Colour, Texture and Solubility:

In general, compounds of transition metals and solutions containing transition metal ions have characteristic colours.

Group I, II, and III metals form white salts and colourless metal ions in solution.

Most metals in the laboratory, with the exception of copper, are usually silvery grey or grey with a shiny lustre.

Ionic compounds which are soluble in water form solutions. The solubility of a sample is usually tested first in cold water. 

Solid/SaltSolubility in

WaterApplication in QA

Nitrates All nitrates are soluble.

This means that any ion, when combined with the nitrate ion in solution, will not precipitate.

Sodium, potassium and ammonium salts

All sodium, potassium and ammonium salts are soluble.

This means that any ion, when combined with the sodium, potassium and ammonium ion in solution, will not precipitate.

Sulphates All sulphates are soluble except for BaSO4,

When a solution containing barium, calcium, or lead (II) ions

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CaSO4 andPbSO4.

are mixed with sulphate ions, a precipitate will be formed.

ChloridesAll chlorides are soluble except for AgCl, PbCl2.

When solutions containing silver or lead(II) ions are mixed with sulphate ions, a precipitate will be formed.

Carbonates

All carbonates are insoluble except for Na2CO3, K2CO3 and (NH4)2CO3.

Any cation except NH4

+and Group I cations e.g. Na+, K+, and will combine with carbonate ions in solution to form a precipitate.

Oxides and Hydroxides

All oxides and hydroxides are insoluble except for Na2O, K2O, NaOH, KOH andCa(OH)2 [slightly soluble].

Most cations except NH4

+ and Group I cations e.g. Na+, K+, and will combine with hydroxide ions in solution to form precipitates.

Colour & Texture of Solid

Inference

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Colourless/white crystals

Gp I, II, III metal salts; no transition metal ions present. Ammonium

compounds e.g. NH4Cl, NH4NO3

Blue crystals

CuSO4.5H2O (hydrated) as shown.

Cu(NO3)2

Green crystalsCuCl2

Pale green crystals

Hydrated iron(II) salts e.g. FeSO4.7H2O

Brown crystalsFe(III) salts e.g. hydrated

FeCl3

Orange crystalsK2Cr2O7

Purple crystalsKMnO4 (purple lustre)

Black crystalsIodine crystals

White powderCaCO3, ZnCO3, ZnO (when cold); Anhydrous CuSO4

Yellow powder

Lead(II) oxide, PbO as shown.

ZnO (yellow when hot and white when cold),

AgI (pale yellow ppt),

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PbI2 (bright yellow ppt).

Bright orange powder

Dilead(II) lead(IV) oxide, Pb3O4

Green powder CuCO3

Reddish brown solid or powder

Cu2O, copper(I) oxide (reddish brown) as shown.

Iron(III) oxide, Fe2O3, Fe(OH)3

PbO2 (dark brown)

Black powderCuO, FeO, MnO2, carbon

powder

Colour of Common Metals

Identity of

metal

Zinc granules

Magnesium ribbon

Copper foil

Colour grey silver grey reddish brown

Identity of

metal

Aluminium powder

Aluminium foil Iron filings

Colour silver silver grey

Colours of Solutions

Colour of Inference

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Unknown solution

Colourless

GpI, II & III cations; most acids and alkalis;hydrogen

peroxide, limewater

BlueCu2+ ions in solution

Pale greenFe2+ ions in solution (pale

green), Cr3+, Ni

GreenCr3+, Ni2+, CuCl2 solution

Yellow/yellow brown

Fe3+ ions in solution

K2CrO4 [potassium chromate(VI)]

Brown Iodine solution

Dark blueComplex ion of Cu2+ :

Cu(NH3)42+

PurpleKMnO4 (potassium

permanganate)

OrangeK2Cr2O7 (potassium

dichromate) in acidic medium

Precipitation Reactions

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Precipitation is a chemical reaction where a solid separates from the solution by combination of ions. The solid formed is called a precipitate.

When two salt solutions are mixed, new salt combinations become possible. The positive metal ion in one salt solution combines with the negative ion in the other and vice versa. If one of these salts is insoluble, then it will precipitate from the solution.

The general formula for a precipitation reaction is:

AB (aq) + XY (aq) AY (s) + XB (aq)

For example, when we mix together potassium iodide solution and lead (II) nitrate solution, a yellow precipitate of lead(II) iodide forms.

Chemical equation: Pb(NO3)2 (aq) + 2KI (aq) PbI2(s) + 2KNO3 (aq)

Ionic equation: Pb2+ (aq) + 2I- (aq) PbI2 (s)

Precipitation Explained

Before mixing After mixing

Ions in Lead(II) nitrate solution:Pb2+, NO3

-

Ions in potassium iodide solution:K+, I-

1. lead(II) ions combine with iodide ions to form lead(II) iodide, which is insoluble.

Pb2+ + 2I-         PbI2 (s)

This is the precipitation reaction.

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2. potassium ions and nitrate ion remain as free, mobile ions in the solution.

K+ (aq) and NO3

- (aq)

Potassium and nitrate ions are called spectator ions because they do not take part in the precipitation reaction.

Thermal Decomposition

Thermal decomposition is a chemical reaction where a single compound breaks up into two or more simpler compounds or elements when heated. It is also an endothermic reaction as heat is required to break chemical bonds in the compound undergoing decomposition. The decomposition reaction is irreversible.

Example 1: Copper (II) carbonate decomposes on heating to form copper(II) oxide and carbon dioxide:

CuCO3 (s) CuO (s) + CO2 (g)

Example 2: Zinc nitrate decomposes on heating to give zinc oxide, nitrogen dioxide and oxygen:

Zn(NO3)2 (s) 2ZnO (s) + 4NO2 (g) + O2(g)

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Summary of Reactivity of Metals and Action of Heat on Their CompoundsGenerally compounds of very reactive metals such as sodium and potassium are more stable to heat than the metals lower down in the reactivity series of metals.

Reactivity Series for

Metal

Action of heat on

hydroxides

Action of heat on

carbonates

Action of heat on nitrates

Most reactive

K

Na

Ca

Mg

Al

Zn

Fe

Pb

Cu

Least reactive

Stable, does not

decompose on heating

Stable, does not

decompose on heating

Decompose to metal nitrite + oxygen

Decompose to metal oxide + steam

Decompose to metal oxide + carbon dioxide

Decompose to metal oxide +

oxygen + nitrogen dioxide

The possible deductions made from the colour changes and colour of the residue after heating

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Observations Possible Cations

white residue orwhite deposits

Compounds of Gp I, II and III metals.Ammonium salts.Residue is likely to be oxides of Gp I, Gp II and Gp III metals formed.

Black powdery residue

Copper(II) compounds, iron(II) compounds.Residue is likely to be CuO, FeO

Yellow residue when hot,white residue when cold.

Zinc compounds; residue is zinc oxide.

Solid changes colour from green to black powder residue.

Solid is Copper(II) carbonate.CuO residue formed.

Black residue from charring, usually with burning odour

Organic compound present; residue is carbon.

Possible Deductions Based On The Gas Evolved During Heating

Observations Possible Deductions

Carbon dioxide produced.

Carbonate is present.But not Na2CO3 or K2CO3 because they are stable to heat.Organic compounds.

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Ammonia produced.

Ammonium salt is present.

Only Oxygen is produced.

NaNO3 or KNO3 is present. Oxidising agents.

Oxygen and Nitrogen dioxide are produced.

All other metal nitrates e.g. Ca(NO3)2 , Zn(NO3)2 , Pb(NO3)2 , Cu(NO3)2.

Sulphur dioxide All sulphites SO32-, eg. iron(II) sulphate

Colourless droplets of liquid condensed on cool walls of test tube.

Water droplets are produced. Water of crystallisation from hydrated crystals e.g. CuSO4.5H2O, FeSO4.7H2O, Na2CO3.10H2O or water from thermal decomposition of metal hydroxides

Examples:1) Thermal decomposition of Copper(II) carbonate: Before heating, copper(II) carbonate is a green powder. When copper(II) carbonate is heated, it undergoes thermal decomposition to produce: i) a black powder, copper(II) oxide, ii) a colourless gas, carbon dioxide, which turns limewater milky.

CuCO3 (s)          CuO (s) + CO2 (g)

2) Thermal decomposition of zinc nitrate: When zinc nitrate crystals are heated, they melt to form a colourless liquid. Stronger heating is required to

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decompose zinc nitrate. On stronger heating, oxygen gas which relights a glowing splint is evolved. Further heating produces a pungent brown gas, nitrogen dioxide, which turns moist blue litmus red. Complete decomposition produces a yellow residue of ZnO which turns white when cooled.

2Zn(NO3)2 (s) 2ZnO (s) + 4NO2(g) + O2 (g)

3) Thermal decomposition of Pb3O4: When dilead(II) lead(IV) powder are heated, they melt to form a dark brown liquid. On stronger heating, oxygen gas which relights a glowing splint is evolved. Complete decomposition produces a yellow residue of lead(II) oxide.

2Pb3O4 (s) 6PbO (s) + O2 (g)

Acid-base Reactions

An acid is a substance that produces hydrogen ions when dissolved in water. Eg. hydrochloric acid, nitric acid and sulphuric acid

A base is a substance that reacts with acids to produce a salt and water. Most insoluble metal oxides and metal hydroxides are bases.

An alkali is a substance that dissolves in water to produce hydroxide ions. An alkali is usually a soluble base. Eg. sodium hydroxide solution, potassium hydroxide solution and aqueous ammonia.

Neutralisation Reaction

In a neutralisation reaction, an acid reacts with an alkali to produce a salt and water. For example:

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HCl (aq) + NaOH (aq)          NaCl (aq) + H2O (l)

Neutralisation Reaction Explained

The chemistry of a neutralisation reaction is essentially the chemistry of hydrogen and hydroxide ions in solution. The ionic equation for neutralisation:

H+ (aq) + OH- (aq)          H2O (l)

Before mixing After mixing

Hydrochloric acid dissociates completely in water to form a solution of hydrogen ions and chloride ions:

HCl (aq)      H+ (aq) + Cl-

(aq)

Sodium hydroxide also dissociates completely in water to form a solution of hydroxide and sodium ions.

NaOH (aq)      OH-

(aq) +Na+ (aq)

1. Hydrogen ions combine with hydroxide ions to form water molecules in the solution.

H+(aq) + OH-

(aq) H2O(l)This is the neutralisation reaction.

2. Sodium ions and chloride ions remain as free, mobile ions in the solution.The sodium and chloride ions are called spectator ions because they do not take part in the neutralisation reaction.

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In general:

Metals below hydrogen in the reactivity series do not react with dilute acids.Metals above hydrogen in the series react to produce hydrogen gas.

The higher the metal in the reactivity series, the more vigorous the reaction.

1)Metal + Acid  Salt + Hydrogen

2)Metal oxide + Acid  Salt + Water

3)Metal hydroxide + Acid  Salt + Water

4)Carbonate + Acid  Salt + Carbon dioxide + Water

Summary of making salts:

Redox Reactions

A redox reaction is a chemical reaction whereby both reduction and oxidation occur simultaneously. Reduction and oxidation always occur together.

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Is the salt soluble?

Use precipitation method.Mix two solutions, one containing the correct positive ion and the other the correct negative ion.

Is it a sodium potassium or ammonium salt?

React an acid with an excess of a solid metal(if suitably reactive), metal oxide, hydroxide or carbonate

Use titration method. React an acid with a solution of sodium or potassium hydroxide or carbonate(or use ammonia solution).

No

Yes No

Yes

Definition in terms of

Oxidation Reaction

Reduction Reaction

Oxygen gain or loss

1. Substance gains oxygen

1. Substance loses oxygen

Hydrogen gain or loss

2. Substance loses hydrogen

2. Substance gains hydrogen

Electron transfer

3. Substance loses electrons to another reactantOILRIGOxygen is loss

3.Substance gains electrons from another reactantOILRIGReduction is gain.

Oxidation number

4. Increase in oxidation number

4. Decrease in oxidation number

Oxidising and Reducing Agents

a. An oxidising agent is a substance that oxidises another substance by

1. adding or donating oxygen to the substance2. removing hydrogen to the substance3. removing electrons from the substance

In the process of oxidising another substance, the oxidising agent itself becomes reduced. In a redox reaction, the substance that is reduced is an oxidising agent.

b. A reducing agent is a substance that reduces another substance by

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1. adding/donating hydrogen to the substance2. removing oxygen from the substance3. adding/donating electrons to the substance

In the process of reducing another substance, the reducing agent itself becomes oxidised. In a redox reaction, the substance that is oxidised is a reducing agent.

Displacement of metals from solution to investigate the reactivity of metals

The displacement of metals from solution is a redox reaction. In general, a more reactive metal will displace a less reactive metal from a solution of its ions.

One common displacement reaction in qualitative analysis is the displacement of copper from a solution of copper(II) salt by zinc metal:

Zn (s) + CuSO4 (aq)          ZnSO4 (aq) + Cu (s)

Zinc displaces copper from the solution because it is more reactive than copper metal. Zinc, being more reactive, loses its electrons readily. The electrons are transferred from the zinc atoms to the copper(II) ions in the solution. Copper(II) ions are reduced to copper metal and zinc atoms become oxidised to zinc ions.

