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24 Dec 2025

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4 pages

How to Make a Buffer Solution for A Level Chemistry - Easy Guide

sheila

@study_withsheila

Buffer solutions are essential in maintaining stable pH levels in... Show more

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21.1- Buffer solutions

Keu facts:
- buffer solution
= a system that
minimises PH
Changes when smau
amounts of an acid /
alkali are added

-

Preparing Buffer Solutions

This section discusses two approaches to prepare buffer solutions in the lab, which is an essential skill for A Level Chemistry students.

Approach 1: Mixing Solutions

The first method involves mixing a solution of a weak acid with a solution containing the salt of the weak acid.

Example: To create an acetate buffer, you might mix a solution of acetic acid (CH3COOH) with a solution of sodium acetate (CH3COONa).

Approach 2: Weak Neutralization

The second method involves adding an aqueous solution of alkali to an excess of a weak acid.

Highlight: In this approach, the weak acid is partially neutralized, resulting in a solution containing both unreacted HA and the A- salt.

Example: To prepare a phosphate buffer, you could add a calculated amount of sodium hydroxide (NaOH) to a solution of phosphoric acid (H3PO4).

Both methods result in a solution containing a weak acid and its conjugate base, which is the essence of a buffer solution.

Vocabulary: Conjugate base - The species formed when an acid donates a proton. In the case of acetic acid (CH3COOH), the conjugate base is the acetate ion $CH3COO-$ .

Calculating pH of Buffer Solutions

Understanding how to calculate the pH of buffer solutions is crucial for A Level Chemistry students, particularly when dealing with buffer calculations in AQA A Level Chemistry questions.

The calculation of buffer pH involves the use of the Henderson-Hasselbalch equation, which is derived from the acid dissociation constant (Ka) expression:

Ka = $H+$ $A-$ / $HA$

Rearranging this equation gives:

$H+$ = Ka × $HA$ / $A-$

Highlight: The value of Ka provides the 'coarse tuning' of the buffer's pH, while the ratio of $HA$ to $A-$ allows for 'fine tuning'.

Key points to remember:

The pH of a buffer depends on the Ka value and the ratio of concentrations of HA and A-
When $HA$ = $A-$ , pH = pKa
For more precise calculations, use the actual concentrations rather than assuming a 1:1 ratio

Example: For a buffer made from a weak acid with pKa = 4.7, if $HA$ = $A-$ , the pH would be 4.7.

Vocabulary: pKa - The negative logarithm of the acid dissociation constant (Ka). It's a measure of the strength of an acid in solution.

Practical Example: Calculating Buffer pH

This section provides a practical example of how to calculate the pH of a buffer solution, which is a common type of question in A Level Chemistry buffer calculations.

Example: Ethanoic acid buffer

Given:

125 cm³ of 0.8 mol dm⁻³ ethanoic acid (CH3COOH)
75 cm³ of 1 mol dm⁻³ sodium ethanoate (CH3COONa)
Ka of ethanoic acid = 1.74 × 10⁻⁵ mol dm⁻³

Step 1: Calculate new concentrations after mixing

$CH3COOH$ = (0.8 × 125) / 200 = 0.5 mol dm⁻³ $CH3COO⁻$ = (1 × 75) / 200 = 0.375 mol dm⁻³

Step 2: Use the equation $H+$ = Ka × $HA$ / $A-$

$H+$ = (1.74 × 10⁻⁵ × 0.5) / 0.375 = 2.32 × 10⁻⁵ mol dm⁻³

Step 3: Calculate pH

pH = -log(2.32 × 10⁻⁵) = 4.63

Highlight: This example demonstrates how to prepare a buffer solution with a specific pH and how to calculate pH of buffer solution given molarity and volume.

Example: If you needed to prepare a pH 4 buffer solution, you would adjust the ratio of acid to conjugate base to achieve this pH.

Understanding these calculations is crucial for mastering buffers in A Level Chemistry OCR and other exam boards.

Understanding Buffer Solutions

Definition: A buffer solution is a system that minimizes pH changes when small amounts of an acid or alkali are added.

The key components of a buffer solution are:

A weak acid (HA)
Its conjugate base $A-$

Highlight: Buffer solutions require a large reservoir of both HA and A- to effectively resist pH changes.

The mechanism of buffer action can be understood in two scenarios:

When acid is added:
- H+ ions are introduced to the solution
- The large reservoir of A- reacts with excess H+
- The equilibrium shifts to the left
- $H+$ and pH are restored
When alkali is added:
- OH- ions are introduced
- H+ + OH- → H2O reaction occurs
- $H+$ decreases
- HA dissociates
- The equilibrium shifts to the right
- $H+$ and pH are restored

Example: In a buffer solution containing acetic acid (CH3COOH) and acetate ion $CH3COO-$ , adding a small amount of HCl would cause the acetate ions to neutralize the added H+ ions, maintaining the pH.

