3. Strong and Weak Acids/Bases

3.1 Strong Acids and Bases

3.1.1 List of Common Strong Acids

1. Hydrochloric acid (HCl)
2. Hydrobromic acid (HBr)
3. Hydroiodic acid (HI)
4. Sulfuric acid (H₂SO₄) – the first proton is strong; the second proton (HSO₄ minus → H plus + SO₄² minus) is considered a strong acid if concentration is high, but at moderate dilutions the second dissociation is weak.
5. Perchloric acid (HClO₄)
6. Nitric acid (HNO₃)

Behavior:

• Virtually 100% of the acid molecules dissociate into H plus and their conjugate base in water.

Strong acids are substances that completely dissociate in water, meaning they release all of their available hydrogen ions (H+) into the solution. The common strong acids listed, such as hydrochloric acid (HCl) and sulfuric acid (H₂SO₄), have this characteristic. In the case of sulfuric acid, while its first dissociation is complete, the second dissociation may not be fully complete depending on the concentration. Since strong acids release a high concentration of protons, they are also very effective at lowering the pH of solutions, making them very acidic.

Think of a strong acid like a very forceful sprayer. When you pull the trigger, it releases all its water immediately, representing the rapid dissociation of strong acids. Conversely, a weak acid would be like a squirt bottle that requires multiple presses to release the same amount of water, illustrating how weak acids do not completely dissociate.

3.1.2 List of Common Strong Bases

1. Sodium hydroxide (NaOH)
2. Potassium hydroxide (KOH)
3. Lithium hydroxide (LiOH)
4. Calcium hydroxide (Ca(OH)₂) – only slightly soluble but highly dissociated in the dissolved portion.
5. Barium hydroxide (Ba(OH)₂) – similar to Ca(OH)₂.
6. Strontium hydroxide (Sr(OH)₂)
7. Cesium hydroxide (CsOH)

Behavior:

• Dissociate completely to give OH minus and the metal cation. (Solubility can be limiting factor for some metal hydroxides.)

Strong bases are substances that completely dissociate in water to produce hydroxide ions (OH-), thereby increasing the basicity of the solution. Common strong bases include sodium hydroxide (NaOH) and potassium hydroxide (KOH). Just as with strong acids, the key feature of strong bases is their ability to provide a high concentration of OH- ions in a solution, resulting in a higher pH, indicating a basic environment.

Imagine a strong base as a powerful cleaning solution that can instantly neutralize acid spills. As it comes into contact with acids, it completely transforms the solution, much like how a strong base completely dissociates into hydroxide ions, effectively neutralizing the acidity.

3.1.3 pH Calculation for Strong Acids/Bases

• For a solution of a strong acid with concentration C, [H plus] = C (assuming negligible contribution from water if C > 10⁻⁶).
• For a solution of a strong base with concentration C, [OH minus] = C (again, ignoring water).
• At 25 °C, pH + pOH = 14.

Example 1: A solution is prepared by dissolving 0.250 mol of HCl in 1.00 L of water.

• [H plus] = 0.250 M → pH = – log₁₀ (0.250) = 0.60.

Example 2: A solution is prepared by dissolving 0.0200 mol of Ca(OH)₂ in 500 mL of water. Ca(OH)₂ dissociates into Ca² plus + 2 OH minus.

• Moles of Ca(OH)₂ = 0.0200 mol, volume = 0.500 L → 0.0400 M Ca(OH)₂.
• [OH minus] = 2 × 0.0400 M = 0.0800 M.
• pOH = – log₁₀ (0.0800) = 1.10 → pH = 14.00 – 1.10 = 12.90.

Calculating the pH for strong acids and strong bases is straightforward because they completely dissociate. For a strong acid, the concentration of hydrogen ions [H+] equals the concentration of the acid. For example, if you have 0.250 M HCl, this means [H+] is also 0.250 M, leading to a pH of 0.60. Similarly, for strong bases, you calculate the concentration of hydroxide ions [OH-] and compute the pH using the relationship that pH + pOH = 14.

Imagine you have a full jar of marbles, each representing a hydrogen ion in a strong acid. If you know exactly how many marbles (H+) there are in the jar (the concentration), you can quickly determine how much acid the jar holds (pH). With strong bases, think of it as having an empty jar fillable by a hose (OH-); the faster you deliver more water, the higher the basicity, allowing you to easily measure how basic the solution is.

3.2 Weak Acids and Bases

3.2.1 Acid Dissociation Constant (Ka)

For a weak acid HA in water: HA ⇌ H plus + A minus

• Ka = ([H plus] × [A minus]) ÷ [HA] at equilibrium. We denote the initial concentration of HA as C₀, and assume initially [H plus] and [A minus] are zero (unless acid or base is already present). At equilibrium, let [H plus] = x, [A minus] = x, and [HA] = C₀ – x. Then: Ka = x × x ÷ (C₀ – x) = x² ÷ (C₀ – x)

Approximations:

• If Ka is significantly smaller than C₀ (for example, Ka < 10⁻² and C₀ > 0.01), then x will be small compared to C₀, so C₀ – x ≈ C₀, and we use x ≈ sqrt(Ka × C₀).
• After finding x, calculate pH = – log₁₀ (x).

Percent Ionization:

• Percent ionization = ([A minus] at equilibrium ÷ C₀) × 100% = (x ÷ C₀) × 100%. As C₀ increases, percent ionization decreases.

Weak acids do not completely dissociate in solution. The degree of dissociation can be quantified using the acid dissociation constant (Ka). To calculate the pH of a weak acid, you first determine the equilibrium concentrations of the ions formed based on the initial concentration and Ka. A key factor in this calculation is the assumption that x (the amount that dissociates) is small compared to C₀. The percent ionization provides insights into how much of the weak acid actually dissociates into ions, which is generally less than 100%.

Consider a weak acid as someone who is shy in a group. They might only participate a little (dissociate partially) rather than jumping into every conversation (completely dissociate). The acid dissociation constant (Ka) represents how likely they are to join in the conversation; a higher Ka means they are more likely to participate, just like a weak acid that can still ionize effectively in solution.

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