Oxidation reaction: Zn (s)       Zn2+ (aq) + 2e-

Reduction reaction: Cu2+ (aq) + 2e-     Cu (s)

Reactions of metals with water

General Summary:

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Metals above hydrogen in the reactivity series:

Metal + Cold Water Metal Hydroxide + HydrogenMetal + Steam Metal Oxide + Hydrogen

Metals below hydrogen in the reactivity series do not react with water.

How do you know when to heat a mixture?

Carbonates and magnesium usually react in the cold

Most other substances usually have to be heated.

Test For Cations

Cations (postively charged ions) are usually identified using two main reagents:

1. Aqueous sodium hydroxide 2. Aqueous ammonia

A cation is usually detected by adding aqueous sodium hydroxide or aqueous ammonia to the unknown solution. 

The solubility table shows that most hydroxides are soluble except for Group I metal hydroxides e.g. NaOH, KOH and Ca(OH)2, sparingly soluble.

All other metal ions react with hydroxide ions in solution to form insoluble hydroxides. The chemical reaction involved in the formation of the insoluble metal hydroxide is called precipitation.

1. Sodium hydroxide

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An aqueous solution of sodium hydroxide solution consists of mainly free, mobile sodium ions, hydroxide ions and water molecules.

NaOH (aq) Na+ (aq) + OH- (aq)

a)Adding NaOH (aq) to salt solutions containing transition metal ions usually form coloured precipitates while salt solutions containing Gp I, II and III metal ions form white precipitates.

b) Metal hydroxides as Al(OH)3 , Zn(OH)2, and Pb(OH)2  react with excess sodium hydroxide to form complex salts that are soluble to form colourless solutions.

c) In the case of ammonium ion, NH4+, no precipitate is

formed with aqueous sodium hydroxide. However, when the mixture of ammonium ions and sodium hydroxide is warmed gently, ammonia gas is liberated

Colour of precipitate

Ion present; Name of precipiate

Ionic Equation

Blue Copper(II) ions;Copper(II)hydroxide

Cu2+(aq) + 2OH- Cu(OH)2(s)

OrangeBrown

Iron (III) ions; Iron(III)hydroxide

Fe3+(aq) + 3OH- Fe(OH)3(s)

Green Iron(II) ions; Iron (II) hydroxide

Fe2+(aq) + 2OH- Fe(OH)2(s)

(NH4)+ ions present; No precipitate but the

NH4+(s/aq) +OH- NH3(g) + H2O(l)

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smell of Ammonia.

2. Aqueous ammonia

Ammonia dissolves in water and dissociates partially to form hydroxide ions:

NH3 (g) + H2O (l) NH4+ (aq) + OH- (aq)

a)Most metal ions (except sodium, potassium and calcium) react with aqueous ammonia to form insoluble hydroxides.

b) Some metal ions e.g. Cu2+ and Zn2+ form complexes with ammonia. Hence, their hydroxides are soluble in excess ammonia.

Al 3+ (aq) ion

Cu2+ (aq) ion

Add dilute NH3 (aq) until no further change is observed

1a) Add dilute NH3 (aq) until a change is observed.

1b) Add an excess of dilute NH3 (aq) to the mixture in test 1(a).

Observation

White precipitate, insoluble in excess

a) blue precipitate formed,

b) ppt. dissolved in excess NH3 (aq) to form a dark blue solution.

Name and Aluminium Copper(II) hydroxide,

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formula of precipitate

hydroxide, Al(OH)3

Cu(OH)2

Type of chemical reaction(s)

Precipitation of Al(OH)3

a) Precipitation of Cu(OH)2

b) Formation of complex salt solution, i.e. Cu(OH)2 reacts with excess ammonia solution to form a soluble salt.

Ionic equation

Al3+ (aq) + 3OH- (aq) Al(OH)3 (s)

Cu2+ (aq) + 2OH- (aq) Cu(OH)2 (s)

Test for Anions

a)Test for Chloride, Bromide and Iodide ions Add dilute nitric acid followed by a few drops of silver nitrate. The dilute nitric acid may be added before or after the addition of silver nitrate solution. It is important to note that dilute hydrochloric acid cannot be used as it introduces chloride ions that interfere with the test.

