AP Chemistry 8.10 Buffer Capacity Study Notes - New Syllabus Effective fall 2024
AP Chemistry 8.10 Buffer Capacity Study Notes- New syllabus
AP Chemistry 8.10 Buffer Capacity Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship between the buffer capacity of a solution and the relative concentrations of the conjugate acid and conjugate base components of the solution.
Key Concepts:
- Buffer Capacity
Buffer Capacity
- Buffer capacity: represents the amount of H+ or OH- the buffer can absorb without a significant change in pH
- The pH of a buffered solution is determined by the ratio of [A-]/[HA]
- Same ratio = same pH
- The buffer capacity is determined by the [A-] & [HA]
- More moles = greater capacity
- Less moles = less capacity bcuz there is less buffer to react with
How to Choose a Buffer
- The most effective buffers (most resistant to change) have a ratio of [A-]/[HA] = 1
- Is true when [A-] = [HA] → This means that pH = pKa (since log 1 = 0)
- When choosing a buffer for a desired pH, choose the buffer system whose pKa is closest to the desired pH
- The log of a larger number is less negative than the log of a smaller number → higher pH
8.10.A.1 Effect of Increasing Buffer Concentration on pH and Capacity:
1. Buffer Definition:
A buffer is a unique solution that resists a change in pH, even when a small quantity of an acid or a base is introduced.
i. How it works:
* A buffer consists of a weak acid and its conjugate base, or a weak base and its conjugate acid.
* The weak acid can neutralize added base.
* The conjugate base will also react with excess acid to neutralize it.
ii. Example:
One of the most familiar buffers is acetic acid (CH₃COOH) and sodium acetate (CH₃COONa)
* If you pour a little strong acid (such as HCl) in, the acetate ion (CH₃COO⁻) reacts with the H⁺ to create more acetic acid, maintaining the pH level.
* When you introduce a minimal amount of concentrated base (such as NaOH), the acetic acid transfers a proton (H⁺) to neutralize the OH⁻, once more maintaining the pH level constant.
iii. Summary:
Buffers are vital in most chemical and biological systems since they ensure that the pH remains constant, which is key to proper chemical reactions, particularly in living organisms.
2. Henderson-Hasselbalch Equation:
This equation shows the relationship between the pH of a buffer solution and the ratio of the concentrations of the conjugate base and the weak acid.
Key Point:
If the ratio stays the same, the pH remains constant, even if the absolute concentrations of the base and acid increase.
However, increasing these concentrations improves the buffer’s capacity, allowing it to neutralize more added acid or base without a significant change in pH.
3. Buffer Capacity: Buffer capacity is a measure of the amount of acid or base that will be neutralized by a buffer before there is any significant change in pH.
Key Point:
Increased concentrations of the buffer’s components (acid and conjugate base) translate to more species available to react with acids or bases added.
* This raises the buffer capacity.
* Notably, as long as the ratio between base and acid remains constant, the pH does not alter, even though the buffer gets stronger (more resistant to pH changes).
8.10.A.2 Buffer Capacity Depends on Component Proportions:
1. Neutralization Reactions:
In a buffer solution, the constituents react with added bases or acids to neutralize the same and sustain the pH. Here’s what happens:
i. When Acid (H⁺) is Added:
* The conjugate base (A⁻) combines with the added H⁺ to neutralize it:
* The conjugate base takes up the H⁺ ions, producing more of the weak acid (HA). This will not make the pH fall too much.
ii. When Base (OH⁻) is Added:
* The weak acid (HA) neutralizes the added OH⁻ as follows:
* The weak acid donates H⁺ ions, neutralizing the OH⁻ and creating additional conjugate base (A⁻). This keeps the pH from increasing substantially.
iii. Result:
* The buffer components react with acids or bases and move the equilibrium to consume the excess H⁺ or OH⁻, maintaining a relatively stable pH.
2. Conjugate Acid-Base Roles:
i. Conjugate Base Neutralizes Acid (H⁺):
* In a buffer solution, the conjugate base (A⁻) neutralizes added acids (H⁺) through the following reaction:
* The conjugate base (A⁻) “sops up” surplus H⁺ ions, creating the weak acid (HA). This keeps the pH from falling too far when acid is added.
ii. Conjugate Acid Neutralizes Base (OH⁻):
* Likewise, the conjugate acid (HA) combines with added bases (OH⁻) to neutralize them:
* The conjugate acid (HA) releases H⁺ ions to neutralize the OH⁻ and produce water (H₂O) and the conjugate base (A⁻). This keeps the pH from rising too high when you add base.
3. Effect of Component Ratio:
i. More Conjugate Acid (HA) → Better at Neutralizing Base (OH⁻):
* If the buffer contains more conjugate acid (HA),it is able to donate more H⁺ ions.
* This facilitates neutralizing added base (OH⁻) to produce water and the conjugate base (A⁻).
* As an example, if in a buffer solution of acetic acid (HA) and acetate (A⁻) you add a base such as NaOH, the acetic acid (HA) will donate H⁺ to neutralize the OH⁻:
ii. More Conjugate Base (A⁻) → Better at Neutralizing Acid (H⁺):
* When the buffer contains more conjugate base (A⁻), it can absorb more H⁺ ions.
* This neutralizes added acid (H⁺) by producing more weak acid (HA).
* For instance, in the same acetic acid/acetate buffer, if you add an acid such as HCl, the acetate (A⁻) will react with H⁺ to neutralize it: