Buffers in Everyday Life

Introduction

Add one drop of hydrochloric acid to a beaker of pure water and the pH can fall from 7 to about 4. Add the same drop to your blood and its pH barely moves from 7.4. Blood is a buffer : a solution that resists changes in pH. Buffers keep living cells working, stop shampoos stinging your eyes and keep food tasting right on the shelf. This page explains, without heavy mathematics, how they manage it.

Core explanation

What a buffer is. A buffer solution keeps its pH almost constant when small amounts of acid or alkali are added, or when it is diluted a little. It cannot resist unlimited additions, but it smooths out small changes.

What a buffer contains. The most common type of buffer contains two partners in similar amounts:

- a weak acid , such as ethanoic acid, CH₃COOH, and - a salt of that weak acid , such as sodium ethanoate, CH₃COONa, which supplies the ethanoate ion, CH₃COO⁻.

The weak acid is a store of H⁺ that can be released; the ethanoate ion is a store of base that can take H⁺ up. Because the acid is weak, the two exist together without simply neutralising each other.

Adding acid. When H⁺ ions are added, ethanoate ions mop them up:

CH₃COO⁻ + H⁺ → CH₃COOH

The added H⁺ becomes part of a weak acid molecule, which hardly ionises, so the concentration of free H⁺ scarcely rises.

Adding alkali. When OH⁻ ions are added, the weak acid neutralises them:

CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O

The added OH⁻ becomes water, so the pH scarcely rises.

In both cases, one partner is used up and the other is made, so the ratio of acid to salt changes only slightly and the pH changes only slightly. Once one partner is nearly used up, the buffer fails and the pH changes rapidly; the amount it can absorb is its buffer capacity .

Buffers in blood. Human blood must stay between about pH 7.35 and 7.45. Outside this narrow range, enzymes and oxygen transport stop working properly, and a pH below about 7.0 or above about 7.8 can be fatal. The main blood buffer is carbonic acid, H₂CO₃ (from dissolved carbon dioxide), with hydrogencarbonate ions, HCO₃⁻:

HCO₃⁻ + H⁺ → H₂CO₃ → H₂O + CO₂

When muscles release lactic acid during exercise, hydrogencarbonate ions remove the extra H⁺, and the carbon dioxide formed is breathed out faster. Proteins such as haemoglobin and phosphate ions act as additional buffers.

Buffers in products.

- Shampoos and skin products are buffered to around pH 5.5 to match skin and avoid irritation. - Foods and soft drinks often contain citric acid with sodium citrate, which keeps flavour and colour stable and helps preservatives work. - Medicines such as some aspirin tablets and eye drops are buffered to reduce irritation. - Swimming pools use hydrogencarbonate to keep pH steady so disinfectants work well. - Laboratories use buffer solutions of known pH to calibrate pH meters.

Step-by-step reasoning

To explain what a buffer does when acid is added:

1. Identify the base partner in the buffer (the ion from the salt). 2. Write the equation for that ion reacting with H⁺. 3. Note that the product is a weak acid, which does not release much H⁺. 4. Conclude that free H⁺ changes only slightly, so pH changes only slightly.

Visual explanation

Picture two graphs side by side, each showing pH against drops of acid added. For pure water, the line plunges steeply from pH 7 after the first drop. For the buffer, the line stays almost flat at pH 4.8 for many drops, then drops sharply once the ethanoate is used up. The flat stretch is the buffering region.

Real-world analogy

A buffer is like a shock absorber on a car. Small bumps (small additions of acid or alkali) are absorbed, and the passengers (the pH) barely feel them. A huge pothole, however, can overwhelm the shock absorber, just as a large addition exceeds a buffer's capacity.

Real-world example

During a sprint, lactic acid pours into the blood. The hydrogencarbonate buffer converts the extra H⁺ into carbonic acid, and you breathe faster to remove the carbon dioxide produced. Blood pH may dip slightly, but it stays within a safe range.

Why?

Why can a strong acid and its salt, such as hydrochloric acid and sodium chloride, not act as a buffer? Chloride ions have almost no tendency to accept H⁺, because HCl ionises fully. There is no base partner to remove added acid, so the pH would change freely.

Common misconception

"A buffer solution always has pH 7." Buffers can be made at almost any pH. An ethanoic acid buffer has a pH near 4.8, blood is 7.4, and an ammonia–ammonium chloride buffer has a pH near 9. A buffer keeps its own pH steady; it does not force neutrality.

Worked example

Question: A buffer contains ethanoic acid and sodium ethanoate. A small amount of sodium hydroxide is added. Write the ionic equation for the reaction and explain why the pH barely changes.

Reasoning: The alkali supplies OH⁻. The acid partner, ethanoic acid, reacts with it. OH⁻ is converted into water and ethanoate ions, so free OH⁻ does not build up.

Answer: CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O; the added hydroxide is removed, so the pH rises only very slightly.

Quick check

1. Name the ion that acts as the main base partner in the blood buffer. Answer: The hydrogencarbonate ion, HCO₃⁻.

Exam focus

State the definition precisely: a buffer resists change in pH when small amounts of acid or alkali are added. Name the two components (a weak acid and its salt) and write one equation for each direction. Avoid saying a buffer keeps pH "exactly" constant.

Advanced insight

The pH of a buffer depends on the strength of the weak acid and the ratio of acid to salt, not on their absolute concentrations. That is why diluting a buffer changes its pH very little. More concentrated buffers, however, have a larger buffer capacity, because they contain more of each partner to react with added acid or alkali.

Summary

A buffer resists pH change when small amounts of acid or alkali are added. Typical buffers contain a weak acid and its salt: the salt's anion removes added H⁺, and the weak acid removes added OH⁻. Blood is buffered near pH 7.4 by carbonic acid and hydrogencarbonate. Buffers are used in shampoos, foods, medicines, pools and pH meter calibration, but every buffer has a limited capacity.

Practice questions

1. Define a buffer solution. Answer: A solution whose pH changes very little when small amounts of acid or alkali are added. 2. Write the equation showing how ethanoate ions remove added hydrogen ions. Answer: CH₃COO⁻ + H⁺ → CH₃COOH 3. Explain why a buffer eventually stops working if a large amount of acid is added. Answer: The base partner (such as ethanoate) is used up, so there is nothing left to remove further H⁺, and the pH then falls sharply. 4. Why are many shampoos buffered to about pH 5.5? Answer: To match the slightly acidic pH of skin and hair, so the product does not irritate skin or eyes and stays stable during storage.