Acids and Bases

Definitions of Acids and Bases

Acids and bases are fundamental chemical species with distinct behaviors in aqueous solutions. Their definitions have evolved to encompass a broader range of chemical reactions.

• Arrhenius Definition:

  • Acid: Produces hydronium ions (H3O+) when dissolved in water.

  • Base: Produces hydroxide ions (OH-) when dissolved in water.

• Brønsted-Lowry Definition:

  • Acid: Proton (H+) donor.

  • Base: Proton (H+) acceptor; must have a lone pair of electrons.

Monoprotic, Diprotic, and Triprotic Acids: Acids are classified by the number of protons they can donate. For example, HCl is monoprotic, H2SO4 is diprotic, and H3PO4 is triprotic.

Acidic hydrogens are typically bonded to electronegative atoms, making them more likely to be donated as protons.

Bases can be neutral molecules (like ammonia) or anions, and must have a lone pair to accept a proton.

Acid-base reactions involve the transfer of a proton from the acid to the base.

Conjugate acid-base pairs differ by one proton and are found on opposite sides of the reaction equation.

Acid and Base Strength

The strength of an acid or base is determined by its tendency to donate or accept protons in water.

• Strong acids: Completely dissociate in water (e.g., HCl, HNO3).

• Weak acids: Only partially dissociate in water (e.g., CH3COOH).

• Strong bases: Completely dissociate to give OH- (e.g., NaOH).

• Weak bases: Partially accept protons (e.g., NH3).

Polyprotic acids donate protons in steps, with each successive dissociation being less favorable.

There is an inverse relationship between acid and conjugate base strength: the stronger the acid, the weaker its conjugate base, and vice versa.

Acid Dissociation Constants (Ka)

The acid dissociation constant (Ka) quantifies the extent of dissociation for a weak acid in water:

$K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• Strong acids: Ka >> 1

• Weak acids: Ka << 1

Each proton donation from a polyprotic acid has a lower Ka value than the previous step.

Water as Both an Acid and a Base (Amphoteric)

Water can act as both an acid and a base, a property called amphoterism. It can donate or accept a proton depending on the reaction partner.

Water self-ionizes to form hydronium and hydroxide ions:

$2H_2O \rightleftharpoons H_3O^+ + OH^-$

The ion-product constant for water at 25°C is:

$K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$

Measuring Acidity: The pH Scale

The pH scale quantifies the acidity or basicity of a solution:

$\mathrm{pH} = -\log_{10}[H_3O^+]$

• Acidic: pH < 7

• Neutral: pH = 7

• Basic: pH > 7

pOH is related to pH by:

$\mathrm{pH} + \mathrm{pOH} = 14.00$

Working with pH and Significant Figures

To convert between pH and [H3O+]:

• $[H_3O^+] = 10^{-\mathrm{pH}}$

• Number of significant figures in [H3O+] equals the number of decimal places in pH.

Acid and Base Equivalents and Normality

One equivalent of acid neutralizes one equivalent of base. Normality (N) is the number of equivalents per liter:

$N = \frac{\text{equivalents of acid or base}}{\text{liters of solution}}$

• For monoprotic acids, N = M (molarity).

• For polyprotic acids, N = M × number of acidic hydrogens.

Common Acid-Base Reactions

Acids react with bases, carbonates, and ammonia in characteristic ways:

• Neutralization: Acid + Base → Salt + Water

• Acid + Carbonate/Bicarbonate: Produces CO2 gas

• Acid + Ammonia/Amines: Produces ammonium or alkylammonium salts

Acidity and Basicity of Salt Solutions

The pH of a salt solution depends on the strengths of the parent acid and base:

• Strong acid + strong base: Neutral solution

• Strong acid + weak base: Acidic solution

• Weak acid + strong base: Basic solution

• Weak acid + weak base: pH depends on relative strengths

Buffer Solutions

Buffers are solutions that resist changes in pH upon addition of small amounts of acid or base. They consist of a weak acid and its conjugate base.

The Henderson-Hasselbalch equation relates buffer pH to the concentrations of acid and conjugate base:

$\mathrm{pH} = \mathrm{p}K_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• Effective buffers have a pKa close to the desired pH and similar concentrations of acid and base.

Titration

Titration is a laboratory technique to determine the concentration of an acid or base by reacting it with a standard solution of known concentration.

• At the endpoint, moles of acid and base are stoichiometrically equivalent.

• For monoprotic acids and bases: $M_{acid} \times V_{acid} = M_{base} \times V_{base}$

• For polyprotic or non-1:1 reactions, use equivalents and normality: $N_{acid} \times V_{acid} = N_{base} \times V_{base}$

Summary Table: Acids and Bases Overview