School Science Lessons
Topic 12
2018-11-01
Please send comments to: J.Elfick@uq.edu.au
12.0 Acids, chemical reactions, chemical equations

Table of contents
12.3.0 Acids

12.2.0 Chemical reactions, types of chemical reactions

12.1.0 Conservation of mass


12.3.0 Acids

12.3.0 Properties of acids

See: Acids (Commercial)

12.3.16 Acid dissociation constant, Ka

12.3.0.1 Amphoteric substances

12.10.0 Boric acid, ionization reaction

12.3.14 Concentrated acids with a non-metal, carbon

12.5.0 Nitric acid

12.3.0.4 pH

12.9.0 Phosphoric acid

12.3.0.2 Polyprotic acids

Experiments
12.3.15 Acids with salts
12.3.12 Concentrated acids with metals, nitric acid with copper
12.3.13 Concentrated acids with metals, sulfuric acid with copper
12.3.5.01 Copper oxide with sodium hydrogen sulfate
12.3.0.6 Ethanoic acid, CH3COOH, acetic acid
Hydrochloric acid
12.3.3.3 Iron with sodium hydrogen sulfate
12.3.3.2 Magnesium with sodium hydrogen sulfate
Nitric acid, HNO3
Nitrous acid, HNO2
12.3.1 Reactions of dilute acids
12.18.5 Reactions of sulfuric acid
12.3.0.3 Strong acids and weak acids, Ka, pKa
12.3.1 Taste of acids, solid acids in the home

12.2.0 Chemical reactions, types of chemical reactions
12.2.0 Chemical reactions, types of chemical reactions
3.80 Energy from chemical reactions
3.12.0 Heat energy from chemical reactions
12.2.1 Oxidation reactions
12.2.2 Reduction reactions
12.2.5.0 Acid-base reactions
12.2.3.0 Decomposition reactions
12.2.4.0 Displacement reactions, substitution reactions
12.2.7.0 Polymerization reactions
10.11.01 Precipitation
12.2.1.0 Precipitation reactions
12.2.6.0 Redox reactions
17.3.1.3 Disproportionation, Hydrogen peroxide with catalase enzyme in raw beef liver
12.2.2.0 Synthesis reactions (combination reactions, direct union of elements)

12.2.3.0 Decomposition reactions
12.2.3.0 Decomposition reactions
Experiments
3.30.0 Decomposition reactions (List)
12.2.3.04 Catalytic decomposition, catalysis
12.2.3.2 Decomposition of copper carbonate, prepare copper oxide
12.2.3.3 Decomposition of zinc carbonate, prepare zinc oxide
12.2.3.02 Electrolytic decomposition, electrolysis
12.2.3.03 Photo decomposition, photolysis

12.2.4.0 Displacement reactions
12.2.4.0 Displacement reactions (substitution reactions)
Experiments
12.2.4.1 Iron displaces copper in copper sulfate solution
12.2.4.3 Magnesium displaces hydrogen in ethanoic acid
12.14.1: Zinc displaces lead from lead nitrate solution

12.5.0 Nitric acid
12.5.0 Nitric acid, aqua fortis, spirits of nitre
12.5.01 Grades of nitric acid
Experiments
12.5.1 Prepare nitric acid (sulfuric acid with sodium nitrate)
12.5.2 Nitric acid with copper
12.3.11.1 Nitric acid with metals
12.5.3 Nitric acid with sulfuric acid

12.7.0 Nitrous acid
12.5.02 Ionization of nitrous acid
16.4.1.4 Nitration
12.7.2 Nitrous acid with amines
Experiments
12.7.3 Heat nitrous acid
12.7.1 Prepare nitrous acid
15.2.5.1 Nitrous acid as oxidizing agent or a reducing agent

12.2.1.0 Precipitation reactions
12.2.1.0 Double replacement reactions, metathesis
12.2.1a Prepare salts by chemical reactions
12.2.1b Ionic equations, double decomposition reactions
12.2.1c Coloured precipitates, double decomposition reactions
Experiments
12.2.1.2 Cobalt chloride with calcium hydroxide
12.2.1.1 Sodium carbonate with calcium hydroxide
12.2.1.4 Sodium carbonate with magnesium sulfate
12.2.1.3 Sodium carbonate with zinc sulfate
12.2.1.11 Sodium chloride with sulfuric acid
12.2.1.8 Sodium hydroxide with alum
12.2.1.6 Sodium hydroxide with cobalt chloride
12.2.1.5 Sodium hydroxide with copper sulfate
12.2.1.7 Sodium hydroxide with iron sulfate
12.2.1.10 Prepare soda lime
12.1.14 Prepare sodium hydroxide
3.71.4 Tests for precipitates when solutions added to lead (II) nitrate
3.71.3 Tests for precipitates, mix salt solutions

12.3.1 Reactions of dilute acids
Experiments
12.3.8 Dilute acids with acidic oxides
12.3.5.1 Dilute acids with amphoteric oxides
12.3.5 Dilute acids with basic oxides (metal oxides), copper (II) oxide
12.3.10.1 Dilute acids with calcium hydrogen carbonate
12.3.9.0 Dilute acids with carbonates, common carbonates
12.3.6 Dilute acids with hydroxides, magnesium hydroxide
12.3.7 Dilute acids with hydroxides, sodium hydroxide
12.10.2.1 Dilute acids with metals
12.3.2 Dilute acids with metals, hydrochloric acid
12.3.2.1 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
12.3.3.1 Dilute sulfuric acid with aluminium
12.3.3 Dilute sulfuric acid with steel wool
12.3.4 Dilute acids with non-metals, carbon, sulfur
12.3.10.0 Dilute acids with sodium hydrogen carbonate
12.3.7.1 Dilute acids with sodium hydroxide
12.3.9.1 Dilute hydrochloric acid with calcium carbonate
12.3.7.2 Dilute hydrochloric acid with hydroxides
12.3.9.2 Dilute hydrochloric acid with sodium carbonate
12.3.11.0 Dilute nitric acid with copper
12.18.5.4 Dilute sulfuric acid as an acid
12.18.5.5 Dilute sulfuric acid as a sulfate
12.3.3.1 Dilute sulfuric acid with aluminium
12.3.9.6 Dilute sulfuric acid with calcium carbonate
12.3.9.4 Dilute tartaric acid with egg shell, soil, wood ash
12.3.9.3 Dilute tartaric acid with sodium carbonate
12.18.5.4 Reactions of dilute sulfuric acid as a acid
12.18.5.5 Reactions of dilute sulfuric acid as a sulfate:

12.2.6.0 Redox reactions
12.2.6.0 Redox reactions (oxidation-reduction reactions, electron transfer reactions)
Experiments
12.2.6.1 Magnesium with dilute hydrochloric acid
12.2.6.2 Oxygen with sulfur dioxide
12.2.6.3 Ammonia with copper oxide
12.2.6.4 Chlorine with water
12.2.6.5 Copper (I) oxide with hot dilute sulfuric acid

12.2.2.0 Synthesis reactions, (combination reactions, direct union of elements)
Experiments
12.2.2.0 Synthesis reactions (combination reactions, direct union of elements)
12.2.2.5 Heat copper sulfate crystals
12.2.2.2 Heat copper wire with iodine crystals
12.2.2.8 Heat copper with sulfur
12.2.2.7 Heat iron with copper
12.2.2.1 Heat iron with sulfur
12.2.2.3 Heat steel wool with iodine crystals
12.2.2.6 Heat zinc with sulfur
12.2.2.4 Zinc with iodine solution

3.71.3 Tests for precipitates, mix salt solutions
Mix 5 mL of each of the different solutions of salts available in your laboratory.
Record observations for each pair of solutions.

3.71.4 Tests for precipitates, acids with lead (II) nitrate
Add lead (II) nitrate to the following: 1. Dilute hydrochloric acid, 2. Dilute sulfuric acid, 3. Sodium hydroxide solution.
A precipitate forms in each test-tube.

12.2.0 Chemical reactions, types of chemical reactions
Chemical reactions involve energy changes.
All chemical reactions involve energy transformations.
The spontaneous directions of chemical reactions are towards lower energy and greater randomness.
In a chemical reaction, a chemical change occurs where elements or compounds (reactants) form new substances (products).
Specific criteria can be used to classify chemical reactions.
Neutralization reactions are reactions producing an aqueous solution with equal concentrations of hydroxyl and hydrogen ions.
Redox reactions (reduction + oxidation), are reactions where one species is oxidized and another species is reduced.
12.2.1 Oxidation reactions
12.2.2 Reduction reactions, reduce
Reversible reactions occur together with their converse to form an equilibrium mixture of reactants and products.

The main types of chemical reactions are as follows:
1. A + B --> AB
A + B --> AB synthesis
A + O --> AO oxidation, combustion
AB + 2O --> AO + BO

2. AB --> A + B
AB --> A + B decomposition, thermal decomposition, reduction.

3. A + BC --> AC + B
A + BC --> AC + B single replacement, precipitation
copper + silver nitrate --> silver + copper nitrate
chlorine + sodium iodide --> sodium chloride + iodine
2A + BC --> BA + CA oxidation.

