Acids, chemical reactions, chemical equations.

School Science Lessons
2024-06-13

Acids, Chemical reactions, Reactions of dilute acids, Prepare salts
(topic12)
Contents
12.3.0 Acids
12.2.0 Chemical reactions
12.4.0 Dilute acids with
12.6.0 Prepare salts

12.3.0 Acids
12.3.1 Properties of acids
12.3.2 Acid dissociation constant, Ka
12.3.3 Amphoteric substances
12.3.4 Boric acid, ionization reaction
12.3.5 Concentrated acids with a non-metal, carbon
12.3.6 pH
12.3.7 Phosphoric acid
12.3.8 Polyprotic acids
Experiments
12.3.9 Acids with salts
12.3.10 Concentrated acids with metals, sulfuric acid with copper
12.3.11 Copper oxide with sodium hydrogen sulfate
12.3.13 Iron with sodium hydrogen sulfate
12.3.14 Magnesium with sodium hydrogen sulfate
12.3.16 Strong acids and weak acids, Ka, pKa
12.3.17 Taste of acids, solid acids in the home

12.2.0 Chemical reactions
12.2.01 Prepare salts by chemical reactions
12.2.02 Prepare boric acid crystals
12.2.03 Prepare soda lime
12.2.04 Types of chemical reactions
12.2.1 Acid-base reactions
12.2.2 Decomposition reactions
12.2.3 Displacement reactions
12.2.4 Double replacement reactions, metathesis
12.2.5 Oxidation reactions
12.2.6 Polymerization reactions
12.2.8 Reduction reactions
12.2.9 Redox reactions
12.2.10 Synthesis reactions
12.2.11 Heat of neutralization reactions, (Experiments)

12.4.0 Dilute acids with
Experiments
12.4.1 Dilute acids with acidic oxides
12.4.2 Dilute acids with amphoteric oxides
12.4.3 Dilute acids with basic oxides (metal oxides), copper (II) oxide
12.4.4 Dilute acids with calcium hydrogen carbonate
12.4.5 Dilute acids with carbonates, common carbonates
12.4.6 Dilute acids with hydroxides, magnesium hydroxide
12.4.7 Dilute acids with hydroxides, sodium hydroxide
12.4.8 Dilute acids with metals, hydrochloric acid
12.4.9 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
12.4.10 Dancing mothballs
12.4.11 Dilute sulfuric acid with steel wool
12.4.12 Dilute acids with non-metals, carbon, sulfur
12.4.13 Dilute acids with sodium hydrogencarbonate
12.4.14 Dilute acids with sodium hydroxide
12.4.15 Dilute hydrochloric acid with calcium carbonate
12.4.16 Dilute hydrochloric acid with hydroxides
12.4.17 Dilute hydrochloric acid with sodium carbonate
12.4.18 Dilute sulfuric acid as an acid
12.4.19 Dilute sulfuric acid as a sulfate
12.4.20 Dilute sulfuric acid with aluminium
12.4.21 Dilute sulfuric acid with calcium carbonate
12.4.22 Dilute tartaric acid with egg shell, soil, wood ash
12.4.23 Dilute tartaric acid with sodium carbonate
12.4.24 Sodium chloride with sulfuric acid

12.6.0 Prepare salts
A salt compound is formed when the hydrogen ions in an acid are replaced by metal ions or ammonium ions.
A salt compound is formed by the combination of an acid radical or positive ions with a basic radical or negative ion.
A salt compound is formed by the replacement of all or part of the hydrogen in an acid by a metallic element.
12.6.1 Prepare salts by neutralization of a soluble acid and a base
12.6.2 Prepare salts by acids + metals
12.6.6 Prepare salts by precipitation reactions
12.6.7 Prepare salts by direct union of elements
12.6.8 Prepare salts by bases + non-metallic oxides
12.6.9 Prepare salts by acids + salts

12.6.1 Prepare salts by neutralization of a soluble acid and a base
Acid-base neutralization: Acid + Base (Alkali) --> Salt + Water
H⁺ (aq) + OH^- (aq) --> H₂O.

