Buffers and Neutralisation Flashcards
What is a buffer solution
a system that minimizes pH changes when small amounts of an acid or a base are added.
What does a buffer solution contain
a weak acid (removes added alkali), and a conjugate base (removes added acid)
Why does a buffer solution only work for small volumes
because there are only a small amount of reserves and once all of the reserves have been reacted then the solution loses its buffering ability
preparation of a buffer from a weak acid and its salt
e. g.CH3COOH and CH3COONa
1. CH3COOH -> CH3COO- H+
2. CH3COONa(s) + (aq) -> CH3COO- + Na+
preparation of a buffer by partial neutralization of the weak acid
an alkali (e.g. NaOH) is added to an excess of a weak acid The weak acid is partially neutralized forming a conjugate base. Resulting solution contains a mixture of the salt and weak acid
How does the conjugate base remove added acid
H+ ions increase and react with the conjugate base. Equilibrium shifts left removing ost of H+
How does the weak acid remove added alkali
[OH-] increases, the small concentration of H+ ions react with OH- ions. HA dissociates shifting equilibrium to the right to restore most of the H+ ions.
When is the buffer most effective
when there are equal concentrations of the weak acid and the conjugate base ([HA]=[A-]
When [HA(aq)] = [A-(aq)] :
- pH of the buffer solution is the same as the pKa value of HA
- the operating pH is typically over about 2 pH units, centered at the pH value of the pKa value
How do you calculate the pH of a buffer solution
- [H+] = Ka x [HA]/[A-]
2. -log[H+]
How do you calculate the pH of a buffer solution made from a weak acid and its salt
- Calculate the amount in mol of the acid and base
- calculate the concentrations
- [H+] = [HA]/[A-]
- -log[H+]
Where is the buffer solution in the body
blood plasma controlling blood pH between 7.35-7.45 for optimum enzyme activity.
What happens if blood pH falls below 7.35
develop a condition called acidosis which may cause fatigue, shortness of breath and in extreme cases death
What happens if blood pH rises above 7.45
a condition called alkalosis develops which can cause muscle spasms, light-headedness and nausea
Equation for the carbonic acid buffer system
H2CO3(aq) -> H+(aq) + HCO3-(aq)
How do you calculate the concentration ratio of HCO3-/H2CO3 in healthy blood
- [HCO3-]/[H2CO3] = ka/[H+]
- convert pH into [H+] and pKa into ka
- Calculate ratio
What is the Henderson - Hasselbalch equation
pH= pka + log([A-]/[HA])
What does a pH meter consist of
an electrode that is dipped into a solution, and connected to a meter that displays a pH meter usually recorded to 2 decimal places
Describe a procedure to monitor the change of pH as an aq base is added to an acidic solution
- using a pipette, add a measured volume of acid to a conical flask
- Place the electrode of the pH meter into the flask
- Add the base to the burette and add to the acid on the conical flask 1cm3 at a time
- After each addition, swirl the contents. Record the pH and the total volume of base added
- repeats steps 3 nd 4 until pH starts to change more rapidly. Then add the base dropwise until pH starts changing less rapidly.
- Add base 1 cm3 at a time again until an excess has been added and the pH has been basic, with little change, for several additions.
Describe the first section of a pH titration curve
Excess of acid - pH increases slightly as basic solution is added
Describe the equivalence point of a pH titration curve
Equivalence point (center of the vertical section) - The volume of one solution that exactly reacts with the volume of the other solution. The solutions have then exactly reacted with one another and the amounts used matching the stoichiometry of the reaction.
Describe the vertical section of a pH titration curve
pH increases rapidly on addition of a very small volume of base. Acid and base concentrations similar.
Describe the final section of a pH titration curve
excess of base - pH increases slowly as basic solution is added
pH titration curve shape of adding an acid to a base
The shape is the same just the other way around
What is an acid-base indicator
a weak acid that has a distinctively different colour from its conjugate base
Acid-base indicator end point
indicator contains equal concentrations of HA and A- and the colour will be in between the 2 extreme colors
What happens to the indicator when a basic solution is added
OH- react with H+ in the indicator. The weak acid dissociates, shifting equilibrium to the right and causing the colour to change with it e.g. methyl orange turns yellow.
What happens to the indicator when an acidic solution is added
H+ ions react with the conjugate base shifting equilibrium left. e.g. methyl orange turns red for the weak acid
At the end point:
- ka = [H+]
2. pka=pH
How do you choose which indicator to use
The indicator must have a colour change which coincides with the vertical section of the pH titration curve. Ideally the end point and equivalence point would coincide. However, this may not be possible and the end point may give a volume that is slightly different to the equivalence point, but any difference will be small.
