Chapter 10: Acids, Bases, and Equilibrium – Study Notes

Chapter 10: Acids, Bases, Equilibrium

10.1 Acids and Bases

This section introduces the fundamental definitions of acids and bases according to the Arrhenius concept, focusing on their behavior in aqueous solutions.

• Acid: A compound that dissociates into H+ (hydrogen ion) when dissolved in water.

• Base: A compound that dissociates into OH- (hydroxide ion) when dissolved in water.

• Example (Acid): $\mathrm{HCl}(g) \rightarrow \mathrm{H^+}(aq) + \mathrm{Cl^-}(aq)$

• Example (Base): $\mathrm{NaOH}(s) \rightarrow \mathrm{Na^+}(aq) + \mathrm{OH^-}(aq)$

10.5 Dissociation of Water

Water undergoes self-ionization, producing hydronium and hydroxide ions through a proton transfer reaction.

• Equation: $\mathrm{H_2O} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^+} + \mathrm{OH^-}$

• Proton transfer reaction: A hydrogen ion (H+) moves from one water molecule to another.

• Hydronium ion: $\mathrm{H_3O^+}$ is formed when water accepts a proton.

• Hydroxide ion: $\mathrm{OH^-}$ is formed when water loses a proton.

• Water is an electrolyte: It conducts electricity due to the presence of ions.

Ion Product of Water

The ion product constant of water ($K_w$) quantifies the equilibrium concentrations of hydronium and hydroxide ions in pure water.

• Equation: $K_w = [\mathrm{H_3O^+}] \times [\mathrm{OH^-}]$

• Value at 25°C: $K_w = 1.0 \times 10^{-14}$

• In pure water: $[\mathrm{H_3O^+}] = [\mathrm{OH^-}] = 1.0 \times 10^{-7}~\mathrm{M}$

Calculations with the Ion Product of Water

Given the concentration of one ion, the concentration of the other can be calculated using $K_w$.

• Example: If $[\mathrm{H_3O^+}] = 0.020~\mathrm{M}$, then $[\mathrm{OH^-}] = \frac{1.0 \times 10^{-14}}{0.020} = 5.0 \times 10^{-13}~\mathrm{M}$

• Practice: Calculate $[\mathrm{OH^-}]$ in a solution with $[\mathrm{H_3O^+}] = 0.00004~\mathrm{M}$

• Practice: Calculate $[\mathrm{H_3O^+}]$ in a solution with $[\mathrm{OH^-}] = 10^{-5}~\mathrm{M}$

• Practice: Calculate $[\mathrm{H_3O^+}]$ in a solution with $[\mathrm{OH^-}] = 2.8 \times 10^{-9}~\mathrm{M}$

Neutral, Acidic & Basic Solutions

The relative concentrations of hydronium and hydroxide ions determine whether a solution is neutral, acidic, or basic.

• Neutral: $[\mathrm{H_3O^+}] = [\mathrm{OH^-}] = 1.0 \times 10^{-7}~\mathrm{M}$

• Acidic: $[\mathrm{H_3O^+}] > [\mathrm{OH^-}]$

• Basic: $[\mathrm{OH^-}] > [\mathrm{H_3O^+}]$

• Example: For $[\mathrm{OH^-}] = 7.2 \times 10^{-5}~\mathrm{M}$ and $[\mathrm{H_3O^+}] = 1.38 \times 10^{-10}~\mathrm{M}$, the solution is basic.

• Example: For $[\mathrm{H_3O^+}] = 2.0 \times 10^{-3}~\mathrm{M}$, the solution is acidic.

10.6 The pH Scale

The pH scale is a logarithmic measure of hydronium ion concentration, used to classify solutions as acidic, neutral, or basic.

• Formula: $\mathrm{pH} = -\log[\mathrm{H_3O^+}]$

• Range: 0 (most acidic) to 14 (most basic)

• Neutral: pH = 7

• Acidic: pH < 7

• Basic: pH > 7

• Example: Seawater with $[\mathrm{H_3O^+}] = 3.67 \times 10^{-9}~\mathrm{M}$ is basic (pH > 7).

• Example: Beer with pH = 5.0 is acidic.

Solution	[H3O+]	[OH-]	pH	Classification
Pure Water	1.0 × 10-7 M	1.0 × 10-7 M	7	Neutral
Seawater	3.67 × 10-9 M	Additional info: Calculate using $K_w$	~8.4	Basic
Beer	Additional info: Calculate using pH	Additional info: Calculate using $K_w$	5.0	Acidic

Additional info: The above notes cover the Arrhenius definitions, water dissociation, ion product, solution classification, and pH scale, which are foundational for understanding acid-base chemistry in General Chemistry.