BASES

CONTENT
⦁ Definition of Bases
⦁ Strength of an Alkali
⦁ Characteristics of Bases
⦁ Preparation of Bases
⦁ Reaction of Bases
⦁ Uses of Bases
⦁ Relative Acidity and Alkalinity (The PH Scale)
⦁ Logarithmic pH Scale
⦁ Worked Examples

Definition of Bases
A base is a substance which will neutralize an acid to yield salt and water only. It is either an oxide or hydroxide of a metal, e.g sodium oxide, magnesium hydroxide, etc, while; an alkaline is a basic hydroxide which is soluble in water. Bases that are soluble in water are referred to as alkalis. Examples are sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca(OH)2)
Oxides of heavy metals like PbO, ZnO and CuO are insoluble in water and are therefore bases not alkalis. CaO and MgO are slightly soluble and are alkalis. Like acids, alkalis may be strong or weak.
Strength of an Alkali
The strength of an alkali is the degree of its ionization in water. Strong alkalis are completely ionized in water. Examples of strong alkalis are potassium hydroxide, sodium hydroxide, calcium hydroxide, strontium hydroxide etc. Weak alkalis are partially ionized in water . example of a weak alkali is aqueous ammonia,.
Characteristics of Bases
⦁ Bases are soapy to touch, e.g. NaOH
⦁ They have bitter taste, e.g. lime water
⦁ They turn red litmus blue
⦁ Concentrated form of the caustic alkalis, NaOH and KOH are corrosive
⦁ They are electrolytes
Preparation of Bases
1. Combustion of a reactive metal in air. When electropositive metals are heated in oxygen, they form metallic oxides.
2Ca  + O2      →         CaO(s)
2. By reaction of metals with water (steam)
Ca +  H2O    →     Ca (OH)2    +   H2
3. Decomposition of metal hydroxides by heating

heat
Ca(OH)2 → CaO + H2O
heat

Cu(OH)2    →    CuO  +  H2O
4. Precipitation or double decomposition reaction
CuSO4   +   2NaOH   →    Cu(OH)2   + Na2SO4
5. Dissolution of metallic oxides in water
Na2O + H2O   →     NaOH
K2O   +  H2O    →     2KOH
EVALUATION
⦁ What is a base? Give three examples.
⦁ (a) Differentiate between a base and an alkaline. (b) Give three uses of base
⦁ State three properties of bases and, mention four methods of their preparation with examples.

Reaction of Bases
1. Reaction with acid: all bases react with acids to form salt and water only
NaOH + HCl    →     NaCl   +   H2O
2. Reaction of metallic hydroxide with heat

heat
Zn(OH)2 → ZnO + H2O
heat

Reaction with ammonium salts: Alkali reacts with ammonium salts in the presence of heat to liberate ammonia gas.
2NH4Cl  + Ca(OH)2     →     CaCl2  +   2NH3   +   2H2O
NaOH   +   NH4NO3    →     NaHNO3    +   H2O   +    NH3

Uses of Bases
⦁ NaOH is the most common base and is very soluble, hence it is used as a drying agent.
⦁ Bases are used in the making of soap.
⦁ Ca(OH)2 is used in the neutralization of soil acidity.
⦁ Mg(OH)2 is used in the production of toothpaste.
⦁ Ca(OH)2 is used in the making of mortar, plaster of Paris, white wash and cement and sugar refining.
⦁ NH4OH is used in weak solution as a common cleanser and grease solvent.
⦁ NH3 is used in making fertilizers and detergents.
EVALUATION
⦁ Describe two chemical properties of bases with examples.
⦁ Give 5 uses of bases with one example each.
Relative Acidity and Alkalinity (The PH Scale)
The term pH denotes hydrogen ion index. It is a number-scale used to express the degree of acidity or alkalinity of a solution and, the number ranges from 0 to 14. A solution with pH value of less than 7 is acidic while that with a value greater than 7 is alkaline. A solution with pH of 7 is neutral i.e. it is neither acidic nor alkaline. The pH of a solution can be measured with an instrument called pH meter.

Colour pH Number Acid/Base
Red 1 – 3 Very acidic
Orange 4 – 5 Weak acid
Yellow 6 Very weak acid
Green 7 Neutral
Blue 8 Very weak base
Indigo 9 – 10 Weak base
Violet 11 – 14 Very basic

pH Range and Colour Changes of Universal Indicator

Logarithmic pH Scale
Sorensen, in 1909, introduced the logarithmic pH scale to eliminate the inconvenience encountered when using negative indices and to give room for wide range of [H+]and [OH–] concentrations that we do come across in acid-base reactions. He defined pH as the negative logarithm of the hydrogen ion concentration to the base 10. For example, if the hydrogen ion concentration of an aqueous medium is 10-5 mol dm-3, the acidity of the solution could be written in terms of pH as follows:
[H+] = 10-5
Log [H+] = log 10-5 = −5
pH = −log [H+] = − (−5) = 5
Thus, if [H+] is 10-x, then pH = x
Proportional to each other:
[H+] [OH−]  = 10-14
PH+ POH = 14, where POH is the hydrogen ion index
pOH = 14 – pH
Note: A high pH value indicates low H+ concentration (weak acidity) and a high OH– concentration (strong alkalinity). At neutrality, [H+] = [OH–] = 10-7
Worked Example
Calculate the pH of 0.005 moldm-3 tetraoxosulphate (vi).
Solution
H2SO4    →     2H+ + SO42-
H+= [2 x 0.005] moldm-3
= 0.001 = 1 x 10-2
pH = −log [H+]
pH = −log [1 x 10-2]
pH = 2
EVALUATION
⦁ (a) Define the term PH. (b) Calculate the  pH of 0.01 M tetraoxosulphate (VI) acid solution.
⦁ Mention 3 different methods by which you could prepare bases in the laboratory. Write an equation to illustrate each method.
⦁ (a). Define pH and POH. (b) What is the relation pH and POH of an aqueous solution? (c) Determine the pH of a solution containing 0.05moldm-3 NaCO3.
⦁ Distinguish between (i) a strong base and a week base. (ii) a Base and an Alkaline.
⦁ A glass cup of orange juice is found to have POH of 11.40. Calculate the concentration of the hydrogen ions in the juice.
⦁ Find the hydrogen ion, H+ and hydroxide ion, OH– concentration in 0.02M solution of H2SO4.