Buffers

Corn Hill, Edward Hopper, 1930

Buffers

A buffer is a chemical system that includes:

• a weak acid and its conjugate base

or

• a weak base and its conjugate acid

Why Buffers?

• Buffers help maintain the pH of a chemical system if external acid or base enters that system.
• Added acid will be neutralized by the base in the buffer.

Why Buffers?

• Buffers help maintain the pH of a chemical system if external acid or base enters that system.
• Added acid will be neutralized by the base in the buffer.
• Added base will be neutralized by the acid in the buffer.

Buffer Example

Blood must maintain a pH of about 7.4 and so contains a buffer composed of H₂CO₃ (a weak acid) and HCO₃^- (its conjugate base).

​

H₂CO₃(aq) + H₂O(l) ⇌ HCO₃^-(aq) + H₃O⁺(aq)

​

1. If acid enters the blood stream (e.g., lactic acid from muscle exertion), what component of the buffer system will neutralize it?

HA(aq) + HCO₃^-(aq) → A^-(aq) + H₂CO₃(aq)

(and then a new equilibrium in the buffer is established)

Buffer Example

Blood must maintain a pH of about 7.4 and so contains a buffer composed of H₂CO₃ (a weak acid) and HCO₃^- (its conjugate base).

​

H₂CO₃(aq) + H₂O(l) ⇌ HCO₃^-(aq) + H₃O⁺(aq)

​

b) If base enters the blood stream, what component of the buffer system will neutralize it?

A^-(aq) + H₂CO₃(aq) → HA(aq) + HCO₃^-(aq)

(and then a new equilibrium in the buffer is established)

Identifying Possible Buffer Solutions

Which of the following combinations form a buffer?

• HF and NaF
• Yes, HF + H₂O ⇌ F^- + H₃O⁺ (forms an equilibrium of an acid and its conjugate base)
• HCl and NaCl
• No, HCl + H₂O → Cl^- + H₃O⁺ (strong acids and bases do not form equilibria and therefore can not form buffer systems)
• HNO₂ and HCl
• No, both are acids
• NH₄Cl and NH₃
• Yes, NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺ (forms an equilibrium of an acid and its conjugate base)

Buffer Capacity

Buffer Capacity is the amount of acid or base that a buffer can absorb without significantly changing the pH.

• Of course, buffers are limited in their ability to neutralize added acid or base by the amounts of the components of the buffer present.
• Higher concentrations of buffer components = higher buffering capacity
• Similar concentrations of buffer components = higher buffering capacity of both added acids and bases

Buffer Capacity

Buffer Capacity

Example:

A 1.0 L buffer solution is 0.10 M in HF and 0.050 M in NaF. Which of the following actions will destroy the buffer?

HF (aq) + H₂O(l) ⇌ F^-(aq) + H₃O⁺(aq)

​

a) Adding 0.050 mol of HCl?

​

Yes, the moles of added acid are equal to the moles of the buffer’s base and will eliminate all of the base present in the buffer.

Buffer Capacity

Example:

A 1.0 L buffer solution is 0.10 M in HF and 0.050 M in NaF. Which of the following actions will destroy the buffer?

HF (aq) + H₂O(l) ⇌ F^-(aq) + H₃O⁺(aq)

​

b) Adding 0.050 mol of NaOH?

​

No, this is fewer moles of NaOH than moles of acid present in the buffer.

Buffer Capacity

Example:

A 1.0 L buffer solution is 0.10 M in HF and 0.050 M in NaF. Which of the following actions will destroy the buffer?

HF (aq) + H₂O(l) ⇌ F^-(aq) + H₃O⁺(aq)

​

c) Adding 0.050 mol of NaF?

​

No, this will simply add more of the conjugate base to the buffer.

Buffer Capacity

Example:

A 1.0 L buffer solution is 0.10 M in HF and 0.050 M in NaF. Which of the following actions will destroy the buffer?

HF (aq) + H₂O(l) ⇌ F^-(aq) + H₃O⁺(aq)

​

d) Does this buffer system have a higher buffering capacity for added acids or added bases?

​

Bases because the buffer contains more of the acid than its conjugate base.

Calculating the pH of a Buffer Solution

The pH of a buffer solution is affected by the common ion effect. Both the acid and the base have the same ion and therefore equilibrium is affected by its presence on both sides of the equilibrium.

Calculating the pH of a Buffer Solution

i) Write the ionization reaction for the acid:

HF(aq) + H₂O(l) ⇌ F^-(aq) + H₃O⁺(aq)

​

Example: What is the pH of a buffer solution that is 0.150 M in HF and 0.130 M in KF? (K_a of HF is 7.2x10^-4)

Calculating the pH of a Buffer Solution

ii) Fill in the RICE table with the given concentrations of each species.

(Can eliminate due to large concentrations and small K_a value)

Example: What is the pH of a buffer solution that is 0.150 M in HF and 0.130 M in KF? (K_a of HF is 7.2x10^-4)

Calculating the pH of a Buffer Solution

iii) Substitute equilibrium concentrations in K_a equation and solve for x.

x = [H₃O⁺] = 8.3x10^-4 M

Example: What is the pH of a buffer solution that is 0.150 M in HF and 0.130 M in KF? (K_a of HF is 7.2x10^-4)

Calculating the pH of a Buffer Solution

iv) Calculate the pH using [H₃O⁺].

pH = -log[8.3x10^-4] = 3.08

x = [H₃O⁺] = 8.3x10^-4 M

Example: What is the pH of a buffer solution that is 0.150 M in HF and 0.130 M in KF? (K_a of HF is 7.2x10^-4)

An easier way…

The Henderson-Hasselbalch Equation:

HA(aq) + H₂O(l) ⇌ A^-(aq) + H₃O⁺(aq)

​

​

Let’s do the previous problem again, this time using the H-H equation:

Notes about using the H-H Equation:

Notes about pK_a

• K_a and pK_a are inversely related
• Larger K_a values = stronger acid
• Larger pK_a values = weaker acid
• A buffer is most effective when the pK_a is near the desired pH.

Example

1. What is the strongest acid in the list of acids to the right?
• HIO₃
2. What acid in the list of acids to the right has the strongest conjugate base?
• HCN
3. What acid in the list of acids to the right would be best for creating a buffer solution with a pH of 6.5?
• carbonic (1)

One last practice

What if two solutions are combined to make a buffer solution?

​

Try this:

What is the pH of a buffer solution that is 0.50 M in HC₂H₃O₂ and 0.40 M in NaC₂H₃O₂?

​