12.6.2 Prepare salts by acids + metals
12.4.9 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
Acid + metal --> salt + hydrogen gas
(Metals that displace hydrogen from an acid may be called "active metals", e.g. Zn, Fe)

12.6.6 Prepare salts by precipitation reactions
Precipitation reactions, double decomposition reactions, double displacement reactions
Mixing two soluble compounds to prepare an insoluble salt.
Solution 1 + Solution 2 --> Insoluble solid 3 + Solution 4
12.2.4 Double replacement reactions, metathesis

12.6.7 Prepare salts by direct union of elements
Direct union of elements to form compounds: 8.0.0
12.2.10 Synthesis reactions

12.6.8 Prepare salts by bases + non-metallic oxides
12.17.3.1 Carbon dioxide with sodium hydroxide solution
CO2 (g) + 2NaOH (aq) --> Na2CO3 (aq)+ H2O (l)

12.6.9 Prepare salts by acids + salts
sulfuric acid + sodium chloride --> sodium sulfate
H2SO4 (aq) + NaCl (s) --> NaHSO4 (s) + HCl (g)
H2SO4 (aq) + MgCl2 (s) --> Mg(HSO4)2(s) + HCl (g)

12.2.01 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!

12.2.02 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.

12.2.03 Prepare soda lime
Dissolve calcium oxide in sodium hydroxide solution and leave to dry so that a granular 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.
It is used in recirculating breathing systems to absorb carbon dioxide.

12.2.04 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 are reactions where one species is oxidized and another species is reduced.
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 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.

12.2.2 Decomposition reactions
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.
Experiments
12.2.2.1 Decomposition of copper carbonate, prepare copper oxide
12.2.2.2 Decomposition of zinc carbonate, prepare zinc oxide
12.2.2.3 Electrolytic decomposition, electrolysis
12.2.2.4 Heat copper sulfate crystals
12.2.2.5 Photo decomposition, photolysis
12.2.2.6 Thermal decomposition of acids
12.2.2.7 Thermal decomposition of sucrose crystals, C12H22O11

12.2.3 Displacement reactions, (substitution reactions)
The reactants are an element and a compound.
In the reaction displacement of Cu²⁺ by zinc
The element replaces part of the compound with the same valence and same sign.
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.
Experiments
Displace a less reactive halogen from halogen compounds: 12.19.2.2
Displacement of copper by zinc: 12.10.3
Iron displaces copper in copper sulfate solution: 12.2.3.1
Magnesium displaces hydrogen in ethanoic acid: 12.2.3.2
Metals, Zn, Fe, displace copper: 12.10.2

12.2.4 Double replacement reactions, metathesis
12.2.4.01 Ionic equations, double decomposition reactions
12.2.4.02 Coloured precipitates, double decomposition reactions
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.
Experiments
12.2.4.1 Calcium hydroxide with cobalt chloride
12.2.4.2 Sodium carbonate with calcium hydroxide
12.2.7.2H Sodium hydroxide with alum
12.2.4.4 Sodium hydroxide with cobalt chloride
12.2.4.5 Sodium hydroxide with copper sulfate
12.2.4.6 Sodium hydroxide with iron (II) sulfate
12.2.4.7 Sodium carbonate with magnesium sulfate
12.2.4.8 Sodium carbonate with zinc sulfate

12.2.4.01 Ionic equationss, 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) + CO3^2- (aq) --> CaCO3 (s).

12.2.4.02 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) + CrO4^2-(aq) --> PbCrO4 (s)
2. silver nitrate (aq) + potassium chromate (aq) --> silver chromate (s) (red precipitate)
Ag⁺ (aq) + CrO4^2- (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) + CO3^2- (aq) --> CuCO3 (s).

12.2.4.1 Calcium hydroxide with cobalt chloride
To half a test-tube of cobalt chloride solution add limewater (calcium hydroxide solution).
Describe the reaction that occurs.
A pink / mauve precipitate of cobalt hydroxide forms.
cobalt chloride + calcium hydroxide --> cobalt hydroxide + calcium chloride.

