School Science Lessons
Topic 12A
2019-07-27
Please send comments to: J.Elfick@uq.edu.au

12A Activity series, buffer solutions, reactions of:

Table of contents
See: Chemicals, (Commercial)

12.10.7.0 Prepare buffer solutions

12.16.0 Reactions of carbonates

12.15.6 Reactions of metals with ligands

12.15.0 Reactions of metals with water

12.15.5.0 Reactions of non-metals with water

12.15.5.1 Heat carbon with steam, water gas

12.17.0 Reactions of oxides

12.10.0 Reactions of salts

12.14.0 Reactivity metals as reducing agents

12.11.5.0 Tests for anions

12.11.4.0 Tests for cations

12.11.6.0 Tests for metallic radicals

12.10.7.0 Prepare buffer solutions
See: Buffer Solutions (Commercial)
12.10.7.0 Prepare buffer solutions
Experiments
12.10.10 Change in pH near the equivalence point
12.10.9 Prepare a buffered salt with sodium acetate
12.10.7.1 Prepare dilute buffer solution
12.10.7.2 Prepare natural buffer
12.10.7.3 Prepare buffered solutions, methyl orange
12.10.8.0 Prepare solutions, pH values 3 to 11, with buffer solutions
12.10.8.1 Prepare solutions, hydrogen ion concentrations 10-3 to 10-6 g ions per litre
12.10.8.2 Prepare solutions, hydrogen ion concentrations 10-7 to 10-11 g ions per litre
12.10.7.4 Salt effect on buffer solutions, methyl orange

12.14.0 Reactivity series of metals as reducing agents (reactivity series, electrochemical series)
12.14.0 Activity series of metals as reducing agents
12.14.2 Metals with copper (II) sulfate solution
12.14.02 Reactions of metals with air or oxygen gas
12.14.03 Reactions of metals with dilute acids
12.14.04 Reactions of metals with concentrated oxidizing acids
Experiments
12.14.2.6 Activity of metals and tendency to form ions
3.80 Exothermic reactions give out heat energy (See 2.)
12.17.2.2 Heat metals with oxides of another metal
12.14.2.2 Iron with copper (II) sulfate solution
12.14.2.3 Iron and zinc with copper (II) sulfate solution
3.72 Magnesium displaces copper from solution of copper ions
12.14.2.1 Magnesium, or zinc, with copper (II) sulfate solution
3.74 Metals displace hydrogen from acids
12.14.1 Zinc displaces lead from lead nitrate solution
12.14.2.4 Zinc with lead nitrate solution, iron with copper (II) sulfate solution
12.14.2.5 Zinc with copper in sulfuric acid

12.16.0 Reactions of carbonates
12.16.0 Carbonates
35.19.2: Carbonates (Geology), CO32-
1.11: List of carbonates
12.16.8 Prepare sodium carbonate, LeBlanc process
12.16.7 Prepare sodium carbonate, Solvay process
12.16.3.5 Smelling salts, ammonium carbonate
Sal volatile
Experiments
12.16.3.3 Ammonium carbonate with acids
12.16.3.2 Ammonium carbonate with alkalis
12.16.3.4 Ammonium carbonate solution precipitates metal carbonates
12.16.1 Carbon dioxide through calcium carbonate suspension
12.16.1.1 Carbon dioxide with calcium hydroxide solution
3.30.1: Decomposition of carbonates
12.3.9.0: Dilute acids with common carbonates:
12.3.10 Dilute acids with sodium hydrogen carbonate
12.3.27 Egg in a bottle
5.0 Expose sodium carbonate decahydrate, washing soda, to the air
12.16.3.1 Heat ammonium carbonate (smelling salts)
12.16.3 Heat carbonates of Cu, Mg, Na, Pb and Zn
12.16.4 Heat sodium hydrogen carbonate
13.7.6: Prepare carbon dioxide, heat carbonates
12.16.2 Prepare sodium hydrogen carbonate
12.16.6 Prepare volcanos with baking soda
12.11.5.7: Tests for carbonates

12.15.0 Reactions of metals with water
12.15.0 Reactions of metals with water
3.73 Reactions of sodium with water
12.15.3 Reactions of metals with steam
12.15.1 Reactions of metals with water, Cu, Zn, Fe, Mg, Al

12.17.0 Reactions of oxides
12.17.0 Oxides, acidic, basic, amphoteric, neutral and mixed oxides
12.17.3 Carbon dioxide, acidic oxides (non-metal oxides)
3.36 Carbon dioxide and photosynthesis
3.37 Carbon dioxide and respiration
3.34.0 Carbon dioxide properties
3.34.3 Carbonic acid, aerated water, carbonated water, soda water
12.17.1.1 Oxides and the periodic table
Experiments
3.38 Carbon dioxide and fermentation for brewing
12.17.3.2 Carbon dioxide with barium hydroxide solution
12.17.3.1 Carbon dioxide with sodium hydroxide solution
12.17.2 Copper (II) oxide (copper oxide), basic oxide (metal oxide)
3.30.5 Decomposition of oxides
3.34.5 Frozen carbon dioxide ("dry ice", "hot ice")
12.17.2.2 Heat metals with oxides of another metal
12.17.2.1 Heat zinc with copper (II) oxide
12.17.1 Properties of oxides
3.34.4 Reduce carbon dioxide with burning magnesium
10.10.1 Reduce metal oxides to metals, red lead to lead and oxygen
3.34.6 Soda acid fire extinguisher
3.34.1.0 Tests for carbon dioxide
3.35.4 Yeast cells convert glucose to carbon dioxide gas and alcohol

12.10.0 Reactions of salts
Experiments
12.10.2.6.1 Artificial gemstones
12.10.1.0 Crystals of different salts, storm glass
12.3.5 Dilute acids with basic oxides, metal oxides, copper (II) oxide
8.0.0 Direct union of elements to form compounds, sodium with chlorine
12.10.3.2 Hydrolysis of ammonium chloride
12.10.5 Hydrolysis of iron (III) chloride
12.10.3 Hydrolysis of sodium carbonate
12.10.3.1 Hydrolysis of sodium hydrogen carbonate
12.10.4 pH of salt solutions
12.10.2 Prepare salts by different methods
12.10.6 Prepare sodium hydrogen sulfate
12.2.4.2 Prepare zinc sulfate crystals
3.75 Reactions of salts with water
12.10.2.6 Salt solutions with another salt

12.11.5.0 Tests for anions
Experiments
12.11.5.0 Tests for acid radicals in solution
12.11.5.1 Tests for acetates
12.11.5.2a Tests for antimonates, borates, oxalates
12.11.5.3 Tests for arsenates
12.11.5.4 Tests for bicarbonates
12.11.5.5 Tests for borates
12.11.5.6 Tests for bromides
12.11.5.7 Tests for carbonates
12.11.5.8 Tests for chlorides
12.11.5.9 Tests for chromates
12.11.5.10 Tests for halides, Cl-, Br-, I-
12.11.5.11 Tests for hydroxides
12.11.5.12 Tests for iodides
12.11.5.13 Tests for nitrates
12.11.5.14 Tests for oxalates
12.11.5.15 Tests for phosphates
12.11.5.16 Tests for sulfates
12.11.5.17 Tests for sulfides
12.11.5.18 Tests for sulfites

12.11.4.0 Tests for cations, prepare a solution for group analysis
Experiments
12.11.4.1 Group 1 tests for Ag+, Pb2+
12.11.4.2 Group 2 tests for Bi3+, Cd2+, Cu2+, Sn2+
12.11.4.3 Group 3 tests for Al3+, Cr3+, Fe2+, Fe3+
12.11.4.4 Group 4 tests for Co2+, Mn2+, Ni2+, Zn2+
12.11.4.5 Group 5 tests for Ba2+, Ca2+, Sr2+
12.11.4.6 Group 6 tests for K+, Mg2+, Na+, NH4+

12.11.6.0 Tests for metallic radicals
12.11.6.0 Tests for metallic radicals
12.11.6.1 Chemistry of group separations.
Experiments
12.11.6.2 Preliminary experiments before separation of Group I metals, silver and lead.
12.11.6.3 Separation into groups.
12.11.7.1 Group I Insoluble chlorides, PbCl2, AgCl (Hg2Cl2 omitted)
12.11.7.2 Group II Sulfides insoluble in dilute hydrochloric acid
12.11.7.2a Group IIa PbS, Bi2S3, CuS, CdS (HgS omitted)
12.11.7.2b Group IIb As2S3, Sb2S3, SnS, SnS2
12.11.7.3 Group III Insoluble hydroxides, Fe(OH)3, Cr(OH)3, Al(OH)3
12.11.7.4 Group IV Insoluble sulfides precipitated by hydrogen sulfide, ZnS, MnS, CoS, NiS
12.11.7.5 Group V Insoluble carbonates, CaCO3, BaCO3, SrCO3
12.11.7.6 Group VI Magnesium, sodium and potassium, Mg, Na, K

12.10.0 Salts, acid salt, sodium chloride, "table salt"
3.71.1 Solubility table and solubility rules
A salt is the product with water of the reaction of an acid with a base.
A salt is a compound formed when the hydrogen ion of an acid is replaced by a metal ion or electropositive complex ion, e.g. NH4.
An acid salt forms when an acid contains more than one replaceable hydrogen ion, e.g. H2SO4 and not all the hydrogen ions are
replaced, e.g. NaH(SO4)2.
Salts are usually crystalline and are composed of positive and negative ions.
You can prepare insoluble salt precipitates from pairs of solutions of salts by using the solubility rules.
Sodium chloride is an ionic solid.
Crystals of sodium chloride contain Na+ and Cl- ions attracted to each other by strong ionic bonds in a crystal lattice.
The crystals are hard and have high melting points and boiling points.
When melted or in solution, sodium chloride conducts electricity, but the solid is a poor conductor of electricity.

