Buffer solutions

what do buffer solutions do

they maintain a reasonably constant pH on the addition of small amount of either acid (H₃O⁺) or base (OH^-)

what are buffer solutions essential for

1. maintaining a constant internal environment in living organisms

2. medical and industrial processes which are pH sensitive eg. producing solutions for delivering drugs and preparing synthetic or processed foods

what do buffer solutions contain similar concentrations of either of the following

• a weak acid and its conjugate base

• a weak base and its conjugate acid

why do buffer solutions maintain a relatively constant pH

because:

• upon addition of a base (OH^- ions) the acid or conjugate acid part of the buffer neutralises the added base by donating protons to the hydroxide ions to form water (pronated)

• upon addition of a base (H₃O⁺ ions)the base or conjugate base part of the buffer neutralises the added acid by accepting protons from the acid - the added hydrogen ions are said to protonate the base or conjugate base

meaning of pronated

the added hydroxide ions in a buffer are said to be pronated

equations for acid and base buffers

base:

HA + OH^- → H₂O + A^-

acid:

A^- + H₃O⁺ → H₂O + HA

what is the buffer zone

buffer solutions are effective only over a pH range of one pH unit above and below the pK_a value for the weak acid in the buffer

• this range of effective pH control (pKa±1) is the buffers zone

example of buffer zone

conjugate acid-base pair:

ethanoic acid and sodium ethanoate

pK_a = 4.76

buffer zone:

3.76-5.76

what are calculations involving buffer solutions limited too

to monoprotic acids eg. ethanoic acid

• which can release one proton only per acid species (molecule or ion) in solution

reaction equation and K_a expression for a weak acid HA

HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)

K_a = [H₃O⁺][A⁻]/[HA]

rearranged formula:

[H₃O⁺]= K_a x [HA]/[A⁻]

in this formula: [H₃O⁺]= K_a x [HA]/[A⁻] what are some of the properties

• [A⁻]≠[H₃O⁺] because the [A⁻] is from the added salt, not from the reaction of HA

• [H₃O⁺] and hence the pH of the buffer solution is dependent on the ratio [HA]/[A⁻]

how

how can this formula [H₃O⁺]= K_a x [HA]/[A⁻] also be written

by using the negative logarithmic formula

pH=pK_a+ log[A⁻]/[HA]

where [HA]= acidic species and [A⁻] = basic species

what happens if the concentrations of both species are equal ([HA]=[A^-])

then [H₃O⁺]= K_a and pH = pKa

what happens to the pH of a buffers solution if the solution is diluted

the pH of a buffer solution is unchanged

how are buffer solutions prepared

by making solutions that contain similar concentrations of a weak acid and its conjugate base or a weak base and its conjugate acid