12.2.4.2 Sodium carbonate with calcium hydroxide
To half a test-tube of sodium carbonate solution add limewater (calcium hydroxide solution).
Describe the reaction that occurs.
A white solid (precipitate) of calcium carbonate forms.
sodium carbonate + calcium hydroxide --> calcium carbonate + sodium hydroxide.

12.2.7.2H 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).

12.2.4.4 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.

12.2.4.5 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 SO4^2-, remain in solution.

12.2.4.6 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.

12.2.4.7 Sodium carbonate with magnesium sulfate
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 SO4^2-, remain in solution.

12.2.4.8 Sodium carbonate with zinc sulfate
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 SO4^2-, remain in solution.

12.2.5H Oxidation reactions
Oxidizing means adding oxygen to a substance.
The addition of oxygen to, or the loss or removal of hydrogen from, a compound.
he increase in the proportion of negative constituents in a molecule or compound.
The loss or removal of an electron from an atom or molecule.
Oxidation, in a cell, oxidation occurs at the anode and reduction occurs at the cathode
Examples include rusting, respiration, combustion, oxidation of ethanol, photosynthesis
Oxidation and reduction, redox reactions: 15.2.0
Oxidation can affect air pressure: 15.3.08
Oxidation and air pressure, steel wool over water: 4.241
Oxidation, Catalytic oxidation of ammonia forms nitrogen monoxide: 13.6.6.1
Oxidation of glucose, blue bottle experiment: 9.8.27
Oxidation of glycerine: 12.1.6
Oxidation of iron: 6.3.2 (Soils)
Oxidation of methanol to methanal: 16.3.2.8
Oxidation reactions: 12.2.5H
Redox reactions: 12.2.9
ORP (Oxidation-Reduction Potential): 18.7.50

12.2.6 Polymerization reactions
See 16.2.4.2.1: Cyanamides, inorganic, CN2^2-, 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.

12.2.8 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
Reduce potassium manganate (VII) by sulfur dioxide: 3.51.2
Reduce red iron oxide, or rust, to iron: 10.10.3.

12.2.9 Redox reactions
12.2.6.01 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
17.7.6.1 Disproportionation, Hydrogen peroxide with catalase enzyme in raw beef liver

12.2.10 Synthesis reactions
Synthesis reactions, combination reactions, direct union of elements, reactions of two 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.
Experiments
12.2.10.2 Heat copper wire with iodine crystals
12.2.10.3 Heat copper with sulfur
12.2.10.4 Heat iron with copper
12.2.10.5 Heat iron with sulfur
12.2.10.6 Heat steel wool with iodine crystals
12.2.10.7 Heat zinc with sulfur

12.2.11 Heat of neutralization reactions
1. 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 carefully stir with a thermometer.
Note the maximum temperature reached.
The increase of temperature should be 13^oC.
The volume of water has been doubled by adding one solution to the other.
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.
2. Put 25 mL of 2.0 M sodium hydroxide solution into a polystyrene cup and measure the temperature of the sodium hydroxide solution.
Stand the polystyrene cup in a beaker.
Put 25 mL of 2.0 M hydrochloric acid in a measuring cylinder and measure the temperature of the acid.
The sodium hydroxide and the hydrochloric acid should have the same room temperature.
Pour the acid into the sodium hydroxide, stir the solution with the thermometer, and note the highest temperature of the mixture.
Record the temperature difference reached in the mixture.
After each temperature measurement rinse the thermometer and dry it with absorbent paper.
Repeat the experiment with 50 mL of 2.0 M sodium hydroxide and hydrochloric acid.
Repeat the experiment with 50 mL of 1.0 M sodium hydroxide and hydrochloric acid.
Compare the temperature differences reached in the 3 experiments.

