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Buffers - what are they?
Buffers - what are they?
- Buffers are solution that resist small changes in pH when small amounts of acid or base are added
- Buffers can be acidic or basic and are generally made up of a weak acid and it's conjugate base(a salt / ion) or a weak base and it's conjugate acid (salt / ion) at equal concentrations
- Midway through the buffer zone/region is halfway to the equivalence point and is where pH = pKa (see below)
- The buffer works either side of the mid point and is most effective within a pH unit of 1 of this point
- A common example of an acidic buffer solution would be ethanoic acid and the ethaonate ion. This could be produced by adding a base such as sodium hydroxide to produce the salt or by dissolving the salt into the acid
- A common example of a basic buffer would be a mix of ammomia and the ammonium ion; this could be from ammonium chloride
- To work buffers must be able to remove extra H+ ions (H₃O+) or OH- ions
example of Acidic and basic buffers
example of Acidic and basic buffers
buffers at equilibrium
buffers at equilibrium
buffers at work
buffers at work
The image below illustrates how the ethanoic acid / ethanoate buffer works. As ethanoic acid is a weak acid it only partially dissociates, meaning there is a significant amount of this acid in the buffer. The salt, in this case sodium ethanoate fully disassociates so there is also a significant amount of the ethanoate ion present in the buffer.
The image below illustrates how the ethanoic acid / ethanoate buffer works. As ethanoic acid is a weak acid it only partially dissociates, meaning there is a significant amount of this acid in the buffer. The salt, in this case sodium ethanoate fully disassociates so there is also a significant amount of the ethanoate ion present in the buffer.
buffers : calculations
buffers : calculations
Let's assume the buffer contains a weak acid and its conjugate base (eg ethanoic acid and ethanoate ion) which can be represented by the following general terms:
Let's assume the buffer contains a weak acid and its conjugate base (eg ethanoic acid and ethanoate ion) which can be represented by the following general terms:
HA for the acid and
HA for the acid and
A- for the conjugate base / salt
A- for the conjugate base / salt
Writing a general dissociation reaction equation with water and putting this into a Ka expression gives:
Writing a general dissociation reaction equation with water and putting this into a Ka expression gives:
Ka will be given, or can be worked out from pKa, or looked up.
Ka will be given, or can be worked out from pKa, or looked up.
You can rearrange, this, to find the [hydronium ion]
You can rearrange, this, to find the [hydronium ion]
If the inverse logarithm for both sides of the Ka equation is taken then you get the Henderson-Hasslebalch equation, which can be used to work out buffer calculations
If the inverse logarithm for both sides of the Ka equation is taken then you get the Henderson-Hasslebalch equation, which can be used to work out buffer calculations
(For alkaline buffers p H = pKb + log salt / base)
calculation examples:
calculation examples:
indicators - what are they?
indicators - what are they?
Indicators are weak acids at equilibrium with its' conjugate base:
Indicators are weak acids at equilibrium with its' conjugate base:
HInd(aq) ⇄ H+(aq) + Ind-(aq)
HInd(aq) ⇄ H+(aq) + Ind-(aq)
Indicators are used to show the equivalence point in acid - base titrations.
Indicators are used to show the equivalence point in acid - base titrations.
Navigate to the titration page, using:
- the side navigation
- the QR code
- titrations information
The colour indicated depends which side of the equilibrium reaction is favoured.
The colour indicated depends which side of the equilibrium reaction is favoured.
For example using the generic indicator above if more acid (H+) is added then the reverse reaction will be favoured as H+ is removed and the colour shown will be due to the indicator.
For example using the generic indicator above if more acid (H+) is added then the reverse reaction will be favoured as H+ is removed and the colour shown will be due to the indicator.
If more base (OH-) is added then the excess OH- will react with the H+, forming H₂O, thereby removing H+.
If more base (OH-) is added then the excess OH- will react with the H+, forming H₂O, thereby removing H+.
In this situation equilibrium will shift, favouring the forward reaction producing more H+, as well as the conjugate base In-, and the colour will change to reflect the conjugate base colour.
In this situation equilibrium will shift, favouring the forward reaction producing more H+, as well as the conjugate base In-, and the colour will change to reflect the conjugate base colour.
