ACIDS, BASES AND SALTS

Acids

An acid is a substance that ionizes in water to give hydrogen ions as the only positively charged ions in solutions.

Or

Acid can be defined as a substance that react with a bass to form a salt and water only.

Examples

Properties of acid

1. All acid turns blue litmus paper red
2. Acid have got a sour , sharp test
3. Acids react with carbonates and hydrogen carbonates to give a salt, carbondixide, and water.

 

Examples

2HCL _[aq] + Na CO_3(aq) → 2NaCL _[aq] + CO₂ + H₂O _[L]

Ca [HCO₃]_2(s) + 2HCL _[aq]  → CaCL_{2 [aq]} + 2H₂O _[l] + 2CO_{2 [g]}

1. Acids react with hydroxides and metallic oxides to give salts and water.

CuO _[S] + 2HCL _[aq] → CuCl_2(aq) + H₂O _(l)

2NaOH _[aq] + H₂SO_{4 [aq]} → Na₂ SO_4(aq) +2H₂O_[l]

1. Acids react with metals to form a salt and hydrogen gas.

Zn(s) + 2HCl(aq)→ ZnCl_{2 [aq]} + H_{2 [g]}

1. Action on indicators

 

Examples

1. Litmus solution blue red
2. Methyl orange red
3. Phenolphthalein colorless pick

An indicator is a substance that has different colours in acid and alkaline solutions

Ph is the concentration of hydrogen ions in a solution.

1                                 7                                14

Neutral

The universal indictor a series of color changes as the acidity of the solution changes as the acidity of the solution changes.

Basicity

It is the number of ionizable hydrogen atoms in one molecule of acid.

There are three types of acids normally;

I) monobasic acids

ii) Dis-basic acids

iii) tri- basic   acid

 Weak and strong acids

A weak acid is one that undergoes partial ionization when dissolved in water while a strong acid is an acid that complete ionization when put in water.

Strong acids

Therefore H_XA                        XH⁺ + A^–

WEAK ACIDS

CH3 COOH

CH₃CH₂ COOH

CH₃ CH₂ CH₂ COOH

Generally all organic acid are weak acid. eg lactic acid

BASES

A base is a substance that reacts with an acid to form water and salt only

 Examples

1. Poltasium hydioxlde- [KOH]
2. Sodium hydroxlde [NaOH]
3. Calcium hydroxide [Ca] [OH₂)
4. Ammonium hydroxide [NH, OH]

Metallic oxides like;

5. Magnesium oxide Mgo
6. Copper [II] OXIDE
7. Sodium oxide [Na₂^o]

Properties of bases

1. Bases turned litmus paper blue.
2. They react with acids to form salts and water.

Eg 2NaaH + HS2O4                        Na2SO4 + 2H2O [L]

This is called neutralization reaction

Basicity of an acid;

It is the  number  of  ionisable  hydrogen  atom  in  one  molecule  of  acid .

Or

it is the  number  of  hydrogen  ion6  released  by molecule  of  the  acid  on  complete  dissociation   in  water.

Monobasic  acid;

(An acid with a basicity of none) Are acids which dissociate to given only one hydrogen ion per molecule of the acid? E.g. HCL _(aq) Hydrochloric acid (strong acid).

Water

1. HCL      H_[aq]⁺ + CL^–_[aq]

One molecule hydrogen iron

2. Nitric acid; HNO₃.

HNO₃                          H⁺ + NO₃^–_[aq]                [strong acid]

1molecule hydrogen iron

3. Ethanoic acid; CH₃

CH₃COOH                             CH₃COO^–[aq] + H⁺ [aq].         [Weak acid]

1 molecule [weak acid]                                        hydrogen ion

The basicity of an acid is not the number of hydrogen atom in one molecule of the acid but it is the number of ionizable hydrogen atoms per molecule o the acid.

Di basic acid;

(Acids with abasicity of two )These contain two ionisable hydrogen atom per molecule of the acid.

Examples

4. Sulphric acid; H₂SO

H₂SO₄                                     2H+     + SO₄^2- _aq                     [strong acid]

[Strong acid] 2 hydrogen iron

1. Carbonic acid H₂CO₃ (weak acid)

H₂CO₃                      2H²⁺ + CO₃^2-

Tribasic acid;

(Acids with a basicity of three)These contain three ionisable hydrogen atoms per molecule of the acid e.g. Phosphoric acid H₃PO₄

                                                [orthophoric acid]

H₃PO₄                                 3H⁺ _[aq] + PO₄^3-_[aq]                     [weak acid]

Strong and weak acid;

Strong acid –

This is an acid which fully ionized in dilute aqueous solution.