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4 pages

How to Make a Buffer Solution for A Level Chemistry - Easy Guide

sheila

@study_withsheila

Buffer solutions are essential in maintaining stable pH levels in various chemical and biological systems. This summary explores the concept of buffer solutions, their preparation, and pH calculations in A Level Chemistry. It covers key aspects of how to... Show more

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Preparing Buffer Solutions

This section discusses two approaches to prepare buffer solutions in the lab, which is an essential skill for A Level Chemistry students.

Approach 1: Mixing Solutions

The first method involves mixing a solution of a weak acid with a solution containing the salt of the weak acid.

Example: To create an acetate buffer, you might mix a solution of acetic acid (CH3COOH) with a solution of sodium acetate (CH3COONa).

Approach 2: Weak Neutralization

The second method involves adding an aqueous solution of alkali to an excess of a weak acid.

Highlight: In this approach, the weak acid is partially neutralized, resulting in a solution containing both unreacted HA and the A- salt.

Example: To prepare a phosphate buffer, you could add a calculated amount of sodium hydroxide (NaOH) to a solution of phosphoric acid (H3PO4).

Both methods result in a solution containing a weak acid and its conjugate base, which is the essence of a buffer solution.

Vocabulary: Conjugate base - The species formed when an acid donates a proton. In the case of acetic acid (CH3COOH), the conjugate base is the acetate ion $CH3COO-$ .

Sign up to see the contentIt's free!

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Calculating pH of Buffer Solutions

Understanding how to calculate the pH of buffer solutions is crucial for A Level Chemistry students, particularly when dealing with buffer calculations in AQA A Level Chemistry questions.

The calculation of buffer pH involves the use of the Henderson-Hasselbalch equation, which is derived from the acid dissociation constant (Ka) expression:

Ka = $H+$ $A-$ / $HA$

Rearranging this equation gives:

$H+$ = Ka × $HA$ / $A-$

Highlight: The value of Ka provides the 'coarse tuning' of the buffer's pH, while the ratio of $HA$ to $A-$ allows for 'fine tuning'.

Key points to remember:

The pH of a buffer depends on the Ka value and the ratio of concentrations of HA and A-

When $HA$ = $A-$ , pH = pKa

For more precise calculations, use the actual concentrations rather than assuming a 1:1 ratio

Example: For a buffer made from a weak acid with pKa = 4.7, if $HA$ = $A-$ , the pH would be 4.7.

Vocabulary: pKa - The negative logarithm of the acid dissociation constant (Ka). It's a measure of the strength of an acid in solution.

Sign up to see the contentIt's free!

Access to all documents

Improve your grades

Join milions of students

By signing up you accept Terms of Service and Privacy Policy

Practical Example: Calculating Buffer pH

This section provides a practical example of how to calculate the pH of a buffer solution, which is a common type of question in A Level Chemistry buffer calculations.

Example: Ethanoic acid buffer

Given:

125 cm³ of 0.8 mol dm⁻³ ethanoic acid (CH3COOH)

75 cm³ of 1 mol dm⁻³ sodium ethanoate (CH3COONa)

Ka of ethanoic acid = 1.74 × 10⁻⁵ mol dm⁻³

Step 1: Calculate new concentrations after mixing

$CH3COOH$ = (0.8 × 125) / 200 = 0.5 mol dm⁻³ $CH3COO⁻$ = (1 × 75) / 200 = 0.375 mol dm⁻³

Step 2: Use the equation $H+$ = Ka × $HA$ / $A-$

$H+$ = (1.74 × 10⁻⁵ × 0.5) / 0.375 = 2.32 × 10⁻⁵ mol dm⁻³

Step 3: Calculate pH

pH = -log(2.32 × 10⁻⁵) = 4.63

Highlight: This example demonstrates how to prepare a buffer solution with a specific pH and how to calculate pH of buffer solution given molarity and volume.

Example: If you needed to prepare a pH 4 buffer solution, you would adjust the ratio of acid to conjugate base to achieve this pH.

Understanding these calculations is crucial for mastering buffers in A Level Chemistry OCR and other exam boards.

Sign up to see the contentIt's free!

Access to all documents

Improve your grades

Join milions of students

By signing up you accept Terms of Service and Privacy Policy

Understanding Buffer Solutions

Buffer solutions play a crucial role in maintaining stable pH levels in various chemical and biological systems. This section explores the fundamental concepts of buffer solutions and their mechanisms.

Definition: A buffer solution is a system that minimizes pH changes when small amounts of an acid or alkali are added.

The key components of a buffer solution are:

A weak acid (HA)

Its conjugate base $A-$

Highlight: Buffer solutions require a large reservoir of both HA and A- to effectively resist pH changes.

The mechanism of buffer action can be understood in two scenarios:

When acid is added:

H+ ions are introduced to the solution

The large reservoir of A- reacts with excess H+

The equilibrium shifts to the left

$H+$ and pH are restored

When alkali is added:

OH- ions are introduced

H+ + OH- → H2O reaction occurs

$H+$ decreases

HA dissociates

The equilibrium shifts to the right

$H+$ and pH are restored

Example: In a buffer solution containing acetic acid (CH3COOH) and acetate ion $CH3COO-$ , adding a small amount of HCl would cause the acetate ions to neutralize the added H+ ions, maintaining the pH.