12.3.1 Properties of acids
An acid is a good electrolyte, reacts with active metals, turns blue litmus red and has a sour taste, (Latin, acidus, sour tasting).
Acids contain hydrogen replaceable by metals.
Acids neutralize alkalis and alkalis neutralize acids.
Acids are sour corrosive, mainly liquids, that can dissolve metals.
Acids with water produce hydrogen ions, H⁺.
An acid is a proton donor (H⁺) (Bronsted-Lowry definition).
Acids can donate protons or accept pairs of electrons.
The "acid test" was originally the nitric acid test for gold, because only gold would not dissolve in it.
To acidify is to add acid to usually a solution.
Acidic substance have pH < 7.
An acid salt is formed by an acid where incomplete exchange of replaceable hydrogen occurs.
An acid dye is a dye that is a metallic salt of an acid and can be applied in an acid medium.
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)]
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.
However, 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^-.

12.3.2 Acid dissociation constant, Ka
Acid dissociation constant, Ka or Acidity constant or Acid-ionization constant or " Dissociation constant"
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.
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, 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 25^oC.
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".

12.3.3 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 --> CO3^2- + H3O⁺
Also, bicarbonate ion can act as a base to accept a proton to form carbonic acid:
HCO3^- + H2O --> H2CO3 + OH^-.

12.3.4 Boric acid, ionization reaction
1. Orthoboric acid, trioxoboric acid (III) acid, boracic acid, sassolite, H3BO3 is a weak acid.
It has white to colourless triclinic crystals, MP 169^oC, and occurs in volcanic steam vents, and is slightly soluble in cold water.
It is used to make borosilicate glass, 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⁺ + HBO3^2-
HBO3^2- --> H⁺ + BO3^2-.

12.3.5 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).

12.3.6 pH
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.

12.3.7 Phosphoric acid
Ionization reaction
H3PO4 + H2O --> H3O⁺ + H2PO4^-
H2PO4^- + H2O --> H3O⁺ + HPO4^2-
HPO4^2-+ H2O --> H3O⁺ + PO4^3-.

12.3.8 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 --> CO3^2- + H3O⁺
(The second proton to be donated to a water molecule.)

12.3.9 Acids with salts
* Add small quantities of sodium chloride, sodium nitrate, sodium acetate, sodium sulfite and iron sulfide to about 5 mL of dilute hydrochloric acid in test-tubes.
Observe what happens when the mixtures are cold and when they are warmed.
* Repeat the procedure using dilute sulfuric acid and then concentrated sulfuric acid.
* 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.
* Acids react with sulfites to produce sulfur dioxide, water and a salt.
* Acids react with sulfides to produce hydrogen sulfide (rotten egg gas) and a salt.
* Concentrated sulfuric acid reacts with chlorides to produce hydrogen chloride and a sulfate.
* Concentrated sulfuric acid reacts with nitrates to produce nitric acid and a sulfate.
* Concentrated sulfuric acid reacts with acetates to produce acetic acid and a sulfate.

12.3.10 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).

12.3.11 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.
CuO + NaHSO4 --> Na2SO4 + CuSO4 + H2O

12.3.13 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.

12.3.14 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.

12.3.16 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 totheir reaction with water and has nothing to do with concentration or the number of moles in a volume.
Strong acids: | perchloric acid (HClO4) | hydrochloric acid (HCl) | hydrobromic acid (HBr) | hydroiodic acid (hydriodic acid), (HI) | nitric acid (HNO3) | sulfuric acid (H2SO4) |
For example, hydroiodic acid, (hydriodic acid), is an aqueous solution of hydrogen iodide, and is completely ionized strong acid in aqueous solution, which reacts with oxygen in air to give iodine.
4 HI + O2 --> 2 H2O + 2 I2 Any other acid is a weak acid, because it does not completely dissociate in water.

12.3.16b 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⁺ + ClO^4-, 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.

12.3.17 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.

12.4.1 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.

12.4.2 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).

12.4.3 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).

12.4.4 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.

12.4.5 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.
CuCO3(s) + H2SO4(aq) --> CuSO4(aq) + CO2(g) + H2O(l) Add 5 mL vinegar or dilute HCl or dilute H2SO4 or dilute HNO3 to pea size amounts of finely divided: sodium hydrogencarbonate, 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 limewater hanging on the tip of a glass rod and finally a burning splinter.
In each case the gas is carbon dioxide.

