School Science Lessons
2016-09-02 SP MF
Please send comments to: J.Elfick@uq.edu.au
14.0 Thermochemistry, heat of reaction
Table of contents
Chemical
14.3.0 Chemiluminescence, bioluminescence

14.3.0a Phosphorescence

14.2.0 Endothermic reactions take in heat energy

14.01 Energy of reactions, heat of reaction

14.1.0 Exothermic reactions give out heat energy

12.8.0 Reactions of iron, Fe
12.19.1.0 Properties of halogens

Chemical
Chemicals, Chemistry
Chemical potential energy, enthalpy
Chemical reactions, chemicals

14.1.0 Exothermic reactions give out heat energy
14.1.0 Enthalpy of reaction, heat of reaction
Experiments
3.80 Exothermic reactions & rise in temperature
14.2.6 Combustion of potassium nitrate, string fuse
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
3.82 Heat of neutralization reactions, HCl, NaOH
14.1.5 Heat of neutralization of HCl with NaOH
14.1.5.1 Heat of neutralization with calorimeter
14.1.9 Heat of reaction, Cr (VI) oxide with ethanol
14.1.8 Heat of reaction, KMNO4 with ethanol
14.1.7 Heat of reaction, KMNO4 with glycerol
14.1.10 Heat of reaction, Kwith diethyl ether
14.1.4 Heat of rusting, steel wool
14.1.1 Heat of solution of copper (II) sulfate
14.1.12 Iron powder heat pack
3.80 Reactions that give out heat energy
14.1.11 Sodium acetate heat pack
14.1.13 Sodium thiosulfate heat pack
14.1.3 Sulfuric acid with water

14.2.0 Endothermic reactions take in heat energy
14.2.0 Endothermic reactions
14.2.4 Ammonium nitrate cold pack
14.2.7 Refrigerants
Experiments
14.2.3 Ammonium carbonate with ethanoic acid
3.81 Endothermic reactions take in heat energy
14.2.8 Methylene chloride
14.2.5 Potassium nitrate with water
3.81 Reactions that take in heat energy
14.2.1 Reactions of NH4 and K salts with water
14.2.2 Urea with water

14.3.0 Chemiluminescence, bioluminescence
14.3.0 Chemiluminescence, bioluminescence
14.3.0a Fluorophores
14.3.1 Luminol tests for blood, Cu, Fe, Cn-

12.8.0 Reactions of iron, Fe
12.8.4 Burn steel wool
12.8.10 Black iron oxide is a mixed base, Fe3O4
12.8.15 Detect iron in fruit juice using black tea
12.8.13 Heat hydrated iron chlorides
12.8.8 Heat iron (II) sulfide, FeS2 (pyrite)
12.8.6 Heat iron filings with powdered sulfur
12.2.2.1: Heat iron with sulfur, synthesis reaction
12.8.12 Iron displaces H from HCl
12.8.11 Iron displace H from H2SO4
12.8.3 Oxidation of iron (II) salts
12.8.7 Prepare iron (II) oxide, FeO
12.8.14 Prepare iron (II) ammonium sulfate
12.8.9 Prepare iron (III) oxide, Fe2O3
12.8.5 Reduce iron (III) salts
12.8.1 Reactions of iron (II) salts & iron (III) salts
15.3.0 Rusting

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 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 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!

14.1.1 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.

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.

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, 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 calorimeter, dilute Hcl and NaOH 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)

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.
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 Heat of neutralization with calorimeter, dilute HCl / 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 laboratory, 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
CH3COOH (aq) + NaOH (aq) --> CH3COONa (aq) + H2O (l)

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 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 × 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 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, (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 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
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 commercial "heat packs".

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 -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 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 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. 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)
Commercial hand warmers contain iron powder, water, soium 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 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 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 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.
Students may 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

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 39.6oC and is used in the drinking bird heat engine.
24.3.7 Drinking bird heat engine, drinking duck, dippy bird, dunking bird

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.
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.
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, 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.

