R3.1.16 Buffer solutions (Higher Level Only)

Buffers: Definition, Components, and pH Calculation

You’re conducting a chemical reaction in the lab, and the reaction environment requires a constant pH.
What happens if you accidentally add a small amount of acid or base to the solution?

Without a buffer, the pH would change dramatically, potentially ruining your experiment.

What Are Buffers?

Definition

Buffers

Buffers are solutions that resist changes in pH when small amounts of acid ( $H ⁺$ ions) or base ( $O H ⁻$ ions) are added.

Example

They play a critical role in maintaining pH stability in chemical reactions, biological systems (like blood), and industrial processes.

To understand how buffers work, consider this:

If you add a strong acid to pure water, the pH drops significantly because there’s nothing in the water to counteract the influx of $H ⁺$ ions.
A buffer, however, "absorbs" most of these $H ⁺$ ions, minimizing the pH change.

Analogy

Think of a buffer as a sponge. Just as a sponge absorbs water to prevent spills, a buffer absorbs excess $H ⁺$ or $O H ⁻$ ions, preventing drastic pH changes.

Components of a Buffer

Every buffer consists of two key components:

A weak acid and its conjugate base(e.g., $C H ₃ C O O H$ / $C H ₃ C O O ⁻$ ).
A weak base and its conjugate acid(e.g., $N H ₃$ / $N H ₄ ⁺$ ).

Example

Acetic Acid/Sodium Acetate Buffer

Weak acid: Acetic acid ( $C H ₃ C O O H$ )
Conjugate base: Acetate ion ( $C H ₃ C O O ⁻$ )

When a small amount of strong acid (e.g., HCl) is added, the acetate ion ( $C H ₃ C O O ⁻$ ) reacts with the $H ⁺$ ions, forming more acetic acid:
$CH₃COO⁻ (aq) + H⁺ (aq) \to CH₃COOH (aq)$

When a small amount of strong base (e.g., NaOH) is added, the acetic acid reacts with the $O H ⁻$ ions, forming water and more acetate ions:
$CH₃COOH (aq) + OH⁻ (aq) \to CH₃COO⁻ (aq) + H₂O (l)$

Example

Ammonia/Ammonium Chloride Buffer

Weak base: Ammonia ( $N H ₃$ )
Conjugate acid: Ammonium ion ( $N H ₄ ⁺$ )

When a strong acid is added, the ammonia reacts with the $H ⁺$ ions:
$NH₃ (aq) + H⁺ (aq) \to NH₄⁺ (aq)$

When a strong base is added, the ammonium ion ( $N H ₄ ⁺$ ) reacts with the $O H ⁻$ ions:
$NH₄⁺ (aq) + OH⁻ (aq) \to NH₃ (aq) + H₂O (l)$

Note

In both cases, the buffer absorbs the added acid or base, keeping the pH relatively constant.

Tip

Buffers work best when the concentrations of the weak acid and its conjugate base (or weak base and its conjugate acid) are roughly equal.

Applications of Buffers

Buffers are used in a wide range of real-world applications:

Biological Systems: Blood is buffered by the carbonic acid/bicarbonate system to maintain a pH of 7.35–7.45, ensuring proper cellular function.
Pharmaceuticals: Buffers stabilize the pH of medicines to enhance their effectiveness and shelf life.
Industrial Processes: Buffers are used in dyeing fabrics, fermentation, and electroplating to maintain optimal pH conditions.

Reflection

Self review

What are the two components of a buffer system?

Theory of Knowledge

In biology, the buffering capacity of blood is vital for homeostasis. How might the failure of this buffering system affect the human body?
What ethical considerations arise when developing artificial blood substitutes to mimic this buffering system?

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R3.1.16 Buffer solutions (Higher Level Only)

Buffers: Definition, Components, and pH Calculation

What Are Buffers?

Components of a Buffer

Acetic Acid/Sodium Acetate Buffer

Ammonia/Ammonium Chloride Buffer

Applications of Buffers

Reflection

You've read 2/2 free chapters this week.

Buffer Solutions

Examples