12.10.1.0 Crystals of different salts, storm glass
1. Dissolve different salts in water.
Slowly evaporate the solution until salt crystals start to form.
Add a crystal of salt to help crystallization.
Describe the colour and shape of different salt crystals.

2. Storm glass
The "storm glass" is a solution of salts that changes to form different types of crystals in different weather conditions, probably caused
by changes in temperature.
The liquid within the sealed glass container contains ethanol, potassium nitrate, ammonium chloride and camphor.
It was invented by Captain Robert Fitzroy, founder of the UK Meteorological Office and formerly the captain of HMS Beagle.
From Australian Geographic:
During Charles Darwin's historic voyage in the Beagle, 27-12-1831 to 2-10-1836, Captain FitzRoy carefully documented how the
storm glass would predict the weather:
If the liquid in the glass is clear, the weather will be bright and clear.
If the liquid is cloudy, the weather will be cloudy as well, perhaps with precipitation.
If there are small dots in the liquid, humid or foggy weather can be expected.
A cloudy glass with small stars indicates thunderstorms.
If the liquid contains small stars on sunny winter days, then snow is coming.
If there are large flakes throughout the liquid, it will be overcast in temperate seasons or snowy in the winter.
If there are crystals at the bottom, this indicates frost.
If there are threads near the top, it will be windy.
In 1859, violent storms struck the British Isles.
In response, the British Crown distributed storm glasses, then known as "FitzRoy's storm barometers," to many small fishing
communities around the British Isles that were to be consulted by ships at port before setting sail.

12.10.2.6 Salt solutions with another salt
This is the only way to prepare an insoluble salt.
In this type of reaction, the needed salt forms a precipitate.
When solutions of two ionic substances are mixed and the ions of an insoluble salt are in this mixture, then a precipitate of the insoluble
salt forms.

Experiments
Make dilute solutions of different salts in separate test-tubes, e.g. barium nitrate, silver nitrate and lead nitrate.
To each add a small quantity of dilute hydrochloric acid from a dropping tube.
Note the colour and appearance of any precipitate that forms.
Repeat the procedure using 1. sodium chloride solution, 2. sodium sulfate solution, 3. dilute sulfuric acid.
silver ions (aq) + chloride ions (aq) --> silver chloride (s) (Silver chloride is insoluble in water)
lead ions (aq) + chloride ions (aq) --> lead chloride (s) (Lead chloride is insoluble in water)
sodium nitrate (aq) + copper (II) sulfate (aq) --> sodium ions (aq) + nitrate ions (aq) + copper ions (aq) + sulfate ions (aq)
(No precipitate because both sodium sulfate and copper nitrate are soluble in water.)

12.10.2.6.1 Artificial gemstones
Half fill a Petri dish with water.
At one side, carefully pour some potassium sulfate solution.
At the other side carefully pour some aluminium sulfate solution (swimming pool flocculent powder).
Leave to allow potassium aluminium sulfate crystals to form in the middle.
Add some lead chromate solution.
The crystals will change colour like artificial gemstones.

12.10.3 Hydrolysis of sodium carbonate
Washing powders contain di-sodium tetraborate (III)-10-water (borax) + sodium carbonate and are alkaline in solution.
Hydrolysis is a chemical reaction of a compound with water.
Hydrolysis of salts is the reverse of neutralization.
Salts of weak acids or weak bases hydrolyse when dissolved in water.
Weak acids with weak alkalis dissociate very slightly.
Solvation occurs when solvent molecules form bonds with a solute particle.
Experiment
Dissolve sodium carbonate in water.
Some hydrogen ions react to form the weak acid carbonic acid leaving excess hydroxyl ions in the solution.
The solution turns red litmus blue.
salt + water --> acid + base
Na2CO3 (aq) <--> 2Na+ (aq) + CO32- (aq)
H2O (l) <--> H+ (aq) + OH- (aq)
2H+ (aq) + CO32- (aq) <--> H2CO3 (aq) carbonic acid
Na2CO3 (aq) + H2O (l) <--> 2NaOH (s) + H2CO3 (aq)

12.10.3.1 Hydrolysis of sodium hydrogen carbonate
Sodium hydrogen carbonate (sodium bicarbonate, baking soda), has a basic reaction and can be used to neutralize acids in fruit or
neutralize bee stings.
Low cost: from supermarkets as baking soda.
Experiment
Dissolve sodium hydrogen carbonate in water.
A sodium hydrogen carbonate solution turns red litmus blue.
NaHCO3 (aq) <--> Na+ (aq) + HCO3- (aq)
H2O (l) <--> H+ (aq) + OH- (aq)
HCO3- (aq) + H+ (aq) <--> H2CO3(aq)

12.10.3.2 Hydrolysis of ammonium chloride
Dilute ammonia solution is only slightly dissociated, because it is a very weak alkali.
The ammonium ions react with hydroxyl ions to form undissociated dilute ammonia solution leaving excess of hydrogen ions.
The solution of ammonium chloride has pH value of about 6.
Experiment
Dissolve ammonium chloride in water.
NH4Cl (aq) <--> NH4+ (aq) + Cl- (aq)
NH4+ (aq) + OH- (aq) <--> NH4OH (s)
NH4Cl (aq) + H2O (aq) <--> NH4OH (s) + H+ (aq) + Cl- (aq)

12.10.4 pH of salt solutions
See: pH (Commercial)
In a "normal salt", the replaceable hydrogen atoms of an acid have been completely replaced by a metal, e.g. sodium chloride, NaCl.
In an "acid salt", the replaceable hydrogen atoms of an acid have not all been replaced, e.g. potassium hydrogen sulfate, KHSO4.
However, a "normal salt" is not necessarily a neutral salt because hydrolysis may occur, e.g. sodium carbonate is alkaline in solution, but
ammonium chloride is acidic in solution.
(In most chemistry curricula nowadays, the term "normal" is no longer used!)

Experiments
(1.) Half fill seven test-tubes with water and add two drops of universal indicator.
Add 1.25 mL of the following salts: sodium carbonate, sodium sulfite, sodium chloride, ammonium chloride, aluminium chloride, borax,
iron (II) sulfate.
Note the pH value according to the colour produced.
(2.) Warm the solutions and note whether this increases the hydrolysis, in some cases producing greater divergence from neutrality.

(3.) Add three drops of universal indicator to 5 mL of 0.2 M a salt solution
Record: Salt, colour, pH
1. NH4Cl, orange-red, pH 5,
2. NaCl, yellow-green, pH 7,
3. Na2HPO4, blue-green, pH 9,
4. KNO2 , blue, pH 9.5, 5. Na2CO3, violet, pH 10,
6. Na2S, red-violet, pH 10.5.

(4.) Test solutions with litmus paper
1. Sodium sulfate solution, neutral
2. Iron sulfate solution, blue litmus paper turns pale mauve, acid solution
3. Sodium hydrogen carbonate solution, alkaline solution

12.10.5 Hydrolysis of iron (III) chloride
Iron chloride exists as anhydrous iron (III) chloride (FeCl3), and hydrated Iron (III) chloride-6-water (FeCl3.6H2O).
Iron (III) chloride is rapidly hydrolysed in moist air and is partially hydrolysed in solution.
Hydrolysis can be suppressed by addition of HCl.
Fe(OH)2, green, is oxidized to Fe(OH)3, brown, in moist air.

Experiment
Dissolve iron (III) chloride in boiling water.
Add drops of dilute ammonia solution.
The reaction forms a red-brown precipitate of iron (III) hydroxide.
FeCl3 (aq) + NH4OH (aq) --> Fe(OH)3 (s) + NH4Cl (aq)
Heat to evaporate some solution.
The reaction forms a red-brown precipitate of iron (III) hydroxide.
FeCl3 (aq) + 3H2O (l) --> Fe(OH)3 (s) + 3HCl (l)
Pour the clear saturated solution into hot water.
The reaction forms a red precipitate of hydrated iron (III) oxide.
2FeCl3 (aq) + 3H2O (l) --> Fe2O3 (s) + 6HCl (l)
Add drops of sodium hydroxide solution.
The reaction forms a red-brown precipitate of iron (III) hydroxide.
FeCl3 (aq) + 3NaOH (aq) --> Fe(OH)3 (s) + NaCl (aq)

12.10.6 Prepare sodium hydrogen sulfate
Sodium hydrogen sulfate is called an "acid salt" because it is the salt of an acid containing more than one acidic hydrogen, e.g. H2SO4,
that has not had all the hydrogen replaced by positive ions.
Experiment
Add drops of 2 M sulfuric acid to 2 M sodium hydroxide.
Count the drops until the solution is neutral to litmus.
Repeat the experiment by adding half the number of drops of acid.
H2SO4 (aq) + NaOH (aq) --> NaHSO4 (aq) + H2O (l).