Chloride ion

Bromide ion

Iodide ion

Observation

White precipitate

Pale cream precipitate

Yellow precipitate

Identity of precipitate formed

Silver chloride, AgCl

Silver bromide, AgBr

Silver Iodide, AgI

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Ionic equation

Ag+(aq) + Cl-(aq) AgCl (s)

Ag+ (aq) + Br- (aq) AgBr (s)

Ag+ (aq) + I- (aq) AgI (s)

b)Test For Sulphate ion, SO42- (aq)

A solution of the suspected sulphate is made in pure water. Enough dilute hydrochloric acid is added to make it acidic and then some barium chloride solution is added. A sulphate will produce a white precipitate of barium sulphate. Both barium chloride and barium nitrate solutions are used to detect anions that form insoluble barium salts.

Ba2+ (aq) + SO42- (aq)         BaSO4 (s)

It is important to note that dilute sulphuric acid cannot be used as it introduces sulphate ions that interfere with the test.

c)Test for carbonate ion in solution, CO32-

In the cold if a dilute acid is added to a solid carbonate, carbon dioxide is produced.

It is best to use dilute nitric acid because some acid-carbonate combinations can produce insoluble salts that coat the solid carbonate, stopping the reaction. As all nitrates are soluble it is best to use dilute nitric acid.

Dilute Nitric acid is added to the solution. Effervescence occurs in the cold of a colourless gas that turns limewater milky, carbon dioxide is produced.

Ionic equation of carobate reacting with any acid:

CO32- (s) + 2H+(aq) CO2 (g) + H2O(l)

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Collecting and Identifying

Tests for Gases

a) Hydrogen: Hydrogen is a colourless and odourless gas which is less dense than air and is almost insoluble in water. It can be collected over water or by upward delivery. Test: Hydrogen ‘pops’ when a lighted splint is held to the mouth of a test tube of hydrogen. The hydrogen combines explosively with oxygen in the air to make water.

2H2(g) + O2(g) 2H2O(l)

b)Oxygen: Oxygen is a colourless and odourless gas which is almost the same density as air and is only slightly soluble in water. It is normally collected over water. Test: Oxygen relights a glowing splint.

c)Carbon Dioxide: Carbon dioxide is a colourless gas and odourless gas which is denser than air and is slightly soluble in water. It can be collected both over water and by downward delivery.

Test: Carbon dioxide turns limewater(Calcium hydroxide solution) milky due to the formation of a white precipitate of Calcium carbonate.

CO2 (g) + Ca(OH)2        CaCO3 (s) + H2O (l)

Limewater         white ppt.

d) Chlorine Gas: Chlorine is a green and pungent smelling gas which denser than air and too soluble in water. It is usually collected by downward delivery. It can also be collected over concentrate salt solution as it is less soluble in the salt solution.

It turns moist blue litmus red and then bleaches

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it. Chlorine gas dissolves in water to form a strong acid and a bleaching agent. Hydrochloric acid dissociates to form hydrogen ions, which turn blue litmus red. Cl2 (g) + H2O (l)    HCl (aq) + HClO (aq)                                                      (bleaching agent)

e)Ammonia Gas: Ammonia gas is a pungent smelling alkaline gas which is less dense than air and extremely soluble in water. It can only be collected by upward delivery. It turns moist red litmus blue. Ammonia gas dissociates partially in water to form hydroxide ions, OH-, which turns red litmus blue.

NH3 (g) + H2O (l)        NH4+ (aq) + OH-(aq)

Tests for water: Using anhydrous copper(II)sulphate: Water turns white anhydrous copper(II) sulphate blue. anhydrous copper(II)sulphate lacks water of crystallization and is white. Dropping water onto it replaces the water of crystallization and turns it blue.

CuSO4(s) + 5H2O(l) CuSO4.5H2O(s)

Using cobalt chloride paper: When cobalt chloride paper is dry it is blue. Adding water to it turns it pink.

Cobalt chloride paper is filter paper that has been dipped into cobalt(II)chloride solution and then dried thoroughly in a desiccator. A dessiccator is a piece of glassware or a small cabinet that contains a tray of some substance which absorbs water.

Testing Water Purity: These tests do not show that water is pure. To check purity, it can be shown that water freezes at exactly 0degrees C and boils at 100degrees C at 1 atmospheric pressure.

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Flame tests:

Potassium (K+) Lilac

Sodium (Na+) Yellow/Orange

Lithium (Li2+) Dark Red (Crimson)

Calcium (Ca2+) Dull Red (Brick red)

Copper (Cu2+) Bluish green

Barium (Ba2+) Apple Green

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