4. AB + CD --> AD + CB
AB + CD --> AD + CB double replacement (displacement, metathesis), neutralization
copper + silver nitrate --> silver + copper nitrate
chlorine + sodium iodide --> sodium chloride + iodine
HCL + NaOH --> H2O + NaCL
lead nitrate + potassium iodide --> lead iodide + potassium nitrate
precipitation
AB (aq) + CD (aq) --> AD (aq) + CB (s)
Combustion reaction
CxHy + O2 --> CO2 + H2O
methane + Oxygen --> CO2 + H2O
Hydrolysis reaction
X + H2O --> HX + OH.

12.2.1 Oxidation reactions
1. Oxidizing means adding oxygen to a substance.
2. The addition of oxygen to, or the loss or removal of hydrogen from, a compound.
3. The increase in the proportion of negative constituents in a molecule or compound.
4. The loss or removal of an electron from an atom or molecule.

12.2.2 Reduction reactions, reduce
1. The use of a chemical reaction to reduce a substance to a simpler form.
2. The loss or removal of oxygen from, or the addition of hydrogen to, an atom or molecule.
3. The decrease in the proportion of electronegative constituents in a molecule or compound.
4. The lowering of the oxidation number of an atom, the charge in negative form of the electron charge of an atom.
5. The conversion of a metallic ore or oxide to a metal by smelting.
Reduce copper oxide with natural gas, methane: 16.5.1.4
Reduce copper (I) oxide (copper oxide) to copper: 10.10.2
Reduce iron (III) chloride with sulfur dioxide: 3.51.3
Reduce metal oxides to metals with hydrogen gas: 3.41.7
Reduce metal oxides to metals, red lead to lead and oxygen: 10.10.1
Separate to metals by reduction of metal oxides, charcoal blocks: 10.10.0
Reduce potassium manganate (VII) by sulfur dioxide: 3.51.2
Reduce red iron oxide, or rust, to iron: 10.10.3.

12.2.1.0 Double replacement reactions, metathesis
Double replacement reactions, double exchange reactions, metathesis, double decomposition reactions, precipitation reactions, neutralization reactions
Metathesis occurs when radicals are exchanged.
A reaction between ions is shown by precipitation of an insoluble salt as a solid.
Use (aq) to show a solution and use (s) to show a precipitate, solid.
AgNO3 (aq) + HCl (aq)--> HNO3 (aq) + AgCl (s)
Metathesis is a chemical reaction between two substances that produces two other substances with either ionic or covalent bonds.
AB + CD ---> AC + BD
The above reaction may also be called a double decomposition reaction (metathesis) because the positive and negative parts of two
compounds swap partners, i.e. exchange radicals.
It can occur if one of the products (substances formed) is insoluble or is a gas.
However, some people think a double decomposition reaction is similar, except one of the substances does not dissolve in the solvent
AB (aq) + CD (s) --> AD (aq) + CB (s)
Precipitation reactions result in the appearance of a solid from reactants in aqueous solution.

12.2.1a Prepare salts by chemical reactions
Salts can be prepared by the action of acids with alkalis, carbonates, metals, metal oxides, and by replacement and double
decomposition reactions.
A salt contains a metal and part of an acid, e.g. copper sulfate from sulfuric acid, sodium chloride from hydrochloric acid.
A salt is a compound formed when the hydrogen of an acid is replaced by a metal.
For example, when zinc reacts with hydrochloric acid it replaces the hydrogen and forms the salt, zinc chloride.
The hydrogen comes away as hydrogen gas.
Zn + 2HCl --> ZnCl2 + H2.

Experiments
1. Add silver nitrate solution to sodium chloride solution.
A silver chloride precipitate forms that can be separated from the sodium nitrate solution.
AgNO3 (aq) + NaCl (aq) --> AgCl (s) + NaNO3 (aq)
silver nitrate + sodium chloride --> silver chloride + sodium nitrate
Be careful! Silver nitrate is expensive!

2. Add silver nitrate solution to potassium chloride solution.
A silver chloride precipitate forms that can be separated from the potassium chloride solution.
AgNO3 (aq) + KCl (aq) --> KNO3 (aq) + AgCl (s)
silver nitrate + potassium chloride --> silver chloride + potassium nitrate
Be careful! Silver nitrate is expensive!

12.2.1b Ionic equations, double decomposition reactions
AgNO3 (aq) + NaCl (aq) --> AgCl (s) + NaNO3 (aq)
An ionic equation that shows all the substances
Ag+ (aq) + NO3- (aq) + Na+ (aq) + Cl- (aq) --> AgCl (s) + Na+ (aq) + NO3- (aq)
1. A net ionic equation that does not contain the "spectator ions" that appear on both sides of the equation but do not form a precipitate
Ag+ (aq) + Cl- (aq) --> AgCl (s).

2. silver nitrate + potassium chloride ---> potassium nitrate + silver chloride (white precipitate)
AgNO3 (aq) + KCl (aq) ---> KNO3 (aq) + AgCl (s)
Ag+ (aq) + Cl- (aq) ---> AgCl (s).

3. calcium chloride + sodium carbonate --> calcium carbonate + sodium chloride
CaCl2 (aq) + Na2CO3 (aq) --> CaCO3 (s) + 2NaCl (aq)
Ca2+ (aq) + CO32- (aq) --> CaCO3 (s).

12.2.1c Coloured precipitates, double decomposition reactions
A reaction that forms a coloured precipitate is a good way to show double decomposition reactions.
1. lead (II) nitrate (aq) + potassium dichromate (aq) ---> lead (II) chromate (IV) (s) (yellow precipitate, chrome yellow)
Pb2+ (aq) + CrO42- (aq) ---> PbCrO4 (s)
2. silver nitrate (aq) + potassium chromate (aq) ---> silver chromate (s) (red precipitate)
Ag+ (aq) + CrO42- (aq) ---> AgCrO4 (s)
3. lead nitrate (aq) + potassium iodide (aq) ---> lead iodide (s) (yellow precipitate)
Pb2+ (aq) + 2I- (aq) ---> PbI2 (s)
4. copper (II) sulfate (aq) + sodium carbonate (aq) ---> copper carbonate (s) (green precipitate)
Cu2+ (aq) + CO32- (aq) ---> CuCO3 (s).

12.2.1.1 Sodium carbonate with calcium hydroxide, double decomposition
To half a test-tube of sodium carbonate solution add lime water (calcium hydroxide solution).
Describe the reaction that occurs.
A white solid (precipitate) of calcium carbonate forms.
sodium carbonate + calcium hydroxide --> calcium carbonate + sodium hydroxide.

12.2.1.2 Cobalt chloride with calcium hydroxide, double decomposition
To half a test-tube of cobalt chloride solution add lime water (calcium hydroxide solution).
Describe the reaction that occurs.
A pink / mauve precipitate of cobalt hydroxide forms.
cobalt chloride + calcium hydroxide --> cobalt hydroxide + calcium chloride.

12.2.1.3 Sodium carbonate with zinc sulfate, double decomposition
Add sodium carbonate solution to the zinc sulfate solution.
Filter the mixture to obtain the precipitate of zinc carbonate.
Allow the filter paper, unfolded, to dry, and scrape off the white powder.
ZnSO4 (aq) + Na2CO3 (aq) --> ZnCO3 (s) + Na2SO4 (aq)
sodium carbonate + zinc sulfate --> sodium sulfate + zinc carbonate
The spectator ions, Na+ and SO42-, remain in solution.

12.2.1.4 Sodium carbonate with magnesium sulfate, double decomposition
Add magnesium sulfate solution to sodium carbonate solution.
Filter the solution to obtain the white precipitate of insoluble magnesium carbonate.
MgSO4 (aq) + Na2CO3 (aq) --> MgCO3 (s) + Na2SO4 (aq)
magnesium sulfate + sodium carbonate --> magnesium carbonate + sodium sulfate
The spectator ions, Na+ and SO42-, remain in solution.

12.2.1.5 Sodium hydroxide with copper sulfate
Copper ions, Cu2+, are precipitated from solution by addition of hydroxide anions, OH-.
Add 5 mL of sodium hydroxide solution to the same volume of copper sulfate solution.
A pale blue precipitate of insoluble copper hydroxide forms.
CuSO4 + 2NaOH --> Cu(OH)2 + Na2SO4
Cu2+ (aq) + 2OH- (aq) --> Cu(OH)2 (s)
sodium hydroxide + copper sulfate > copper hydroxide + sodium sulfate.
The spectator ions, Na+ and SO42-, remain in solution.

12.2.1.6 Sodium hydroxide with cobalt chloride
Add 5 mL of sodium hydroxide solution to the same volume of cobalt chloride solution.
A blue-green precipitate of cobalt hydroxide forms.
2NaOH + CoCl2 --> Co(OH)2 + 2NaCl
Sodium hydroxide + cobalt chloride > cobalt hydroxide + sodium chloride.
A double replacement reaction.
The spectator ions, Na+ and Cl -, remain in solution.

12.2.1.7 Sodium hydroxide with iron (II) sulfate
Add 5 mL of sodium hydroxide solution to a dilute solution of iron sulfate.
Stopper the test-tube, shake well, and leave to stand.
Both solutions are colourless.
A light green precipitate of iron (II) hydroxide, ferrous hydroxide, forms that turns orange-brown on standing
The green iron hydroxide first forms, but it soon reacts with oxygen gas to form a different type of iron hydroxide, which is brown.
FeSO4 (aq) + 2NaOH (aq) --> Fe(OH)2 (s) + Na2SO4 (aq)
Pure iron (II) hydroxide is white, but in the presence of oxygen it forms a green rust as Iron (III ions form.

