STRONG BASE + WEAK ACID

When equal moles of sodium hydroxide, NaOH, and acetic acid CH₃COOH, were mixed they would also react completely

However {CH₃COO}‾ can itself participate in an equilibrium reaction with water, behaving like a base to accept a proton from water.

It is known as the conjugate (or corresponding) base of acetic acid. So there will be a little acetic acid left no matter what. It is not possible to achieve a stoichiometric point. The final pH of this solution would be slightly basic (highly than 7). For such a titration it is best to use an indicator which changes colour at a pH higher than 7. Like phenolphthalein.

WEAK BASE and STRONG ACID

With equal moles of ammonium hydroxide, NH₄OH, and hydrochloric acid,

It is also not possible to achieve a stoichiometric point. This solution would be slightly acidic (pH less than 7). The best indicator to use would be methyl red.

WEAK BASE AND WEAK ACID

For reaction for equal moles of ammonium hydroxide and acetic acid

both conjugate acid and conjugate base {NH₄}⁺ + {CH₃COO}‾ will be produced. There is no suitable indicator for such titrations and it is best that they be avoided.

SUMMARY

• Titration with weak acids will produce conjugate bases.
• Titration with weak will produce conjugate acids.
• The end point will not be the stoichiometric point.

WATER

Water has both acidic and basic properties; or amphiprotic.

The concentration of {H₃O}⁺ in pure water was found to be 1.004x10‾⁷ moles per litre. So the concentration of {OH}‾ should also be the same. That is water when pure is neutral in terms of acidity. The pH for this solution would then be;

So when the pH of an aqueous solution is 7 the solution is said to be a neutral solution (in terms of acidity). pH less than 7 is considered acidic and greater than 7 alkaline (or basic).

Just to make this topic complete it should be mentioned that a term autoprotolysis constant of water, K_w, is defined as;

BUFFERS

Henderson-Hasselbatch Equation

Based on the equilibrium constant of the Bronsted equilibrium, the expression for pH would be,

This is known as the Henderson-Hasselbatch Equation.

Let us see what happens in a solution containing a weak acid (HA) and the salt of its conjugate base (M⁺A‾). We will choose a MA that dissociate fully in water (meaning very soluble salts). Let us prepare the solution such that [HA]₀ = [M⁺A‾]₀.

The subscribe "0" denote the value at time=0, or initial value. Such values are always known value. Once the two are mixed, two equilibriums would come into play

HA is a weak acid so the equilibrium is very much to the left and for M⁺A‾ it will be very much to the right. The concentration of A‾ would push the equilibrium for the dissociation of HA even further to the left. So we can assume that at anytime the concentration of HA is [HA]₀ and the concentration of A‾ is [MA]₀. Using this value for the Henderson-Hasselbatch Equation we will get

So we can prepare a solution with known pH using this technology. The best is yet to come.

Let us take the case of phosphoric acid / potassium biphosphate; H₃PO₄ / KH₂PO₄, where [H₃PO₄]₀ = [KH₂PO₄]₀ = 0.200M. The pK_a value for phosphoric acid is 2.12. So the pH of the solution would be 2.12.

Suppose we accidentally drop some HCl so that its concentration in the mixture is 0.001M. The equilibriums that come into play would be;

Consequently [HA] = (0.200 + 0.001) M and [MA] = (0.200 - 0.001) M, the Henderson-Hasselbatch Equation will then be;

The pH hardly changes. For a 0.001M HCl solution the pH should have been 3. Such phosphoric acid / potassium biphosphate solution is known as a buffer.

We can prepare various buffers (for various pH value) by using the combination of a weak acid and the salt of its conjugate base (or weak base and the salt of its conjugate acid).

pH buffers are widely used in industry if the reaction is pH sensitive. One good example will be the detergent industry. Most of the detergents have built in pH buffers to keep the pH at a specific value. The next time you see a facial cleansing soap pay attention to this.