Examples  

Sulphric acid

A dilute solution of a strong and contains many ions and very few or no molecule of the acid

ii) Hydrogen acid (HCL) _(aq)

Water

HCL                                     H⁺_[aq] + CL^–_(aq)

One

Molecule

iii) Nitric acid HNO₃

HNO₃                                  H⁺ [aq] + NO₃^–[aq]

Weak acid;

Are acids which are only slightly ionized (slightly dissociated) in dilute aqueous solution?

Such acids contains many molecules of the acid and few ions (hydrogen ions)

Examples

1. i) Carbonic acids H₂CO _(aq)

H₂CO₃ [aq]                              2H⁺ _[aq] + CO^2-

Many molecules                                  few ions

Of the acid

ii) Organic acids;

e.g. Ethanoic acid [acetic], CH₃COOH

CH₃COOH _[L]                          CH₃COO^– + H⁺ _[aq]

Preparation of acids

Method of preparation acids

General methods

1. By reaction between acid and hydrate. (Acidic oxide of a non metal) and water.

Examples of acids prepared by this method are;

1. sulphurous acid

H₂O _[L] + SO_{2 [g]}                       H₂SO_{[g] [aq]}

Sulphur [iv] oxide

2. Carbonic acid;

H₂O _[L] + CO_{2 [g]}                      H₂CO_{[g] [aq]}

Carbondioxide [carbon [iv] oxide]

3. Phosphoric acid;

P₄O_{2 [aq]} + 6H₂O _[L]                                4H₂PO_{4 [aq]}

1. ii) By displacement of more volatile acid by less volatile acids.

Sulphuric acid being less volatile, hydrochloric acid sulphric acid can displace hydrogen chloride from metallic chlorides.

Conc H₄SO_{2 [aq]} + NaCL _[s]                               NaHSO_{4 [s]} + HCL_[g]

Higher B.P metallic                                                     lower B.P

Less volatile chloride                                   more volatile

Similarly sulphuric acid, less volatile acid displaces more volatile acid from metallic nitrates

Conc H₂SO + NaHO _[aq]                                     NaHSO _[aq] + HCL _[ag]

iii) By perspirations of insoluble sulphides from metallic salt using hydrogen sulphides.

[CH₃COO]₂Pb + H2S_[g]          PbS_[s]   + CH₃ COOH_[aq]

Lead [111] ethanoate                          ethanoic acid

Bases and Alkaline (additions)

Abase is a hydrogen ion accepter. (Proton accepter

An acid can be defined as hydrogen ion donor.

(Proton donor)

Alkalis

An alkali is a soluble base. The commonest solvent is water

 Examples of alkali are;

1. Sodium oxide (Na₂O)

Na₂O_[s] + H₂O _[L]                     2NaOH _[aq]

Sodium hydroxide

1. Pottosium oxide (K₂O)

K₂O_[s] + H₂O _[L]                       2KOH

2. magnesium oxide MgO

MgO_[s] + H₂O _[L]                      Mg [OH] _2[aq]

CaO_[s] +H2O_[g]                        2CaOH _[aq]

Note

Insoluble bases are not alkali

Examples of insoluble bases

1. i) copper (ii) oxide (CuO)
2. ii) Lead (ii) oxide (PbO)

iii) Iron (ii) oxide [Fe₂O₃]

1. iv) Copper (ii) hydroxide Cu [OH] ₂
2. v) Lead (ii) hydroxide Pb [OH]₂
3. vi) Zinc (iii) hydroxide Zn [OH]₂

Strong and weak alkalis

1. a) Strong alkali

eg sodium hydroxide

Potassium hydroxide

Strong alkalis are electrovalent and are completely ionized in aqueous solution and in the solid state.

Water

NaOH_[s]                          Na⁺ _[aq] + OH^–_[aq]

                                Water

KOH_[s]                               K⁺ _[aq] + OH^–_[aq]

Weak alkalis; bonding in these alkalis is covalent and they exists an molecule. They are only slightly ion sable in dilute aqueous solution and their ionization is irreversible

Example

Ammonia        NH₃

Properties of Alkalis;

• Bitter taste
• Soapy fill when touched
• Change colours of indictors;

Turn litmus blue, methyl orange to yellow, phenophlen from colourless to pink.