As indictors are weak acids at equilibrium we can write a Ka expression:
As indictors are weak acids at equilibrium we can write a Ka expression:
From this we can determine that Kind = [H+] when the concentrations of the indicator (HInd) and its' conjugate base (Ind-) are equal:
From this we can determine that Kind = [H+] when the concentrations of the indicator (HInd) and its' conjugate base (Ind-) are equal:
In other words
In other words
From this we can see we want to. Ideally, an indicator has a pKa = pH at equivalence point of the titration. In reality the indicator's pKa needs to fall within +/- 1 of pH at equivalence.
From this we can see we want to. Ideally, an indicator has a pKa = pH at equivalence point of the titration. In reality the indicator's pKa needs to fall within +/- 1 of pH at equivalence.
pKa (Ind) = +/- 1 pH (equivalence point)
pKa (Ind) = +/- 1 pH (equivalence point)
In other words
In other words
* for a strong acid - strong base titration where pH at equivalence is 7 the indicator should have a pKa of 7. Within a range of +/- 1 of pH at equivalence works, ie the pKa of the indicator for this titration should fall between 6 - 8.
* for a strong acid - strong base titration where pH at equivalence is 7 the indicator should have a pKa of 7. Within a range of +/- 1 of pH at equivalence works, ie the pKa of the indicator for this titration should fall between 6 - 8.
*for a weak acid - strong base titration where pH at equivalence, is for example 9.24 the indicator should have a pH of 9.24. Within a range of +/- 1 of pH at equivalence works, ie the pKa of the indicator for this titration should fall between 8.24 - 10.24.
*for a weak acid - strong base titration where pH at equivalence, is for example 9.24 the indicator should have a pH of 9.24. Within a range of +/- 1 of pH at equivalence works, ie the pKa of the indicator for this titration should fall between 8.24 - 10.24.
indicator chart
indicator chart
This chart shows a range of indicators, their pKa and their pH at equivalence point use, along with the associated colour changes.
This chart shows a range of indicators, their pKa and their pH at equivalence point use, along with the associated colour changes.
indicator use - examples.
indicator use - examples.
Phenolphthalein has a pKa of 9.5
Phenolphthalein has a pKa of 9.5
It would be used in a titration which has a pH at equivalence point of between 8.5 - 10.5; in other words a weak acid - strong base titration.
It would be used in a titration which has a pH at equivalence point of between 8.5 - 10.5; in other words a weak acid - strong base titration.
Phenolphthalein as the weak acid is colourless. When more OH- is added removing the H+ equilibrium shifts and favours the forward reaction. This in turns means there is more of the conjugate base present (phph-) which is pink - this is the colour change as the end point is reached.
Phenolphthalein as the weak acid is colourless. When more OH- is added removing the H+ equilibrium shifts and favours the forward reaction. This in turns means there is more of the conjugate base present (phph-) which is pink - this is the colour change as the end point is reached.
The images below shows where this colour change lies and that it is the appropriate indicator for a weak acid-strong base titration, as opposed to methyl orange which changes colour around a pH of 4, indicating it would be a better choice for a weak base - strong acid titration.
The images below shows where this colour change lies and that it is the appropriate indicator for a weak acid-strong base titration, as opposed to methyl orange which changes colour around a pH of 4, indicating it would be a better choice for a weak base - strong acid titration.
A good choice of indicator for a strong acid - strong base titration would be bromothymol blue.
A good choice of indicator for a strong acid - strong base titration would be bromothymol blue.
pH at equivalence point for this titration would be 7.
pH at equivalence point for this titration would be 7.
Looking at the chart above bromothymol blue has a pKa of 7.3 so which is within +/- 1 of a pH at 7, which is the equivalence point for such a titration.
Looking at the chart above bromothymol blue has a pKa of 7.3 so which is within +/- 1 of a pH at 7, which is the equivalence point for such a titration.
This means this indicator would change colour, from yellow to blue around the equivalence point as shown below.
This means this indicator would change colour, from yellow to blue around the equivalence point as shown below.
CHOOSING THE CORRECT INDICATOR:
CHOOSING THE CORRECT INDICATOR:
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