12.2.1.8 Sodium hydroxide with alum
Alum, aluminium potassium sulfate, Al2(SO4)3.K2(SO4).24H2O
Add 5 mL of sodium hydroxide solution to a dilute solution of alum.
A faint white precipitate forms.
The part of alum that reacts with the alkali is aluminium sulfate.
Al2(SO4)3 (s) + 6NaOH --> 2 Al(OH)3 (s) + 6Na+ (aq) + 3(SO4)2- (aq).

12.2.1.10 Prepare soda lime
Dissolve calcium oxide in sodium hydroxide solution and leave to dry so that a mixture of sodium hydroxide and calcium oxide forms,
NaOH + CaO.
Leave the product to dry.
The CaO with NaOH keeps the hygroscopic NaOH dry, to aid fusion of a dry product, soda lime.
Soda lime is toxic if ingested and is corrosive.
Sold as: Soda lime, CaO / NaOH, granular, pellets.

12.2.1.11 Sodium chloride with sulfuric acid
3.71.1 Solubility table and solubility rules
1. Concentrated sulfuric acid with solid sodium chloride BE CAREFUL!
The reactions contain no water.
Two reactions occur and both go to completion if heated.
The reactions occur because hydrogen chloride has a lower boiling point than sulfuric acid.
2NaCl (s) + H2SO4 (l) --> Na2SO4 (aq) + 2HCl (g)
NaCl (s) + H2SO4 (l) --> NaHSO4 (aq) + HCl (g)
2. Dilute sulfuric acid with solid sodium chloride: The reaction does not go to completion because the hydrochloric acid dissolves in the
water.
One product of the reaction is a slightly ionized substance, e.g. water.
In neutralization reactions HOH is forming, so the reaction can almost go to completion.
One product of the reaction is a precipitate.
An insoluble substance leaves the solution.
The solubility rules state that all chlorides are soluble except Ag+, Hg2+ and Pb2+ (slightly).
Predict whether the following reaction occurs.
The reaction occurs because insoluble silver chloride precipitates.
NaCl (aq) + AgNO3 (aq) --> NaNO3 (aq) + AgCl (s)
NaOH (aq) + HCl (l) --> NaCl (aq) + H2O (l).

12.2.2.0 Synthesis reactions, (combination reactions, direct union of elements)
Reactions of mixtures of two elements, iron with sulfur, copper with sulfur, zinc with sulfur, zinc with iodine to form compounds.
Be careful! The following reactions are vigorous.
Do not use large quantities of the chemicals.
Use eye protection.
Do not get close to the fumes from the reaction.
Elements or simple molecules combine to form a new compound.

12.2.2.1 Heat iron with sulfur
Heat iron filings with sulfur powder (synthesis reaction)
See diagram 12.2.1: FeS
S8 (s) + 8Fe (s) --> 8FeS (s)
Be careful! The following reactions are vigorous.
Do not use large quantities of the chemicals.
Be careful! The reaction of iron (II) sulfide with hydrochloric acid will form the poisonous gas, hydrogen sulfide, with an odour of rotten eggs.

1. Mix half a metal bottle top of powdered sulfur with the same volume of iron filings.
Heat a small portion of the mixture on the metal bottle top with the cork removed or in a hard glass test-tube.
When the reaction begins, i.e. the mixture starts to glow, stop heating by moving the Bunsen burner to the side.
If the glow stops, heat the test-tube again.
The reaction of a mixture of iron with sulfur gives out so much heat that the mixture becomes red hot.
Note the following different properties of powdered sulfur, iron filings and the product, the compound iron (II) sulfide: 1.1 appearance,
1.2 colour, 1.3 hardness, 1.4 magnetism.
Iron is magnetic so is easily removed from a mixture of iron and sulfur but iron (II) sulfide is not magnetic.
Fe + S ---> FeS (s).

2. Use < 10 g total material iron with sulfur in a fume cupboard.
Heat the mixture to start the reaction.
However, be aware that unreacted sulfur may catch fire and produce sulfur dioxide gas to irritate the lungs.

3. Mix uniformly reduced iron powder and powdered sulfur in a weight ratio of seven to four.
See diagram 12.2.1: FeS
Carve the word "FeS" on a red coloured brick with a knife.
Spread the iron sulfur mixture throughout the word groove and press the powdered mixture solid.
Heat one tip of a glass rod until red hot with an alcohol burner and then immediately dig the hot tip into the mixture at one end of the
word groove.
A chemical reaction is starts immediately.
The reaction continues violently to release a large amount of heat and meanwhile to develop rapidly a red glow, which looks like a
small "fiery dragon".
The heat lost by the reaction is more than the heat needed to start the reaction.
The reaction produces a new black solid substance, iron (II) sulfide, that has different properties from the two reactants, iron and sulfur.
Compare iron powder, powdered sulfur and iron (II) sulfide.
Note their appearance.
Test them respectively with a magnet.
Add in drops hydrochloric acid solution to them respectively.

3. Mix equal amounts of iron filings and powdered sulfur.
Heat the mixture in a crucible or a small tin with sand in the bottom.
The sand prevents the bottom of the tin from melting by spreading the heat.
Heat the mixture strongly until you see a red glow spreading through the mass.
The heat lost by the chemical reaction is more than the heat needed to start the reaction.
The reaction forms a new substance iron (II) sulfide that has different properties from the two elements used to make it.
Compare iron filings, powdered sulfur, and iron (II) sulfide.
Note their appearance.
Test with a magnet.
Add drops of hydrochloric acid.

4. Make a mixture of 7 parts of iron filings with 4 parts of sulfur powder in a sealed plastic bag.
Hold a magnet over the plastic bag to show that the iron filings can be easily separated from the mixture.
Quarter fill an ignition tube with the mixture.
Near an open window or in a fume cupboard heat the end of the ignition tube with a Bunsen burner.
When the mixture glows move the Bunsen burner away but when the glow stops move the Bunsen burn back again until all the mixture
reacts.
Leave the ignition tube to cool
then move the magnet near it.
The magnet can no longer attract iron filings or the iron sulfide in the ignition tube.

12.2.2.2 Heat copper wire with iodine crystals (synthesis reaction)
1. Heat a mixture of iodine crystals and copper wire in a hard glass test-tube.
Stop heating when you hear a hissing noise.
Heat again to make sure all the copper reacts with the iodine.
Excess iodine sublimes and solidifies up the test-tube.
Let the test-tube cool then scrape out the product of the reaction.
Compare the crushed product with the reactants copper and iodine.
The reaction forms a new substance, copper (I) iodide.
Cu2+ + 2I- --> CuI2.

2. Use a fume cupboard.
Dip a strip of sheet copper into iodine crystals in a hard-glass test-tube.
Use a Bunsen burned to heat the copper touching the iodine crystals.
Yellow copper iodide forms on the copper strip.

12.2.2.3 Heat steel wool with iodine crystals
Use a fume cupboard.
Put iodine crystals in a test-tube and then push in a plug of steel wool.
Clamp the test-tube at an angle and heat the steel wool with a Bunsen burner.
The steel wool glows red and the iodine evaporates.
The purple iodine vapour reacts with the hot steel wool to form iron (II) iodide, by direct synthesis.
Fe (s) + I2 (g) --> FeI2 (s).

12.2.2.4 Zinc powder with iodine solution
1. Do not use a Bunsen burner in this experiment.
Dissolve iodine crystals in 5 mL of ethanol then note the temperature of the solution.
Add zinc powder to the solution then stir it with thermometer and note any rise in temperature.
When the solution becomes clear, filter it into a beaker leaving behind any excess zinc.
Place a watch glass over a beaker of hot water from an electric jug.
Pour the filtrate into the watch glass and observe the zinc iodide, a white salt formed by direct synthesis as the alcohol evaporates.
zinc + iodine --> zinc iodide.
Zn + I2 --> ZnI2.

2. Dissolve 2 g of iodine crystals in 25 mL of ethanol and note the temperature of the solution.
Add 2 g of zinc powder to the solution, stir and note the temperature.
The colour of the iodine is lost during the reaction and some of the zinc remains.
Filter off the unreacted zinc and heat the filtrate in an evaporating basin on an electric hot plate, not over a Bunsen burner.
A white odourless crystals of zinc iodide remain.
Anhydrous zinc iodide is a white hygroscopic substance used as an X-ray opaque substance to improve the fine details on X-ray plates.
Zn + I2 --> ZnI2
Repeat the experiment using antimony with iodine.

12.2.2.5 Heat copper sulfate crystals
1. Prepare white copper sulfate by heating a finger width of blue copper sulfate crystals in an evaporating basin, while stirring with a
glass rod.
Do not overheat, and stop heating as soon as the chemical has turned white.
Leave to cool.
Add drops of water until the powder is blue but still dry.
The white compound has combined chemically with the water to form the blue compound, because all the added water has
disappeared.
This reaction is the combination of an anhydrous salt with water to form a hydrate.
CuSO4.5H2O (s) <--> CuSO4 (s) + 5H2O
blue crystals <--> white powder.

2. Heat 1.66 g of blue copper sulfate crystals in a hard-glass test-tube.
Steam from the crystals condenses in the upper part of the test-tube.
The blue crystal turn into white, or blue-white, powder, weight 1.06 g.