• Alkalis react with an acid to form salt and water only. (neutralization)

Acid + Alkali – salt + water

HCL _[aq] + Na OH_[aq]                NaCL + H₂O

• Alkalis precipitate insoluble metal hydrogen in solution of salts.

PH scale of acidity or Alkalinity.

The pH scale is a scale of number from 0to 14 to express acidity or alkalinity. PH is related to hydrogen ion concentration.

PH of 7 represents neutral point – This is the pH of distilled water.

Solutions which have pH value below 7 are acidic; most fruit juices are weak acid have value of about 6 to 5.

PH above 7 represents alkalis-The higher the pH value the stronger the alkali and the lower the pH value the weaker the alkali.

Universal indictors

Indictors like methyl orange and phenopthaline

Do not show weather the acid is strong or weak .A universal indictor, enables us to density solutions as neutral, weak and strong bases or acids. It has a chart with different coloure which match with the number or pH values.

Classification of oxides

There are four classes of oxides

• Acid oxides (acidic hydrants)

Are oxides of non metal which dissolve in water to give acidic solution?

1. sulphur dixide.

sulphur trioxide.

Oxides of phrosphous dissolve in water to give acidic solution

• Neutral oxide – These are neutral to litmus eg water

1. ii) carbondioxide.

Basic oxide – These are oxides of metal which reacts with to form salts and water only

Alkalis are bases which are soluble in water eg sodium oxide.

• Amphoteric oxides – metallic oxides which have properties of acid and bases.

Lead (ii) oxide.

Aluminum (ii) oxide.

Zinc oxide.

When they dissolve in acids they behave as bases and in alkalize as acids.

SALTS

It is a compound containing a negative ion from acid and metallic c or ammonium group (radical)

Types of salts

Normal salt;

It is a salt in which all the ioniseble hydrogen atoms of the acid have been replaced by a metal or ammonium group. Normal salt don’t contain hydrogen from the acid

Examples of normal salts

1. salts from hydro chlochloric acid HCL;

NH₄CL, Amonium chloride

NaCL, Sodium chloride.

KCL, Potasium chloride.

LiCL, Lithium chloride.

1. Salts from Nitric acid, HNO₃;

LiNO₃  Lithium Nitrate,

NaNa₃ Sodium nitrate,

KNO₃ potasium nitrate

Normal salts from sulphiric acid,

H₂ SO₄; Na₂so₄ Sodium sulphiric,

K₂ SO₄ Potassium sulphate,

Li₂ SO₄  Lithium sulphate,

Zn SO₄  Zinc sulphate,

Mg SO₄  Magnesium sulphate,

[NH₄]₂ SO₄ ammonium sulphate

NB

Monobasic acids cannot form salts because they contain only one ionisable hydrogen atoms.

Acid salts;

An acid salt is one which contains some ion able hydrogen atoms from the acid.

An acid salt is one which is capable of further ionization in aqueous solution to give hydrogen ions.

Examples of acid salts

Preparation of salts

The method used in the preparation of a given salt depends on whether the salt is soluble or insoluble in water.

Table of soluble in soluble salts

Methods of preparation of salts

Method 1

Synthesis [direct combination of elements

 Examples

1. Iron [ii] sulphide, Fes[s] is prepared by heating iron powder with sulphur powder

Zinc[ii] sulhide Zn S[s], by heating zinc powder with sulphur power

Zn[s] + S[s]                 ZnS[s]

1. Iron [iii] chloride FeCL3

2Fe[s] + 3CL2[aq]                  2Fe CL3

[Black crystals]

By heating iron in dray chloride

1. Iron [ii] chloride, FeCL2 + H2

Prepared by heating iron metal in one atmosphere of hydrogen chloride gas

NOTE

Sodium chloride, NaCL and magnesium chloride, Mg CL2 can be prepared by heating the respective metal in the atmosphere of chlorine gas.

Mg_[s] + CL_{2 [g]}                         MgCL_{2 [s]}

2No_[s] + CL_{2 [g]}                                     2NaCL

[white crystals]

Method 11

Reacting an acid with a metal or an insoluble oxide, hydroxide, carbonate.

Examples

In the preparation of soluble salts of copper, lead, iron and zinc. The general procedure is;-

1. Add metal or metal hydroxide, oxide or carbonate, to the appropriate acid if the solid is in excess, heating if necessary

1. Filter off excess solid
2. Saturate the filtrate by evaporation [crystals can only form from concentrated filtrate to crystallize.