12.4.6 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).

12.4.7 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.

12.4.8 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.
Mg (s) + 2HCl (aq) --> H2 (g) + MgCl2 (aq)
Dilute hydrochloric acid with zinc:
Zn (s) + 2HCl (aq) --> 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 limewater 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 | heat energy is produced (exothermic) | heat energy is 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, most active to the least active, is: | 1. magnesium | 2. aluminium | 3. zinc | 4. iron | 5. lead and copper, 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 limewater 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|.

12.4.9 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.

12.4.10 Dancing mothballs
Vinegar contains up to 8% acetic acid. Mothballs consist of naphthalene, which is flammable, so newer mothballs contain other chemicals that are not flammable 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 hydrogencarbonate forms bubbles of carbon dioxide on the mothballs.
The mothballs rise to the surface, lose the bubbles and sink again, so this experiment may be called "dancing mothballs". NaHCO3 + CH3COOH --> CH3COONa + H2O + CO2 (g).

12.4.11 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).

12.4.12 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.

12.4.13 Dilute acids with sodium hydrogencarbonate
Alka-Seltzer, (Experiments)
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 hydrogencarbonate, bicarbonate of soda, is used in baking soda, baking powder, self raising flour, effervescent fruit salts, soda acid fire extinguishers.
It is also used for treatment for acid burns.
Some people swallow sodium hydrogencarbonate to counteract excess acid in the stomach.
However, it would be better to use magnesium oxide or magnesium hydroxide that does not react with acids to produce carbon dioxide, (and burping).
better.
1. Add sodium hydrogencarbonate, 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.

12.4.14 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.

12.4.15 Dilute hydrochloric acid with calcium carbonate
See diagram 9.154: Limewater 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 limewater.
The surface of the limewater turns milky.
Shake the test-tube containing the limewater.
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 limewater add water to the first test-tube and quickly replace the stopper.
The limewater 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.

12.4.16 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).

12.4.17 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 limewater 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: CO3^2- + 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 limewater.
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.

12.4.20 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.

12.4.23 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 limewater add water to the first test-tube and quickly replace the stopper.
The limewater 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.

12.4.22 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 limewater, at the end of a glass tube, in the mouth of the test-tube.

12.4.21 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 limewater add water to the first test-tube and quickly replace the stopper.
The limewater 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 calcium sulfate formed is only slightly soluble.
It deposits on the carbonate and prevents it from reacting with the acid.
So the reaction with hydrochloric acid above is much better.
sulfuric acid + calcium carbonate --> calcium sulfate + carbon dioxide + water.

12.4.24 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).

12.2.10.2 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 burner to heat the copper touching the iodine crystals.
Yellow copper iodide forms on the copper strip.

12.2.10.3 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).

12.2.10.4 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.

12.2.10.5 Heat iron with sulfur
Heat iron filings with sulfur powder (synthesis reaction).
See diagram 12.2.1: Iron (II) sulfide, FeS
S8 (s) + 8Fe (s) --> 8FeS (s)
Be careful! The following reactions are vigorous.
Do not use large quantities of the chemicals.
Be careful! The reaction of iron (II) sulfide with hydrochloric acid will form the poisonous gas, hydrogen sulfide, with an odour of rotten eggs.
Experiments
1. Mix half a metal bottle top of powdered sulfur with the same volume of iron filings.
Heat a small portion of the mixture on the metal bottle top with the cork removed or in a hard glass test-tube.
When the reaction begins, i.e. the mixture starts to glow, stop heating by moving the Bunsen burner to the side.
If the glow stops, heat the test-tube again.
The reaction of a mixture of iron with sulfur gives out so much heat that the mixture becomes red hot.
Note the following different properties of powdered sulfur, iron filings and the iron (II) sulfide: | appearance | colour | hardness | magnetism|.
Iron is magnetic so is easily removed from a mixture of iron and sulfur, but iron (II) sulfide is not magnetic.
Fe + S --> FeS (s).
2. Use < 10 g total material iron with sulfur in a fume cupboard.
Heat the mixture to start the reaction.
However, be aware that unreacted sulfur may catch fire and produce sulfur dioxide gas to irritate the lungs.
3. Mix uniformly reduced iron powder and powdered sulfur in a weight ratio of seven to four.
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.
4. Mix equal amounts of iron filings and powdered sulfur.
Heat the mixture in a crucible or a small tin with sand in the bottom.
The sand prevents the bottom of the tin from melting by spreading the heat.
Heat the mixture strongly until you see a red glow spreading through the mass.
The heat lost by the chemical reaction is more than the heat needed to start the reaction.
The reaction forms a new substance iron (II) sulfide that has different properties from the two elements used to make it.
Compare iron filings, powdered sulfur, and iron (II) sulfide.
Note their appearance.
Test with a magnet.
Add drops of hydrochloric acid.
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.

