7.2.11 Buffer Solutions - Theory

What is a Buffer?

Definition: Buffer

A buffer solution is a solution that resists changes in pH when small amounts of acid or alkali are added.

Key Terminology

• “Resists Change”: It does NOT keep pH constant. The pH changes slightly, but drastically less than if the buffer was not there.
• “Small amounts”: Every buffer has a “limit”. If you dump too much acid/alkali in the buffer, it will be overwhelmed.

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Types of Buffers

There are two main ways to make a buffer solution, depending on whether you want to stabilise an acidic pH or an alkaline pH.

A. Acidic Buffer (pH < 7)

Made from a Weak Acid and its Conjugate Base (Salt).

• Recipe: Ethanoic Acid () + Sodium Ethanoate ().

B. Basic Buffer (pH > 7 )

Made from a Weak Baseand its Conjugate Acid (Salt).

• Recipe: Ammonia () + Ammonium Chloride ().

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Acidic Buffer Mechanism

Setup

Consider a buffer containing a high concentration of Ethanoic Acid () and Ethanoate ions ():

• Ethanoic acid () is a Weak Acid and partially ionises in solution to form a relatively low concentration of ethanoate ions ().

• Sodium Ethanoate () is a salt which fully ionises in solution to form a relatively high concentration of ethanoate ions ().

These form the reserve supplies of the acid () and its conjugate base ().

A buffer solution contains:

• Relatively high concentrations of acid (), due to the partial ionisation of the weak acid.
• Relatively high concentrations of the acid’s conjugate base (), due to the complete ionisation of the salt ().

This leads to an equilibrium between the acid () and its conjugate base ():

Scenario 1: Adding Acid ()

• Disturbance: is added to the buffer solution, concentration of rises.
• Response: Position of equilibrium shifts to the left.
  • The excess ions react with the to form until equilibrium is re-established.

• Result: The added is consumed and turned into weak acid. The pH drops only very slightly.

Scenario 2: Adding Alkali ()

• Disturbance: is added to the buffer solution, the reacts with the to form water, concentration of drops.
• Response: Position of equilibrium shifts to the right.
  • The acid ionises to form more and until equilibrium is re-established.

• As there is a large supply of , the change in concentration of is negligible when it dissociates to form more ions.
  
• As there is a large supply of , the change in concentration of is negligible when more is formed by the dissociation of .
  
• Result: The concentration is restored. The pH rises only very slightly.

Warning: pH does not remain the same

• When a strong acid or alkali is added to a buffer solution, the pH does not come back exactly the same as it was before.
• This is because a buffer solution does not stop changes in pH, it resists.

Scenario 1: is added

• Position of equilibrium shifts to the left:
• Concentration weak acid () increases.
• Concentration of conjugate base () decreases.
• According to the Henderson-Hasselbalch Equation:
  • remains constant for a given weak acid.
  • decreases.
  • increases.
  • The logarithmic term will be smaller than it previously was, as the ratio becomes smaller.
  • Therefore, the pH will slightly decrease.

Scenario 2: is added

• Position of equilibrium shifts to the right:
• Concentration weak acid () decreases.
• Concentration of conjugate base () increases.
• According to the Henderson-Hasselbalch Equation:
  • remains constant for a given weak acid.
  • increases.
  • decreases.
  • The logarithmic term will be larger than it previously was, as the ratio becomes larger.
  • Therefore, the pH will slightly increase.

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Alkaline Buffer Mechanism

Setup

Consider a buffer containing a high concentration of Ammonia () and ammonium ions ():

• Ammonia is a Weak Base and partially ionises in solution to form a relatively low concentration of ammonium ions ().

• Ammonium chloride () is a salt which fully ionises in solution to form a relatively high concentration of ammonium ions ().

These form the reserve supplies of the base () and its conjugate acid ().

A buffer solution contains:

• Relatively high concentrations of base () due to the partial ionisation of the base.
• Relatively high concentrations of the base’s conjugate acid (), due to the complete ionisation of the salt ().

This leads to an equilibrium between the base () and its conjugate acid ():

Scenario 1: Adding Acid ()

• Disturbance: is added to the buffer solution, the concentration of rises.
• Response: Position of equilibrium shifts to the right.
  • The excess ions react with the to form until equilibrium is re-established.

• Result: The added is consumed and turned into the conjugate acid of the weak base (). The pH drops only very slightly.

Why doesn't the react with instead?