12.10.7.0 Prepare buffer solutions
39.0 Prepare standard buffer solutions
7.9.11 Buffer
A buffer solution is a mixture of substances that tend to hinder large changes in acid or basic properties of a solution.
The term is used in a more general sense outside chemistry.
The pH of a buffer solution is not greatly changed by the addition of an acid or an alkali.
Most buffer solutions are a mixture of a weak acid or base with one of its salts.
In body fluids, the buffers include H2CO3 with HCO3-.
Buffers based on weak acids work because the dissociation of a weak acid is an equilibrium reaction.
If a weak acid, HA, dissociates in solution as follows:
HA (aq) <==> H+ (aq) + A- (aq), both [H+ (aq)] and [A- (aq)] are small because HA is a weak acid
If OH- ions are added, the OH- ions react with H+ to form water to disturb the equilibrium to cause more HA to dissociate until the
original pH value is restored.
If soluble salt, Na+A-, is added to increase the A-, more H+ will be removed.
So the weak acid in a buffer solution is a source of H+ to react with added OH-, and the salt component reacts with added H+.
Acidic buffer, e.g. sodium hydrogen carbonate with carbonic acid solutions, the salt of the weak acid is completely dissociated into ions
but the weak acid is only partly dissociated.
Basic buffer, e.g. ammonium chloride in ammonia solutions
The pH value of buffer solutions changes very little when acids or alkalis are added or when diluted with water.
Although the salts of weak acids are completely dissociated into ions, weak acids do not dissociate completely.
A buffer solution contains a weak acid and the salt of the weak acid,
e.g. H2CO3 / HCO3- (carbonic acid / sodium hydrogen carbonate).
By mixing an acid with its conjugate base, definite hydrogen ion concentrations, within a certain range depending on the dissociation
constant of the acid, are obtainable.
Such solutions have the advantage that evaporation will not affect the value of (H+), because the ratio (acid) / (base) remains constant.
Contamination by small quantities of acidic or basic impurities will not affect the pH.
If an acid is added to a buffer solution, the H+ added reacts with the HCO3-.
If a base is added to a buffer solution, the OH- reacts with the undissociated H2CO3 to form the salt and water.
Natural body fluids are buffered.
Examples of buffer solutions:
1. Hydrochloric acid with ammonia in excess, HCl with NH3 in excess, i.e. strong acid with weak base in excess.
2. Hydrochloric acid with sodium acetate in excess, HCl with CH3COONa in excess, i.e. base of a weak acid with strong acid.
3. Sodium hydroxide with acetic acid in excess, NaOH with CH3COOH in excess, i.e. strong base with weak acid in excess.
4. Sodium acetate with acetic acid, CH3COONa with CH3COOH, i.e. base of a weak acid with weak acid.
5. Sodium hydroxide with ammonium chloride in excess, NaOH with NH4Cl in excess.
6. Ammonium chloride with ammonia, NH4Cl with NH3.
NH4+ (aq) <=> NH3 (aq) + H+ (aq)
If add OH-: H+ (aq) + OH- (aq) --> H2O (l)
If add H+: NH3 (aq) + H+ (aq) --> NH4+ (aq)
7. Blood is buffered to pH 7.4, mainly from the following equation:
H+ (aq) + HCO3- (aq) <=> CO2 (aq) + H2O (l)
8. Skin products, shampoos and detergents may be buffered using phosphoric acid and sodium phosphate.

12.2.4.2 Prepare zinc sulfate crystals
Add pieces of granulated zinc to sodium hydrogen sulfate solution in a test-tube, and heat.
When there is no further reaction and all the blue colour has disappeared, filter the mixture.
The filtrate is zinc sulfate solution.

12.10.7.1 Prepare dilute buffer solutions
Add 1 mL of 0.01M HCl to 1 mL of water.
The pH value changes from 7 to 5.

12.10.7.2 Prepare natural buffer
Add 1 mL of 0.01M HCl to one cube of beef soup (beef cube infusion).
Almost no pH change occurs because of buffering action.

12.10.7.3 Prepare buffered solutions
Methyl orange: pH 2.5 (red), pH 3.5 (straw colour), pH 4.5 (orange).
Add a drop of methyl orange to the following:
(1.) deionized water.
It turns yellow.
(2.) deionized water + 5 drops ethanoic acid (acetic acid).
It turns pink.
(3.) deionized water + 5 drops ethanoic acid + crystals of sodium acetate-3-water.
It turns yellow.
The (3.) solution is buffered, so it does not turn pink as in the (2.) solution.

12.10.7.4 Salt effect on buffer solutions
Add drops of methyl orange to:
1. deionized water.
The solution turns yellow.
2. Dilute hydrochloric acid.
The solution turns red.
3. Dilute ethanoic acid (acetic acid).
The solution turns slightly red.
4. Very dilute acetic acid.
The solution turns red.
The very dilute acetic acid is red as with dilute hydrochloric acid.
5. Half the very dilute acetic acid solution + sodium chloride crystals.
The solution turns pale red.
The salt effect prevents reformation of molecular acetic acid.

12.10.8.0 Prepare solutions, pH values 3 to 11, with buffer solutions
The pH value of a buffer solution does not alter for small additions of acid or alkali, e.g. a mixture of highly ionized sodium acetate.
CH3COONa, and partly ionized acetic acid, CH3COOH or HAc.
1. If add hydrogen ions to the solution, the HAc that forms is undissociated and so H+ are removed from the solution.
H+ + Ac- --> HAc
2. If add alkali to the solution, more HAc dissociates to form hydrogen ions that combine with the hydroxyl ions to form H2O that is
undissociated, and so OH- ions are removed from the solution.
HAc --> H+ + Ac-
H+ + OH- --> H2O.

12.10.8.1 Prepare solutions, hydrogen ion concentrations 10-3 to 10-6 g ions per litre
Use the following solutions:
1. 0 1M acetic acid solution
2. 0.1 M sodium acetate solution (13.6 g of crystalline sodium acetate, CH3COONa.3H2O per litre)
1.1 Hydrogen ion concentration 10-3: 1 litre 0.1 M acetic acid and 18 mL 0.1 M sodium acetate
1.2 Hydrogen ion concentration 10-4: 1 litre 0.1 M acetic acid and 180 mL 0.1 M sodium acetate
1.3 Hydrogen ion concentration 10-5: 555 mL 0.1 M acetic acid and 1 litre 0.1 M sodium acetate
1.4 Hydrogen ion concentration 10-6: 55 mL 0.1 M acetic acid and 1 litre 0.1 M sodium acetate.

12.10.8.2 Prepare solutions, hydrogen ion concentrations 10-7 to 10-11 g ions per litre
Use the following solutions:
1. Disodium phosphate solution (Na2HPO4): Dissolve 0.1 mole of the crystalline salt Na2HPO4, 35.8 g
2. 0.1 M hydrochloric acid
3. 0.1 M sodium hydroxide.
4.1 Hydrogen ion concentration 10-7: 1 litre Na2HPO4 solution and 322 mL 0.1 M HCl solution
4.2 Hydrogen ion concentration 10-8: 1 litre Na2HPO4 solution and 47 mL 0.1 M HCl solution
4.3 Hydrogen ion concentration 10-9: 1 litre Na2HPO4 solution and 5 mL 0.1 M HCl solution
4.4 Hydrogen ion concentration 10-10: 1 litre Na2HPO4 solution and 3.6 mL 0.1 M NaOH solution
4.5 Hydrogen ion concentration 10-11: 1 litre Na2HPO4 solution and 3.6 mL 0.1 M NaOH solution.

12.10.9 Prepare a buffer salt with sodium acetate, universal indicator
A buffer salt is essentially a highly ionized salt of a weak acid.
1. Add two drops of universal indicator to 10 mL of 0.1 M sodium hydroxide solution.
Titrate the mixture with 0.1 M hydrochloric acid.
Note the colour changes that indicate the rapid change of pH about the equivalence point.
2. Add 5 g of sodium acetate to 10 mL of 0.1 M sodium hydroxide solution, then two drops of indicator.
Titrate the mixture with 0.1 M hydrochloric acid.
Note hydrogen ions are added, but the green colour of the indicator persists because the pH remains constant over a long period of
addition of hydrogen ions.
The buffer salt, sodium acetate, is highly ionized and gives acetate ions.
The hydrogen ions from the hydrochloric acid form molecular acetic acid instead of increasing the hydrogen ion concentration in the
solution.
NaAc --> Na+ + Ac-
H+ + Ac- <--> HAc
When a large excess of hydrogen ions is added, the pH of the solution decreases.
Adding a strong alkali to a highly ionized salt of a weak base does not at first increase the pH of the mixture.
Ammonium chloride solution gives ammonium ions that react with the added hydroxide ions of a strong alkali to form molecular
"ammonium hydroxide".
(Not "ammonium hydroxide, NH4OH".
Ammonia solution is shown as NH3 (aq) because "NH4+" ions and "OH-" ions can be detected, but "NH4OH" cannot be detected.)
The pH of the solution rises only after an excess of alkali is added.

12.10.10 Change in pH near the equivalence point
See: pH (Commercial)
1. Add two drops of universal indicator to 10 mL of sodium hydroxide solution.
Titrate the mixture with hydrochloric acid.
Note the rapid change of colour from blue-green at pH about 8.5 to orange-red at pH about 4.
When a strong alkali is titrated against a strong acid, the indicator indicates the equivalence point with negligible error.
Repeat the experiment with a low pH indicator, e.g. methyl orange, and a high pH indicator, e.g. phenolphthalein, and note the slight
difference.
2. Add two drops of universal indicator to 10 mL of sodium hydroxide solution and titrate the mixture with acetic acid.
Note that when the equivalence point is reached, the pH is about 8.5. Note also the considerable excess of acid necessary to approach
the orange colour of pH 4, showing that only a high pH indicator is efficient in the titration of a strong alkali with a weak acid.
Repeat the experiment using phenolphthalein and methyl orange.

