick" at night fairs or for Halloween activities.
The "light stick" contains dilute hydrogen peroxide dissolved in a phthalic ester solvent and contained in a very thin glass ampoule surrounded by a phenyl oxalate ester solution (Cyalume) and the fluorescent dye, fluorophore, e.g. 9, 10-bis (phenyl ethynyl) anthracene, (BPEA).
When you break the ampoule by agitating the light stick, the hydrogen peroxide and phenyl oxalated ester react to form phenol and peroxyacid ester.
The ester forms carbon dioxide and transfers energy to the dye molecule that produces a green-yellow cold light as it returns to the energy ground state.
See diagram 14.3.0: Luminol structural formula and equations

14.03 Fluorophores
Fluorophores are fluorescent chemical compounds.
(1.) 1-chloro-9,10-bis(phenyl ethynyl) anthracene emits green-yellow light in 30-minute high-intensity Cyalume sticks,
(2.) 2-chloro-9,10-bis(phenyl ethynyl) anthracene emits green light in 12-hour low-intensity Cyalume sticks.
(3.) 1,8-dichloro-9,10-bis(phenyl ethynyl) anthracene emits yellow light in Cyalume sticks.
(4.) 9,10-diphenylanthracene (DPA), C26H18, emits blue light in light sticks, yellow powder.
(5.) 1-chloro-9,10-diphenyl anthracene (1-chloro(DPA)) emits blue-green light.
(6.) 2-chloro-9,10-diphenyl anthracene (2-chloro(DPA)) emits blue-green light.
(7.) 9,10-bis(phenyl ethynyl) anthracene (BPEA), C30H18 emits "ghostly" green light in light sticks.
(8.) 2,4-di-tert-butylphenyl 1,4,5,8-tetracarboxynaphthalene diamide emits deep red light, but with DPA emits white or hot-pink light.
(9.) Fluorescein isothionate, FITC, emits green light.
(10.) Rhodomine tetramethyl isothionate, TRITC, "Rhodominr Red-X", emits orange light.
(11.) Hydroxycoumarin, emits blue light.
(12.) Cyanine dyes, Cy2 cyanine, Cy3 indocarbocyanine, Cy5 indodicarbocyanine.
(13.) Biological fluorphores, e.g. green fluorescent protein, GFP, occurs from a jellyfish.
(14) Quantum dot nanocrystals.

14.3.1 Luminol tests for blood
"Cool Blue Light Kit", luminol, chemiluminescence, forensic science, (toy product)
"Cool Chemical Light Powder", luminol (toy product)
Luminol is used to detect copper, iron, peroxides and cyanides.
The chemical luminol, 5-amino-2,3-dihydro-1,4-phthalazinedione, 3-aminophthalhydrazide, C8H7N3O2, reacts with oxygen to produce an intermediate molecule, metal chelate, that releases energy as blue-green luminescence when oxidized in alkaline solution.
It has low solubility in water and is a yellow grainy substance.
When luminol is placed in a basic solution such as a permanganate, hypochlorite or hydrogen peroxide, with a metal catalyst, e.g. cobalt, the luminol is oxidized.
The two nitrogen atoms are replaced by two oxygen atoms and nitrogen gas is discharged, leaving the luminol in an excited state with additional energy that is then released as light amino acids, and serum albumin can also react with luminol to produce blue-green light.
So luminol is used in biology and biochemistry for testing and for the detection of blood for forensic science.
Blood is slightly alkaline and contains haemoglobin, which contains iron.
Luminol can detect very small amounts of blood, even if many years old.
In a television story, detectives used luminol to find blood stains, although the murderer had tried to wipe clean all the blood at the crime scene.
In the television story, the detective applied the luminol, turns out the lights, the glow appears and the case is solved!
However, further testing is required to decide if the reactant is blood, because luminol may glow when in contact with other substances.
Luminol destroys genetic markers in the blood so it is used as a last resort in crime scenes investigations.
Dissolve luminol in sodium hydroxide solution then oxidize it with hydrogen peroxide or sodium hypochlorite to form the unstable disodium salt of 3-aminophthalic acid.
Experiments 1. Solution A: Add 100 mL of 5% NaOCl, household bleach to 900 mL water.
Solution B: Dissolve 0.4 g luminol and 40 g sodium hydroxide in 1 litre of water.
The luminol will not dissolve completely.
Record the temperature of both solutions.
Dim the lights and pour both solutions simultaneously into a larger beaker.
Note the pale glow that lasts for a few seconds.
Measure the temperature of the mixture to show that no heat came from the reaction.
2. Solution A: Dissolve 0.1 g luminol and 5 mL 5% NaOH in 100 mL water.
Solution B: Add 10 mL 3% hydrogen peroxide + 0.25 g potassium ferricyanide, K3[Fe(CN)6] + 1 litre of water.
Pour the two solutions together into a larger beaker.
3. Oxalyl chloride mixed with hydrogen peroxide and a fluorescent dye produces chemiluminescence.
Also, phenyl oxalate ester mixed with hydrogen peroxide and a dye, gives a brighter light, but not as efficient as a firefly.
Some oxalate esters react with hydrogen peroxide with the help of a salicylates catalyst to form a peroxyacid ester and phenol.
The peroxyacid ester decomposes to form more phenol and a high energy intermediate compound that gives up its energy to a dye molecule, which then fluoresces.
Most light sticks use the dye molecule 9,10-bis(phenylethynyl) anthracene to make green, and 9,10-diphenylanthracene to make blue.

14.3.2 Enthalpy of reaction, heat of reaction
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!.

14.3.3 Heat of reaction, anhydrous copper (II) sulfate with water
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)

14.3.4 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!
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).

