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1. Introduction to Acids, Bases and Buffers

Acids and bases are among the most important classes of chemical substances. Their reactions occur in laboratories, biological systems, industrial processes, environmental chemistry and analytical chemistry.

The terms acid and base were originally associated with characteristic properties such as sour taste, reaction with metals, neutralization and changes in indicator colour. Modern theories provide a more useful explanation based on hydrogen ions, proton transfer and electron-pair transfer.

A buffer solution is a solution that resists a large change in pH when a small amount of acid or base is added.

2. Arrhenius Theory of Acids and Bases

The Arrhenius theory explains acids and bases mainly in aqueous solutions.

Arrhenius Acid

An Arrhenius acid is a substance that increases the concentration of hydrogen ions, H⁺, in aqueous solution.

In water, the hydrogen ion does not exist as a completely free particle. It is associated with water molecules, commonly represented as hydronium ion.

Arrhenius Base

An Arrhenius base is a substance that increases the concentration of hydroxide ions in aqueous solution.

Neutralization

An acid and a base can react to form salt and water.

Limitations

• It is mainly applicable to aqueous solutions.
• It defines bases in terms of producing OH⁻.
• It does not adequately explain many acid-base reactions in non-aqueous solvents.
• It cannot explain substances such as BF₃ as bases or acids in the broader sense.

3. Brønsted-Lowry Theory

The Brønsted-Lowry theory defines acid-base behaviour in terms of proton transfer.

Example: Reaction of HCl with Water

HCl donates a proton to water, so HCl is the acid. Water accepts the proton, so water is the base.

After donating a proton, HCl becomes Cl⁻. After accepting a proton, H₂O becomes H₃O⁺.

Another Example

NH₃ accepts a proton from water and therefore behaves as a Brønsted-Lowry base.

4. Lewis Theory of Acids and Bases

Lewis theory gives a broader definition of acid-base reactions based on electron pairs.

Example: BF₃ and NH₃

NH₃ has a lone pair of electrons and donates it. Therefore, NH₃ is the Lewis base.

BF₃ accepts the electron pair and therefore acts as the Lewis acid.

Comparison of the Three Theories

5. Conjugate Acid-Base Pairs

When a Brønsted-Lowry acid donates a proton, the species formed is its conjugate base.

When a Brønsted-Lowry base accepts a proton, the species formed is its conjugate acid.

Here:

• HA = acid
• A⁻ = conjugate base
• H⁺ = transferred proton

Example

NH₃/NH₄⁺ is one conjugate acid-base pair.

H₂O/OH⁻ is the other conjugate acid-base pair.

6. Strength of Acids and Bases

Acid strength describes how readily an acid donates a proton. Base strength describes how readily a base accepts a proton.

Strong Acids

Strong acids ionize almost completely in aqueous solution.

Common examples include:

• HCl
• HBr
• HI
• HNO₃
• HClO₄
• H₂SO₄ for its first ionization

Weak Acids

Weak acids ionize only partially and establish an equilibrium with their ions.

Examples include:

• CH₃COOH
• H₂CO₃
• HF
• HCN

Strong and Weak Bases

NaOH and KOH are common strong bases because they dissociate extensively in water.

NH₃ is a common weak base because it reacts only partially with water.

7. Ionization of Weak Acids and Bases

Weak Acid

Consider a weak monoprotic acid HA.

If the initial concentration of HA is C and x mol L⁻¹ ionizes, the equilibrium concentrations are:

Degree of Ionization

The fraction of the original electrolyte that ionizes is called the degree of ionization.

If percentage ionization is required:

8. pH and pOH

The hydrogen ion concentration of a solution may vary over a very large range. The pH scale provides a convenient way to express hydrogen ion concentration.

Similarly:

Relationship Between pH and pOH

This relationship applies at 25°C because the ionic product of water is 1.0 × 10⁻¹⁴ at this temperature.

Interpretation

• pH < 7: acidic solution
• pH = 7: neutral solution at 25°C
• pH > 7: basic solution

9. Acid Dissociation Constant (K_a) and Base Dissociation Constant (K_b)

Acid Dissociation Constant

For a weak acid:

A larger K_a generally indicates a stronger acid.

Base Dissociation Constant

For a weak base:

A larger K_b generally indicates a stronger weak base.

pK_a and pK_b

10. Ionic Product of Water

Water undergoes slight self-ionization.

More accurately, the proton is associated with water to form hydronium ion, but the simplified equation is commonly used for calculations.

At 25°C:

Therefore:

Relationship with pK

At 25°C:

11. Common Ion Effect

The common ion effect is the suppression of the ionization of a weak electrolyte by adding a strong electrolyte containing an ion common to the weak electrolyte.

Example: Acetic Acid and Sodium Acetate

Sodium acetate provides CH₃COO⁻ ions.

The increased concentration of acetate ion shifts the ionization equilibrium of acetic acid toward the left.

Importance

• Preparation of buffer solutions
• Control of ionization
• Selective precipitation
• Qualitative analysis
• Control of solution composition

12. Buffer Solutions

A buffer solution resists a significant change in pH when a small amount of strong acid or strong base is added.

Types of Buffer

1. Acidic buffer
2. Basic buffer

Acidic Buffer

An acidic buffer usually consists of a weak acid and its salt containing the conjugate base.

