Acids, Bases, and the pH Scale: Study Notes for General Chemistry

Acid and Base Chemistry

Bronsted-Lowry (B/L) Acids and Bases

The Bronsted-Lowry definition is one of the most widely used concepts for acids and bases in general chemistry. It focuses on the transfer of protons (H+) between chemical species.

• Acid: A substance that donates a proton (H+).

• Base: A substance that accepts a proton (H+).

• Example: In the reaction HCl + H2O → Cl- + H3O+, HCl acts as an acid (donates H+), and H2O acts as a base (accepts H+).

Conjugate Acid-Base Pairs

Every acid has a conjugate base, and every base has a conjugate acid. These pairs differ by one proton.

• Conjugate Acid: Formed when a base gains a proton.

• Conjugate Base: Formed when an acid loses a proton.

• General Reaction:

$HA(aq) + H_2O(l) \rightarrow A^-(aq) + H_3O^+(aq)$

• Example: For HSO4-, the conjugate acid is H2SO4, and the conjugate base is SO42-.

• Additional info: Water is amphoteric, meaning it can act as either an acid or a base depending on the reaction partner.

Strong Acids and Bases

Strong acids and bases dissociate completely in water, making them strong electrolytes.

• Six Strong Acids:

  • HCl (hydrochloric acid)

  • HBr (hydrobromic acid)

  • HI (hydroiodic acid)

  • HNO3 (nitric acid)

  • H2SO4 (sulfuric acid)

  • HClO4 (perchloric acid)

• Strong Bases: Alkali and alkaline earth metal hydroxides, such as:

  • NaOH (sodium hydroxide)

  • KOH (potassium hydroxide)

  • LiOH (lithium hydroxide)

  • Ca(OH)2 (calcium hydroxide)

  • Sr(OH)2 (strontium hydroxide)

  • Ba(OH)2 (barium hydroxide)

The Dissociation of Water

Water can self-ionize to form hydronium and hydroxide ions. This process is fundamental to understanding acid-base chemistry.

• Equation:

$H_2O(l) + H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$

• Equilibrium Constant (Kw): The product of the concentrations of H+ and OH- in pure water at 25°C.

$K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$

• Neutral Solution: $[H^+] = [OH^-] = 1.0 \times 10^{-7}$ M

• Acidic Solution: $[H^+] > [OH^-]$

• Basic Solution: $[OH^-] > [H^+]$

pH and the pH Scale

The pH scale is a logarithmic measure of the concentration of hydrogen ions in solution. It is used to quantify acidity and basicity.

• Definition: $\text{pH} = -\log[H^+]$

• Neutral pH: At 25°C, pure water has $\text{pH} = 7$

• Relationship: $\text{pH} + \text{pOH} = 14$

• Example Calculation: If $[H^+] = 1.8 \times 10^{-5}$ M, then $\text{pH} = -\log(1.8 \times 10^{-5}) = 4.74$

Acidity of Rainwater

Rainwater is naturally slightly acidic due to the dissolution of atmospheric carbon dioxide, forming carbonic acid.

• Reaction: $CO_2(g) + H_2O(l) \rightleftharpoons H_2CO_3(aq)$

• Effect: In colder climates, rainwater has a lower pH due to increased solubility of CO2.

Summary Table: Strong Acids and Bases

Strong Acids	Strong Bases
HCl (hydrochloric acid)	NaOH (sodium hydroxide)
HBr (hydrobromic acid)	KOH (potassium hydroxide)
HI (hydroiodic acid)	LiOH (lithium hydroxide)
HNO3 (nitric acid)	Ca(OH)2 (calcium hydroxide)
H2SO4 (sulfuric acid)	Sr(OH)2 (strontium hydroxide)
HClO4 (perchloric acid)	Ba(OH)2 (barium hydroxide)

Key Equations and Relationships

• $K_w = [H^+][OH^-]$

• $\text{pH} = -\log[H^+]$

• $\text{pOH} = -\log[OH^-]$

• $\text{pH} + \text{pOH} = 14$

Examples and Applications

• Example: If a solution has a pH of 9.26, then $[OH^-] = 5.5 \times 10^{-5}$ M (using $K_w$ and pOH calculations).

• Application: Ranking solutions by acidity or basicity using pH and ion concentrations.

Additional info:

• Concept checks and practice problems are included to reinforce understanding of acid-base equilibria and pH calculations.

• Students should be familiar with logarithmic relationships and how to manipulate equations involving pH, pOH, and $K_w$.