Biology · The Chemistry of Life

Acidity, pH and buffering in living systems

Two solutions can begin at the same pH and respond differently when a small amount of acid is added. The starting reading describes one property; resistance to change describes another. A buffer helps us see why a single number does not describe an entire chemical environment.

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By the end, you can

Interpret pH changes and explain why equal pH does not mean equal resistance to added acid or base.

Acid and base describe a transfer

In the proton-transfer description used here, an acid can donate a proton and a base can accept one. H⁺ is shorthand for hydrogen ions in water, where protons are hydrated rather than bare particles floating independently. H₃O⁺ is one useful representation.

Hydrogen-ion conditions can affect the charge state of molecular groups. A group that accepts or loses a proton may interact differently with water or neighbouring groups. This connects acidity to molecular behaviour without implying that every cell or compartment has the same pH.

Acid strength and concentration are different properties. Calling an acid weak refers to its tendency to donate protons in the stated medium, not to having only a small amount present. A pH reading therefore cannot, by itself, identify the acid, its strength or the total acid concentration.

A pH step represents a factor of ten

Formally, pH is the negative base-ten logarithm of hydrogen-ion activity. Activity accounts for chemical behaviour beyond a simple particle concentration. In a sufficiently dilute ideal approximation, we use pH ≈ −log₁₀(c_H/(1 mol L⁻¹)), where c_H is hydrogen-ion concentration.

You do not need logarithms to read our three-row model. A concentration of 10⁻⁶ mol L⁻¹ corresponds to pH 6; 10⁻⁷ to pH 7; 10⁻⁸ to pH 8. Each upward pH step means a tenfold lower activity, or approximately tenfold lower concentration under this model.

In dilute aqueous conditions at 25 °C, neutrality corresponds approximately to pH 7. The neutral condition relates hydrogen-ion and hydroxide-ion conditions, and its numerical pH changes with temperature. 'Neutral always means pH 7' leaves out a real condition.

Equal pH steps correspond to multiplicative changes

Concentration correspondence is an ideal dilute approximation. The formal pH quantity uses activity.

Read the diagram in words

At the same stated conditions, pH 6 to 7 corresponds approximately to a tenfold fall in hydrogen-ion concentration; pH 6 to 8 to a hundredfold fall. pH has no concentration unit.

Equal pH steps correspond to multiplicative changes
pH	mol L⁻¹
6	10⁻⁶ mol L⁻¹
7	10⁻⁷ mol L⁻¹
8	10⁻⁸ mol L⁻¹

One pH step upward: activity divided by ten

A buffer responds rather than staying perfectly unchanged

A common buffer contains appreciable amounts of a weak acid, written HA, and its conjugate base, A⁻. These names refer to two forms differing by a proton. Added hydrogen ions can be taken up by A⁻ to form HA; HA can supply protons when added base removes them.

The mixture resists a pH change within limits. Its pH can still move, and sufficient acid or base can overwhelm its useful capacity. Buffering is not a guarantee of perfect constancy and does not make a solution automatically neutral.

Equal starting pH does not establish equal buffering capacity. Solutions with different total amounts of the buffer components can begin at similar pH but tolerate different additions. To compare them, specify composition, volume, temperature and what is added.

The two buffer forms respond to different additions

HA is an uncharged monoprotic weak acid in this schematic, and H⁺ is hydrated-ion shorthand. The arrows describe net proton transfer without giving an exact final pH.

Read the diagram in words

The conjugate base can take up added hydrogen ions. The weak acid can donate a proton to neutralise added hydroxide. The supply of these forms is finite, so buffering has limits.

Added acid

A⁻ + H⁺ → HA

Some conjugate base becomes acid

Added base

HA + OH⁻ → A⁻ + H₂O

Some acid becomes conjugate base

Worked example

Compare pH 6 with pH 8

Under the stated dilute approximation, pH 6 corresponds to 0.000001 mol L⁻¹ and pH 8 to 0.00000001 mol L⁻¹. Divide the first by the second: the ratio is 100. A two-unit pH increase corresponds to a hundredfold lower hydrogen-ion concentration in this model, not a doubling of some general quantity called 'acidity'. This calculation tells us neither the total amount of all acids nor the buffering capacity.

Check your understanding

In the dilute model, how do hydrogen-ion concentrations at pH 7 and pH 6 compare?
Show reasoning
The pH 7 concentration is one tenth of the pH 6 concentration.
Two solutions have the same pH. Must they resist added acid equally?
Show reasoning
No. Buffer composition and the amounts present can differ. A pH reading alone does not measure buffering capacity.
Does a buffer have to keep pH exactly fixed or at seven?
Show reasoning
Neither. It resists change over a useful range; its pH may be acidic or basic, and its capacity is finite.

What to carry forward

Separate hydrogen-ion activity, total composition and resistance to pH change.

References

Resources supporting this lesson’s explanations and examples.