The reaction of an acid with a base is therefore characterised by the transfer of a proton from the acid to the base. Consider for example the hydrolysis of acetic acid. IN this reaction, CH₃COOH donates a proton to water, this indicates that CH₃COOH is an acid in the Bronsted sense and H₂O functions as a base.

------- H+ -----à

CH₃COOH + H₂O CH₃COO- + H₃O+

Acid base

General Acid / Base Reaction

i) HA + B BH⁺ + A^-

acid 1 base 1 acid 2 base 2

e.g.

HCl + NH₃ Cl^- + NH₄⁺

acid 1 base 1 base 2 acid 2

NOTE

These reactions are shown as reversible and in the reverse reaction, acid 2 acts as the acid donating the proton.

i.e. HCl and NH₄⁺ are both acids (HCl is a much stronger acid and the reaction is mainly forward) and NH₃ and Cl^- are both bases. HCl and Cl^- differ only by one proton and they are called a conjugate pair. So, Cl^- is the conjugate base of the acid HCl. Similarly, NH₃and NH₄⁺are also a conjugate pair with NH₄⁺being the conjugate acid of the base NH₃

Acid-Base Reactions in Aqueous Solution

Water may react either as an acid by forming OH^- or as a base by forming H₃O⁺.

e.g.

i) as an acid

H₂O + NH₃ NH₄⁺ + OH^-

Acid conj base

ii) as a base

H₂O + HCl H₃O⁺ + Cl^-

Base conj acid

Water can also react with itself, but only very slightly

-----(-)--------à

H₂O + H₂O H₃O⁺ + OH^-

Acid 1 base 2 acid 2 base 1

Another example of such a substance is HCO₃^-

As a base

HCO₃^- + H₂O H₂CO₃ + OH^-

As an acid

HCO₃^- + H₂O CO₃ ^2- + H₃O⁺

Entities that can act as both an acid and a base are called amphiprotic

Some molecules can donate more than one proton. An example would be sulphuric acid, H₂SO₄

i.e.

H₂SO₄+ H₂O H₃O⁺+ HSO₄ ^-

HSO₄ ^- + H₂O H₃O⁺+ SO₄^2- (incomplete)

The sulphuric acid molecule donates two protons to the base, H₂O. Sulphuric acid is called a diprotic acid.

In a similar way phosphoric acid can donate protons, i.e.

H₃PO₄ + H₂O H₃O⁺+ H₂PO₄ ^- (partial)

H₂PO₄ ^- + H₂O H₃O⁺+ HPO₄ ^{2 -} (slight)

HPO₄ ^{2 -} + H₂O H₃O⁺+ PO₄ ^{3 -} (very slight)

Such an acid is said to be triprotic

Distinguish between strong and weak acids on the basis of completeness of ionization.

Strength of Acids

The strength of an acid depends on the degree to which it ionises in aqueous solution. Strong acids such as HCl, HNO₃ , H₂SO₄ ionize virtually to completion whereas weak acids such as CH₃COOH, H₂CO₃ and H₃PO₄ ionize only partly or very little at all. Sulphurous acid, H₂SO₃ can be referred to as an acid with moderate strength. It is a stronger than the weaker acids listed here, but it is not as strong at the stronger acids mentioned.

The strength of am acid is also greatly influenced by the relative strength and stability of its conjugate base. Strong acids generally have very stable conjugate bases, which are consequently weak bases. The reverse reaction is there for negligible, e.g, Cl^- is a stable and weak base. Weak acids relatively unstable conjugates which are quite strong bases and the reverse reaction is considerable, so CH₃COO^- is a strong base compared to Cl^-

In general strong acids have weak conjugate bases

weak acids have strong conjugate bases

Distinguish between strong and weak bases on the basis of completeness and ionization.

Strength of Bases

The strength of a base primarily depends on its ability to absorb a proton. IN the case of bases soluble aqueous solutions (alkalis), the strength of the base depends on the degree of dissociation of the molecule in water. So bases with a high degree of dissociation in water will be strong bases while those bases with a low degree of dissociation in water will be weak acids

NaOH Na⁺ + OH^-complete à strong base

KOH K⁺ + OH^-complete à strong base

Ca(OH)₂ Ca ²⁺ + 2OH^-slight à weak base

Mg(OH)₂ Mg ²⁺ + 2OH^-slight à weak base

NH₃ + H₂O NH₄⁺ + OH^-slight à weak base

The pH System

To obtain a numerical system for comparing the strengths of acids and bases a system has been established, the pH system.

