Acids, Bases, and Equilibrium Notes

Binary Acids

  • A binary acid contains only two elements: hydrogen and a more electronegative element (typically a halogen).

  • Examples: HF, HCl, HBr

Binary Acid Nomenclature

  1. Begins with the prefix "hydro-".

  2. Followed by the root name of the second element.

  3. Ends with the suffix "-ic acid".

Oxyacids

  • An oxyacid contains hydrogen, oxygen, and a third element (usually a nonmetal).

Oxyacid Nomenclature

  • The root of the name comes from the third element in the compound.

  • Acids with fewer oxygen atoms end with the suffix "-ous acid."

  • Acids with more oxygen atoms end with the suffix "-ic acid."

Strong vs. Weak Acids

  • Strong acids ionize completely in water; no reactant molecules remain.

  • Example: HCl+H<em>2O→H</em>3O++Cl−\text{HCl} + \text{H}<em>2\text{O} \rightarrow \text{H}</em>3\text{O}^+ + \text{Cl}^-

  • Weak acids partially ionize in water; an equilibrium is established.

  • Example: CH<em>3COOH+H</em>2O⇌CH<em>3COO−+H</em>3O+\text{CH}<em>3\text{COOH} + \text{H}</em>2\text{O} \rightleftharpoons \text{CH}<em>3\text{COO}^- + \text{H}</em>3\text{O}^+

Strong vs. Weak Bases

  • Strong bases dissociate completely in water, producing many hydroxide ions.

  • Example: NaOH→Na++OH−\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-

  • Weak bases produce few hydroxide ions in solution.

  • Example: NH<em>3+H</em>2O⇌NH<em>4OH⇌NH</em>4++OH−\text{NH}<em>3 + \text{H}</em>2\text{O} \rightleftharpoons \text{NH}<em>4\text{OH} \rightleftharpoons \text{NH}</em>4^+ + \text{OH}^-

Neutralization Reactions

  • Neutralization is a reaction between an acid and a base.

  • It always produces water and a salt.

  • Example: HCl+NaOH→NaCl+H2O\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}

  • Applications of Neutralization:

    • Antacid medications (to reduce stomach acid).

    • Pesticide treatment and cleanup.

pH and pOH

  • pH is a measure of the acidity or alkalinity of a solution.

  • pH=−log⁡[H+]\text{pH} = -\log[\text{H}^+] (where [H+] is the hydrogen ion concentration)

  • pOH measures the hydroxide ion concentration in a solution.

  • pOH=−log⁡[OH−]\text{pOH} = -\log[\text{OH}^-] (where [OH-] is the hydroxide ion concentration)

  • The Kw expression is the equilibrium constant for water.

  • pKw=pH+pOH=14\text{pKw} = \text{pH} + \text{pOH} = 14

  • pKw can be used to measure the pH of a basic solution or the pOH of an acidic solution.

  • To measure the pH of a basic solution: pH=pKw−pOH\text{pH} = \text{pKw} - \text{pOH}

  • The lower the pH, the greater the concentration of acid.

Practice Problems:

  • Calculate the pH for the following acids:

    • 0. 01 M of HCl

    • 0. 00025 M of acetic acid

    • 0. 05 M H2SO4

  • Calculate the pOH for the following bases:

    • 0. 01 M of NaOH

    • 0. 00025 M of Baking soda

    • 0. 05 M Ammonia