12.2.2.6 Heat zinc with sulfur
Heat a mixture of sulfur in the bottom of a test-tube and a strip of zinc half way up the test-tube to form
the compound zinc sulfide
Zn (s) + S (s) --> ZnS (s)
zinc + sulfur --> zinc sulfide
Be careful!
Zinc powder and sulfur react violently with a yellow-green flame to form yellow zinc sulfide.
The reaction has been used to propel rockets
Zn (s) + S (s) --> ZnS (s).

12.2.2.7. Heat iron with copper
Heat a mixture of iron filings and copper turnings in a hard glass test-tube.
No reaction of iron filings with copper is observed because no new compound forms.

12.2.2.8 Heat copper with sulfur
Do this experiment in a fume cupboard.
Copper sulfide is an irritant.
Put sulfur powder and 10 cm of copper wire in a hard-glass test-tube.
Heat the mixture with a Bunsen burner.
The sulfur melts and then forms yellow vapour which reacts with the copper wire to form dark grey copper sulfide.
2Cu (s) + S (s) --> Cu2S (s).

12.2.3.0 Decomposition reactions
A compound breaks down into simpler compounds or into elements, usually caused by heat, the opposite of a synthesis reaction.
All compounds decompose on heating to a high enough temperature to form elements or simple molecules.
In a decomposition reaction a compound is broken into smaller chemical species.
A decomposition reaction may be thought of as the breakdown of a single phase, a molecule or a reaction intermediate into two or
more phases.

12.2.3.02 Electrolytic decomposition, electrolysis
Electrolytic decomposition, electrolysis, occurs when electric current passes through an aqueous solution of a compound.
2.1 Electrolysis of water, decomposition of water, Hofmann voltameter, electrochemical coulometer: 15.5.4
2.2 Hydrogen peroxide decomposition, with different catalysts: 17.3.1.0 (List)
2.3 Electrolysis of sodium chloride solution: 15.5.12
2.4 Decomposition of molten sodium chloride, to sodium and chlorine.

12.2.3.03 Photo decomposition, photolysis
Photo decomposition, photolysis, occurs when a substance is broken down by light, photons
Photolysis: 7.9.42
3.1 Silver chloride precipitate in photography: 7.8.7.1
3.2 Decomposition of silver bromide
3.3 Hydrogen peroxide decomposition, with different catalysts: 17.3.1.0 (List)
Decomposition of hydrogen peroxide: In the presence of light, hydrogen peroxide decomposes into water and oxygen.

12.2.3.04 Catalytic decomposition, catalysis
Catalytic decomposition, catalysis, occurs when a catalyst assists in the decomposition
4.1 Decomposition of potassium chlorate with manganese dioxide catalyst: 3.30.11
4.2 Hydrogen peroxide decomposition, with different catalysts: 17.3.1.0 (List)
Decomposition of dilute solutions of hydrogen peroxide (H2O2) into water and oxygen, with powdered manganese dioxide, to form
oxygen gas and water.

12.2.3.2 Decomposition of copper carbonate, prepare copper oxide
1. Prepare copper carbonate by mixing sodium carbonate solution and copper sulfate solution.
Pour off the liquid when the copper carbonate has settled in the test-tube.
Heat to evaporate remaining liquid and heat more strongly to form the oxide.
The oxide could be purified by washing with water, using a filtration apparatus.
Most carbonates decompose to form a metal oxide and carbon dioxide, e.g. copper carbonate.

2. Prepare copper carbonate by adding half a test-tube of copper sulfate solution to the same quantity of sodium carbonate solution.
Leave the insoluble copper carbonate to settle then pour off the liquid above the precipitate.
Gently heat the carbonate just enough to drive off the remaining water as steam.
Fit a stopper and bent tube to the test-tube.
With the end of the bent glass delivery tube dipping into the lime water solution, heat the copper carbonate more strongly.
The copper carbonate turns black and the lime water turns milky.
The copper carbonate has been decomposed by the heat into black copper oxide and carbon dioxide gas, which turns lime water milky.

12.2.3.3 Decomposition of zinc carbonate, prepare zinc oxide
Heat part of the zinc carbonate, from experiment 163, in an evaporating basin, or stand fairly strongly, and until the white powder, zinc
oxide turns yellow.
Allow it to cool and heat it again.
Note the colour changes.
Zinc oxide is yellow when hot, white when cold, so it is said to be thermochromic.
When white zinc oxide is heated it loses some oxygen to cause a yellow colour, but it returns to a white colour when cooled.
The carbonate decomposed on heating to form zinc oxide and carbon dioxide.
Most carbonates decompose to form a metal oxide and carbon dioxide, e.g. zinc carbonate.
ZnCO3 (s) --> ZnO + CO2 (g).

12.2.4.0 Displacement reactions (substitution reactions)
See 12.14.1: Zinc displaces lead from lead nitrate solution
The reactants are an element and a compound.
The element replaces part of the compound with the same valence and same sign, e.g. displacement of Cu2+ by zinc
Zn (s) + CuSO4 (aq) --> ZnSO4 (aq) + Cu (s)
zinc + copper sulfate --> zinc sulfate + copper
The copper precipitates as the element and the zinc metal goes into solution as zinc ions.

12.2.4.1 Iron displaces copper in copper sulfate solution
1. Prepare iron sulfate crystals.
Add iron filings to sodium hydrogen sulfate solution in a test-tube, and heat.
When there is no further reaction, filter the mixture and pour a finger width of sulfate solution of the filtrate into a evaporating basin.
Pale green crystals of iron sulfate form.

2. Add iron filings to half a test-tube of copper sulfate solution.
Leave until the colour of the solution changes from blue to pale green and the iron metal turns pink-brown as displaced copper is
deposited on it.
Filter the solution.
Pour part of the filtrate into an evaporating basin.
Pale green crystals of iron sulfate form no different from the crystals formed in the previous experiment.
Fe (s) + CuSO4 (aq) --> FeSO4 (aq) + Cu (s).

3. Place copper sulfate crystals in the test-tube and add a quarter of a test-tube of water.
Shake to get a blue solution.
Drop in the iron nail, which must not be rusty.
Leave for ten minutes.
Take out the nail.
The iron nail has turned a pinkish colour, due to a deposit of copper on it.
Iron + copper sulfate > copper + iron sulfate
Repeat the experiment but use a finger width of iron filings instead of the nail.
Leave the iron nail test-tube and contents for a few hours.
The blue liquid turns a pale green sulfate colour.
The iron has completely displaced the copper in the copper sulfate, forming a solution
of iron sulfate that is pale green.
Repeat the experiment using a centimetre of magnesium ribbon instead of the iron nail.
Observe the metal a few moments after it has been in the copper sulfate solution.
The metal turns pink or copper coloured.
Note whether iron or magnesium metal reacts more quickly.
Magnesium is a more reactive metal than iron.

12.2.4.3 Magnesium displaces hydrogen in ethanoic acid
1. All acids contain hydrogen, and many metals can displace it, thus setting the hydrogen gas free.
The acid in vinegar is ethanoic acid, acetic acid.
Add 3 cm of magnesium to half a test-tube of vinegar.
As soon as bubbles of hydrogen gas are coming well, hold your thumb or finger over the mouth of the test-tube for half a minute to trap
a quantity of hydrogen gas.
Mg (s) + 2CH3COOH (aq) --> Mg (CH3COO)2 (aq) + H2 (g).

Quickly hold the open test-tube to the spirit burner flame.
Describe what happens.
A small pop or squeak occurs.
It is a minor explosion.
If you got no result, repeat the procedure of trapping the gas and igniting it in the flame.
When hydrogen gas mixes with air it explodes, i.e. combines extremely rapidly with the oxygen gas in the air.

2. Repeat the experiment using tartaric acid, CHOHCOOH)2
Magnesium displaces hydrogen in tartaric acid.
Repeat the experiment with iron filings instead of magnesium.
You may have to heat the mixture.
Tartaric acid exists in rhubarb, so some people have experienced unpleasant symptoms, e.g. abnormal heart rhythm, after eating
magnesium and rhubarb.

12.2.5.0 Acid-base reactions
Acid-base reactions involve transfer of protons from donors to acceptors.
An acid dissociates in water to produce positive hydrogen ions, H+, that is solvated to produce hydronium ions (hydroxonium ions,
oxonium ions) H3O+, by transferring a proton (H+) to a water molecule.
HCl (g) + H2O (l) --> H3O+ (aq) + Cl- (aq)
A base dissociates in water to produce negative hydroxide ions, OH-.
NaOH --> Na+ (aq) + OH- (aq)
Acids react with bases to from salts and water.
The products are neither acidic nor basic so this reaction is called neutralization.
HCl + NaOH --> NaCl + H2O
hydrochloric acid + sodium hydroxide --> sodium chloride + water
The ionic equation that shows all the substances
H3O+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) --> 2H2O + Na+ (aq) + Cl- (aq)
The net ionic equation
H3O+ (aq) + OH- (aq) --> 2H2O.

12.2.6.0 Redox reactions (oxidation-reduction reactions, electron transfer reactions)
Redox reactions involve a transfer of electrons and a change in oxidation number.
Electrons move from one atom to another.
Oxidation is loss of electrons.
Reduction is gain of electrons.
Oxidation reactions and reductions reactions must occur together.
The same number of electrons are gained in the reduction as are lost in the oxidation.