12.10.12 pH values of oxides
See: pH (Commercial)
See 12.17.0: Oxides, acidic, basic, amphoteric, neutral and mixed oxides
Add 3 drops of Universal Indicator to 2 CC of the following oxides and note the colour change, pH value and state whether the
oxides are acid, alkali or neutral
1. 0.2 M Nitric acid (nitrogen oxide and water)
2. 0.2 M Sodium hydroxide (sodium oxide and water)
3. 0.2 M Potassium hydroxide (potassium oxide and water)
4. 0.2 M Phosphoric acid (phosphorus (V) oxide and water)
5. 0.2 M Calcium hydroxide (calcium oxide and water)
The soluble oxides of metals are alkaline and the oxides of non-metals are acidic oxides.

12.11.5.0 Tests for acid radicals in solution
Before testing a solution for acidic radicals remove heavy metals that may interfere with the tests, leaving only sodium, potassium or
ammonium in solution, e.g. to test for a sulfate radical in solution, add dilute hydrochloric acid and barium chloride solution.
A white precipitate of barium sulfate indicates the presence of a sulfate.
Ba2+ + SO42- --> BaSO4 (s)
However, if the solution already contains the silver ion, the white precipitate is silver chloride.
Ag+ + Cl- --> AgCl (s).

Experiment
Boil 1 g of the finely divided unknown solid with sodium carbonate solution to precipitate heavy metals as carbonates, or as hydroxides
by hydrolysis.
Filter off the precipitates.
Copper may rarely form a soluble double carbonate.
The acidic radicals, originally combined with the heavy metals, are now in the filtrate as the sodium salts if double decomposition has
occurred, e.g. a mixture containing barium chloride and calcium nitrate:
BaCl2 + Na2CO3 --> BaCO3 (precipitate) + 2NaCl (solution)
Ca(NO3)2 + Na2CO3 --> CaCO3 (precipitate) + 2NaNO3 (solution)
The filtrate is alkaline with excess sodium carbonate and now must be made acidic, e.g. barium chloride use hydrochloric acid, with
silver nitrate use nitric acid so you do not add the radical you are testing for.
If the solutions are not made acid, the sodium carbonate precipitates the metal of the testing reagent as a heavy metal carbonate.

12.11.5.1 Tests for acetates
1. Add to 5 drops of original solution drops of dilute HCl, or HNO3 if using a Pb salt.
If effervescence occurs, pass the gas through lime water.
A milky precipitate indicates (CO3)2-.
If effervescence does not occur, heat the solution.
The odour of vinegar indicates CH3COO-.

2. Neutralize with dilute nitric acid and ammonia, then add iron (III) chloride solution.
A blood red colour, lost by adding hydrochloric acid, indicates an acetate.

3. Add an equal volume of alcohol and then drops of concentrated sulfuric acid.
Heat gently and smell the vapour.
The fruity smell of ethyl acetate indicates the presence of an acetate.
CH3COONa + C2H5OH + H2SO4 --> CH3COOC2H5 + NaHSO4 + H2O.

12.11.5.2a Tests for antimonates, borates, oxalates
Add excess dilute nitric acid, followed by silver nitrate solution.
Filter off the precipitate.
Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a white precipitate forms that indicates
antimonate, borate, or oxalate in the filtrate.

12.11.5.3 Tests for arsenates
1. Add dilute nitric acid and excess ammonium molybdate solution.
Heat to boiling.
A yellow precipitate of ammonium arsenomolybdate (NH4)3AsO4.12MoO3, indicates arsenate.
2. Add excess dilute nitric acid, followed by silver nitrate solution.
Filter off the precipitate.
Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a brick-red precipitate forms that indicates
arsenate in the filtrate.

12.11.5.4 Tests for bicarbonates
Add magnesium sulfate solution.
A white precipitate in the cold indicates the presence of carbonate.
No precipitate in the cold, but a white precipitate on boiling, confirms bicarbonate.
If the original solid is insoluble in water, an aqueous suspension of it may be boiled.
A solution that produces carbon dioxide indicates the presence of bicarbonate.

12.11.5.5 Tests for borates
1. Dissolve 1g of boric acid in 10 mL of ethanol.
Use a trigger pump-operated spray bottle, e.g. window cleaner spray bottle, to spray the solution onto a roaring Bunsen burner flame.
A green flame indicates borates.
2. Add concentrated sulfuric acid to the unknown substance then pour into methylated spirit into an evaporating dish while stirring with
a glass rod.
Heat the evaporating dish and light the vapour rising it.
A green colour in the flame produced by the volatile compound, ethyl borate, indicates borate radical.
Na2B4O7 + H2SO4 + 5H2O --> Na2SO4 + 4H3BO3
H3BO3 + 3C2H5OH --> B(OC2H5)3 + 3H2O
The test may not work for a few minerals containing boron, e.g. borosilicates.
3. To confirm borate, acidify the solution and test with turmeric paper.
Dry the paper over a small flame.
The change of colour from yellow to brown, which becomes blue or blue-black in caustic soda solution indicates a borate.
4. Tests for borate, oxalate, antimonate
Add excess dilute nitric acid, followed by silver nitrate solution.
Filter off the precipitate.
Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a white precipitate forms that indicates the
presence of antimonate, borate, or oxalate in the filtrate.

12.11.5.6 Tests for bromides
1. Add excess dilute nitric acid, followed by silver nitrate solution.
A pale yellow precipitate of silver bromide, sparingly soluble in ammonia, indicates the presence of the bromide radical.
Ag+ + Br- --> AgBr (s)
2. To confirm the bromide radical, heat the solid with manganese dioxide and concentrated sulfuric acid and observe the dark red
vapour of bromine.

12.11.5.7 Tests for carbonates
Add magnesium sulfate solution.
A white precipitate in the cold confirms carbonate.
No precipitate in the cold, but a white precipitate on boiling, confirms bicarbonate.
If the original solid is insoluble in water, an aqueous suspension of it may be boiled.
If the solution produces carbon dioxide, a bicarbonate is indicated.

12.11.5.8 Tests for chlorides
1. Add excess dilute nitric acid, followed by silver nitrate solution.
A white precipitate of silver chloride, soluble in ammonia, indicates the presence of chloride radical.
Ag+ + Cl- --> AgCl (s)
AgCl + 2NH3 --> Ag(NH3)2Cl (soluble silver amine)

12.11.5.9 Tests for chromates
Most chromates are only slightly soluble or insoluble so the tests are mainly for sodium, potassium or ammonium chromate (VI) ions.
A solution with a bright yellow colour indicates that it is worth testing for chromate (VI) ions.
Oxidation reactions involve the reduction of solutions of chromate or dichromate ions that cause colour changes from yellow or orange
to pale green or colourless solutions.
The reactions with the formation of an insoluble metal chromate give brightly coloured precipitates.
Do not attempt to isolate these precipitates because they are carcinogenic.
Prepare these precipitates in the smallest quantities that allow them to be seen.

Experiments
1. Acidify with dilute nitric acid, add ammonia solution, NH3 (aq) ("ammonium hydroxide"), until just alkaline.
Heat to boiling then divide intro 2 parts.
To one part add the solution.
silver nitrate solution.
A crimson red precipitate, soluble in dilute nitric acid indicates chromate.
2Ag+ + CrO42- --> Ag2CrO4 (s)
To the other part add barium chloride solution.
A yellow precipitate soluble in hydrochloric acid confirms chromate (VI) ions
Ba2+ (aq) + CrO42- (aq) --> BaCrO4 (s)
2. Acidify the sodium carbonate extract with dilute sulfuric acid.
Add drops of amyl alcohol then hydrogen peroxide solution.
Shake then leave to stand.
A blue colour in the alcohol confirms chromate.
3. Add excess dilute nitric acid, followed by silver nitrate solution.
Filter off the precipitate.
Add ammonia
to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a crimson red precipitate
forms that indicates chromate in the filtrate.
4. Add lead nitrate (II) solution to a solution of chromate (VI) ions to form bright yellow precipitate lead (II) chromate (VI), the
"chrome yellow paint pigment.
Pb2+ (aq) + CrO42- (aq) --> PbCrO4 (s)
5. If dilute sulfuric acid is added to a solution containing chromate (VI) ions, the orange colour of dichromate (VI) ions appears.
However, this is not a reliable test for chromate (VI) ions because the colour may be caused by an acid-base indicator in the solution.

12.11.5.10 Tests for halides, Cl-, Br-, I-
Mix 1 g of unknown solid with 1 g of MnO2 add concentrated H2SO4 then heat.
Orange-red gas indicates Br-.
Violet layer of gas indicates I-.
Yellow-green gas that turns KI / starch paper blue to indicate Cl-.

12.11.5.11 Tests for hydroxides
Add one drop of sodium hydroxide solution to ten drops of the unknown solution.
1. A white or glassy precipitate indicates Al3+, Bi3+, Cd2+, Mg2+, Mn2+, Pb2+, Zn2+, Sn2+.
2. A green precipitate indicates Fe(OH)2, Ni2+, Cr3+.
3. A brown precipitate indicates Ag+ and Fe(OH)3.
4. A blue precipitate indicates Cu2+ and Co2+.
5. The reaction with Ca2+ forms a slightly soluble white precipitate.
If the reaction forms no precipitate, heat the solution to identify the presence of NH4+ from the odour of ammonia.