14.3.0a Fluorophores
(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), C30 H18 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, green light
(10.) Rhodomine tetramethyl isothionate, TRITC, "Rhodominr Red-X" orange light
(11.) Hydroxycoumarin, blue light
(12.) Cyanine dyes, Cy2 cyanine, Cy3 indocarbocyanine, Cy5 indodicarbocyanine
(13.) Biological fluorphores, e.g. green fluorescent protein, GFP, from a jellyfish
(14) Quantum dot nanocrystals

14.3.1 Luminol tests for blood, Cu, Fe, Cn-
"Cool Blue Light Kit", luminol, chemiluminescence, forensic science, (commercial)
"Cool Chemical Light Powder", luminol, (commercial)
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.
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)63-.
So a thiocyanate solution can be used as a test for iron (III) ions because iron (II) ions do not cause a colour change.

6. 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.

7. 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)

8. 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)

9. 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. Pass chlorine gas through an iron salt solution.
Brown iron (IIII) 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.

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, FeOFe2O3, 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 + + S2-
S2- - 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
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.)
4FeS2 (s) + 11O2 --> 2Fe2O3 (s) + 8SO2 (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 FeSO47H2O, 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 FeCl24H2O.
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 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-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.

12.19.1.0 Properties of halogens
Be careful! Chlorine and bromine are harmful when inhaled or when they contact the skin.
Fluorine, F, is a yellow gas.
Chlorine, Cl, is a green gas.
Bromine, Br, is a red-brown liquid.
Iodine, I, is a grey-black crystal.
All halogens are slightly soluble in water to form weak acidic solutions that are bleaches.
Chlorine is the most soluble and the strongest bleach.
Test each solution with universal indicator.
Universal indicator turns red then is bleached.

1. Compare the colours and states of the elements at room temperature.
At room temperatures fluorine is a pale yellow gas, chlorine is a green-yellow gas, bromine is a red-brown liquid that gives a brown
vapour and iodine is a grey-black solid.

2. Compare the colours, states and solubility in water of sodium fluoride, chloride, bromide and iodide.
They are all soluble, white, crystalline solids.

3. Compare the activity of fluorine, chlorine, bromine and iodine by investigating, which will displace another from their compounds.
Prepare, sodium fluoride solution, sodium chloride solution, sodium bromide solution and sodium iodide solution.
Add chlorine solution to each solution.
Chlorine has no visible effect on sodium fluoride solution or sodium chloride solution.
Chlorine turns sodium bromide solution brown-yellow.
Chlorine turns sodium iodide solution deep brown.
So chlorine displaces bromine from sodium bromide and chlorine displaces iodine from sodium iodide.
However, chlorine does not displace fluorine from sodium fluoride.
So chlorine is more active than bromine and iodine but chlorine is less active than fluorine.

4. Compare the colours and solubility of silver fluoride, chloride, bromide and iodide by adding silver nitrate solution to sodium fluoride
solution, sodium chloride solution, sodium bromide solution and sodium iodide solution.
Silver fluoride is soluble.
Silver chloride, silver bromide and silver iodide are insoluble.
Silver chloride is white.
Silver bromide is very pale yellow.
Silver iodide is a deep yellow.
Table 12.19.1.0
Element ->
Fluorine Chlorine Bromine Iodine
Description ->
-
yellow gas,
irritating smell
green-yellow gas
irritating smell
red liquid,
irritating smell
black solid,
irritating smell
12.19.1.0
Prepare
halogens
electrolysis
of KHF2
-
heat mixture of
chloride, MnO2
& conc. H2SO4
heat mixture of
bromide, MnO2
& conc. H2SO4
heat mixture of
iodide, MnO2
& conc. H2SO4
Activity
-
-
very reactive, combines
with most  metals
& non-metals
very reactive, combines
with most metals
& non-metals
very reactive, combines
with most metals
& non-metals
reactive, combines
with most metals
& few non-metals
Replacing action
-
replaces all
halogens
replaces Br2 & I2 from
bromides & iodides
replaces iodine from
iodides
-
-
Oxidizing action
-
very powerful
oxidizing agent
very powerful
oxidizing agent
powerful
oxidizing agent
weak
oxidizing agent
Action with alkalis
forms fluoride
forms hypochlorite
forms hypobromite
forms hypoiodite
Halogen aciid
-
fuming gas,
weak acid
fuming gas,
strong acid
fuming gas,
strong acid
fuming gas,
strong acid