12.2.6.1 Magnesium with dilute hydrochloric acid
In the reaction of dilute hydrochloric acid on magnesium ribbon, each magnesium atom loses two electrons to two hydrogen atoms.
Mg (s) + HCl (aq) --> MgCl2 (aq) + H2 (g)
Mg (s) + 2H3O+ (aq) + 2Cl- (aq) --> H2 (g) + Mg2+ (aq) + 2Cl- (aq) + 2H2O.

12.2.6.2 Oxygen with sulfur dioxide
In reactions where no ions form, use the idea of oxidation number (oxidation state) to show the "apparent charge" on an atom.
In the reaction between gases:
2SO2 (g) + O2 (g) --> 2SO3 (g)
Give the oxygen atom a net charge of -2, but give O2 a net charge of zero because the oxygen atom is in the elemental form.
Then the sulfur atom in SO2 has an oxidation number +4 and the sulfur atom in SO3 has an oxidation number +6.
The sulfur atoms have been oxidized because the oxidation number has increased and the oxygen gas atoms in O2 have been reduced
because the oxidation number has decreased.

12.2.6.3 Ammonia with copper oxide
Similarly, in the following equation:
NH3 + CuO --> Cu + H2O + N2
The oxidation number of hydrogen atom in NH3 is +1 and in H2 is zero, because the hydrogen atom is in the elemental form.
The oxidation number of the nitrogen atom has increased from -3 in NH3 to 0 in N2, because N in N2 is in the elemental form.
The oxidation number of the copper has decreased from +2 in CuO to zero in Cu, because the Cu atom is in elemental form.
The nitrogen atom has been oxidized and the copper atom has been reduced.

12.2.6.4 Chlorine with water
The disproportionation process occurs when a chemical species is oxidized and reduced simultaneously.
When chlorine has dissolves in water a disproportionation occurs because the chlorine becomes both oxidized, when HClO is formed,
and reduced, when HCl is formed.
Chlorine dissolving in water
Cl2 (g) + H2O (l) <--> HClO (aq) + Cl- (aq) + H+ (aq)
So the chlorine is both oxidized and reduced.

12.2.6.5 Copper (I) oxide with hot dilute sulfuric acid
Copper (I) ions in solution form a precipitate of copper and copper (II) ions, because disproportionation occurs.
So for 2Cu+ (aq), one ion is reduced to copper,
Cu+ (aq) --> Cu (s)
the other ion is oxidized.
Cu+ (aq) --> Cu2+ (aq)
In the reaction of copper (I) oxide with hot dilute sulfuric acid, a solution of copper (I) sulfate and water does not form.
Instead a brown precipitate of copper and a blue solution of copper (II) sulfate forms because of disproportionation.
2Cu+ (aq) --> Cu2+ (aq) + Cu (s).

12.2.7.0 Polymerization reactions
See 16.2.4.2.1: Cyanamides, inorganic, (CN22-), ionization reaction of methylamine, cyanic acid, melamine
Polymerization reactions produce large molecules, polymers with repeating units, monomers.
The physical properties of addition and condensation polymers are related to their structure described by the terms thermoset,
thermoplastic, elastomer, vulcanization, amorphous, crystalline.
Polymer properties depend on the chain length, side branches and cross-linking.

12.3.0 Properties of acids
Acids are good electrolytes, react with active metals, turn blue litmus red, and have a sour taste.
Dilute acids contain hydrogen ions in aqueous solution.
You can represent the hydrogen ion, which is really a proton, in different ways to show how it is related to the water molecules in the
solution.
You can show it as the hydrated hydrogen ion, [proton, H+ (aq)] or as the hydronium ion [oxonium ion, H3O+ (aq)] but, for
convenience, use H+ (aq).
Concentrated sulfuric acid exists mainly as H2SO4 molecules.
Hydrochloric acid and nitric acid dissociate into ions even in concentrated solution.
Weak acids, e.g. ethanoic acid (acetic acid, CH3COOH) carbonic acid and sulfurous acid dissociate very little in aqueous solution,
but their salts, e.g. potassium acetate (CH3COOK) are completely dissociated into ions.
Using the Bronsted-Lowry definition of acids and bases an acid donates a proton (H+) to another substance and a base accepts a
proton from another substance.
When sulfuric acid dissociates in water it donates a proton (H+) to the water molecule.
So in this reaction the water molecule acts as a base.
H2SO4 + H2O --> HSO4- + H3O+
When ammonia dissolves in water, ammonia accepts a proton and so it is the base.
So in this reaction the water molecule acts as an acid.
NH3 + H2O <--> NH4+ + OH-.

12.3.0.1 Amphoteric substances
Amphoteric substances can act as an acid or a base.
In the above reactions water is acting as a base with sulfuric acid and is acting as an acid with ammonia.
Similarly, bicarbonate ion can act as an acid to donate a proton to form carbonate ion:
HCO3- + H2O <--> CO32- + H3O+
Also, bicarbonate ion can act as a base to accept a proton to form carbonic acid:
HCO3- + H2O <--> H2CO3 + OH-.

12.3.0.2 Polyprotic acids
Polyprotic acids can donate more than one proton, e.g. carbonic acid.
H2CO3 + H2O <--> HCO3- + H3O+ (The first proton to be donated to a water molecule.)
HCO3- + H2O <--> CO32- + H3O+ (The second proton to be donated to a water molecule.)

12.3.0.3 Strong acids and weak acids, Ka, pKa
A strong acid completely dissociates into ions, e.g. nitric acid has almost complete dissociation, 93%
HNO3 (aq) + H2O --> H3O+ (aq) + NO3- (aq)
A weak acid only partly dissociates into ions, e.g. acetic acid.
CH3COOH + H2O <--> CH3COO- + H3O+
So describing acids and bases as strong or weak only refers to their reaction with water and has nothing to do with concentration or
the number of moles in a volume.
The strong acids are perchloric acid (HClO4), hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (hydriodic acid),
(HI), nitric acid (HNO3), and sulfuric acid (H2SO4).
Any other acid is a weak acid because it does not completely dissociate in water.

12.3.0.3a Acid dissociation constant, Ka
The acid dissociation constant, Ka of the acid HB:
HB (aq) <--> H+ (aq) + B- (aq)
Ka = [H+][B-] / [HB]
Ka is a measure of the degree to which an acid or base will dissociate in water.
Stronger acids have a larger Ka and a smaller pKa than weaker acids.
The greater the value of Ka, the more the formation of H+ is favoured, and the lower the pH of the solution.

12.3.0.3b Acid dissociation constant at logarithmic scale, pKa
(pK is the negative logarithmic scale of any rate constant)
pKa = -log10Ka
Strong acids have pKa value < 2
When the pH of solution is at the value of pKa for a dissolved acid, that acid will be 50% dissociated.
Sulfuric acid, H2SO4 --> HSO4-, pKa -10
Hydroiodic acid, HI,
HI (g) + H2O (l) --> H3O+ (aq) + I- (aq), pKa -9
Hydrobromic acid, HBr
HBr (g) + H2O (l) --> H3O+ (aq) + Br- (aq), pKa -8
Perchloric acid, HClO4
HClO4 + H2O --> H3O+ + ClO4- , pKa -10
Hydrochloric acid, HCl,
HCl (g) + H2O (l) --> H3O+ (aq) + Cl- (aq), pKa -7
Hydronium ion, H3O+
H2O + H2O <--> H3O+ + OH-, pKa -1.74
Nitric acid, HNO3
HNO3 + H2O --> H3O+ + NO3-, pKa - 1.3
Chloric acid, HClO3, pKa -1.0
Weak acid has pKa value 2 to 12 in water
Acetic acid, CH3COOH, pKa 4.75.

12.3.0.4 pH
See: pH (Commercial)
Water can transfer a proton from one molecule to another, autionization.
2H2O <--> H3O+ + OH-
and
H2O <--> H+ + OH-
The product of hydrogen ion concentration, [H+] and hydroxide ion concentration, [OH-] = the constant, Kw
Kw = [H+] × [OH-] = 1.00 × 10-14
So [H+] = 10-7 and [OH-] = 10-7
The hydrogen ion concentration is very small in pure water so the concentration is describes in terms of its negative log.
pH is the negative log of the hydrogen ion concentration, pH = -log[H+], so hydrogen ion concentration, [H+] = 10-pH.
So acidic solutions have a high [H+] and low pH values .
Basic solutions have low [H+] and high pH values.
A solution that is neither acidic nor basic, a neutral solution, has [H+] = [OH-], so pH = 7.
A more acid solution has pH approaching 1.
A more basic solution has pH approaching 14.

12.3.0.6 Ethanoic acid, CH3COOH, acetic acid
Ethanoic acid, vinegar
Bone, calcium hydroxyapatite, Ca10(PO4)6(OH)2, + other ions
A weak acid can dissolve the calcium in a bone or egg shell
1. The rubber chicken bone experiment.
Select a chicken bone, e.g. a "drumstick (tibia and fibula) or a "wishbone", furcula, the forked bone.
Clean any tissue off the bone , wash it with warm salt water and dry it.
Try to bend the bone between your fingers.
Cover a chicken bone with vinegar or dilute acid hydrochloric acid.
Change the solutions each day for 2 - 5 days.
The vinegar reacts with the calcium in the bone to form soluble calcium acetate.
Dry the bone with absorbent paper.
Try to bend the bone between your fingers.
The bone can be bent and even tied into a knot because all the calcium has been removed.
CaCO3 + 2CH3COOH --> Ca(CH3COO)2 + H2O + CO2
2. The bouncy egg experiment.
Cover a fresh egg with vinegar or dilute acid hydrochloric acid.
Change the solutions each day for 7 days.
Pick up the decalcified egg and drop it to show that it will bounce and not break.