12.2.10.6 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).

12.2.10.7 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).

12.2.2.1 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 limewater solution, heat the copper carbonate more strongly.
The copper carbonate turns black and the limewater turns milky.
The copper carbonate has been decomposed by the heat into black copper oxide and carbon dioxide gas, which turns limewater milky.

12.2.2.2 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).
12.2.3.1 Iron displaces

12.2.3.1 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.

12.2.2.3 Electrolytic decomposition, electrolysis
Electrolytic decomposition, electrolysis, occurs when electric current passes through an aqueous solution of a compound.
1. Electrolysis of water, decomposition of water, Hofmann voltameter, electrochemical coulometer: 15.5.4
2. Hydrogen peroxide decomposition, with different catalysts: 17.3.1 (List)
3. Electrolysis of sodium chloride solution: 15.5.12
4. Decomposition of molten sodium chloride, to sodium and chlorine.

12.2.2.4 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 crystals turn into white, or blue-white, powder, weight 1.06 g.

12.2.2.5 Photo decomposition, photolysis
Photo decomposition, photolysis, occurs when a substance is broken down by light, photons
Photolysis: 7.9.42
1. Silver chloride precipitate in photography: 7.8.7.1
2. Decomposition of silver bromide
3. Hydrogen peroxide decomposition, with different catalysts: 17.3.1 (List)
Decomposition of hydrogen peroxide: In the presence of light, hydrogen peroxide decomposes into water and oxygen.

12.2.2.6 Thermal decomposition of acids
1. Decomposition of boric acid
Boric acid, H3BO3, loses water until, above 170^oC, it decomposes to the anhydride, metaboric acid, B2O3, a white, cubic crystalline solid.
H3BO3 --> HBO2 + H2O
Above 300^oC, metaboric acid, loses more water and forms tetraboric acid (pyroboric acid), H2B4O7:
4HBO2 --> H2B4O7 + H2O.
2. Decomposition of carbonic acid
Carbonic acid + heat --> carbon dioxide + water
H2CO3 + heat --> CO2 + H2O.
3. Decomposition of nitric acid
Nitric acid + heat --> nitric oxide + oxygen + water
(Decomposition occurs at 83^oC, catalysed by light, but below 83^oC the colourless pure nitric acid is brown if nitric oxide is already dissolved in it.)
4HNO3 + heat --> 4NO2 + O2 + 2H2O.
4. Decomposition of oxalic acid
Oxalic acid begins to sublime at 100^oC, becomes anhydrous at 189^oC and when heated rapidly decomposes into carbon dioxide, carbon monoxide, formic acid and water.
5. Decomposition of sulfuric acid
Sulfuric acid + heat --> sulfur trioxide + water
H2SO4 + heat --> SO3 + H2O.
6. Decomposition of tartaric acid
Heat tartaric acid (CHOHCOOH)2
Tartaric acid contains the elements carbon, hydrogen, and oxygen, so a residue of carbon is left on heating, and steam forms.

