An acid is a substance that ionizes in water to give hydrogen ions as the only positively charged ions in solutions.
Acid can be defined as a substance that react with a bass to form a salt and water only.
Properties of acid
- All acid turns blue litmus paper red
- Acid have got a sour , sharp test
- Acids react with carbonates and hydrogen carbonates to give a salt, carbondixide, and water.
2HCL [aq] + Na CO3 2NaCL [aq] + CO2 + H2O [L]
Ca [HCO3]2 + 2HCL [aq] CaCL2 [aq] + 2H2O [L] + 2CO2 [g]
- Acids react with hydroxides and metallic oxides to give salts and water.
CuO [S] + 2HCL [aq] CuCL2 + H2O [L]
2NaOH [aq] + H2SO4 [aq] Na2 SO4 +2H2O[l]
- Acids react with metals to form a salt and hydrogen gas.
Zn[z] + 2HCL [aq] ZnCL2 [aq] + H2 [q]
- Action on indicators
- Litmus solution blue red
- Methyl orange red
- 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
The universal indictor a series of colour changes as the acidity of the solution changes as the acidity of the solution changes.
It is the number of ionisable 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
- HNO3 H+ + NO3 – 1
- HCL H+ + CL – 2
- H2SO4 2H+ + SO42- – 2
- H2SO4 3H+ + PO43- -3
- H2CO3 2H+ + CO32- – 2
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.
Therefore HXA XH+ + A–
CH3 CH2 CH2 COOH
Generally all organic acid are weak acid. eg lactic acid
A base is a substance that reacts with an acid to form water and salt only
- Poltasium hydioxlde- [KOH]
- Sodium hydroxlde [NaOH]
- Calcium hydroxide [Ca] [OH2)
- Ammonium hydroxide [NH, OH]
Metallic oxides like;
- Magnesium oxide Mgo
- Copper [II] OXIDE
- Sodium oxide [Na2o]
Properties of bases
- Bases turned litmus paper blue.
- 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 .
it is the number of hydrogen ion6 released by molecule of the acid on complete dissociation in water.
(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).
- HCL H[aq]+ + CL–[aq]
One molecule hydrogen iron
- Nitric acid; HNO3.
HNO3 H+ + NO3–[aq] [strong acid]
1molecule hydrogen iron
- 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.
- Sulphric acid; H2SO
H2SO4 2H+ + SO42- aq [strong acid]
[Strong acid] 2 hydrogen iron
- Carbonic acid H2CO3 (weak acid)
H2CO3 2H2+ + CO32-
(Acids with a basicity of three)These contain three ionisable hydrogen atoms per molecule of the acid e.g. Phosphoric acid H3PO4
H3PO4 3H+ [aq] + PO43-[aq] [weak acid]
Strong and weak acid;
Strong acid –
This is an acid which fully ionized in dilute aqueous solution.
A dilute solution of a strong and contains many ions and very few or no molecule of the acid
ii) Hydrogen acid (HCL) (aq)
HCL H+[aq] + CL–(aq)
iii) Nitric acid HNO3
HNO3 H+ [aq] + NO3–[aq]
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)
- i) Carbonic acids H2CO (aq)
H2CO3 [aq] 2H+ [aq] + CO2-
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
- By reaction between acid and hydrate. (Acidic oxide of a non metal) and water.
Examples of acids prepared by this method are;
- sulphurous acid
H2O [L] + SO2 [g] H2SO[g] [aq]
Sulphur [iv] oxide
- Carbonic acid;
H2O [L] + CO2 [g] H2CO[g] [aq]
Carbondioxide [carbon [iv] oxide]
- Phosphoric acid;
P4O2 [aq] + 6H2O [L] 4H2PO4 [aq]
- 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 H4SO2 [aq] + NaCL [s] NaHSO4 [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  ethanoate ethanoic acid
Bases and Alkaline (additions)
Abase is a hydrogen ion accepter. (Proton accepter
An acid can be defined as hydrogen ion donor.