12.3.1 Taste of acids, solid acids in the home
BE CAREFUL! NEVER TASTE ACIDS IN THE LABORATORY!
Citric acid, C6H8O7
Acetic acid, CH3COOH, ethanoic acid, vinegar
Do NOT taste these acids in the laboratory.
Each acid has a sour taste that is a characteristic of acids.
Lemon juice contains the white crystalline citric acid.
Vinegar contains ethanoic acid (acetic acid, CH3COOH).
Moisten your finger with a very dilute solution of hydrochloric acid.
Rub your fingers together and then lick them.
Repeat the procedure with very dilute solutions of acetic acid and citric acid.
Do not taste any other acids because they may damage living tissues.

12.3.2 Dilute acids with metals, hydrochloric acid
Reactions of acids with metals are exothermic.
The higher the metal is in the reactivity series the greater the heat liberated.
Dilute hydrochloric acid with zinc:
Zn (s) + 2HCl (l) --> H2 (g) + ZnCl2 (aq)
The order of activity of metals with acids is similar to the order of activity with water.
Evolution of hydrogen occurs
Table 12.3.2
Metal 2M Hydrochloric acid 2M Sulfuric acid
Magnesium very rapid rapid
Aluminium slight none
Zinc moderate Slight
Iron very slight very slight
Tin none none
Lead none none
Copper none none

1. Use different cleaned metals, e.g. calcium pieces, iron nail, lead sinker, magnesium ribbon, copper wire, aluminium sheet and zinc
granules.
Rub them with emery paper to make surfaces clean of oxides.
Put each metal into a separate test-tube.
Add 10 mL of 2 M hydrochloric acid to test-tubes.
Observe the properties of any gas liberated and name it.
Test it with moist pieces of red and of blue litmus paper, with a drop of lime water hanging from a glass rod and with a lighted splint.
Compare the rate at which hydrogen gas evolves by noting the rate and size of the hydrogen gas bubbles from the reaction.
Describe the rate of reaction as nil, very slow, slow, moderately fast, very fast, and whether energy, in the form of heat, is produced
(exothermic) or absorbed (endothermic).
List the acids in order of their activity towards metals and state whether the same gas was liberated during each reaction and whether
a salt may be isolated when the acids react with a metal.

2. Make up a reactivity series by listing the elements in approximate order of their activity with respect to acids, from the most active
to the least active.
Compare the results with the table of the reactivity series of some metals.
The order of activity of the metals used, from the most active to the least active, is: magnesium, aluminium, zinc, iron with lead and
copper displaying no noticeable reaction.
When reaction did occur, the gas liberated was hydrogen gas.
The reactions of these acids with metals are exothermic.
The order of activity of the acids is that dilute hydrochloric and dilute sulfuric acids are about equal in activity but that they are more
reactive than acetic acid.
The order of activity of the metals with respect to acids is similar to that with respect to water.
Magnesium ribbon forms most rapid bubbles of hydrogen gas then zinc then iron.
Tin forms few bubbles of hydrogen gas.
Copper forms no bubbles of hydrogen gas.
Lead forms some lead chloride precipitate on the surface of the lead.
Aluminium develops a layer of aluminium oxide that obstructs further chemical reactions.

3. Note the properties of any gas that forms.
Test the gas with moist litmus paper a lighted splint and a hanging drop of lime water on a glass rod.
4. Feel the test-tube to note whether heat energy is released or absorbed.
The reactions of these acids with metals are exothermic.
5. List the elements in approximate order of their activity with respect to hydrochloric acid from the most active to the least active.
The order of activity is: magnesium, aluminium, zinc, iron, lead (no noticeable reaction), copper (no noticeable reaction).

12.3.2.1 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
Dilute hydrochloric and dilute sulfuric acids are about equal in activity, but that they are more reactive than ethanoic acid (acetic acid).
Note the slower production of hydrogen gas with the weak acetic acid.
The reaction with sulfuric acid forms insoluble sulfates on the surface of calcium and lead that obstructs or stops reactions.
List the acids in order of their activity on metals.
2CH3COOH (aq) + Mg (s) --> Mg(CH3COO)2 (aq) + H2 (g)
ethanoic acid + magnesium --> magnesium ethanoate + hydrogen.

12.3.3 Dilute sulfuric acid with steel wool
Add dilute sulfuric acid to steel wool in a test-tube.
Test the gas that forms with a lighted taper.
BE CAREFUL! THE GAS IS HYDROGEN GAS!
Heat the mixture in a beaker of hot water until all the steel wool has dissolved.
Add more acid when necessary.
Filter the hot solution then leave it to cool.
Crystals form on cooling.
If no crystals form, add alcohol because the salt is less soluble in it.
Dry the green crystals of iron (II) sulfate-7-water between absorbent paper.
Fe (s) + H2SO4 (aq) --> H2 (g) + FeSO4 (aq).

12.3.3.1 Dilute sulfuric acid with aluminium
Heat dilute sulfuric acid with pieces of aluminium foil in a test-tube.
Some effervescence occurs but sometimes not enough to test for hydrogen gas with a lighted taper.
After heating for 5 minutes, decant the solution that contains aluminium sulfate into another test-tube and add ammonia solution.
A white jelly-like precipitate of aluminium hydroxide forms.

12.3.3.2 Magnesium with sodium hydrogen sulfate
Add 3 cm of magnesium ribbon to 3 cm of sodium hydrogen sulfate solution in a test-tube.
The metal reacts with the sulfuric acid in the solution.
Tests for hydrogen gas.
Remove any magnesium that has not reacted from the solution, pour part of the liquid into an evaporating basin and leave for magnesium
sulfate crystals to form.

12.3.3.3 Iron with sodium hydrogen sulfate
Add a finger width of iron filings to a finger width of sodium hydrogen sulfate solution in a test-tube.
Heat the mixture to speed up the reaction.
The metal reacts with the sulfuric acid in the solution.
Tests for hydrogen gas.
Leave to stand until all bubbles have ceased to appear.
Pour part of the liquid into an evaporating basin and leave for magnesium sulfate crystals to form.
Test the liquid with universal indicator paper.
(The indicator changes colour to red, orange, or yellow for acids and green, or violet for alkalis.)
Pale green is the colour for neutral substances.
Before testing, make the paper this colour by dipping it into neutral tap water for a few moments.
The Universal indicator turns yellow indicating the presence of an acid.
Filter the liquid, and pour part of the clear solution into the evaporating basin and leave for pale green crystals of iron sulfate to form
(FeSO4.7H2O, green vitriol).
A solution of a salt is not necessarily neutral because some salts, like iron sulfate, form acids when dissolved in water.

12.3.4 Dilute acids with non-metals, carbon, sulfur
Add a piece carbon and sulfur to dilute hydrochloric acid, dilute sulfuric acid and dilute ethanoic acid (acetic acid) in separate test-tubes.
Heat the test-tubes.
No reaction occurs.
Non-metals do not react with dilute acids.

12.3.5 Dilute acids with basic oxides
Basic oxides are mostly metal oxides, e.g. copper oxide.
acid + basic oxide ---> salt + water
1. Heated dilute acids react with metal oxides to form a salt and water:
Pour dilute sulfuric acid into a Pyrex test-tube and heat in a beaker of boiling water until the sulfuric acid is nearly boiling.
BE CAREFUL!
Add pieces of copper (II) oxide one by one while stirring until some remains unreacted with the acid.
Filter the undissolved copper oxide from the hot solution.
Leave the filtrate in a watch glass to cool and form crystals.
Blue crystals of copper (II) sulfate-5-water form with water.
Remove the crystals and dry them by pressing between absorbent paper.
H2SO4 (aq) + CuO (s) --> CuSO4 (aq) + H2O (l)
acid + basic oxide ---> salt + water
2. Repeat the experiment with dilute nitric acid.
2HNO3 (aq) + CuO (s) --> Cu(NO3)2 (aq) + H2O (l).

12.3.5.1 Dilute acids with amphoteric oxides
(Greek, amphoteros: compare, both)
Oxides of Sn, Al, Zn, Pb, and Sb are amphoteric.
They have both acidic and basic properties,
Acid + amphoteric oxide --> salt + water
Amphoteric oxides react with bases to form a salt + water.
Amphoteric oxides react with acids to form a salt + water.
Add dilute hydrochloric acid to zinc oxide.
2HCl (aq) + ZnO (s) --> ZnCl2 (aq) + H2O (l)
2NaOH (aq) + ZnO (s) --> Na2ZnO2 (aq) + H2O (l).

12.3.5.01 Copper oxide with sodium hydrogen sulfate
Add half a test-tube of sodium hydrogen sulfate solution to copper oxide in a test-tube.
Heat the solution slowly until it turns blue.
Be careful of spurting from the test-tube.
Some copper oxide may remain after the reaction.
Filter the solution obtain the filtrate of copper sulfate solution.