12.2.2.7 Thermal decomposition of sucrose crystals, C12H22O11
1. Heat sugar on a tin lid or in an old spoon.
Note if any gases evolve and if any colours change.
Note the residue left after much heating.
Steam forms, and a black residue of charcoal (carbon), forms.
Sugar is a carbohydrate, a compound of carbon, hydrogen, and oxygen.
The last two elements are usually in the ratio of two to one as in water.
So when sugar is heated, water as steam forms, leaving a residue of black carbon.
2. Heat sucrose to form carbon.
Heat sugar on a metal lid.
The substance melts to form a liquid, which soon turns brown.
If it is cooled at this stage, the brown solid obtained is called caramel.
When heated more strongly the sugar decomposes, giving off inflammable vapours and leaving a black mass of sugar charcoal on the lid.

12.2.3.04 Catalytic decomposition, catalysis
Catalytic decomposition, catalysis, occurs when a catalyst assists in the decomposition
1. Decomposition of potassium chlorate with manganese dioxide catalyst.
Potassium chlorate, KClO3, decomposes above 368oC into potassium perchlorate and oxygen gas. 2 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 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.

12.2.6.01 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.
Experiments
12.2.3.2 Magnesium displaces hydrogen in ethanoic acid
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

12.2.6.1 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 increased and the oxygen gas atoms in O2 reduced, because the oxidation number 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:
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).

12.4.0 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 20^oC C. 36% "ANALAR" solution
D. Sold as: Muriatic acid, for use in the building trades.

12.10.2 Metals, Zn, Fe, displace copper
See diagram 3.2.83: Temperature rise of the reacting solution.
1. Put 25 mL of 0.2 M copper (II) sulfate solution in a 100 mL plastic bottle fitted with a one-hole stopper and thermometer.
Replace the stopper, invert the bottle and shake it gently.
Record the temperature of this solution.
Turn the bottle the right way up, remove the stopper and add 0.5 g of zinc dust.
The quantity of zinc powder is in excess to ensure that all the copper (II) sulfate is used up in the reaction, so some zinc will remain when the reaction stops.
Replace the stopper, invert the bottle, and shake gently.
Record the highest temperature reached.
Calculate the rise of temperature.
This rise of temperature in not affected by the volume of 0.2 M copper (II) sulfate used for the experiment.
For a 1 M solution, multiply the rise in temperature by 5 (5 × 0.2M = 1.0 M).
The reactants lost energy to the solution.
The temperature change is usually between 9^oC and 10^oC.
Zn (s) + Cu²⁺ (aq) --> Zn²⁺(aq) + Cu (s)
2. Repeat the experiment with 0.5 g of iron powder or iron filings.
This amount is again in excess so that all the copper (II) sulfate will be used up in the reaction.
The temperature change is usually between 6^oC and 7^oC.
The zinc metal became zinc ions and copper ions became copper metal, because of 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.
3. Fix two lead foil strips in a beaker and add 200 mL of 1 mol per litre sulfuric acid.
Connect the lead electrodes to a power pack set at 2 V and switch it on for two minutes.
The lead strip connected to the positive terminal becomes covered with brown lead dioxide.
Disconnect the power pack and connect the lead strips to a torch battery.
The battery glows, but the brown leads dioxide on the positive terminal does not disappear.
Repeat the experiment with increasing charging times.
The time the battery glows increases with charging time up to 30 seconds then hardly changes.
Repeat the experiment with different charging voltages.
Different charging voltage makes hardly any difference in the time the battery glows.
However, at high charging voltages hydrogen is produced at the negative electrode and oxygen at the positive electrode.
Charging
At the positive electrode: Pb (s) + 2H2O (l) --> PbO2 (s) + 4H⁺ (aq) + 4e^-
At the negative electrode: 2H⁺(aq) + 2e^- --> H2(g)
Also, lead reacts with the sulfuric acid to produce lead sulfate
At the positive electrode: PbSO4 (s) + 2H2O (l) --> PbO2 (s) + 4H⁺ (aq) + SO4^2- + 2e^-
At the negative electrode: PbSO4 (s) + 2e^- --> Pb (s) + SO4^2-(aq)
So sulfuric acid is produced during charging and is consumed during discharging.
As sulfuric acid has about twice the density of water, the density of the electrolyte shows the state of charge of the battery.
4. When the battery is fully charged, the specific gravity = 1.280, electrode A is lead and electrode, B is lead dioxide.
When the battery discharges, electrode A changes from lead to lead sulfate, electrode B changes from lead dioxide to lead sulfate, concentration of H2SO4 decreases.
When the battery is being charged, these processes are reversed.
The concentration of sulfuric acid suggests the state of charge of the battery so this concentration can be measured with a battery hydrometer.
Electrode A: Pb + SO4^2- --> PbSO4 + 2e^-
Electrode B: PbO2 + 4H3O⁺ + SO4^2- + 2e^- --> PbSO4 + 6H2O
In a motor car battery, the electrodes have a coat of lead (II) oxide (PbO) and lead powder (Pb).
In the electrolyte, electric current converts the PbO to Pb on the negative plate, and the PbO to lead (IV) oxide (lead peroxide) PbO2 on the positive plate.
Discharging -->
PbO2 + 2H2SO4 + Pb < = > 2PbSO4 + 2H2O
<-- Charging
If you pass electricity through the battery after it is fully charged, "gassing" occurs, i.e. water is decomposed into hydrogen and oxygen gas.
Never smoke or allow a naked flame near a charging battery.