Example:

Here, CH₃COOH is the weak acid and CH₃COONa provides its conjugate base, CH₃COO⁻.

Basic Buffer

A basic buffer usually consists of a weak base and its salt containing the conjugate acid.

Example:

NH₃ acts as the weak base and NH₄Cl supplies NH₄⁺, its conjugate acid.

13. Mechanism of Buffer Action

Acidic Buffer: CH₃COOH / CH₃COO⁻

Consider a buffer containing acetic acid and sodium acetate.

The weak acid establishes the equilibrium:

When a Small Amount of Acid is Added

Added H⁺ is consumed by the acetate ion.

Therefore, the increase in free hydrogen ion concentration is much smaller than it would be in the absence of the buffer.

When a Small Amount of Base is Added

The added OH⁻ reacts with acetic acid.

Thus, the added base is largely consumed and the pH changes only slightly.

14. Henderson-Hasselbalch Equation

For an acidic buffer consisting of a weak acid HA and its conjugate base A⁻:

Rearranging:

Taking negative logarithms gives:

For a buffer made from a weak acid and its salt, the equation is commonly written:

For a Basic Buffer

The corresponding expression can be written using K_b:

At 25°C:

15. Buffer Capacity

Buffer capacity describes the ability of a buffer to resist a change in pH when acid or base is added.

A buffer with greater capacity can neutralize a larger amount of added acid or base before its pH changes significantly.

Factors Affecting Buffer Capacity

• Total concentration of buffer components
• Ratio of acid to conjugate base
• Amount of strong acid or base added
• pH relative to pK_a

Buffer capacity is generally greatest when the weak acid and its conjugate base are present in comparable amounts.

Therefore, a buffer is especially effective near the pK_a of its weak acid.

16. Acid-Base Indicators

An acid-base indicator is a weak organic acid or base whose different forms have different colours.

A general representation for an indicator is:

The two forms may have different colours.

Indicator Transition Range

An indicator changes colour over a particular range of pH rather than at one exact pH value.

Choice of Indicator

The indicator should be selected so that its transition range falls within the steep pH change near the equivalence point of the titration.

17. Acid-Base Neutralization

Neutralization is a reaction in which an acid reacts with a base.

Strong Acid and Strong Base

The net ionic equation is:

Weak Acid and Strong Base

The nature of the resulting solution depends on the relative strengths of the acid and base and the extent of hydrolysis of the resulting salt.

Equivalent Point and End Point

The equivalence point is the point at which chemically equivalent amounts of reacting substances have reacted according to the stoichiometry of the reaction.

The end point is the experimentally observed point at which the indicator changes colour.

18. Worked Numerical Problems

19. Common Mistakes

• Confusing strong and concentrated: strength refers to ionization, while concentration refers to amount per unit volume.
• Using pH = 14 − pOH at every temperature: this relationship has the value 14 at 25°C.
• Forgetting the negative sign: pH = −log[H⁺].
• Mixing up acid and conjugate base: an acid loses H⁺ and becomes its conjugate base.
• Using the wrong ratio in Henderson-Hasselbalch: for an acidic buffer, use conjugate base/weak acid.
• Thinking a buffer keeps pH perfectly constant: it only resists pH change within its effective range and capacity.
• Confusing equivalence point with end point: the first is based on reaction stoichiometry; the second is observed experimentally, usually with an indicator.
• Ignoring units: concentrations used in equilibrium calculations should be expressed consistently.
• Applying the weak-acid approximation without checking it: the approximation [H⁺] ≈ √(K_aC) should be checked when accuracy matters.

Practical and Laboratory Connections

1. Preparation of an Acetate Buffer

An acetate buffer can be prepared using acetic acid and sodium acetate.

The final pH depends mainly on the pK_a of acetic acid and the ratio of acetate ion to acetic acid.

2. pH Measurement

pH may be estimated using indicators, universal indicator paper or measured more accurately using a pH meter.

3. Titration

Acid-base titration is used to determine the concentration of an unknown acid or base using a solution of known concentration.

20. TU Examination Preparation

Quick Check

Try these questions before looking at the answers.

21. Quick Revision

• Arrhenius acid produces H⁺ in water.
• Arrhenius base produces OH⁻ in water.
• Brønsted acid donates H⁺.
• Brønsted base accepts H⁺.
• Lewis acid accepts an electron pair.
• Lewis base donates an electron pair.
• Conjugate acid-base pairs differ by one proton.
• Strong acids ionize extensively.
• Weak acids ionize partially.
• pH = −log[H^+].
• pOH = −log[OH⁻].
• At 25°C, pH + pOH = 14.
• K_w = [H⁺][OH⁻].
• pK_a = −log K_a.
• pK_b = −log K_b.
• Common ion effect suppresses ionization.
• Acidic buffer = weak acid + conjugate base salt.
• Basic buffer = weak base + conjugate acid salt.
• Acidic buffer equation: pH = pK_a + log(salt/acid).
• Basic buffer can be treated using the corresponding pOH relationship.
• Buffer capacity measures resistance to pH change.
• Indicator colour changes over a pH transition range.
• Equivalence point is based on reaction stoichiometry.
• End point is the experimentally observed indicator change.