An alkali is a soluble base. The commonest solvent is water
Examples of alkali are;
- Sodium oxide (Na2O)
Na2O[s] + H2O [L] 2NaOH [aq]
- Pottosium oxide (K2O)
K2O[s] + H2O [L] 2KOH
- magnesium oxide MgO
MgO[s] + H2O [L] Mg [OH] 2[aq]
CaO[s] +H2O[g] 2CaOH [aq]
Insoluble bases are not alkali
Examples of insoluble bases
- i) copper (ii) oxide (CuO)
- ii) Lead (ii) oxide (PbO)
iii) Iron (ii) oxide [Fe2O3]
- iv) Copper (ii) hydroxide Cu [OH] 2
- v) Lead (ii) hydroxide Pb [OH]2
- vi) Zinc (iii) hydroxide Zn [OH]2
Strong and weak alkalis
- a) Strong alkali
eg sodium hydroxide
Strong alkalis are electrovalent and are completely ionized in aqueous solution and in the solid state.
NaOH[s] Na+ [aq] + OH–[aq]
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
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.
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?
- sulphur dixide.
Oxides of phrosphous dissolve in water to give acidic solution
- Neutral oxide – These are neutral to litmus eg water
- 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.
When they dissolve in acids they behave as bases and in alkalize as acids.
It is a compound containing a negative ion from acid and metallic c or ammonium group (radical)
Types of salts
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
- salts from hydro chlochloric acid HCL;
NH4CL, Amonium chloride
NaCL, Sodium chloride.
KCL, Potasium chloride.
LiCL, Lithium chloride.
- 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
Monobasic acids cannot form salts because they contain only one ionisable hydrogen atoms.
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
Synthesis [direct combination of elements
- 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]
- Iron [iii] chloride FeCL3
2Fe[s] + 3CL2[aq] 2Fe CL3
By heating iron in dray chloride
- Iron [ii] chloride, FeCL2 + H2
Prepared by heating iron metal in one atmosphere of hydrogen chloride gas
Sodium chloride, NaCL and magnesium chloride, Mg CL2 can be prepared by heating the respective metal in the atmosphere of chlorine gas.
Mg[s] + CL2 [g] MgCL2 [s]
2No[s] + CL2 [g] 2NaCL
Reacting an acid with a metal or an insoluble oxide, hydroxide, carbonate.
In the preparation of soluble salts of copper, lead, iron and zinc. The general procedure is;-
- Add metal or metal hydroxide, oxide or carbonate, to the appropriate acid if the solid is in excess, heating if necessary
- Filter off excess solid
- Saturate the filtrate by evaporation [crystals can only form from concentrated filtrate to crystallize.
- Filter off the crystals and wash them with disliked water.
- Dry the crystals between filter paper or in a desiccators Or under sun rays
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 + HSO4 ZnSO4 + H2[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.
Preparation of copper  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 SO4 [aq] CuSO4 [aq] + H2O[l]
Preparation of lead  nitrate crystals by reaction dilute nitric acid and insoluble lead  carbonate
Lead  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 + CO2 [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  nitrate form, they are flitted off, wasted with distilled water and they are dried.
Copper  sulphate crystal, copper  nitrate crystals. Magnetism sulphate crystal, zinc sulphate crystal, calcium chloride crystals and calcium nitrate crystals can be prepared using this method.
Calcium chloride and calcium nitrate are deliquescent and they do not form crystals. Their solutions must be evaporated to dryness.
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.
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  iodide, PbI2 lead  nitrate solution is added potassium iodide solution. A yellow precipitate of potassium iodide is formed.