12.11.5.12 Tests for iodides
1. Add excess dilute nitric acid, followed by silver nitrate solution.
A yellow precipitate of silver iodide, insoluble in ammonia, indicates the presence of the iodide radical.
Ag+ + I- --> AgI (s)
2. To confirm the iodide radical, heat the solid with manganese dioxide and concentrated sulfuric acid and observe the violet vapour of
iodine.
3. Add 6 M HCl to 3 mL of test solution, then boil, then add 3 mL 0.1 M FeCl3.
Add 1 m L of hexane and shake the solution.
A purple colour of the hexane indicate the presence of I-.

12.11.5.13 Tests for nitrates
1. First test: When the cation is not a salt of Na+, NH4+ or K+, remove it as insoluble carbonate.
Add 10 mL Na2CO3 solution to 1 g of the solid salt, boil, filter and prepare up to 2 mL with deionized water.
Add to 5 drops of unknown solution, 5 drops of water, 5 drops concentrated H2SO4 and Cu foil.
Brown fumes of nitrogen dioxide and a blue-green solution indicate NO3-.
2. Second test: Add to 5 drops of unknown solution in an evaporating basin, 3 drops of concentrated sulfuric acid and a crystal of iron
(II) sulfate.
A purple colour on the crystal indicates NO3-.
3. This test is called the brown ring test.
Add excess of cold dilute sulfuric acid to the unknown solution then add excess freshly prepared iron (II) sulfate solution.
Transfer the solution to a boiling tube to a depth of 2 cm.
Fix the boiling tube in a sloping position then very carefully pour concentrated sulfuric acid down the sloping side of the tube to form a
separate 2 cm layer beneath the solution.
Observe a brown ring at the junction of the acid and unknown solution.
The nitrate and the concentrated sulfuric acid first form nitric acid to be reduced by iron (II) sulfate to nitric oxide.
The nitric oxide reacts with more iron (II) sulfate to form the brown compound NO.2FeSO4.
Carefully shake the boiling tube to spread the brown colour.
The solution becomes warm as the acid and water mix and the brown colour disappears as the unstable brown compound decomposes.
2FeSO4 + 2NaNO3 + 5H2SO4 --> 2NaHSO4 + 3Fe2(SO4)3 + 4H2O + 2NO (g)
NO + 2FeSO4 --> NO.2FeSO4 (brown colour forms)
NO.2FeSO4 --> NO + 2FeSO4 (brown colour disappears)
4. If a bromide or iodide is in the unknown solution, a ring due either to free bromine or to free iodine forms and the iron (II) sulfate is
not part of this reaction.
However, if bromide or iodide is already known to be in the unknown solution, add silver sulfate solution to precipitate the bromide or
iodide as a silver salt and then test the filtrate for the nitrate ion.
5. If a nitrite is in the unknown solution, a diffuse brown ring forms.
To eliminate nitrite, add a concentrated solution of urea, then dilute sulfuric acid and warm until effervescence of nitrogen stops.
Then test for nitrate.
6. Heat a mixture of the original solid with copper and drops of concentrated sulfuric acid.
The nitrate radical reacts with concentrated sulfuric acid to form nitric acid, which reacts with copper to produce brown nitrogen
dioxide gas.
The brown gas indicates the nitrate radical.
Cu + 4HNO3 ---> Cu(NO3)2 + 2H2O + 2NO2 (g).

12.11.5.14 Tests for oxalates
1. Dissolve the unknown substance in water, add excess calcium chloride solution and heat to boiling.
Decant and wash the remaining precipitate of calcium oxalate with warm dilute sulfuric acid.
Add a drops of potassium permanganate solution, which is decolorized.
2KMnO4 + 3H2SO4 + 5H2C2O4 --> K2SO4 + 2MnSO4 + 8H2O + 10CO2
2. Tests for oxalate, antimonate, borate.
Add excess dilute nitric acid, followed by silver nitrate solution.
Filter off the precipitate.
Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a white precipitate forms that indicates
antimonate, borate, or oxalate in the filtrate.

12.11.5.15 Tests for phosphates
1. Add dilute nitric acid and excess ammonium molybdate solution.
Heat but do not boil.
A yellow coloration, with precipitate of ammonium phosphomolybdate on standing (NH4)3PO4.12MoO3, indicates phosphate.
2. Add excess dilute nitric acid, followed by silver nitrate solution.
Filter off the precipitate.
Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a yellow precipitate forms that indicates
phosphate in the filtrate.

12.11.5.16 Tests for sulfates
1. Add to 5 drops of unknown solution 2 drops of hydrochloric acid, heat then add 3 drops of barium chloride solution.
A white precipitate indicates SO42-.
Ba2+ + SO42- --> BaSO4 (s)
2. Add excess dilute hydrochloric acid, and then barium chloride solution.
A white precipitate of barium sulfate shows the presence of the sulfate radical.
3. To confirm the presence of sulfates, heat the unknown with fusion mixture on a charcoal block and test the residue on a wet silver
surface.
A black stain of silver sulfide indicates a sulfide formed by partial reduction of the sulfate.
This test is not applicable if sulfide is in the unknown substance.

12.11.5.17 Tests for sulfides
Add lead acetate solution.
A black precipitate indicates sulfide.

12.11.5.18 Tests for sulfites
Add barium chloride solution.
A white precipitate, soluble in hydrochloric acid, indicates sulfite.

12.14.0 Activity series of metals as reducing agents
The activity series is also called reactivity series or electrochemical series.
Decreasing activity from left to right: potassium, sodium, barium, calcium, magnesium, aluminium, zinc, iron, tin, lead (hydrogen) copper,
mercury, silver, platinum, gold.
Metals above lead in the activity series react with acids with liberate hydrogen gas.
However, nitric acid and concentrated sulfuric acid react with metals above platinum but do not produce hydrogen gas.
Reactions of acids with metals are exothermic and the higher the metal in the activity series, the greater the heat liberated in its reaction
with an acid.
1a = reaction with cold water to give the oxide and hydrogen gas
1b = reaction with hot water to give the oxide and hydrogen gas
1c = reaction with steam to give the oxide and hydrogen gas
2a = reaction with air (when heated form peroxides)
2b = reaction with air (when heated as powders form oxides)
3a = react with dilute hydrochloric acid or sulfuric acid to form hydrogen gas and metal ions and react with concentrated nitric acid or
sulfuric acid to produce metal ions and nitrogen dioxide or sulfur dioxide
3b = react with concentrated nitric acid or sulfuric acid to produce metal ions and nitrogen dioxide or sulfur dioxide
3c = react with aqua regia (concentrated nitric acid and hydrochloric acid)
Table 12 14 0
K 1a 2a 3a Zn 1c 2b 3a .
Hg 2b 3b .
Ba 1a 2a 3a Fe 1c 2b 3a .
Ag .
3b .
Sr 1a 2a 3a Ni 1c 2b 3a .
Pt .
.
3c
Na 1a 2a 3a Sn .
2b 3a .
Au .
.
3c
Ca 1a 2a 3a Pb .
2b .
3b .
.
.
.
Mg 1b 2b 3a H .
.
.

.
.
.
.
Al 1c 2b 3a Cu .
2b .
3b .
.
.
.


12.14.02 Reactions of metals with air or oxygen gas
All elements except Ag, Au and Pt react with air.
K, Na and Ca form peroxides.
The other elements form oxides, when heated as powders.

12.14.03 Reactions of metals with dilute acids
Pb, Cu, Hg, Ag, AU and Pt do not react with dilute HCl or HNO3.
Pt and Au react with aqua regia.
Metals react with dilute acids to form hydrogen gas and the metal ion.

12.14.04 Reactions of metals with concentrated oxidizing acids
Au and Pt do not react with concentrated HNO3 or H2SO4.
Reactions form the metal ions of high oxidation number and sulfur dioxide if H2SO4. Reactions form nitrogen dioxide if HNO3,
e.g. copper has two oxidation numbers, number 1 (Cu+1), and number 2 (Cu2+).

12.14.1 Zinc displaces lead from lead nitrate solution
A metal displaces a metal lower in the activity series from its salt solutions.
The more active metal atoms lose electrons more easily to go into solution as ions.
The less active metal ions attract electrons more easily to leave the solution as metal atoms.
The position of the metal in the activity series represents its relative ease of oxidation, i.e. ease of losing electrons to form ions.
The most active metals replace hydrogen from water.
Metals that replace hydrogen from dilute acids are placed above hydrogen.
Metals that do not replace hydrogen from such acids are placed below hydrogen.
These metals may be oxidized by the oxidizing acids nitric acid and hot concentrated sulfuric acid.
Gold and platinum do not react with the oxidizing acids, but do react with aqua regia (a mixture of concentrated hydrochloric acid and
concentrated nitric acid in ratio 3:1 by volume).

Experiment
Put a piece of granulated zinc in a test-tube containing lead (II) nitrate solution.
The zinc becomes covered with metallic lead solution.
The zinc granule becomes corroded.
Zinc displaces lead from lead salt solutions.

12.14.2 Metals with copper (II) sulfate solution
A metal higher in the activity order is needed to displace copper metal from copper ions solutions.

