topic14

School Science Lessons
Topic 14 Thermochemistry, heat of reaction, iron, chemiluminescence
2012-05-11 SPw
Please send comments to: J.Elfick@uq.edu.au

Table of contents
14.0.0 Thermochemistry
14.3.0 Chemiluminescence, bioluminescence
14.2.0 Endothermic reactions, reactions that take in heat energy
14.01 Energy of reactions, enthalpy, thermal capacity, heat of reaction, Hess's law
14.1.0 Exothermic reactions, reactions that give out heat energy
12.8.0 Iron, Fe

12.8.0 Iron, Fe12.8.10 Black iron oxide is a mixed base, Fe₃O₄
13.3.3 Burn steel wool in oxygen, burn iron filings
12.8.15 Detect iron in fruit juice using back tea
12.8.13 Heat hydrated iron chlorides
12.8.8 Heat iron (II) sulfide, (FeS₂, pyrite, fool's gold)
12.8.6 Heat iron filings with powdered sulfur
12.8.11 Iron displaces hydrogen from sulfuric acid to form iron (II) sulfate
12.8.12 Iron displaces hydrogen from hydrochloric acid to form iron (II) chloride
12.8.3 Oxidation of iron (II) salts
12.8.14 Prepare iron (II) ammonium sulfate
12.8.7 Prepare iron (II) oxide
12.14.3 Prepare iron (III) ammonium alum
12.8.9 Prepare iron (III) oxide
12.8.1 Reactions of iron (II) salts and iron (III) salts, Prussian blue
12.8.5 Reduce iron (III) salts

14.1.0 Exothermic reactions, reactions that give out heat energy
14.1.0 Exothermic reactions, reactions that give out heat energy
3.80 Exothermic reactions, Reactions that give out heat energy,
14.2.6 Combustion of potassium nitrate, fire line on paper, string fuse
15.3.12 Conditions necessary for rusting
14.1.2 Copper (II) sulfate solution with magnesium
3.83 Heat of reaction when metals displace copper
14.1.6 Heat of displacement reaction, zinc with copper (II) sulfate solution
3.82 Heat of neutralization reaction
14.1.5 Heat of neutralization of hydrochloric acid with sodium hydroxide
14.1.5.1 Heat of neutralization with a simple calorimeter, dilute hydrochloric acid / ethanoic acid with sodium hydroxide solution
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, steel wool
14.1.1 Heat of solution of anhydrous copper (II) sulfate
14.1.12 Iron powder heat pack
3.52.1 Mass of iron and its temperature increases during rusting
3.52.3 Metals can prevent rusting
3.52.2 Oxygen gas combines with iron during rusting
14.1.11 Sodium acetate heat pack
14.1.13 Sodium thiosulfate heat pack
14.1.3 Sulfuric acid with water
5.03 Supersaturated solution of sodium ethanoate-3-water, "heat-pack"

14.2.0 Endothermic reactions, reactions that take in heat energy
Order online: Ice Melting Blocks, endothermic heat flow
14.2.3 Ammonium carbonate with ethanoic acid
14.2.4 Ammonium nitrate cold pack
14.2.0 Endothermic reactions, reactions that take in heat energy
14.2.5 Potassium nitrate with water
3.81 Reactions that take in heat energy, endothermic reactions
14.2.1 Reactions of ammonium salts and potassium salts with water
14.2.7 Refrigerants
14.2.2 Urea with water

14.3.0 Chemiluminescence, bioluminescence
14.3.0 Chemiluminescence, bioluminescence
14.3.0.1 Luminol

14.01 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. Ionic bonding occurs when positive and negative ions are held together in a crystal lattice by electrostatic forces. When chemical bonds, whether ionic or covalent, form between different elements, a chemical compound is obtained, which can be represented by a chemical formula.
3. Forces weaker than covalent bonding exist between molecules. The structure of a metal involves positive ions embedded in a sea of electrons.
4. 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.
5. 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).
6. 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.
7. 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).
8. 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.
9. 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

14.1.0 Exothermic reactions
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!

14.1.1 Heat of solution of anhydrous copper (II) sulfate
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.

14.1.2 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.

14.1.3 Sulfuric acid with water
Be careful! When diluting strong acids always slowly add ACID to WATER. Never add water to acid.
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. Pass hydrogen chloride gas into water. Add acetic acid to water. Acetic acid, a weak acid, produces less heat than the strong acids sulfuric acid and hydrochloric acid.
1. 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.
2. 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. 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, steel wool
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 simple calorimeter, dilute hydrochloric acid and sodium hydroxide solution
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, which is 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 hydrochloric acid is therefore equal to 005 kg, i.e. m₁ = 005 kg, so is the mass of 50 mL of 1 mol / L sodium hydroxide 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)
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. Wait until the initial temperatures are the same, then add the 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 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.