Pb [ NO3]2[aq] + 2KI PbI2[s] + 2KNO3[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  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 SO4 [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 + H2SO4 [aq] BoSO4[S] + 2H NO3 [aq]
The white precipice is fettered off, washed in distilled water and dried
Lab preparation of lead  chloride crystals
Lead  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 CL2[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
- Ammonium chloride
When it is heated it sublimes to form gaseous ammonium chloride. On further heating the gas dissociates into ammonia and hydrogen chloride
NH4CL [s] NH4CL [g]
NH4CL [s] NH3 [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 in group I [sodium and potassium] decompose on heating to give the nitrite of the metal and oxygen gas.
2KNO3[s] 2KNO2 + O2
2NaNO3[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  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] + 4NO2 [g] + O2 [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  oxide is yellow when hot and white on cooling
2Zn [NO3]2 [s] 2ZnO[s] + 4NO2 [g] + O2 [g]
A solution is a uniform mixture of a solvent and a solute
This is the substance which dissolves in a solvent. It can be a solid, liquid or gases
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
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.
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.
Qualities of a good graph
- Should have a little
- X and y axis should be drown and well labeled
- Divisions should be equal
- Scale for x and the y axis should be stated
- Points should be clearly plated
- 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
Scale for the y axis – 1:5 grammes
Range = 107. 2
Scale for x – axis -1cm: 5oc
Range = 60 – 0 = 60
18cm : 60
1 : 3.33oc
Cont IV; 6 :0
17 :1 grammes
17.1g of potassium chloride crystallises out
75 grammes of a saturated solution salt x contained 30 gms of salt
- Mass of water in the solution
75g – 30g
- The percentage of water in the solution
= x 100%
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] + 4NO2 [g] + O2 [g]
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
2AgNO3 2Ag[s] + 2NO2 [g] + O2 [g]
Effect of heat on carbonates
- 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 2NH3 [g] + CO2 [g] H2O[g]
It is a white powder, on heating it decomposes to give off a colourless gas which turns lime water milky [carbondioxide]
ZnCO3[S] ZnO[s] + CO2 [g]
Cont copper  carbonate, Cu CO3- it is a green powder, it decomposes on heating to give a black residue [of copper  oxide, Cuo] and a colourless gas which turns lime water milky [carbodioxide, CO2]
CuCO3[s] CuO[s] + CO2 [g]
Effect of heat on lead  carbonate, PbCO3 – it is a white powder. On heating, it decomposes to give reddish – brown a residue which turns yellow on cooling [lead  oxide, Pbo] and a colourless gas which turns lime water milky [ carbondioxide; CO2]
PbCO3[s] PbO[s] + CO2 [g]
Potassium and sodium decompose on heating to give the corresponding carbonates, water and carbon dioxide gas.
Sodium hydrogen carbonate
2NaHCO3[s] Na2CO3[s] + H2O[g] + CO2 [g]
Potassium hydrogen carbonate
2KHCO3 K2CO3[s] + H2O[g] + CO2 [g]
Calcium hydrogen carbonate
Ca [HCO3]2[s] CaO[s] + H2O[g] + 2CO2 [g]
It decomposes on heating to give calcium oxide, water and carbondioxide gas
Effects of heat on sulphates
- Copper  sulphate crystals, CuSO4. SH2O [Copper  sulphate – water [copper  sulphate pentshydrate
Blue crystals on gentle heating, they loss water of crystallization to form a white anhydrous solid.
Cu SO4.SH2O[s] CuSO4[s] + 5H2O
Blue crystals white powder
On strong heating the white anhydrous solid decomposes to give a black residue [of copper  oxide] and white fumes of sulphur trioxide gas
CuSO4 CuO[s] + SO3 [g]
White powder black black sulphur trioxide
Residues [sulphur [v1] oxide
Iron  sulphate crystals, FeSO4. 7H2O – Iron  sulphate – 7 water
–iron  sulphate
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
On strong heating sulphur trioxide gas and sulphur dioxide gas are given off. The residue is reddish – brown [iron  oxide].
2Fe SO4[s] Fe2O3[s] + SO2 [g] + SO3 [g]