Weak Acid and Base
Table of Contents
Acids are substances that donate protons (H+), while bases are substances that accept them. In aqueous solutions, their behavior can vary greatly depending on their strength. Weak acids and bases only partially dissociate in water. This means that only a small fraction of their molecules participate in chemical reactions, while the majority remain undissociated. These reactions are reversible and reach a dynamic equilibrium between reactants and products.
Weak acids and bases behave differently from strong acids and bases, which dissociate almost completely in water.
Dissociation Equations
The reversible dissociation of weak acids and bases is represented using equilibrium arrows (⇌), highlighting that these reactions do not proceed to completion. [1-6]
Weak Acid
A generic weak acid (HA) donates a proton (H+) to a water (H2O) molecule, forming a hydronium ion (H3O+) and a conjugate base (A–):
HA (aq) + H2O (l) ⇌ H3O+ (aq) + A– (aq)
Example
Acetic acid (CH3COOH), found in vinegar, partially donates a proton to water, forming acetate ions (CH3COO–) and hydronium ions:
CH3COOH (aq) + H2O (l) ⇌ H3O+ (aq) + CH3COO– (aq)
Most of the acetic acid remains undissociated, illustrating its weak nature.
Weak Base
A weak base (B) accepts a proton from water, resulting in the formation of a conjugate acid (BH+) and a hydroxide ion (OH⁻):
B (aq) + H2O (l) ⇌ BH+(aq) + OH– (aq)
Example
Ammonia (NH3) acts as a weak base by accepting a proton from water to form ammonium ions (NH4+) and hydroxide ions:
NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH– (aq)
Again, only a small proportion of ammonia molecules participate in this reaction.
Strength
The strength of a weak acid or base is determined by how much it dissociates in solution. This is quantified using equilibrium constants: [1-4]
Acid Dissociation Constant (Ka)
Ka = [H3O+][A–] / [HA]
Base Dissociation Constant (Kb)
Kb = [BH+][OH–] / [B]
Lower Ka and Kb values indicate that the corresponding acids and bases are weak. It also means that the dissociation in these compounds is less and not complete.
Alternative Measurement Scales
The Ka and Kb values can vary widely, making it inconvenient for chemists to compare the weak acids and the weak bases. Therefore, their logarithmic forms, pKa and pKb, are often used for easier comparison:
pKa = – log10 (Ka)
pKa = – log10 (Kb)
Weak acids and bases have higher pKa and pKb values, respectively.
Examples of pKa and pKb Values (at 25 °C)
- Acetic acid (CH3COOH): pKa = 4.74
- Formic acid (HCOOH): pKa = 3.74 (stronger than acetic acid)
- Ammonia (NH3): pKb = 4.75
- Methylamine (CH2NH2): pKb = 3.34 (stronger than ammonia)
Acidity and Basicity
The pH scale measures the acidity or basicity of a solution. However, calculating the pH for weak acids and bases is more complex than for strong ones because they do not fully dissociate. Chemists use the equilibrium conditions to estimate the concentrations of H3O+ and OH⁻, along with the Ka or Kb values. The pH is then calculated as follows:
Weak Acids
pH = – log10 [H3O+]
Weak Bases
pOH = – log10 [OH⁻] and pH = 14 – pOH
Example Problem
Calculate the pH of a 0.10 M acetic acid (CH3COOH) solution. Given: Ka = 1.8 x 10-5.
Solution
Dissociation Equation:
CH3COOH ⇌ H3O+ + CH3COO–
Let [H3O+] = x
Then,
Ka = [H3O+][CH3COO—] / [CH3COOH]
=> 1.8 x 10-5 = x2/0.1
=> x = √ (0.1) (1.8 x 10-5) ≈ 1.34 x 10-3
pH is given by:
pH = – log10 [H3O+]
=> pH = – log10 (1.34 x 10-3) ≈ 2.87
The pH of the 0.10 M acetic acid solution is approximately 2.87.
Examples [1-4]
There are several weak acids and bases in chemistry. The following tables list the most common ones.
Weak Acids
| Name | Chemical Formula | Common Uses and Applications |
|---|---|---|
| Acetic acid | CH3COOH | Found in vinegar; used in cooking, food preservation, and cleaning |
| Citric acid | C6H8O7 | Found in citrus fruits; used in food flavoring, soft drinks, and cosmetics |
| Carbonic acid | H2CO3 | Present in carbonated beverages; helps regulate blood pH |
| Formic acid | HCOOH | Found in ant stings; used in leather production and as a preservative |
| Phosphoric acid | H3PO4 | Used in soft drinks, rust removal, and fertilizers |
| Boric acid | H3BO3 | Used as an antiseptic, eyewash, and in insecticides |
| Hydrocyanic acid | HCN | Used in chemical synthesis |
| Benzoic acid | C6H5COOH | Used as a food preservative and in pharmaceuticals |
Weak Bases
| Name | Chemical Formula | Common Uses and Applications |
|---|---|---|
| Ammonia | NH3 | Used in household cleaners, fertilizers, and refrigeration systems |
| Methylamine | CH3NH2 | Used in the manufacture of pharmaceuticals and pesticides |
| Aniline | C6H5NH2 | Used in dyes, rubber processing, and pharmaceuticals |
| Pyridine | C5H5N | Used as a solvent and in the production of vitamins and agrochemicals |
| Trimethylamine | (CH3)3N | Found in fish odor; used in chemical manufacturing |
| Aluminum hydroxide | Al(OH)3 | Used as an antacid and in water purification |
| Magnesium hydroxide(Milk of magnesia) | Mg(OH)2 | Used as an antacid and laxative |
Water as the Weakest Acid and Base
Water (H2O) is a unique substance because it can act as both an acid and a base, a property known as amphoteric behavior. However, it is extremely weak in both roles and barely ionizes, making it one of the weakest acids and the weakest bases in chemistry. [2]
As a weak acid, water donates a proton (H+) to form hydroxide ions (OH–):
H2O ⇌ H+ + OH−
As a weak base, it accepts a proton to form hydronium ions (H3O+):
H2O + H+ ⇌ H3O+
This extremely low activity places pure water at a neutral pH of 7, balancing between acidic and basic properties.
From everyday products like vinegar and antacids to industrial chemicals and biological buffers, weak acids and bases play vital roles in our daily lives. By studying dissociation and calculating its effect on pH, students acquire valuable knowledge about real-world chemical applications.