The concentration of H₃O⁺ion in solution is taken as the standard. Pure water is taken as the zero point or point of neutrality, where the concentrations of both H₃O⁺and OH^- are exactly equal

Note The greater the concentration of H₃O⁺, the stronger is the acidity of the solution.

The symbol [ ] indicates the concentration of a solution in molarity

The [H₃O⁺] of 0.1 M HCl is 10^-1

The [H₃O⁺] of pure water is 10^-7

The [H₃O⁺] of 0.1 M NaOH is 10^-13

As these values of [H₃O⁺] are difficult to work with, a further definition of the acidity of a solution is made – the pH of a solution

So pH = – log₁₀ [H₃O⁺]

So 0.1 M HCl, pH = 1 pure water, pH =7 0.1 NaOH, pH = 13

Distinguish between acidic. basic and neutral solutions using the pH system.

Distinguishing between acid, basic and neutral solutions using the pH system becomes quite a simple matter.

1) If the pH of a solution is less than 7, the solution is acidic, the lower the pH of the solution, the stronger the acidity

2) If the pH of a solution is equal to 7, the solution is neutral [OH^-] = [H₃O⁺] = 10^-7

3) If the pH of a solution is greater than 7, the solution is basic, the higher the pH of a solution, the stronger the bascisity

In pure water

2 H₂O H₃O⁺ + OH^-

[H₃O⁺] = [OH^-] = 10^-7 (pH = 7)

[H₃O⁺] X [OH^-] = 10^-14

so pH + pOH = 14

Calculate pH from [H+] or [OH-]

Calculating the pH of a Solution

Using [H⁺] = [H₃O⁺]

To determine the pH of a solution knowing it’s [H₃O⁺] is simple : pH = – log₁₀ [H₃O⁺]

So example : calculate the pH of a solution of 0.01M HCl

\ pH = – log₁₀ (0.01)

= – log₁₀ (10^-2)

= 2

Using [OH^-]

To determine the pH of a solution knowing its [OH^-] is also quite simple, but it requires 2 steps.

So Example : calculate the pH of a solution of 0.01 M NaOH

\ pOH = – log₁₀ (0.01)

= – log₁₀ (10^-2)

pOH = 2

since pH + pOH = 14

pH = 14 – pOH

= 14 – 2

pH = 12

Measurement of pH of a Solution

There are three basic methods of measuring the pH of a solution

pH Paper

this is specially treated paper which will change colour as soon a sit is placed in a solution containing H₃O⁺. Depending on the colour change the pH of the solution can be determined approximately. Using pH paper has the benefit of being quick and easy, however, it is only an approximate value of the pH of the solution (~ 1 to 0.5 of a pH unit) and each piece of paper can only be used for one measurement.

pH Meter

this is an electronic device which will give a quick and very accurate reading of the pH of a solution containing H₃O⁺ , as soon as the electrodes connected to the meter are placed in the solution. Using a pH meter has many advantages, it is quick, accurate provided the machine and its electrodes are calibrated correctly and cared for will last for many years of service

indicators

during an acid / base titration we want to know exactly when the reaction between the acid and the base is finished, i.e. when the acid and the base have just neutralized on another. To do this we use acid / base indicators, which are substances that undergo definite colour changes within a fairly narrow range of pH. That is, the indicators change colour just as soon as the reaction between the acid and the base is finished

Some Common Indicators

Indicator	Ph Range for Colour Change	Colour at Lower pH	Colour at Higher pH
Thymol Blue	1.2 – 2.8	Red	Yellow
Methyl Orange	3.1 – 4.4	Red	Yellow
Methyl Red	4.2 – 6.3	Red	Yellow
Litmus	5 – 9	Red	Blue
Bromothymol Blue	6.0 – 7.6	Yellow	Blue
Thymol Blue	8.0 – 9.6	Yellow	Blue
Phenylphthalon	8.3 – 10.0	Colourless	Red

A universal indicator is really a mixture of several indicators chosen to display a variety of colours over a particular range of pH. Universal indicators are useful in the determination of the approximate pH of solutions.

Write Balanced Ionic and Molecular Equations for the Neutralisation of Common Acids with Common Bases

Neutralization Reactions of Acids and Bases

In general, when an acid and a base react with one another, the products are water and a salt

e.g.

HCl + NaOH NaCl + H₂O

Acid base salt

H₂SO₄ + 2KOH K₂SO₄ + 2 H₂O

Acid base salt

Equations such as these in which no charged particles (ions) are shown are called molecular equations for the neutralization of an acid with a base.

When equations in which charged particles (ions) are used, the equations are called ionic equations for the neutralization of an acid with a base.

e.g.