12.14.2.1 Magnesium, or zinc, with copper (II) sulfate solution
Magnesium or zinc displaces copper that is lower in the activity series from its salt copper (II) sulfate.
Experiment
Use magnesium ribbon or zinc dust in a test-tube of copper (II) sulfate solution.
The reaction can be vigorous with the magnesium.
Copper metal deposits and the blue colour gradually disappear as the copper ion is displaced by the more reactive metal that is higher
in the activity series.
The reaction loses heat.
When the solution is colourless, decant the solution leaving red copper powder at the bottom of the test-tube.
Mg (s) + CuSO4 (aq) --> MgSO4 (aq) + Cu (s)
Mg loses electrons: Mg --> Mg2+ + 2e- (oxidation)
Cu gains electrons: Cu2+ + 2e- --> Cu (reduction).

12.14.2.2 Iron with copper (II) sulfate solution
Clean a large iron nail with emery cloth.
Put it in a test-tube containing copper (II) sulfate solution.
The reaction forms a coating of copper on the iron nail as copper leaves the solution.
The colour of the solution changes from blue to green iron enters the solution as ions.
The iron nail is corroded.
Iron displaces copper from copper salt solutions.

12.14.2.3 Iron and zinc with copper (II) sulfate solution
1. Add 10 g of copper (II) sulfate solution to 50 mL of water in two beakers.
Add shiny iron nails to beaker 1. Add shiny pieces of zinc metal to beaker 2. Leave to stand and after 2 hours note any change in
colour of the solution and any precipitate.
3. Add iron nails to the solution containing the zinc and add shiny pieces of zinc to the solution containing the iron nails.
Notice any further reactions that take place.
CuSO4 + Zn --> ZnSO4 (aq) + Cu (s)
CuSO4 + Fe --> FeSO4 (aq) + Cu (s)
FeSO4 + Zn --> ZnSO4 (aq) + Fe (s)
ZnSO4 + Fe --> no reaction

12.14.2.4 Zinc with lead nitrate solution, iron with copper (II) sulfate solution
Clean a small strip of zinc and an iron nail with emery cloth.
Make separate solutions of lead (II) nitrate and copper (II) sulfate.
Put the zinc in the lead nitrate solution and put the iron in the copper (II) sulfate solution.
After a few minutes remove the metal strips and observe the appearance of each.
Note a copper coating on the iron nail.
Note the crystals of metallic lead on the zinc.
After leaving the metals in the solution for a longer time you will notice that the original metal has corroded.
The copper (II) sulfate solution the blue colour will be gradually replaced by a dirty green colour.

12.14.2.5 Zinc with copper in sulfuric acid
1. Hold a clean strip of zinc in dilute sulfuric acid.
If the zinc is very pure, few bubbles of hydrogen gas will form from its surface.
Remove the zinc and hold a strip of copper in the acid.
No gas forms.
2. Put both metal strips in the acid so that an edge of the zinc is in contact with the copper.
Copious bubbles of gas are given off from the copper plate and practically none from the zinc.

12.14.2.6 Activity of metals and tendency to form ions
See diagram 33.3.1: Simple cell
Dip pairs of strips of zinc, copper, iron, lead and magnesium, into sodium chloride solution.
Connect the metals to a 0-3 V voltmeter, or galvanometer, and note the direction of current flow.
The more reactive metal forms the negative pole and so electrons flow from it.
Test zinc with copper, lead, iron and magnesium.
Test copper with lead, magnesium and iron.
Test lead with iron and magnesium.
Test iron with magnesium.
For each pair of metals, note which metal forms the positive terminal, which metal forms the negative terminal and the voltage for each
combination.
The more active metal becomes the negative pole of the cell from which electrons flow.
Metals high in the activity series, e.g. zinc, tend to release electrons to form ions.
Metals low in the activity series, e.g. copper, do not readily form ions and these
ions readily form metal atoms.
Zn (s) --> Zn2+ (aq) + 2e-
Cu2+ (aq) + 2e- --> Cu (s)
The metals in order of activity are (most active) Mg, Zn, Fe, Pb, Cu (least active).

12.15.0 Reactions of metals with water
All metallic elements except Sn, Pb, Cu, Hg, Ag, Au and Pt react with cold water or hot water or steam.
1. Metals act as reducing agents in displacing hydrogen from water.
2. K, Ba and Na displace hydrogen from cold water.
3. K reacts violently and forms hydrogen gas that catches alight and burns with a pink flame.
4. Ca reacts slowly and the solution turns milky because of the formation of calcium hydroxide.
5. Mg reacts slowly with cold water and quickly with hot water.
6. Al, Zn, Fe and Ni react with steam to produce oxide and hydrogen gas.
7. Sn, Pb, Cu, Hg, Ag, Au and Pt do not react with water.

12.15.1 Reactions of metals with water, Cu, Zn, Fe, Mg, Al
See diagram 12.15.1B: Metals with water
If metals are not pure, some reactions may be caused by the impurity.
1. Boil deionized water for 5 minutes to remove dissolved air leave it to cool then pour into test-tubes.
Put in the test-tubes pieces of freshly polished: copper, zinc, iron, magnesium, aluminium.
Leave for 10 minutes.
Observe any changes in the metal or water.
When you see bubbles on the metals, put the metal under an inverted test-tube of water and leave for two days to collect the gas.
Test the collected gas with litmus paper, lime water, and a lighted splint.
The bubbles are hydrogen gas.
Calcium reacts slowly then sinks.
Magnesium reacts very slowly in cold water, but reacts vigorously in steam.
Ca (s) + 2H2O (l) --> Ca(OH)2 (aq) + H2 (g)
2. Use test-tubes containing deionized water or demineralized water.
Boil the water then leave to cool.
Put small pieces of freshly polished copper, zinc, iron, magnesium and aluminium in the boiled water and leave for 10 minutes.
Note any change in the metal or water.
Boil the water + metals for 5 minutes.
Note any changes.
If you see any bubbles on the metals, put the metal in a small basin of water and invert a test-tube of water over it.
Leave for a few days to see whether larger quantities of the gas in the bubbles may be collected.
Test any gas collected with litmus, lime water and a lighted splinter.
The purpose of boiling the water before placing the metal into it is to remove any dissolved air that might react with the metal.

12.15.3 Reactions of metals with steam
See diagram 12.15.3: Metals with steam
Put wet cotton wool or glass wool at the bottom of a test-tube.
Put another small piece of cotton wool or glass wool half way up the tube.
Clean and polish a piece of magnesium ribbon and put it on the upper plug.
Insert a one-hole stopper fitted with a glass tube.
Use a Bunsen burner to heat the lower cotton wool or glass wool until steam comes off.
Use a second Bunsen burner to heat the magnesium ribbon.
Tests for hydrogen gas with a lighted splint.
Repeat the experiment with cleaned aluminium, copper wire,
and iron wire.
When heated in steam, magnesium, aluminium and iron react, but not copper.
Mg (s) + H2O (g) --> MgO (s) + H2 (g)
Al (s) + H2O (g) --> Al2O3 (s) + H2 (s)
Fe (s) + H2O (g) --> Fe2O3 (s) + H2 (s).

12.15.5.0 Reactions of non-metals with water
1. Shake small quantities of sulfur, carbon and iodine separately with water.
Are there any indications of solution or chemical reaction? Filter each mixture.
Test a little of the filtrate from the mixture containing iodine by pouring a little of it on to a piece of starch.
Evaporate each filtrate to dryness and residue remains.
The slight blue colour with iodine shows that iodine is slightly soluble in water.
Sulfur and carbon are insoluble in water.
2. Pass some chlorine into water in a test-tube and shake the test -tube.
Drop small pieces of red and blue litmus paper into the chorine and water.
The blue litmus paper turns red then white as the chlorine and water bleaches it.
Chlorine dissolves in water to produce hydrochloric acid and hypochlorous acid.

12.15.5.1 Heat carbon with steam, water gas
Carbon is insoluble in water, but carbon heated to 1000oC reacts with steam to produce the fuel "water gas" that can be added to coal
gas.
carbon (s) + water (g) --> carbon monoxide (g) + hydrogen (g)
C (s) + H2O (g) <--> CO (g) + H2 (g) water gas

12.16.0 Carbonates
K, Na, Ca, Mg, Zn, and Pb carbonates are white.
Fe carbonate is brown.
Cu carbonate is blue-green.
Only K and Na carbonates are soluble in water and are not decomposed by heat.
Ammonium carbonate is a white powder.

12.16.1 Carbon dioxide through calcium carbonate suspension
Pass carbon dioxide through a suspension of calcium carbonate then boil the mixture.
The calcium carbonate suspension disappears because the reaction forms soluble calcium hydrogen carbonate.
Note that the reaction is reversible.
Calcium hydrogen carbonate easily decomposes when heated.
CaCO3 (s) + CO2 (g) <--> CaHCO3 (aq)

12.16.1.1 Carbon dioxide with calcium hydroxide solution
Whitewash is a suspension of calcium hydroxide in water used as marker on grass and a cheap paint.
Carbon dioxide in the air slowly changes the slightly soluble calcium hydroxide to insoluble calcium carbonate.
Ca(OH)2 (aq) + CO2 (g) --> CaCO3 (s) + H2O (g)
Whiting is pulverised fine chalk used for whitewashing, cleaning, filling cracks.