12.3.6 Dilute acids with hydroxides, magnesium hydroxide
Basic hydroxides are insoluble in water and react with acids to form a salt and water.
Many metallic hydroxides react with acids to form a salt and water.
Add magnesium hydroxide to dilute sulfuric acid until the reaction stops.
Filter the mixture.
Test the filtrate with litmus paper.
Evaporate the filtrate to dryness so that crystals form.
Mg(OH)2 (s) + H2SO4 (aq) --> MgSO4 (aq) + H2O (l).

12.3.7 Dilute acids with hydroxides, sodium hydroxide
Acids react with (neutralize) alkalis to form a salt and water.
Pour 5 mL of dilute sodium hydroxide solution into a watch glass.
Test with litmus paper.
Red litmus turns blue.
Add dilute hydrochloric acid drop by drop.
Stir as each drop is added.
Test the mixture with the litmus paper until the litmus paper is neither red nor blue, but between these colours.
Evaporate the solution to dryness by heating the watch glass over a beaker of boiling water.
Crystals of sodium chloride (common salt) form.

12.3.7.1 Dilute acids with sodium hydroxide
Repeat the previous experiment with: dilute sulfuric acid, dilute nitric acid, ethanoic acid (acetic acid).
HCl (aq) + NaOH (aq) --> NaCl (aq) + H2O (l)
hydrochloric acid + sodium hydroxide --> sodium chloride + water.

12.3.7.2 Dilute hydrochloric acid with hydroxides
[NH3 (aq) is used because while "NH4+" ions and "OH-" ions can be detected, "NH4OH" cannot be detected, so ammonia solution is
shown as "NH3 (aq) + H2O (l)"]
Repeat the experiment with dilute solutions of: potassium hydroxide, calcium hydroxide, aqueous ammonia solution.
acid + (base) alkali --> salt + water
HCl (aq) + NaOH (aq) --> NaCl (aq) + H2O (l)
HNO3 (aq) + NaOH (aq) --> NaNO3 (aq) + H2O (l)
HCl (aq) + KOH (aq) --> KCl (aq) + H2O (l)
HCl (aq) + NH3 (aq) + H2O (l) --> NH4Cl (aq) + H2O (l).

12.3.8 Dilute acids with acidic oxides
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Note any reaction for five minutes then evaporate to dryness.
In each case, no reaction occurs.
In each experiment there is no precipitate.
If you evaporate a sample of a remaining solution to dryness in a fume cupboard, no residue remains.
Pass carbon dioxide through hydrochloric acid or ethanoic acid (acetic acid) solution.
Pass sulfur dioxide through hydrochloric acid or ethanoic acid (acetic acid) solution.

12.3.9.0 Dilute acids with carbonates, common carbonates
Dilute acids react with metal carbonates to form a salt, carbon dioxide and water.
Geologists use this reaction to identify calcium carbonate in rock.
Drops of hydrochloric acid cause bubbles to form.
1. Make a chemical egg peeler.
Put an egg in vinegar (contains acetic acid, ethanoic acid).
Note the bubbles forming on the outside of the egg.
Leave overnight then, the next day, pick up the egg with your fingers. The egg has become soft.
Leave to stand for a few days and the egg shell disappears completely.
You can now see through the raw egg.
acetic acid in the vinegar + calcium carbonate in the egg shell --> calcium acetate in solution + bubbles of carbon dioxide + water.

2. Add 5 mL vinegar or dilute hydrochloric acid or dilute sulfuric or dilute nitric acid to pea size amounts of finely divided common
carbonates: sodium hydrogen carbonate, sodium carbonate, calcium carbonate, magnesium carbonate, nickel carbonate, limestone,
lime, oyster shells, egg shell, snail shell, coral.
Continue to add the solid until no further reaction occurs.
Filter and evaporate the filtrate to dryness.
Note any visible changes.
Test any gas liberated by inserting in the mouth of the tube first damp pieces of red and of blue litmus paper then a drop of lime water
hanging on the tip of a glass rod and finally a burning splinter.
In each case the gas is carbon dioxide.

12.3.9.1 Dilute hydrochloric acid with calcium carbonate
See diagram 9.154: Lime water test for carbon dioxide in the breath
1. Put calcium carbonate in a test-tube.
Add 2 mL 1.0 M hydrochloric acid.
Tilt the test-tube so that its mouth is touching a second test-tube containing 5 mL of lime water.
The surface of the lime water turns milky.
Shake the test-tube containing the lime water.
The milky colour on the surface disappears.
CaCO3 (s) + 2HCl (aq) --> CO2 (g) + CaCl2 (aq) + H2O (l)
carbonate + acid --> carbon dioxide + salt + water.

2. Put 5 g of marble chips (calcium carbonate) and the same quantity of dilute hydrochloric acid in a test-tube fitted with a one-hole
stopper and delivery tube.
With the end of the delivery tube dipping into a second test-tube of lime water add water to the first test-tube and quickly replace the
stopper.
The lime water turns milky.
The acid reacts with calcium carbonate to form a salt, carbon dioxide, and water.
hydrochloric acid + calcium carbonate --> calcium chloride + carbon dioxide + water.

12.3.9.2 Dilute hydrochloric acid with sodium carbonate
1. Put sodium carbonate in a test-tube and add drops of dilute hydrochloric acid.
Test any gases formed from the reaction with moist litmus paper, a lighted splint, and a drop of lime water on a glass rod.
The reaction forms carbon dioxide.
Add more carbonate until no more reaction occurs.
Filter and evaporate the filtrate to dryness.
Repeat the experiment with dilute nitric acid.
Repeat the experiment with magnesium carbonate.
Na2CO3 (s) + 2HCl (aq) --> 2NaCl (aq) + H2O (l) + CO2 (g)
Na2CO3 (s) + 2HNO3 (aq) --> 2NaNO3 (aq) + H2O (l) + CO2 (g).

2. Sodium carbonate with hydrochloric acid
Stage 1. Na2CO3 + HCl ---> NaHCO3 + NaCl
Stage 2. NaHCO3 + HCl ---> NaCl + H2O + CO2
Overall equation: Na2CO3 + 2HCl ---> 2NaCl + H2O + CO2
Net ionic equation: CO32- + 2H+ --> H2O + CO2.

3. Shake different solid acids in separate test-tubes half filled with water.
Divide the solutions in the test-tubes into three different test-tubes:
Test-tube A: Add small pieces of red and of blue litmus paper.
Test-tube B: Add three drops of methyl orange solution.
Test-tube C: Add three drops of phenolphthalein solution.
Observe any changes in the solutions.
Add solid sodium carbonate to each acid solution.
Observe any changes in the solutions.
Pass some gas given off into a test-tube containing lime water.
Shake the test-tube for thorough mixing.
Note how milky the solution is because carbon dioxide was produced when the acids reacted with sodium carbonate.

12.3.9.3 Dilute tartaric acid with sodium carbonate
Put 5 g of sodium carbonate and the same quantity of tartaric acid in a test-tube fitted with a one-hole stopper and delivery tube.
With the end of the delivery tube dipping into a second test-tube of lime water add water to the first test-tube and quickly replace the
stopper.
The lime water turns milky.
The acid reacts with sodium carbonate to form a salt, carbon dioxide, and water.
tartaric acid + sodium carbonate --> sodium tartrate + carbon dioxide + water.

12.3.9.4 Dilute tartaric acid with egg shell, soil, wood ash
Many common substances, such as mortar, egg shell, most soils, contain calcium carbonate and wood ashes contain potassium
carbonate.
Observe the action of tartaric acid on these substances in a test-tube.
Tests for carbon dioxide by holding a drop of lime water, at the end of a glass tube, in the mouth of the test-tube.

12.3.9.6 Dilute sulfuric acid with calcium carbonate
Put 5 g of marble chips (calcium carbonate) and the same quantity of dilute sulfuric acid in a test-tube fitted with a one-hole stopper
and delivery tube.
With the end of the delivery tube dipping into a second test-tube of lime water add water to the first test-tube and quickly replace the
stopper.
The lime water turns milky.
The acid reacts with calcium carbonate to form a salt, carbon dioxide, and water.
The reaction of sulfuric acid with calcium carbonate proceeds only for a few moments because the salt formed, calcium sulfate, is only
slightly soluble and deposits on the carbonate, preventing this compound from reacting with the acid.
So the reaction with hydrochloric acid above is much better.
sulfuric acid + calcium carbonate --> calcium sulfate + carbon dioxide + water.

12.3.10 Dilute acids with sodium hydrogen carbonate
Alka-Seltzer
See 3.34.6: Soda-acid fire extinguisher
Sold as: Mighty Seltzer Rocket, uses Alka Seltzer tablets, Be careful!
The only stable hydrogen carbonates are KHCO3 and NaHCO3.
Sodium hydrogen carbonate, bicarbonate of soda, is used in baking soda, baking powder, self raising flour, effervescent fruit salts,
e.g. Alka-Seltzer, and soda acid fire extinguishers and treatment for acid burns.
Some people swallow sodium hydrogen carbonate to counteract excess acid in the stomach but using magnesium oxide or magnesium
hydroxide that does not react with acids to produce carbon dioxide is better.
1. Add sodium hydrogen carbonate, or other hydrogen carbonates, to acids to form carbon dioxide, water and a salt.
NaHCO3 + HCl --> CO2 + H2O + NaCl
hydrogen carbonate + acid --> carbon dioxide + water + salt.