• Filter off the crystals and wash them with disliked water.

1. Dry the crystals between filter paper or in a desiccators Or under sun rays

Cont. examples

Preparation of zinc sulphate crystals

Dilute sulphric acid is poured in a glass beaker and zinc granules are added to the acid. Effervescence occurs.

If the reaction is a low a little copper [ii] sulphate solution is added as a catalyst and the reactants are warmed.

Zn_[2] + HSO₄                           ZnSO₄  + H_2[g]

Metal + acid                 + gas

When the reaction steps, more zinc is added to make sure that the acid is not left in considerable amount excess zinc granules and solid impurities are filtered off. The filtrate is gently heated to concentrate it.

The concentrate filtrate is then called. White crystals of zinc sulphate form. They are filtered off, wasted with distilled water they are dried between filter papers

Magnesium sulphate crystals and iron [ii] sulphate crystals can be prepared in the same way using magnesium metal and iron fillings respectively.

Example 2

Preparation of copper [11] sulphate crystals

Copper [ii] oxide; CuO, is added a little a a time to worm dilute sulphate sulphric acid in a glass beaker until no more dissolves. Excess copper [ii] oxide + solide impurities are filtered off. The filtrate is evaporated to concentrate it.

The concentrated filtrate is then cooled. Blue crystals of copper [ii] sulphate -5- water, CuSO₄. 5H20 form. The crystals are filtered, washed with distilled water then dried in desiccators or sunshine.

Zinc sulphate crystals and lead [ii] nitrate crystals can be prepared in the same way.

CuO_[s] + H₂ SO_{4 [aq]}                  CuSO_{4 [aq]} + H₂O_[l]

Example 3

Preparation of lead [11] nitrate crystals by reaction dilute nitric acid and insoluble lead [1] carbonate

Lead [11] carbonate is added a little at a time to dilute nitric acid in a beaker. Effervescence occurs as carbondioxide is given off. More carbonate is added until no more reacts showing that old the acid has reacted.

Pb CO₃ + 2HNO₃                   Pb [NO₃]₂ + CO_{2 [g]} + H₂O _[L]

Carbonate + acid                         salt + carbondixide + water

Excess carbonate is filtered off and the filtrate is evaporated until crystal begins to form when it cools.The concentrated filtrate is cooled. White crystals of lead [11] nitrate form, they are flitted off, wasted with distilled water and they are dried.

Copper [11] sulphate crystal, copper [11] nitrate crystals. Magnetism sulphate crystal, zinc sulphate crystal, calcium chloride crystals and calcium nitrate crystals can be prepared using this method.

Note:

Calcium chloride and calcium nitrate are deliquescent and they do not form crystals. Their solutions must be evaporated to dryness.

Method 111

Preparation of salt by action of an acid a soluble hydroxide or carbonate

Salts of sodium;

Potassium and ammonium can be prepared by this method from solutions of sodium hydroxide, potassium hydroxide and ammonium solution respectively using the appropriate acid.

The set up for the experiment is shown in the figure below

A known volume sodium hydroxide solution is piped into a conical flask and phenolphthalein indicator is added to give a pink liquid.

Dilute hydrochloric acid is added from the burette to sodium hydroxide solution little at a time until the solution turns calourless. The volume of the acid used is noted.

The resulting solution is discarded because it contains an indicator. Equivalent volume of the acid and alkali are now added, this time without using indicator.

Na OH [aq] + HCL [aq]                     NaCL[aq + H2O[L]

The resulting mixture is evaporated to dryness using a water bath to recover sodium chloride crystals.

Insoluble salts

They are prepared by precipitation method or double decomposition method. In a double decom position reaction, anions and cat are exchanged.

For example in the preparation of lead [11] iodide, PbI₂ lead [11] nitrate solution is added potassium iodide solution. A yellow precipitate of potassium iodide is formed.

Pb [ NO₃]_2[aq] + 2KI                            PbI_2[s] + 2KNO_3[aq]

The yellow precipitate is filtered off, washed with distilled water and is dried.

Normally, in the preparation of an insoluble salt, two solution of soluble salt and insoluble salt. The insoluble slot is filter off, washed with distilled water and dried

Preparation of lead [11] sulphate

Dilute sulphic acid added to lead [ii] nitrate solution in a breaker. A white precipitate of lead [ii] sulphate is formed.