12.10.3 Displacement of copper by zinc
See diagram 33.3.2a: Copper and zinc foil in a voltmeter, a simple cell.
1. Put concentrated copper (II) sulfate solution in a beaker.
Connect copper foil to the positive terminal, red wire, of a voltmeter and a zinc foil to the negative terminal, black wire.
Simultaneously dip the two metals briefly into the copper sulfate solution.
Record the readings on the voltmeter.
The voltage falls to zero after a short time, because black copper deposited on the zinc and caused the reaction to stop.
When copper deposits on the zinc electrode, it prevents more zinc from entering the solution.
This causes the voltage to fall to zero after a short time and the cell becomes "dead".
Separate the electrolytes to prevent the voltage fall by using 1. a Daniell Cell that has a porous pot, or 2. a salt bridge.
2. Pour concentrated copper (II) sulfate solution into a beaker.
Connect a copper rod to the positive terminal of a voltmeter and a zinc rod to the negative terminal.
Dip the two metals briefly into the copper (II) sulfate solution.
Zinc dissolves and hydrogen bubbles form on the surface of the copper.
The voltmeter reads 1.1 V, so electrons are moving from the zinc to the copper.

12.4.18 Dilute sulfuric acid as an acid
1. Add 2 cm of dilute sulfuric acid to 1 cc of zinc powder.
Close the test-tube with the thumb until enough hydrogen gas forms to give a mild explosion when the mouth of the test-tube is held in a flame.
2H⁺ + Zn (s) --> Zn²⁺ + H2 (g).
2. Add 2 cm of sodium carbonate to 1 cc of zinc powder.
Tests for carbon dioxide by passing the gas given off to pass into limewater that turns milky, because of the fine precipitate of calcium carbonate.
2H⁺ + CO3^2- --> H2O +CO2 (g)
Ca(OH)2 + CO2 (g) --> CaCO3 (s) +H2O.

12.4.19 Dilute sulfuric acid as a sulfate
Add an equal volume of barium chloride solution to 3 cm of dilute sulfuric acid.
Note the white precipitate of barium sulfate.
Allow the precipitate to settle, filter, wash and leave to dry.
SO4^2- + Ba²⁺ --> BaSO4 (s).

12.19.2.2 Displace a less reactive halogen from halogen compounds.
More reactive halogen displaces less reactive halogen from its compound.
1. Add iodine solution to colourless potassium bromide solution.
No reaction, because iodine is less active than bromine.
Pass chlorine gas through colourless potassium bromide solution.
The more active chlorine displaces the less active bromine and the solution turns orange.
2KBr (aq) + Cl2 (g) --> 2KCl (aq) + Br2 (aq).
2. Pass chlorine gas through potassium iodide solution.
The more active chlorine displaces the less active iodine and the solution turns deep brown.