Acid and Base Chemistry Notes

Objectives

• Review the concept of “acidity and basicity”.

• Define the terms “acid” and “base”.

• Illustrate the types of acids.

• Define the Brønsted-Lowry Theory of conjugate acid-base pairs.

• Illustrate strength of acids and bases.

Acids and Bases

• Acids and bases are two extremes that describe chemicals, similar to how hot and cold describe temperature.

• Mixing acids and bases can neutralize their extreme effects, analogous to mixing hot and cold water to even out the temperature.

pH Scale

• The pH scale measures how acidic or basic a substance is.

• It ranges from 0 to 14.

• A pH of 7 is neutral.

• A pH less than 7 is acidic.

• A pH more than 7 is basic.

• The pH of pure water is 7 (neutral).

• When chemicals are mixed with water, the mixture can become either acidic or basic.

Acid and Base Definitions

• Arrhenius Definition:

  • An acid produces or ionizes to produce $H^+$ ions in water (aqueous solution).

  • A base produces $OH^-$ ions in water (aqueous solution).

• Lewis Definition:

  • An acid is an electron pair acceptor.

  • A base is an electron pair donor.

• Brønsted-Lowry Definition:

  • An acid is a proton donor.

  • A base is a proton acceptor.

Classification of Acids

• Monoprotic: Yields one proton.

  • Example: $CH<em>3COOH(aq) \leftrightarrow CH</em>3COO^-(aq) + H^+(aq)$

• Diprotic: Yields two protons.

  • Example:

    • $H<em>2SO</em>4(aq) \leftrightarrow HSO_4^-(aq) + H^+(aq)$

    • $HSO<em>4^-(aq) \leftrightarrow SO</em>4^{2-}(aq) + H^+(aq)$

• Triprotic: Yields three protons.

  • Example:

    • $H<em>3PO</em>4(aq) \leftrightarrow H<em>2PO</em>4^-(aq) + H^+(aq)$

    • $H<em>2PO</em>4^-(aq) \leftrightarrow HPO_4^{2-}(aq) + H^+(aq)$

    • $HPO<em>4^{2-}(aq) \leftrightarrow PO</em>4^{3-}(aq) + H^+(aq)$

Brønsted-Lowry Theory

• Generally regarded as the classical theory of acids and bases.

  • An acid is a substance that is a 'proton donor'.

    • Examples: hydrochloric acid (HCl), sulphuric acid ($H<em>2SO</em>4$), and nitric acid ($HNO_3$).

  • A base is a substance which is a 'proton acceptor'.

    • Examples: ammonia ($NH<em>3$) and sodium hydrogencarbonate ($NaHCO</em>3$ or baking soda).

• Water as a Solvent:

  • Water is made up of $H_2O$ molecules and can dissolve many substances.

  • An acidic substance, when dissolved in water, donates a hydrogen ion (proton) ($H^+$) to water to form hydronium ions ($H_3O^+$).

    • $HCl + H<em>2O \rightleftharpoons H</em>3O^+ + Cl^-$

  • A basic substance, when dissolved in water, accepts a proton from water to form hydroxide ions ($OH^-$).

    • $NH<em>3 + H</em>2O \rightleftharpoons NH_4^+ + OH^-$

Acid-Base Conjugate Pairs

• The removal of a proton ($H^+$) from an acid produces its conjugate base.

• The reception of a proton by a base produces its conjugate acid.

• Example: $HCl + H<em>2O \rightleftharpoons H</em>3O^+ + Cl^-$

  • The removal of $H^+$ from HCl produces the chloride ion ($Cl^−$), the conjugate base of the acid.

Strength of Acids

• A strong acid completely ionizes in water to give a high concentration of protons and a low pH.

  • In a 6M solution of HCl, 99.996% of the HCl molecules react with $H<em>2O$ to form $H</em>3O^+$ and $Cl^-$ ions.

    • $HCl(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + Cl^-(aq)$

• A weak acid does not readily transfer $H^+$ ions to water.

  • In a 1M solution of acetic acid, less than 0.4% of the $CH<em>3CO</em>2H$ molecules react with water to form $H<em>3O^+$ and $CH</em>3CO_2^-$ ions.

    • $CH<em>3CO</em>2H(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + CH<em>3CO</em>2^-(aq)$

Strength of Acids - $K_a$

• The relative strengths of acids is described in terms of the acid-dissociation equilibrium constant, $K_a$.

• Generic equation: $HA(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + A^-(aq)$

• The value of $K<em>a$ is calculated as: $K</em>a = \frac{[H^+][A^-]}{[HA]}$

• For strong acids, K_a > 1, indicating that only a small concentration of HA remains in solution.

  • Example: HCl has a $K_a$ of roughly $1 \times 10^6$.

• For weak acids, Ka < 1, indicating that the product of the concentrations of $H</em>3O^+$ and $A^-$ ions is smaller than the concentration of the residual HA molecules.

  • Example: Acetic acid has a $K_a$ of only $1.8 \times 10^{-5}$.

$K<em>a$ and $pK</em>a$

• The larger the $K_a$, the stronger is the acid.

• $pK<em>a = -\log{K</em>a}$

• Therefore, the larger the $pK_a$, the weaker is the acid.

Strength of Bases - $K_b$

• A strong base completely ionizes in water to give a high concentration of hydroxide ions ($OH^-$) and a high pH.

• Weaker bases will not completely ionize to give high concentrations of $OH^-$ ions and will give a lower (still >7) pH.

• The relative strengths of bases is described in terms of a base-dissociation equilibrium constant, $K_b$.

• For strong bases: K_b > 1

• For weak bases: K_b < 1

$K<em>b$ and $pK</em>b$

• The larger the $K_b$, the stronger is the base.

• $pK<em>b = -\log{K</em>b}$

• Therefore, the larger the $pK_b$, the weaker is the base.

Relative Strengths of Conjugate Acid-Base Pairs

• Strong acids have a weak conjugate base.

  • Example: HCl is a strong acid; thus, $Cl^-$ is a weak base.

    • $HCl(g) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + Cl^-(aq)$

• Strong bases have a weak conjugate acid.

  • Example: Ammonia ($NH<em>3$) is a reasonably good base; thus, $NH</em>4^+$ is a weak acid.

    • $NH<em>4^+(aq) + H</em>2O(l) \rightleftharpoons H<em>3O^+(aq) + NH</em>3(aq)$

Strong and Weak Acids

• List of acids and their conjugate bases, illustrating increasing acid strength.

Direction of Proton-Transfer Reactions

• The stronger acid will transfer a proton to the stronger base, yielding the weaker acid and weaker base as favored species at equilibrium.

• $HA + B \rightleftharpoons A^- + HB^+$

  • Strong acid donates a proton to a strong base, resulting in a weak base and a weak acid.

  • Good proton donors and acceptors lead to poor proton donors and acceptors.