Experiment
Add water to cool freshly made calcium oxide (quicklime) in an evaporating basis to form calcium hydroxide.
The reaction is exothermic and forms steam.
CaO (s) + H2O (l) --> Ca(OH)2 (s)
Mix 1 mL of the solid calcium hydroxide with 10 mL of water.
Test this with an indicator to show that it is a base.
Leave the solution to stand.
Decant the clear liquid that is lime water, a diluted solution of calcium hydroxide
Tests for carbon dioxide
See diagram 9.154: Lime water test for carbon dioxide in the breath
Pass carbon dioxide through the clear liquid, the lime water.
The reaction forms a white precipitate of calcium carbonate.
Ca(OH)2 (aq) + CO2 (g) --> CaCO3 (s) + H2O (l)
Continue to pass carbon dioxide through the solution.
Soluble calcium hydrogen carbonate forms and the solution becomes clear again.
CaCO3 (s) + CO2 (g) + H2O (l) --> Ca(HCO3)2 (aq)
CO2 (g) + H2O (l) --> H2CO3 (aq) carbonic acid
H2CO3 (aq) + 2OH- (aq) --> CO32- (aq) + 4H2O
Ca2+ (aq) + CO32- (aq) --> CaCO3 (s)
CaCO3 (s) + H2CO3 (aq) --> Ca2+ (aq) + 2HCO3- (aq) bicarbonate ion
This reaction also occurs when the wet cement mortar used in bricklaying sets hard to hold the bricks together.
The water evaporates leaving the solid calcium carbonate.

12.16.2 Prepare sodium hydrogen carbonate
Pass carbon dioxide through sodium carbonate solution to form sodium hydrogen carbonate sodium bicarbonate.
If you heat dry sodium hydrogen carbonate the reverse reaction occurs.
Na2CO3 (aq) + CO2 (g) + H2O (l) --> 2(NaHCO3) (aq)

12.16.3 Heat carbonates of Cu, Mg, Na, Pb and Zn
Carbonates, except Na2CO3 and K2CO3, decompose on heating to form carbon dioxide and the oxide.
Experiment
Heat powdered calcium carbonate with a strong burner.
The calcium carbonate decomposes to form calcium oxide (quicklime) and carbon dioxide.
CaCO3 (s) --> CaO (s) + CO2 (g)
Heat different carbonates in a test-tube, e.g. carbonates of Cu, Mg, Na, Pb and Zn.
Test the gases that form with: moist litmus paper, a drop of lime water on a glass rod, a lighted splint.
The reaction forms carbon dioxide.
PbCO3 (s) --> PbO (s) + CO2 (g)

12.16.3.1 Heat ammonium carbonate (smelling salts)
Formerly this chemical was used to revive young ladies who had fainted by heating the container by hand to give off ammonia.
To make smelling salts, coarsely powdered ammonium carbonate was moistened with a mixture of oil of orris root, oil of lavender
flowers, extract of violet, and ammonia water.
Ammonium carbonate is a white powder fairly soluble in water forming a weak alkali.
Experiment
Heat ammonium carbonate.
Heat ammonium carbonate in a dry test-tube held sloping downwards.
Observe the steam and condensed water on the cooler rim.
Tests for ammonia gas by smell and hold damp red litmus at the mouth of the test-tube.
It turns blue.
Ammonium carbonate decomposes to form three gases or vapours: 1. steam, 2. ammonia, 3. carbon dioxide, leaving no residues.
Note the smell of the ammonia given off.

12.16.3.4 Ammonium carbonate solution precipitates metal carbonates
Add ammonium carbonate solution to solutions of copper (II) sulfate, iron (II) sulfate, magnesium sulfate, zinc sulfate and lime water.
Note the colours of the precipitated metal carbonates.

12.16.3.5 Smelling salts, ammonium carbonate
Ammonium carbonate
Ammonium carbonate is sold in small bottles as "smelling salts" to revive a fainting person by sniffing the ammonia gas formed from the
ammonium carbonate in the bottle.
Common names: bakers' ammonia.
Smelling salts is ammonium carbonate and scent.
The warmth of hand causes ammonium carbonate to break down to form ammonia, carbon dioxide and water.
The ammonia revives people who have fainted.
Sal volatile, solution in alcohol of ammonium carbonate, double salt: [ammonium hydrogen carbonate, ammonium
aminomethanoate (carbamate)], [NH4HCO3.NH2COONH4], in a small bottle, to be sniffed as a restorative for faintness.

12.16.4 Heat sodium hydrogen carbonate
This reaction is used in baking powder.
1. Heat a hydrogen carbonate in a test-tube.
Test gases that form with: moist litmus paper, a drop of lime water on a glass rod, a lighted splint.
The reaction forms carbon dioxide.
2. Heat sodium hydrogen carbonate (baking soda).
Solid sodium hydrogen carbonate begins to decompose at 100oC and is completely decomposed at 200oC.
The solution in water starts to decompose at room temperature.
2NaHCO3 (s) --> CO2 (g) + H2O (g) + Na2CO3 (s).

12.16.6 Prepare volcanos with baking soda
1. Make a heap of sand to represent the volcano and push a test-tube or long thin jar down into the heap of sand.
Put baking soda or baking powder, food colouring, detergent and even glitter into the glass container.
Carefully pour vinegar into the glass container.
BE CAREFUL! DO NOT LOOK DOWN INTO THE GLASS CONTAINER!
2. Mix 2 parts pf flour (plain white flour), 2 parts of sodium chloride, 0.5 parts of cooking oil, 2 parts of water and mix thoroughly to
form a dough.
Mould the dough around the bottle without putting any inside or covering the mouth of the bottle.
Stand a glass bottle on a baking pan.
Fill the bottle with warm water.
Add food colouring.
Add 5 drops of liquid detergent.
Add 0.25 parts baking soda (sodium hydrogen carbonate).
Slowly add vinegar to the bottle.
BE CAREFUL! DO NOT LOOK DOWN INTO THE BOTTLE!
3. Create a simulated underground explosion.
Pour a tablespoon (15 -20 mL) of baking soda into a balloon.
Add 150 mL of vinegar and quickly tie a knot in the neck of the balloon.
The balloon swells and may burst.
CH3COOH + NaHCO3 --> CH3COONa + H2CO3
Acetic acid (vinegar) + sodium hydrogen carbonate --> sodium acetate + carbonic acid
H2CO3 --> H2O + CO2
carbonic acid --> water + carbon dioxide

12.16.7 Prepare sodium carbonate, Solvay process
Soda ash is used to produce glass, detergents for metal refining, and for water purification.
In nature sodium carbonate decahydrate can be formed by the action of concentrated salt solutions on limestone.
2NaCl (aq) + CaCO3 (s) --> Na2CO3.10H2O (s) + CaCl2 (aq)
In the laboratory sodium carbonate solution precipitates calcium carbonate from an aqueous solution of calcium chloride, but in nature
the reaction may be very slowly reversed in evaporating deposits because of the very high concentration of sodium chloride.
In the Solvay process, soda ash is produced by the reaction:
CaCO3 + 2NaCl --> Na2CO3 + CaCl2
The natural direction of this reaction is backwards but the reaction can be moved forward by various reactions including forcing carbon
dioxide is forced under pressure into a concentrated cold brine solution saturated with ammonia adding ammonium ions and bicarbonate
ions to the sodium and chloride ions already present.
NH3 (g) + CO2 (g) + NaCl (aq) + H2O (l) --> NaHCO3 (s) + NH4Cl (aq)
The least soluble combination of ions is sodium bicarbonate, which precipitates.
This anhydrous product is called light soda.
The liquor is fed to the ammonia recovery plant where it is liberated with lime to leave calcium chloride.
Lime kilns produce both lime and carbon dioxide for the process.
Sodium bicarbonate is decomposed to sodium carbonate and the carbon dioxide released is recycled.
The ammonia is regenerated and recycled by decomposing the ammonium chloride formed.
Sodium carbonate solid is
hydrated to monohydrate crystals for easier handling.
Washing soda is produced by recrystallization, using the monohydrate from water to form the decahydrate.
Washing soda, Na2CO3.10H2O, will dehydrate spontaneously by efflorescence back to the monohydrate under dry conditions.
Some of the waste concentrated calcium chloride liquor is used as a drilling mud for the oil industry and as an ice and snow melting salt
in cold climates.

12.16.8 Prepare sodium carbonate, LeBlanc process
This process was discovered by Nicolas LeBlanc in 1791.
It was an important source of sodium carbonate, but is no longer used.
Step 1. Sea salt boiled in sulfuric acid
2NaCl + H2SO4 --> Na2SO4 + 2HCl (HCl discharged into the atmosphere to cause air pollution)
Step 2. Sodium carbonate mixed with limestone and coal, then burnt to form "black ash", sodium carbonate dissolved out of black ash
with water
Na2SO4 + 2C --> Na2S + CO2
(carbon in coal oxidized to carbon dioxide, sulfate reduced sulfide)
Na2S + CaCO3 --> Na2CO3 + CaS
(calcium and sodium swap ligands to form thermodynamically-favourable sodium carbonate and calcium sulfide)
Na2SO4 + CaCO3 + 2C --> Na2CO3 + 2CO2 + CaS
sodium sulfate + calcium carbonate --> sodium carbonate + carbon dioxide + calcium sulfide
The sodium carbonate was leached out of the products of the reactions to leave "galligu", a mixture of calcium sulfide, unburnt coal, coal
ash and sodium sulfide, which caused widespread land pollution.

12.17.0 Oxides, acidic, basic, amphoteric, neutral and mixed oxides
Oxides are formed by direct combination of elements, addition of oxygen by oxidation, decomposition by heat of carbonates,
hydroxides and some nitrates.
Oxides can be reduced back to the element with reducing agents, e.g. hydrogen, carbon, carbon monoxide.
Metal oxides act as bases.
Non-metal oxides act as acids.
Oxygen gas reacts with metals to form basic oxides.
Oxygen gas reacts with non-metals to form acidic oxides.
Metal oxides on the left of the period form alkaline solutions in water.
Non-metal oxides on the right of the period form acidic solutions in water.
Oxides of metals (semi-metals) in the middle of the period, e.g. SiO2, show amphoteric behaviour.
Elements lower in a group have more basic oxides.