2. Mix vinegar with bicarbonate of soda in a glass jar.
Drop some naphthalene mothballs into the solution.
The carbon dioxide formed by the reaction of the vinegar (acetic acid) with the sodium hydrogen carbonate forms bubbles of carbon
dioxide on the mothballs in the bottom of the jars.
The mothballs rise to the surface, lose the bubbles and sink again.
NaHCO3 + CH3COOH --> CH3COONa + H2O + CO2 (g).

12.3.10.1 Dilute acids with calcium hydrogen carbonate
Put powdered calcium carbonate into a test-tube containing about 10 mL of water.
Bubble carbon dioxide through the suspension until no further change takes place.
Soluble calcium hydrogen carbonate forms.
Boil the mixture for 10 minutes.
Add acids to form carbon dioxide, water and a salt.

12.3.11.0 Dilute nitric acid with copper
Very dilute nitric acid may react with very active metals, e.g. magnesium to form hydrogen gas.
When nitric acid reacts with most metals, it oxidizes the hydrogen to water.
Add drops of dilute nitric acid to copper.
Nitrogen monoxide forms, which immediately reacts with oxygen gas in the air to form nitrogen dioxide.
3Cu (s) + 8HNO3 (aq) --> 3Cu(NO3)2 (aq) + 4H2O (l) + 2NO (g)
2NO (g) + O2 (g) --> 2NO2 (g).

12.3.11.1 Nitric acid with metals
Add slowly small pieces of copper, magnesium and zinc to small amounts of dilute nitric acid in separate test-tubes.
If no change is taking place, gently heat the mixture.
Repeat the procedure 1. with concentrated nitric acid, 2. with concentrated sulfuric acid, and 3. with concentrated hydrochloric acid.
The reactions of metals with nitric acid and concentrated sulfuric acid are different from reactions of metals with hydrochloric acid,
dilute sulfuric acid and dilute acetic acid.
Although copper does not react with dilute acids or with concentrated hydrochloric acid, it does react with both dilute and concentrated
nitric acids and with hot concentrated sulfuric acid but does not produce hydrogen gas in reaction with them.
The residual mixtures contain solutions of salts but writing equations for the reactions is difficult because more than one reaction can
occur simultaneously between copper or magnesium or zinc and nitric acid.
For example, when zinc reacts with nitric acid the reaction may produce nitrogen dioxide, nitric oxide, nitrous oxide, zinc nitrate and
ammonium nitrate.

12.3.12 Concentrated acids with metals, nitric acid with copper
Nitric acid reacts with metals above platinum in the reactivity series, but does not form hydrogen gas.
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Pour drops of concentrated nitric acid on pieces of copper in a test-tube. Put a stopper on the test-tube immediately because brown
nitrogen dioxide gas forms.
The nitric acid acts as an oxidizing agent and is reduced to nitrogen dioxide and water.
The reaction is exothermic.
Cu (s) + 4HNO3 (aq) --> Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g).

12.3.13 Concentrated acids with metals, sulfuric acid with copper
Concentrated acids should be handled only by experienced science teachers.
Concentrated sulfuric acid reacts with metals above platinum in the reactivity series, but does not form hydrogen gas.
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Add hot concentrated sulfuric acid to a piece of copper foil.
Brown nitrogen dioxide gas forms.
The sulfuric acid acts as an oxidizing agent.
Cu (s) + 2H2SO4 (aq) --> CuSO4 (aq) + 2H2O (l) + SO2 (g).

12.3.14 Concentrated acids with a non-metal, carbon
DO NOT DEMONSTRATE THIS EXPERIMENT!
Hot sulfuric acid and nitric acid can react as oxidizing agents with carbon.
Carbon is oxidized to carbon dioxide and nitric acid is reduced to nitrogen dioxide and water.
C (s) + 4HNO3 (aq) --> CO2 (g) + 4NO2 (g) + 2H2O (l).

12.3.15 Acids with salts
1. Add small quantities of sodium chloride, sodium nitrate, sodium acetate, sodium sulfite and iron sulfide to about 5 mL of dilute
hydrochloric acid in separate test-tubes.
Observe what happens when the mixtures are cold and when they are warmed.
2. Repeat the procedure using dilute sulfuric acid and then concentrated sulfuric acid.
3. Dilute acids do not react with chlorides, nitrates, sulfates, or acetates unless the metal ions in the salt can form an insoluble salt with
the ions in the acid.
4. Acids react with sulfites to produce sulfur dioxide, water and a salt.
5. Acids react with sulfides to produce hydrogen sulfide (rotten egg gas) and a salt.
6. Concentrated sulfuric acid reacts with chlorides to produce hydrogen chloride and a sulfate.
7. Concentrated sulfuric acid reacts with nitrates to produce nitric acid and a sulfate.
8. Concentrated sulfuric acid reacts with acetates to produce acetic acid and a sulfate.

12.3.16 Acid dissociation constant, Ka
Acid dissociation constant, Ka, acidity constant, acid-ionization constant, dissociation constant
1. The acid dissociation constant, Ka, measures the strength of an acid in solution.
2. An acid, HA, dissociates into A-, conjugate base, and H+, hydrogen ion (proton).
The equilibrium equation when concentrations do not change is: HA <--> A- + H+.
3. Dissociation refers to the break up of a molecule into smaller molecules, atoms or ions.
In a buffer solution of the salt of a weak acid with a weak acid, the dissociation of the weak acid is negligible, but a salt may be
dissociated completely into ions.
4. The dissociation constant, Ka (acid dissociation constant, acidity constant, acid-ionization constant), is the equilibrium constant of
reversible dissociation including the ionization reactions of acids and bases in water.
The dissociation constant Ka = [A-] [ H+] / [HA] in mol / litre.
5. However, dissociation is usually expressed as a logarithmic constant, pKa, where pKa = -log10 (1/Ka)
It is the quotient of the equilibrium concentrations, in mol/L for ionization reactions at 25oC.
For pKa, the larger the value the weaker the acid, so strong acids have pKa < 2, and weak acids have pKa >2, < 12.
6. Confusion occurs because both Ka and pKa are both called "acid dissociation constant".

12.4.0 Hydrochloric acid
Hydrochloric acid is an aqueous solution of hydrogen chloride gas.
Hydrochloric acid dissolves most metals to form chlorides and hydrogen gas.
Hydrochloric acid is available as:
A. 5.0 M, 4.0 M, 2.0 M, 1.0 M and 0.5 M volumetric solutions
B. Minimum assay 36% solution density 1.17 g cm-3 at 20oC C. 36% "ANALAR" solution
D. Sold as: Muriatic acid, for use in the building trades.

12.9.0 Phosphoric acid
Ionization reaction
H3PO4 + H2O <--> H3O+ + H2PO4-
H2PO4- + H2O <--> H3O+ + HPO42-
HPO42- + H2O <--> H3O+ + PO43-.

12.10.0 Boric acid, ionization reaction
1. Orthoboric acid, trioxoboric acid (III) acid, boracic acid, sassolite, H3BO3 is a weak acid.
White to colourless triclinic crystals, m.p. 169oC, occurs in volcanic steam vents, slightly soluble in cold water, used to make
borosilicate glass, used in buffer solutions, detergents and in pharmacy, e.g. "boracic powder" for eye infections.
Action of continuous heat: boric acid, H3BO3 --> metaboric acid + water, H2B4O4 --> tetraboric acid
(pyroboric acid) H2B4O7 --> boric oxide (anhydrous boron (III) oxide) B2O3.
Boric oxide is an intermediate oxide, as is aluminium oxide, with weak acidic and basic properties.
Borax is hydrated sodium borate.
When heated it fuses to form clear glass that can dissolve metal oxides to give characteristic colours of the borax bead test.
Ionization reaction, Ka = 6.0 × 10-10
H3BO3 + H2O <--> H3O+ + H2BO3-
H3BO3 <--> H+ + H2BO3-
H2BO3- <--> H+ + HBO32-
HBO32- <--> H+ + BO32-.

12.10.1 Prepare boric acid crystals
Use 5 g of boric acid crystals.
Pour some into 2 cm boiling water in a test-tube and leave to dissolve.
Continue adding crystals and heat to boiling until all crystals dissolved.
Leave to cool to see fine white crystals form.

12.10.2.1 Dilute acids with metals
The reactions with K and Na are too vigorous.
No reaction for metals below hydrogen in the activity series.

Sigma-Aldrich, Nitric Acid Grades
See: Sigma-Aldrich (Commercial)
Nitric acid puriss.
p.a., ACS reagent, fuming, 99.5%
Nitric acid reagent grade, fuming, >90%
Nitric acid red, fuming, HNO3 >90 %
Nitric acid reagent grade, >90%
Nitric acid ACS reagent, 90.0%
Nitric acid, ACS reagent, 70%
Nitric acid 70%, AR grade
Nitric acid 70%, purified by redistillation, 99.999% trace metals basis
Nitric acid 65%, TraceSELECT Ultra, for ultratrace analysis
Nitric acid 69.0%, TraceSELECT, for trace analysis
Nitric acid puriss. p.a., ACS reagent, ISO, 69%
Nitric acid puriss., 64-66%
Nitric acid purum p.a., 65% (T)
Nitric acid concentrate for 1L standard solution, 1.0 M HNO3 (1.0N).