NaOH_(s) + (aq) Na⁺ _(aq) + OH^- _(aq)

HCl _(l) + (aq) H⁺ _(aq) + Cl^- _(aq)

Ionic equation H⁺ _(aq) + OH^- _(aq) H₂O

Determine the molarity and concentration of common strong acids and common strong bases.

Example 1 A 20ml aliquot of 1.00M NaOH is titrated to end point with 15.2ml of a H₂SO₄ solution of unknown concentration. Calculate the molarity of the H₂SO₄ solution.

H₂SO₄+ 2NaOH 2 H₂O + Na₂SO₄

1 mole 2 moles

x mole 2x moles

no. of moles of NaOH present =

= 0.02 mole

\ no. of mole of H₂SO₄present is 0.01 ( )

\ in 15.2 ml of H₂SO₄, there are 0.01 moles of H₂SO₄

\ the molarity of the H₂SO₄ solution = X 0.01

= 0.66M

example 2 A 20ml aliquot of a 0.5M H₂SO₄solution is titrated to end point with 25 ml of a KOH solution of unknown concentration. Calculate the molarity of the KOH solution

H₂SO₄ + 2 KOH 2 H₂O + Na₂SO₄

1 mole 2 mole

in 20 ml of 0.5M H₂SO₄solution there are = 0.01 moles of H₂SO₄

\ in 25ml of KOH solution there are 0.02 moles of KOH (0.01 X 2)

\ the molarity of KOH solution is = 0.8M

b. Experimentally determine the concentration of a common strong acid or base using a common strong base or acid.

State why aqueous solutions of only some salts of neutralization reactions have a pH = 7

Only some aqueous solutions of salts of neutralization have a pH = 7. Normal salts do not have a replaceable H⁺, so they cannot donate H⁺

The aqueous solutions of salts of neutralization reactions resulting from the neutralization of a strong acid and a strong base are neutral, i.e. pH = 7

The aqueous solutions of salts of neutralization reactions resulting from the neutralization of a strong acid and a weak base are acidic, i.e. pH < 7

The aqueous solutions of salts of neutralization reactions resulting from the neutralization of a weak acid and a strong base are basic, i.e. pH > 7

The pH of the solution of aqueous solutions of salts of neutralization treactions varies greatly in the area of the equivalance or end point

Example Titration of NaOH with HCl

This occurs because in the vicinity of the end point (where the acid and base have nearly neutralized one another) the addition of very minute volumes of titrant (NaOH) caused large changes in the pH of the solution.

Because of the fact that aqueous solutions of only some salts of neutralization reactions have a pH = 7, a wide variety of acid / base indicators are required, to show when a certain acid / base has reached its equivalence point.

The table of acid / base indicators mentioned earlier showing the pH range in which they work is often referred to.

Name all the Ions and Molecules which Aqueous Solutions in CO₂, SO₂, NH₃

Carbon Dioxide

i) CO₂+ H₂O H₂CO₃ (Carbonic Acid)

This reaction is only slight, about 1%

Carbonic acid is only a weak acid. The evidence for this lies in its poor conductivity and is supported by the litmus test. Carbonic Acid is only slightly ionised in the excess water.

H₂CO₃ + H₂O H₃O⁺ + HCO₃^– (slight)

Hydronium Hydrogen Carbonate

Ion Ion

The hydrogen carbonate ion is capable of acting as an acid, although it is a much weaker acid than carbonic acid.

HCO₃^- + H₂O H₃O⁺ + CO₃^2– (very slight)

Sulphur Dioxide

ii) SO₂+ H₂O H₂SO₃ (Sulphurous Acid)

The solution formed is only weakly acidic and is probably similar to carbon dioxide solution in that it contains molecules of the gas which are dissolved but have not reacted with the water. Sulphurous acid will ionise only to a slight extent in the excess water.

H₂SO₃ + H₂O HSO₃^– + H₃O⁺

Hydrogen Sulphite Hydronium

Ion Ion

Further ionisation of the HSO₃^– ions will occur but only to a very slight extent.

HSO₃^– + H₂O SO₃^2– + H₃O⁺

Sulphite Hydronium

Ion Ion

The whole solution is called sulphurous acid, and because it contains several different molecules and ions, it cannot be simply represented by any one formula. However as a matter pf convenience, the formula H₂SO₃is usually used to represent the solution.

Ammonia

iii) NH₃+ H₂O NH₄⁺ + OH^-

Ammonium Hydroxide

Ion Ion

A solution of ammonia inn water is alkaline to litmus. The solution is a poor conductor of electricity and so it can be said, it contains very few ions.

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