About this experiment
Blood, sea water and many medicines must keep their pH almost constant, even when acids or bases are added. They do it with buffers: mixtures of a weak acid and its conjugate base. Add a little strong acid and the base mops it up; add a little strong base and the weak acid neutralises it. Drop acid and base into a buffer and into pure water side by side, watch the indicator colours, and find out when a buffer works well and when it is used up.
Learning goals
- Define a buffer as a solution that resists changes in pH when small amounts of acid or base are added
- Explain how a weak acid and its conjugate base remove added H⁺ and OH⁻ ions
- Calculate the pH of a buffer with the Henderson-Hasselbalch equation
- Explain how the concentration and the base to acid ratio affect buffer capacity
- Choose a buffer pair for a target pH using its pKa
How to use it
- 1Add 1 mL of acid five times. Compare the pH changes in the buffer and in the water.
- 2Press Start again and add 10 mL of base. How far does each pH move?
- 3Keep adding acid until the buffer pH falls quickly. How much acid had been added? Compare with the moles of base in the buffer.
- 4Start again with a buffer concentration of 0.02 mol/L, then 0.4 mol/L. Which one resists change better?
- 5Change the ratio of base to acid to 1 : 4 and 4 : 1. Check the starting pH with pH = pKa + log([base]/[acid]).
- 6Which buffer pair would you use to keep a solution near pH 7.4, like blood?
The model
Assumptions
- Each beaker starts with 100 mL at 25 °C. The acid added is 0.1 mol/L HCl and the base is 0.1 mol/L NaOH, and the volume increases as they are added.
- Concentrations are used instead of activities, and only one acid dissociation step of each buffer pair is included (pKa 4.76, 7.21 and 9.25).
- The pH is found by solving the charge balance exactly, so it stays correct when the buffer runs out. The Henderson-Hasselbalch value is shown for comparison.
- The water starts at pH 7.00: dissolved carbon dioxide from the air is ignored.
Try this in class
- Q1Why does adding 1 mL of 0.1 M HCl change the pH of water by about 4 units, but a buffer by only about 0.01?
- Q2A buffer contains 0.05 mol/L ethanoic acid and 0.05 mol/L sodium ethanoate. What is its pH?
- Q3Why is a buffer most effective when the concentrations of acid and base are equal?
- Q4Explain why diluting a buffer with water hardly changes its pH but does reduce its capacity.
- Q5The carbonic acid and hydrogencarbonate pair keeps blood near pH 7.4. What would happen to blood pH without it during hard exercise?
Frequently asked questions
Is a buffer the same as a neutral solution?
No. A buffer can hold any pH close to the pKa of its weak acid. An ethanoate buffer stays near pH 4.8 and an ammonia buffer near pH 9.3. What makes it a buffer is that the pH changes very little, not what the pH is.
What is buffer capacity?
It is how much strong acid or base the buffer can take before its pH changes noticeably. It is largest when the acid and base are at equal concentrations and grows with the total concentration.
Why does the Henderson-Hasselbalch estimate disappear when the buffer is used up?
The equation needs both the acid and the base to be present. Once all of one of them has reacted, the logarithm cannot be calculated and the pH is decided by the excess strong acid or base instead.