Acid & Base Theories Notes

Acid-Base Theories

• Arrhenius Acid-Base Theory
• Bronsted Acid-Base Theory
• Lewis Acid-Base Theory
• Auto Ionization of Water

Properties of Acids

• Sour taste.
• Aqueous solutions are strong or weak electrolytes.
• Turns blue litmus paper to red.
• Reacts with metals to produce $H_2$.

Properties of Bases

• Soapy texture.
• Bitter taste.
• Slippery feel.
• Changes red litmus paper to blue.
• Aqueous solutions are strong or weak electrolytes.

Arrhenius Acids

• Produce hydrogen ions ($H^+$) in water.
• Examples:
  • Hydrochloric acid: $HCl$
  • Nitric acid: $HNO_3$
  • Sulfuric acid: $H<em>2SO</em>4$
  • Phosphoric acid: $H<em>3PO</em>4$
  • Ethanoic acid: $CH_3COOH$
  • Carbonic acid: $H<em>2CO</em>3$
• Hydrogen ions join to water molecules as hydronium ions ($H_3O^+$).
• Acids vary in the number of hydrogens that can form hydrogen ions.
• Ethanoic acid ($CH_3COOH$) is a monoprotic acid.

Arrhenius Bases

• Produce hydroxide ions ($OH^−$) in water.
• Examples:
  • Sodium hydroxide: $NaOH$ (high solubility)
  • Potassium hydroxide: $KOH$ (high solubility)
  • Calcium hydroxide: $Ca(OH)_2$ (very low solubility)
  • Magnesium hydroxide: $Mg(OH)_2$ (very low solubility)

Limitations of Arrhenius Theory

• Aqueous solutions of $SO<em>2$, $SO</em>3$, $CO_2$ are acidic but do not produce $H^+$.
• Aqueous solutions of $CaO$, $Ca(CO<em>3)</em>2$, $Na<em>2CO</em>3$, $NH_3$ are basic but do not produce $OH^−$.

Brønsted-Lowry Theory

• Acid: Hydrogen-ion donor.
• Base: Hydrogen-ion acceptor.

Conjugate Acids and Bases

• Conjugate acid: Formed when a base gains a hydrogen ion (e.g., $NH<em>4^+$ is the conjugate acid of $NH</em>3$).
• Conjugate base: Remains after an acid loses a hydrogen ion (e.g., $OH^−$ is the conjugate base of $H_2O$).
• A conjugate acid-base pair consists of two ions or molecules related by the loss or gain of one hydrogen ion.

Amphoteric Substances

• Can act as either an acid or a base (e.g., water).
• Water acts as a base with $HCl$ and as an acid with $NH_3$.

Lewis Theory

• Acid: Accepts a pair of unbonded electrons.
• Base: Donates a pair of unbonded electrons.
• Broader definition than Brønsted-Lowry.

Acid-Base Definitions Comparison

• Arrhenius: Acid = $H^+$ producer, Base = $OH^−$ producer.
• Brønsted-Lowry: Acid = $H^+$ donor, Base = $H^+$ acceptor.
• Lewis: Acid = electron-pair acceptor, Base = electron-pair donor.

Key Definitions

• Hydronium ion ($H_3O^+$): Formed when a water molecule gains a $H^+$ ion.
• Conjugate acid: Particle formed when a base gains a $H^+$ ion.
• Conjugate base: Particle remaining when an acid donates a hydrogen ion.
• Conjugate Acid-Base Pair: two substances that are related by the loss or gain of a single hydrogen ion.
• Amphoteric: Substance that can act as both an acid and a base.
• Lewis acid: Substance that can accept a pair of electrons.
• Lewis base: Substance that can donate a pair of electrons.

The Auto-Ionization of Water and the pH Concept

• Water reacts with itself in an acid-base reaction: $2H<em>2O \rightleftharpoons H</em>3O^+ + OH^−$
• Dissociation constant: $K<em>c = \frac{[H</em>3O^+][OH^-]}{[H_2O]^2}$
• $K<em>w = [H</em>3O^+][OH^-] = [H^+][OH^-]$
• At 25°C: $[H<em>3O^+] = [OH^-] = 1.0 \times 10^{-7}$, thus $K</em>w = 1.0 \times 10^{-14}$
• Only Temperature can change $K_w$.

Types of Solutions

• Neutral: $[H^+] = [OH^-]$
• Acidic: [H^+] > [OH^-]
• Basic: [H^+] < [OH^-]

pH Scale

• $pH = -log[H_3O^+] = -log[H^+]$
• $10^{-pH} = [H^+]$
• Inverse relationship between pH and $H^+$ concentration.

pH and pOH Relationship

• $pOH = -log[OH^-]$
• $pH + pOH = 14.00$

Buffers

• Resist change in pH when a small amount of acid or base is added.
• Contain a common ion; important in biochemical and physiological processes.

Buffer Systems

• Human blood: Buffered system with pH = 7.4 ± 0.4.

The Hunderson-Hesselbalch Equation

• $pH = pK_a + log \frac{[A^-]}{[HA]}$

Characteristics of Buffered solutions

• Contain relatively large concentrations of a weak acid and its conjugate base.
• When acid is added, it reacts with the conjugate base
• When base is added, it reacts with the acid.
• pH is determined by the ratio of the base and acid. $pH = pK_a + log \frac{[A^-]}{[HA]}$

Buffer Capacity

• Amount of protons or hydroxide ions that can be absorbed without a significant change in pH.
• A good buffer is the one whose pH is near pKa
• pH is determined by the ratio of [A–]/[HA] and pKa
• Capacity is determined by the magnitudes of [HA] and [A–].
• $pH = pK_a + log \frac{[A^-]}{[HA]}$