14.1.5.1 Heat of neutralization with a simple calorimeter, dilute hydrochloric acid / ethanoic acid with sodium hydroxide solution
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 lab, the following simple apparatus can be used instead.
1. 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 existing in molecular form in aqueous solution. So some energy released by the neutralization must be used to ionize the ethanoic acid molecules.
2. 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
3. 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

14.1.6 Heat of displacement reaction, zinc with copper (II) sulfate solution
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 or a glass bottle, 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 when the reaction stops. Replace the stopper, invert the bottle, and shake gently. Record the highest temperature reached. Calculate the rise of temperature. This rise of temperature in not affected by the volume of 0.2 M copper (II) sulfate used for the experiment. For a 1 M solution, multiply the rise in temperature by 5 (5 × 0.2M = 1.0 M). The reactants lost energy to the solution. The temperature change is usually between 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 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 it is reduced to green chromium (III) oxide.
2CrO₃ (s) + C₂H₅OH (l) + 1½ O₂ (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 Sodium acetate heat pack
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 54ºC in an open container. However, when this solution is in a sealed container, the solution can be cooled below this temperature, as low as -10ºC. 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 its freezing point. This super-cooled 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 so the heat pack returns back to a liquid state and then is allowed to cool below its freezing temperature. It is then ready to be activated again.

14.1.12 Iron powder heat pack
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. Observe the heat energy given off when iron rusts in a 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. Record what happens.
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 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.

14.2.0 Endothermic reactions
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

14.2.1 Reactions of ammonium salts and potassium salts with water
See 3.81: Endothermic reactions take in heat energy
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 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.

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

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

14.2.5 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.

14.2.6 Combustion of potassium nitrate, fire line on paper, string fuse
Combustion of potassium nitrate, fire line on paper, string fuse
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. Some people use potassium nitrate solution to write their own name or the name of their school on paper then see the name burst into flames.
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.

14.2.7 Refrigerants
1. Sodium chloride + ice
2. Potassium nitrate + ammonium chloride + water
3. Potassium nitrate + ammonium chloride + sodium disulfate + water
4. Ammonium nitrate + water
5. Sodium sulfate + dilute hydrochloric acid
6. Sodium sulfate crystals + dilute sulfuric acid
7. Pass a current of air through methylene chloride

14.3.0 Chemiluminescence, bioluminescence
See diagram 14.3.0: Luminol structural formula and equations
Order online: Photo Luminescent Paint, phosphorescence, photo-electric effect
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. 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). Students may have experience of a
chemiluminescence if they have seen or purchased 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, fluorophor, 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.
Other fluorophors include:
(1.) 1-chloro-9,10-bis(phenyl ethynyl) anthracene emits green-yellow light in Cyalume sticks,
(2.) 2-chloro-9,10-bis(phenyl ethynyl) anthracene emits green light in Cyalume sticks,
(3.) 1,8-dichloro-9,10-bis(phenyl ethynyl) anthracene emits yellow light in Cyalume sticks,
(4.) 9,10-diphenylanthracene (DPA) emits blue light,
(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) emits green light,
(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.
Phosphorescence is the green glow from the slow oxidation of white phosphorus. It is an example of chemiluminescence.

14.3.0.1 Luminol
Order online: Cool Blue Light Kit, luminol, chemiluminescence, forensic science
Order online: Cool Chemical Light Powder,
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 light energy. 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. Students may have seen a television series where detectives use luminol to find blood stains, although the murderer tried to wipe clean all the blood at the crime scene. In the television story, the detective applies 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. In a laboratory you can 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.

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.

12.8.1 Reactions 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.
Prussian blue as a dye is made by adding iron (II) sulfate to potassium ferrocyanide, with the later addition of iron (III) chloride. Prussian blue can be distilled to yield prussic acid, hydrocyanic acid, HCN, which is very poisonous.
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)₂·3H₂O 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.

12.8.3 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. Repeat the experiment by substituting other oxidizing materials, e.g. chlorine, bromine, potassium permanganate or hydrogen peroxide, for nitric acid in the above experiment.

12.8.4 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, FeO·Fe₂O₃, that is weakly magnetic.

12.8.5 Reduce 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⁺

12.8.6 Heat iron filings with powdered sulfur
See 12.2.2.1: Heat iron with sulfur
Grey iron (II) sulfide forms, FeS. It is ferrimagnetic.
8Fe + S8 --> 8FeS (direct union of elements to form compounds)

12.8.7 Prepare iron (II) oxide, FeO
Close it 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.

12.8.8 Heat iron (II) sulfide, FeS₂ (pyrite) fool's gold
Iron (III) oxide and sulfur dioxide forms. (FeS₂ is not iron (IV) sulfide.)
4 FeS₂ (s) + 11 O₂ --> 2 Fe₂O₃ (s) + 8 SO₂ (g)

12.8.9 Prepare iron (III) oxide, Fe₂O₃
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. Boil some oxide in concentrated hydrochloric acid and show that it is a base.

12.8.10 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

12.8.11 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 FeSO₄·7H₂O, green vitriol. In the air, iron (II) salts are oxidized to iron (III) salts, so brown iron (III) hydroxide and iron (III) sulfate may form on the blue-green crystals.

12.8.12 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 FeCl₂·4H₂O. In the air, the iron (II) is oxidized to FeCl₃ and Fe₂O₃

12.8.13 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

12.8.14 Prepare iron (II) ammonium sulfate (NH₄)₂SO₄.FeSO₄.6H₂O
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.

12.8.15 Detect iron in fruit juice using back 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-haeme iron from the fruit juice with the tannins in the black tea. The non-haeme 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⁺, haeme iron comes mainly from haemoglobin and myoglobin in red meat.