H₂SO _4[aq] + Pb [NO₃]_{2 [aq]}                                Pb SO_{4 [S]} + 2HNO _3[aq]

The precipitate is filtered off washed with distilled water to remove soluble impurities and it is dried

Laboratory preparation of barium sulphate

Barium nitrate is added to dilute sulphiric acid. A white precipitate of barium sulphate is formed

Ba [NO₃]_2[aq + H2SO_{4 [aq]}                     BoSO_4[S] + 2H NO_{3 [aq]}

The white precipice is fettered off, washed in distilled water and dried

Lab preparation of lead [11] chloride crystals

Lead [11] nitrate solution is added to dilute hydrochloric acid in a beaker a white precipitate of lead[s] chloride is formed.

Pb[NO₃]_{2 [aq]} + 2HCL_[aq]                                  Pb CL_2[S] + 2H[NO₂]_2[aq]

The white precipitate is filtered off, washed with distilled water

Good crystals are obtained by dissolving the washed crystals in a minimum amount of hot water and cooling to recrytalize.

Effects of heat on some salts

1. Ammonium chloride

When it is heated it sublimes to form gaseous ammonium chloride. On further heating the gas dissociates into ammonia and hydrogen chloride

Heat

NH₄CL _[s]                    NH₄CL _[g]

Cool

White solid

Strong heat

NH₄CL _[s]                    NH_{3 [g]} + HCL _[g]

Ammonia and hydrogen chloride recombine to form dense white fumes of ammonium chloride. On cooling, the white fumes condense into a white solid [white sublimate]

Nitrates

Nitrates in group I [sodium and potassium] decompose on heating to give the nitrite of the metal and oxygen gas.

Heating

2KNO_3[s]                                  2KNO₂ + O₂

Potassium                    potassium

Nitrate                         nitrite

Heating

2NaNO_3[s]                                   2NaNO₂ + O₂

ii) Nitrates of metals lower in the reactively seems decompose to give the oxide of the metal, reddish – brown nitrogen dioxide gas and oxygen gas

Effect of heat on lead [11] nitrate crystals

On heating a crackling noise is heard, the white crystals melt evolving reddish – brown nitrogen dioxide and oxygen gas.

2Pb [NO₃]_2[S]                                2PbO _[aq] + 4NO_{2 [g]} + O_{2 [g]}

The residue [lead [ii] oxide] is reddish – brown when hot and yellow on cooling

Effect of heat on zinc nitrate

Zinc nitrate is a white crystalline solid, on heating a reddish – brown gas [nitrogen dioxide] and oxygen gas are given off. The residue [zinc [11] oxide is yellow when hot and white on cooling

2Zn [NO3]_{2 [s]}                         2ZnO_[s] + 4NO_{2 [g]} + O_{2 [g]}

Redishh brown

Solutions

A solution is a uniform mixture of a solvent and a solute

Solute

This is the substance which dissolves in a solvent. It can be a solid, liquid or gases

Solvent

Substance in excess] the solvent can be solid, liquid, or gas.

Note that air is a solution in which solvent is nitrogen and other gases are solutes

Saturated solution

A saturated solution of a solute at a particular temperature is one which can dissolve no more solute at room temperature in the presence of undissolved solute.

Super saturated solution

This one that contains more solvent than it can hold at the same temperature in the presence of undissolved solute.

Solubility curves

A solubility curve of a substance is a group that shows how it solubility varies with temperature.

The solubility of a solute in a solvent at a particular temperature is the number of grammes or moles of the solute required to saturate 100 grammes of solvent at that temperature.

Units are in moles/ 10grammes of water or grammes per 100 grammes of water at given temperature.

Plotting graphs

Qualities of a good graph

1. Should have a little
2. X and y axis should be drown and well labeled
3. Divisions should be equal
4. Scale for x and the y axis should be stated
5. Points should be clearly plated
6. The graph should be a straight line or a curve

Suitable scales for plotting graphs

1:1                   1:10                 1:100

1:2                   1:20                 1:200

1:5                   1:50                 1:500

1:10                 1:100               1:1000

1:0.1

1:0.2

1:0.5

1:1

Scale for the y axis – 1:5 grammes

Range = 107. 2

-12.9

95.3

27cm: 95:3

1cm: 95:3

27

   1: 3.52

Scale for x – axis -1cm: 5^oc

Range = 60 – 0 = 60

18cm   :           60

1          :

1          :           3.33^oc

Cont IV; 6       :0

43        :9

17        :1 grammes

17.1g of potassium chloride crystallises out

75 grammes of a saturated solution salt x contained 30 gms of salt

1. Calculate;
2. Mass of water in the solution

75g – 30g

= 45grammes

1. The percentage of water in the solution

=  x 100%

= 60%

Determine the solubility of the salt.

