School Science Lessons
Topic 14 Thermochemistry, heat of reaction, chemiluminescence
Updated 2009-09-17
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites
Table of contents
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
14.2.0 Endothermic reactions, reactions that take in heat energy
14.3.0 Chemiluminesce, bioluminescence
14.1.0 Exothermic reactions, reactions that give out heat energy
3.80 Reactions that give out heat energy, exothermic reactions
14.1.1 Measure the heat of solution of anhydrous copper (II) sulfate
14.1.2 Reactions of copper (II) sulfate solution with magnesium
14.1.3 Reactions of concentrated sulfuric acid with water
14.1.4 Heat from rusting of steel wool
14.1.5 Measure heat of neutralization of hydrochloric acid with sodium hydroxide
3.82 Heat of a neutralization reaction
14.1.5.1 Measure heat of neutralization with a simple calorimeter, dilute hydrochloric acid / ethanoic acid with sodium hydroxide solution
14.1.6 Measure heat from a displacement reaction, zinc with copper (II) sulfate solution
3.83 Heat of a copper displacement reaction
14.1.7 Test heat of reaction, potassium permanganate with glycerol
14.1.8 Test heat of reaction, potassium permanganate with ethanol
14.1.9 Test heat of reaction, chromium (VI) oxide with ethanol
14.1.10 Test heat of reaction, potassium with diethyl ether
14.1.11 Sodium acetate heat pack
14.1.12 Iron powder heat pack
14.1.13 Sodium thiosulfate heat pack
5.01.2 Supersaturated solution of sodium ethanoate-3-water
3.52 Conditions necessary for rusting
3.52.1 The mass of iron and its temperature increases during rusting
3.52.2 Oxygen combines with iron during rusting
3.52.3 Metals can prevent rusting
14.2.0 Endothermic reactions, reactions that take in heat energy
3.81 Reactions that take in heat energy, endothermic reactions
14.2.1 Reactions of ammonium salts and potassium salts with water
14.2.2 Reactions of urea with water
14.2.3 Reactions of ammonium carbonate with ethanoic acid
14.2.4 Ammonium nitrate cold pack
14.2.5 Reactions of potassium nitrate with water
14.2.6 Refrigerants
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 Co).
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 react ants 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 (DeltaH) 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 (DeltaH) is positive. For exothermic reactions, DH (DeltaH) is negative.
A --> B + xkJ
i.e A --> B, DH =- xkJ / mol A
Be careful! The reactions may be vigorous!
14.1.1 Measure the 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 Reactions of copper (II) sulfate solution with magnesium
1. Pour concentrated copper (II) sulfate solution into the test-tube. Add very small pieces of magnesium ribbon until the blue colour disappears. BE CAREFUL! The reaction is vigorous. Record the change in temperature of the solution.
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 Reactions of concentrated sulfuric acid with water
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 from rusting of steel wool
See 3.52.2: Oxygen combines with iron during 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 Measure 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 X specific heat X change in temperature.
Q = (005 + 005) X 42 X (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 13oC. 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 Measure 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. Therefore, 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 Measure heat from a displacement reaction, zinc with copper (II) sulfate solution
See diagram 3.1.6: 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 10oC.
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 X 0.2M = 1.0 M). The reactants lost energy to the solution. The temperature change is usually between 9oC and 10oC.
Zn (s) + Cu2+ (aq) --> Zn2+ (aq) + Cu (s)
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 6oC and 7oC.
The zinc metal became zinc ions and copper ions became copper metal due to transfer of electrons from zinc metal to the copper ion. To get electrical energy, these electrons must flow in an external conductor, e.g. a wire, from the zinc to the copper. The potential or voltage will reflect the greater activity of zinc over copper. The current flowing will depend on the extent and rate of the reaction
14.1.7 Test 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 Test heat of reaction,  with 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 Test 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) + 3/2O2 (g) ---> Cr2O3 (s) + 2CO2 (g) + 3H2O (l)
14.1.10 Test heat of reaction,  potassium with diethyl ether
This experiment is too dangerous for schools but see the animation from Chemistry Department, University of Leeds, in "Interesting websites".
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 80oC. 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. Instant hot packs: 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. Rusting experiment: Label three test-tubes 1, 2 and 3. Roll up some iron wool, steel wool, to make three little balls that will fit inside the test-tubes. Put one ball in each test-tube. Leave test-tube 1 as a control. Add water to test-tube 2 until the ball is half covered. Shake the tube to wet all the steel wool. Add salt water to test-tube 3 until the steel wool is half covered. Shake the tube to wet all the steel wool 5 Put the three tubes where they won't be disturbed. Observe them after a week.
4. 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.
5. 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: Reactions that take in heat energy, endothermic reactions
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 Reactions of urea with water
Urea (carbamide, H2NCONH2) is a crystalline solid that is very soluble in water.
Use a test-tube containing 10 mL of water. Put a thermometer in the test-tube and record the initial temperature. Put 5 g of urea in the test-tube. Record the changes in temperature of the solution.
14.2.3 Reactions of 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 Reactions 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 90oC. 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 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
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 yellow-green cold light as it returns to the energy ground state.
Other fluorophors include:
(1.) 1-chloro-9,10-bis(phenyl ethynyl) anthracene emits yellow-green 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.
14.3.0.1 Luminol
The chemical luminol, 3-aminophthalhydrazide (C8H7N3O2) reacts with oxygen to produce an intermediate molecule 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 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 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 salicylate 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.