1. Acidic oxides are oxides of non-metals that react with water to form acids or react with bases and alkalis to form salts + water, at
room temperature are usually gases.
CO2 --> carbonic acid, H2CO3
CO2 + H2O --> H2CO3
SO2 --> "sulfurous acid", Sulfurous does not exist in solution but as a vapour
SO2 + H2O <--> H+ + HSO3- (hydrogen sulfide, hydrogen bisulfide)
SO2 (g) + H2O (l) --> H2SO4 (aq) --> H+ (aq) + HSO3- (aq)
A solution of SO2 in water is commonly called "sulfurous acid".
SO2 + NaOH --> NaHSO3 (sodium bisulfite, sodium hydrogen sulfite)
SO3 --> sulfuric acid, H2SO4
N2O3 --> nitrous acid, HNO2
N2O5 --> nitric acid, HNO3
P2O3 (P4O6) --> phosphoric acid, H3PO4
B2O3, boron oxide --> boric acid, H3BO3
SiO2 does not react with water, but reacts with molten sodium hydroxide at high temperature and pressure and is an important reaction
in the geological origin of silicates.
SiO2 + 2NaOH --> H2O + Na2SiO3, sodium silicate

2. Basic oxides are oxides of metals that react with acids to form a salt and water only, do not react with bases, most basic oxides are
insoluble in water but some dissolve to form alkaline solutions, i.e. Na2O, K2O and CaO.
The oxides of feebly acidic cations react exothermically with water to form the hydroxide.
Na2O (s) + H2O (l) --> 2NaOH (aq)
K2O (s) + H2O (l) --> 2KOH (aq)
CaO (s) + H2O (l) <--> Ca(OH)2 (aq) "slaked" lime"
MgO (s) + H2O (l) --> Mg(OH)2 (s) <--> Mg2+ + 2OH-
(a slight reaction, nothing appears to happen, but pH changes)
Cu2O, CuO, FeO, Fe2O3, PbO do not react with water, but some may react with steam.
PbO + 2HNO3 --> Pb(NO3)2 + H2O.

3. Amphoteric oxides behave as acidic oxides and basic oxides, e.g. Al2O3, PbO, SnO, ZnO react with both acids and bases to form
salt and water.
ZnO + 2HCl --> ZnCl2 + H2O
ZnO + 2NaOH + H2O --> Na2[Zn(OH)4] sodium zincate
Zn(OH)2 + 4HCl -> ZnCl2 + 2H2O 2NaOH + Zn(OH)2 -> Na2[Zn(OH)4] sodium zincate.

4. Neutral oxides, e.g. carbon monoxide CO, dinitrogen oxide (nitrous oxide) N2O, nitrogen monoxide (nitric oxide) NO and water
H2O have neutral pH.
Hydrogen peroxide is an example of a higher oxide that forms oxygen gas when heated.

5. Mixed oxides contain more than one oxide
Examples:
The anticorrosive pigment red lead oxide, dilead (II) lead (IV) oxide, Pb3O4(2PbO.PbO2)
The iron ore mineral magnetite, iron (II) iron (III) oxide, Fe3O4(FeO.Fe2O3).

6. Hydroxides refers to "hydrated oxides", OH.

12.17.1 Properties of oxides
All elements except the noble gases (inert gases), form oxides.
1. Different oxides, e.g. magnesium oxide, calcium oxide, aluminium oxide, carbon dioxide, sulfur dioxide, and nitrogen dioxide.
2. Describe the appearance.
3. Describe the odour.
BE CAREFUL! DO NOT INHALE GASES DIRECTLY FROM THE TEST-TUBE!
Fan the gas towards the nose with the hand and sniff cautiously.
If no odour is detected, move closer and try again.
4. Add different oxides to water and shake.
Note the relative solubility.
5. Test the acidity where solution has occurred.
6. Add drops of dilute sulfuric acid to each oxide.
Note any reactions.
Heat if no reaction occurs.
7. Add drops of sodium hydroxide solution to each oxide.
Heat if no reaction occurs.
8. List the oxides in order of increasing acidic character.

12.17.1.1 Oxides and the periodic table
All elements except inert gases form oxides.
The oxides of metals in Group II were thought to be "like earth" and they form alkaline solutions, so the metals were called "alkaline
earth" metals.
Their oxides and hydroxides react with acids but not with alkalis.
The oxide ion reacts with water to form the hydroxide (hydroxyl) ion.
O2- + H2O --> 2OH-
With acids, the oxide ion reacts with the hydronium ion
O2- + 2H3O+ --> 3H2O
The metallic properties become less to the right of the periodic table, e.g. aluminium oxide is insoluble in water, and reacts with both
acids and alkalis to form water and salts, so is called an amphoteric oxide.
Farther to the right of the periodic table, the elements are non-metals.
They may react with water to form acid solutions.
Example 1. Carbon dioxide dissolves in water to form carbonic acid
CO2 (aq) + H2O (l) --> H2CO3 (aq)
Example 2. Phosphorus pentoxide (phosphorus (V) oxide) reacts violently with water to form phosphoric acid.
P4O10 (s) --> H2O (l) + 4H3PO4 (aq)

12.17.2 Copper (II) oxide (copper oxide), basic oxide (metal oxide)
A basic oxide reacts with hydrogen ion to give water and a salt
CuO (s) + 2H+ (aq) --> H2O (l) + Cu2+ (aq)
copper (II) oxide + hydrogen ion --> water + copper ion
CuO (s) + 2HCl (aq) --> CuCl2 + H2O (l)
copper oxide + hydrochloric acid --> copper (II) chloride + water
Basic oxides do not usually react with alkalis.
Experiment
Put copper (II) oxide, calcium oxide, magnesium oxide and iron oxide in separate test-tubes.
Add drops of alkali solution to each.
Heat the mixture.

12.17.2.1 Heat zinc with copper (II) oxide
Weigh 2 g (0.025 mol) copper (II) oxide powder and 1.6 g (0.025 mol) zinc powder, zinc dust.
Mix the powders to a uniform grey colour.
Pour the mixture in the shape of a horizontal cylinder on a coffee tin lid.
Heat one end of the mixture cylinder with a Bunsen burner until the mixture begins to glow.
Stop heating and let the glow move along the cylinder of powder to the end leaving a white-grey mixture.
Heat the coffee tin lid over a Bunsen burner to show that the white powder, zinc oxide, is yellow when hot
and white when cool (because of change in the crystal structure of zinc oxide).
Put the cooled residue in a beaker and add dilute hydrochloric acid to dissolve the zinc oxide and any remaining copper oxide and zinc,
leaving red-brown copper.
Heat the red-brown powder with concentrated nitric acid to give a blue solution of copper nitrate.
Repeat the experiment using coarse magnesium powder instead of zinc powder.

12.17.2.2 Heat metals with oxides of another metal
12.1.5 Thermite reaction, thermit welding
Competition for oxygen
Heat 10 mL of the following mixtures in a crucible.
Put the crucible on a pipe clay triangle on a tripod.
Heat the mixture with a Bunsen burner, slowly then strongly.
Use tongs to remove the crucible from the tripod and leave the mixture to cool.
Examine the contents of the crucible for evidence of a chemical change.
1. Lead oxide with iron filings --> iron oxide + lead
2. Magnesium oxide with iron filings --> no reaction
3. Lead oxide with zinc dust --> zinc oxide + lead
4. Iron oxide with zinc dust --> zinc oxide + iron
The metal are showing competition for oxygen in the sequence of activity of the reactivity series.

12.17.3 Carbon dioxide, acidic oxides (non-metal oxides)
Acidic oxides dissolve in water to form an acid
CO2 (aq) + H2O (l) <--> H2CO3 (aq)
carbon dioxide + water <--> carbonic acid, that dissociates:
H2CO3- <--> CO32- + H+
carbonic acid <--> carbonate ion + hydrogen ion

12.17.3.1 Carbon dioxide with sodium hydroxide solution
Alkalis react with acidic oxides to form salt and water.
Pass carbon dioxide through sodium hydroxide solution.
Note the reduction in the size of the bubbles, which shows that a reaction with carbon dioxide probably occurs.
Stop the flow of carbon dioxide.
Add drops of dilute hydrochloric acid.
Test gases that form from the reaction with: moist litmus paper, a lighted splint, .
The gas is carbon dioxide.
NaOH (aq) + CO2 (g) --> H2O (l) + Na2CO3 (aq)
Na2CO3 (aq) + HCl (aq) --> NaCl (aq) + CO2 (g)

12.17.3.2 Carbon dioxide with barium hydroxide solution
Ionization of barium hydroxide, K2 = 1.4 × 10-1
Ba(OH)2 <--> BaOH+ + OH-
BaOH+ <--> Ba2+ + OH-
Pass carbon dioxide through barium hydroxide solution.
The reaction forms a white precipitate.
Filter off the precipitate.
Add dilute hydrochloric acid to the precipitate.
Test the gas that forms with a lighted splint and moist litmus paper.
The gas is carbon dioxide.
Ba(OH)2 (aq) + CO2 (g) --> BaCO3 (s) + H2