45g – 30g = 15g

45 grammes of water was saturated by 30 grammes of salt

1 gram of water was saturated by  grammes

100grammes of water are saturated by  x 100 gms

= 6 grammes per 100gramms of H₂O

Effects of heat magnesium nitrate crystals

On heating, white crystal decompose to give a reddish brown gas [nitrogendioxide and oxygen the residue is white

2mg [NO₃]_2[s]                           2MgO[s] + 4NO_{2 [g]} + O_{2 [g]}

White crystals

Nitrates of metals at the bottom of reactivity series e.g. silver and mercury decompose on heating to give the metal, nitrogendioxide and oxygen gas

Examples

                                    Heat

2AgNO₃                                  2Ag_[s] + 2NO_{2 [g]} + O_{2 [g]}

White                                                              grey

Crystals

Effect of heat on carbonates

1. Carbonates of group 1 metals e.g. sodium carbonate and potassium carbonate do not decompose on heating.

Effects of heat on ammonium carbonate

The salt is very unstable on heating it decomposes to give ammonia gas, carbon dioxide + water vapour

[NH₄]₂CO₂                              2NH_{3 [g]} + CO_{2 [g]} H₂O_[g]

White crystals

Zinc carbonate

It is a white powder, on heating it decomposes to give off a colourless gas which turns lime water milky [carbondioxide]

ZnCO_3[S]                                  ZnO_[s] + CO_{2 [g]}

Cont copper [11] carbonate, Cu CO_3- it is a green powder, it decomposes on heating to give a black residue [of copper [11] oxide, Cuo] and a colourless gas which turns lime water milky [carbodioxide, CO₂]

Heat

CuCO_3[s]                                  CuO_[s] + CO_{2 [g]}

Effect of heat on lead [11] carbonate, PbCO3 – it is a white powder. On heating, it decomposes to give reddish – brown a residue which turns yellow on cooling [lead [11] oxide, Pbo] and a colourless gas which turns lime water milky [ carbondioxide; CO2]

Heat

PbCO_3[s]                                  PbO_[s] + CO_{2 [g]}

Hydrogen carbonates

Potassium and sodium decompose on heating to give the corresponding carbonates, water and carbon dioxide gas.

Sodium hydrogen carbonate

Heat

2NaHCO_3[s]                                         Na₂CO_3[s] + H2O_[g] + CO_{2 [g]}

Potassium hydrogen carbonate

Heat

2KHCO₃                                 K₂CO_3[s] + H2O_[g] + CO_{2 [g]}

Calcium hydrogen carbonate

Ca [HCO₃]_2[s]                                      CaO_[s] + H2O_[g] + 2CO_{2 [g]}

It decomposes on heating to give calcium oxide, water and carbondioxide gas

Effects of heat on sulphates

4. Copper [11] sulphate crystals, CuSO₄. SH₂O [Copper [11] sulphate – water [copper [11] sulphate pentshydrate

Blue crystals on gentle heating, they loss water of crystallization to form a white anhydrous solid.

Cu SO₄.SH₂O_[s]                                   CuSO_4[s] + 5H₂O

Blue crystals                                        white powder

Hydrated                                             [anhydrous]

On strong heating the white anhydrous solid decomposes to give a black residue [of copper [11] oxide] and white fumes of sulphur trioxide gas

CuSO₄                                    CuO_[s] + SO_{3 [g]}

White powder             black black sulphur trioxide

Residues [sulphur [v1] oxide

Iron [11] sulphate crystals, FeSO₄. 7H₂O – Iron [11] sulphate – 7 water

–iron [11] sulphate

Heptahydrate

The crystals are pole green. On gentle heating, the crystals lose water of crystallization to form a dirty white anhydrous solid.

Fe SO₄. 7H2O                         FeSO₄ [s] + 7H₂0_[l]

Pale green                                dirty white

Crystale

On strong heating sulphur trioxide gas and sulphur dioxide gas are given off. The residue is reddish – brown [iron [111] oxide].

2Fe SO_4[s]                    Fe₂O_3[s] + SO_{2 [g]} + SO_{3 [g]}