Some of the does react with the , however:

• The concentration of is negligible compared to concentration of .
  • For example, while the concentration of is usually around 0.1 to 1.0 M, at a pH of 9 (weak alkali), the pOH is 5, meaning the concentration of is about M.
• It is statistically more likely for the to find a molecule to react with than a .
• When some of the do react with the , the equilibrium shifts to replace them by the dissociation of .

Scenario 2: Adding Alkali ()

• Disturbance: is added to the buffer solution. The concentration of rises.
• Response: Position of equilibrium shifts to the left.
  • The conjugate acid () acts as an acid, it donates a proton () to the incoming , turning it into water.

• Result: The added is turned into water. The pH rises only very slightly.

Warning: pH does not remain the same

• When a strong acid or alkali is added to a buffer solution, the pH does not come back exactly the same as it was before.
• This is because a buffer solution does not stop changes in pH, it resists.

Scenario 1: is added

• The weak base reacts with the protons:
• Concentration of conjugate acid () increases.
• Concentration of weak base () decreases.
• According to the Henderson-Hasselbalch Equation:
  • remains constant for the ammonium ion.
  • () decreases.
  • () increases.
  • The logarithmic term will be smaller than it previously was, as the ratio becomes smaller.
  • Therefore, the pH will slightly decrease.

Scenario 2: is added

• The conjugate acid neutralizes the hydroxide:
• Concentration of conjugate acid () decreases.
• Concentration of weak base () increases.
• According to the Henderson-Hasselbalch Equation:
  • remains constant for the ammonium ion.
  • () increases.
  • () decreases.
  • The logarithmic term will be larger than it previously was, as the ratio becomes larger.
  • Therefore, the pH will slightly increase.

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Use of Buffers: Blood

Context

The pH of human blood plasma must be tightly maintained between 7.35 and 7.45.

• If pH < 7.35: Acidosis (leads to fatigue, shortness of breath, shock).
• If pH > 7.45: Alkalosis (leads to muscle spasms, nausea).

To survive, the body uses the Hydrogencarbonate ion () as the conjugate base and Carbonic Acid () as the weak acid.

Warning: Blood uses an acidic buffer system and not alkaline.

• Given that blood pH need to be between 7.35 and 7.45, which are both above 7, you would assume that an alkaline buffer would be needed.

• However an important distinction needs to be made:

Acidic buffer does not always mean pH less than 7

• An acidic buffer is simply a buffer made from a Weak Acid + Conjugate Base.
• It can buffer any pH depending on the of the acid and the Ratio of .

The Mathematical Proof (Henderson-Hasselbalch Equation)

• The pH of a buffer is determined by:

For the blood buffer (Carbonic Acid):

• : The of Carbonic Acid () is roughly 6.1.
• If [Base] = [Acid]: , so (Too acidic for life!)

How does the blood pH get up to 7.4?

• The body maintains a massive imbalance in the acid and conjugate base concentrations.
• It keeps the concentration of Bicarbonate (base) roughly 20 times higher than Carbonic Acid.

Conclusion An acidic buffer system can maintain an alkaline pH (7.4) simply by having a huge excess of the conjugate base relative to the weak acid.

Formation of the Buffer Mixture

The buffer components are naturally generated from cellular respiration.

1. Respiration: Cells produce as a waste product.
2. Dissolution: dissolves in blood plasma and reacts with water.
   • This reaction is catalysed by the enzyme Carbonic Anhydrase

3. Dissociation: The carbonic acid is a weak acid, so it partially dissociates to establish the buffer equilibrium.

The "Effective" Equilibrium

Since is unstable, the combined equilibrium equation is often preferred:

Regulation of Blood pH

The system relies on the large reservoir of (to soak up acid) and the ability of the lungs to remove (to remove acid).

Scenario 1: Excess Acid (e.g. Lactic Acid from exercise)

• Disturbance: Concentration of in blood rises.
• Response: Position of equilibrium shifts to the left.
  • The conjugate base accepts protons:
• Regulation: The excess produced is transported to the lungs and exhaled. This prevents the acid form () from building up and shifting the equilibrium back.

Scenario 2: Excess Alkali (rare, e.g. ingesting alkaline drugs)

• Disturbance: Excess reacts with to form water, concentration of in blood drops.
• Response: Position of equilibrium shifts to the right.
  • Carbonic acid dissociates to restore the lost .
• Regulation: The kidneys help remove the excess in urine if necessary.