14.1.3 Heat of reaction, acids with water
"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 (20^o C) → H3O⁺+ HSO4^– (130^o 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.

14.1.4 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.

14.1.5 Heat of neutralization with a calorimeter
See diagram 3.1.5: 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).

14.1.6 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.

14.1.7 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.

14.1.8 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.
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.

14.1.9 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.
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).

14.1.10 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!
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).

14.1.11 Heat of reaction, crystallization of sodium acetate, sodium acetate heat pack
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 54C 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.

14.1.12 Heat of reaction, iron with oxygen gas, iron powder heat pack, 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.

14.1.13 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.

14.8.1 Reaction of iron (II) salts and iron (III) salts, Prussian blue
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)6^3-.
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.

14.8.31 Prepare iron (II) sulfate crystals with iron filings
When iron is treated with dilute sulfuric acid hydrogen gas forms and a solution of iron (II) sulfate forms.
Heat half a test-tube of dilute sulfuric acid over a flame, but do not boil the liquid.
Remove the test-tube from the flame and add iron filings on the end of a metal spatula.
A vigorous effervescence occurs, because of the formation of hydrogen.
Test with a glowing splint.
When the action dies down, add more iron filings, and then more, until a total of 4 mL is added.
Put the test-tube on one side for 15 minutes until effervescence has nearly ceased.
Filter the liquid into an evaporating basin.
Make sure that acid is left, by testing with blue litmus paper.
The presence of acid prevents the solution from oxidizing to brown iron (III) sulfate.
To obtain large crystals of iron (II) sulfate, leave the solution undisturbed for a day or two to crystallize out.
Small crystals can form more quickly by evaporating the solution over a flame until only one third of it remains.
If a brown colour appears in the liquid during the evaporation add a drop or two of dilute sulfuric acid.
When the evaporation finishes, leave the remaining liquid to cool.
Many small crystals of iron (II) sulfate are deposited.

14.8.32 Prepare iron (III) hydroxide and iron (III) oxide
Add 4 cm of dilute ammonia solution to 2 cm of ammonium iron (III) sulfate solution in a test-tube.
Shake the test-tube to mix the liquids.
A brown jelly-like precipitate of iron (III) hydroxide forms.
Filter off the precipitate of iron (III) hydroxide.
Put of the red-brown jelly left in the filter paper on to a clean metal lid or into a metal screw cap.
Hold the lid or cap in a pair of pliers and heat it carefully above a flame.
Steam is produced and a red powder formed.
This is iron (III) oxide, a very pure form of rust.
The same substance forms by heating iron (II) sulfate crystals.

14.8.14 Prepare iron (II) ammonium sulfate
Iron (II) ammonium sulfate, ammonium iron (II) sulfate, Mohr's salt, (NH4)2Fe(SO4)2(H2O)6
Add 4 mL of concentrated sulfuric acid to 30 mL of deionized water in a conical flask.
Slowly add 5 g of iron then heat to boiling.
Add 10 g of ammonium sulfate and evaporate to two thirds of the original volume.
Add a loose stopper loosely and leave the double salt to crystallize.
This salt is not an alum.

14.8.15 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.

14.8.16 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

14.8.17 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.

14.8.18 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.

14.8.19 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⁺
Test1: 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.

14.8.3 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, potassium permanganate or hydrogen peroxide, for nitric acid in the above experiment.

14.8.4 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.

14.8.22 Reduction of iron (III) salts
1. Put 2 cm of iron (III) chloride solution in a test-tube.
Pass hydrogen sulfide through the solution until there is no further precipitate of sulfur occurs.
Filter the solution and note the pale green solution.
Test the filtrate with potassium ferricyanide for proof of iron (II) salt.
2FeCl3 + H2S --> 2FeCl2 + 2HCl + S (s)
Iron (III) ions are reduced to iron (II) ions by electron gain.
Fe³⁺ + e^- --> Fe²⁺Sulfide ions are oxidized by electron loss.
H2S <--> 2H⁺ + S^2-
S^2- - 2e^- --> S.
2. Add an equal volume of concentrated hydrochloric acid and pieces of granulated zinc to 3 cm of iron (III) salt solution.
Leave for half an hour then filter.
Test the filtrate with excess of sodium hydroxide solution to show that reduction to iron (II) is complete.
In the presence of acid, zinc atoms ionize and the electrons are accepted by iron (III) ions that are reduced to iron (II) ions.
Zn --> Zn³⁺ + 2e^-
2Fe³⁺ + 2e^- --> 2Fe²⁺.

14.8.23 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).

14.8.24 Prepare iron (II) oxide
Close with a plug of wool a dry test-tube containing 3 cm of iron (II) oxalate.
Heat gently then strongly to convert all the yellow oxalate to black iron (II) oxide.
Remove the plug of wool and sprinkles the iron (II) oxide into an evaporating basin.
The iron (II) oxide spontaneously ignites as it oxidizes to red iron (III) oxide.
FeC2O4 --> FeO + CO + CO2
Dissolve the particles left in the test-tube in hydrochloric acid.
Test the solution for iron (II) ions.
Iron (II) oxide is a base, but iron (II) salts are prepared with metallic iron and acid.

14.8.25 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).

14.8.26 Prepare iron (III) oxide
Add excess ammonia solution, NH3 (aq) ("ammonium hydroxide") to an iron (III) salt and filter off the iron (III) hydroxide.
Heat the filter paper and contents in a crucible to leave red iron (III) oxide, Fe2O3 .
Boil some oxide in concentrated hydrochloric acid and show that it is a base.

14.8.27 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.

14.8.28 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.

14.8.29 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.

14.8.30 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.

14.8.33 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.