Introduction to Acid-Base Equilibria

Acids and bases are fundamental chemical substances that interact through proton transfer. Their reactions create equilibria that follow predictable patterns based on their strengths.

When studying acid-base chemistry, we need to understand several key definitions. An acid can donate protons, while a base accepts protons. When these substances interact, they form products called salts plus water in a process called neutralization.

Strong acids and bases dissociate completely in water, whereas weak acids and bases only partially dissociate. This difference in behavior creates varying pH levels that we can measure and predict mathematically.

Remember: Water itself can act as both an acid and a base - this dual nature makes it amphoteric and is essential to understanding acid-base chemistry!

Arrhenius Acids and Bases

In 1884, Svante Arrhenius provided one of the earliest modern definitions of acids and bases. According to his theory, an acid is a substance that produces hydronium ions (H₃O⁺) in water, while a base produces hydroxide ions (OH⁻) in water.

When an acid like HCl dissolves in water, it reacts to form hydronium ions:

H₂O(l) + HCl(aq) → H₃O⁺(aq) + Cl⁻(aq)

Bases work slightly differently. Many bases are metal hydroxides like NaOH or KOH that directly release OH⁻ ions when dissolved:

NaOH(s) → Na⁺(aq) + OH⁻(aq)

Some bases like ammonia (NH₃) don't initially contain hydroxide ions but produce them by reacting with water:

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

Chemistry in Action: When ammonia reacts with water, a proton transfers from water to ammonia - this is shown using curved arrows in reaction mechanisms to track electron movement!

Brønsted-Lowry Acids and Bases

The Brønsted-Lowry definition expands our understanding of acids and bases beyond the Arrhenius theory. In this model, an acid is a proton donor, and a base is a proton acceptor.

This definition introduces the important concept of conjugate pairs. When an acid donates a proton, what remains is called the conjugate base. Similarly, when a base accepts a proton, it forms a conjugate acid.

For example, in the reaction:

HCl(aq) + H₂O(l) ⇌ H₃O⁺(aq) + Cl⁻(aq)

HCl is the acid, H₂O is the base, H₃O⁺ is the conjugate acid, and Cl⁻ is the conjugate base.

In another example:

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

NH₃ acts as the base, H₂O as the acid, NH₄⁺ is the conjugate acid, and OH⁻ is the conjugate base.

Quick Tip: To identify conjugate pairs, look for species that differ by just one H⁺. For example, NH₃/NH₄⁺ and H₂O/OH⁻ are conjugate pairs.

Lewis Acids and Bases

The Lewis definition provides the most general way to understand acids and bases. A Lewis base is a substance that can donate a pair of electrons, while a Lewis acid accepts a pair of electrons.

This definition broadens our understanding beyond just proton transfer. A Lewis acid-base reaction involves the formation of a coordinate covalent bond, where both electrons in the new bond come from the same atom.

For example, when boron trifluoride (BF₃) reacts with ammonia (NH₃):

F₃B + :NH₃ → F₃B-NH₃

The nitrogen atom in ammonia donates its lone pair of electrons to form a bond with boron.

This definition explains reactions that don't involve hydrogen ions at all, making it the most comprehensive acid-base theory. It can describe interactions between metal ions and ligands, which is crucial in understanding coordination chemistry.

Why This Matters: The Lewis definition explains reactions in organic chemistry, biochemistry, and materials science that wouldn't be considered acid-base reactions under earlier models!

The Acid-Base Properties of Water

Water is special because it can act as both an acid and a base - a property called amphoteric. It can donate protons (acting as an acid) or accept them (acting as a base) depending on what it's reacting with.

Water undergoes a process called autoionization:

2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

This equilibrium has a constant called Kw:

Kw = [H₃O⁺][OH⁻] = 1 × 10⁻¹⁴ at 25°C

In pure water, $H₃O⁺$ = $OH⁻$ = 1 × 10⁻⁷ M, which creates a perfect balance. This balance determines whether a solution is:

• Acidic: when $H₃O⁺$ > $OH⁻$
• Basic: when $H₃O⁺$ < $OH⁻$
• Neutral: when $H₃O⁺$ = $OH⁻$

The relationship between these ions is always governed by Kw, meaning if one concentration changes, the other must change to maintain their product at 1 × 10⁻¹⁴.

Think About It: Even in highly acidic solutions, there are still some OH⁻ ions present, just at very low concentrations!

The pH Scale

The pH scale provides a convenient way to express the acidity or basicity of a solution without writing out scientific notation. The pH is defined as:

pH = -log[H⁺] or pH = -log[H₃O⁺]

This logarithmic scale means that each unit change in pH represents a tenfold change in hydrogen ion concentration. For instance, a solution with pH 3 has ten times more hydrogen ions than a solution with pH 4.

On this scale at 25°C:

• Neutral solutions have pH = 7
• Acidic solutions have pH < 7
• Basic solutions have pH > 7

Common substances have characteristic pH values: stomach acid (~1), orange juice (~3.5), pure water (7), blood (~7.4), and household ammonia (~11.5).

We can also express hydroxide concentration using pOH:

pOH = -log[OH⁻]

At 25°C, pH and pOH are related by:

pH + pOH = 14

Real-World Application: Your body maintains blood pH between 7.35-7.45 - even small deviations can be life-threatening, showing how critical pH balance is to biological systems!

Strong Acids and Bases

Strong acids and bases completely dissociate in water, meaning essentially 100% of their molecules break apart into ions. This dissociation is so complete that it's not considered an equilibrium process.

Common strong acids include:

• Hydrochloric acid (HCl)
• Nitric acid (HNO₃)
• Sulfuric acid (H₂SO₄) - first ionization only
• Perchloric acid (HClO₄)

Common strong bases include:

• Group 1A hydroxides (LiOH, NaOH, KOH)
• Group 2A hydroxides (Ca(OH)₂, Ba(OH)₂)

For strong acids, the hydronium ion concentration $H₃O⁺$ equals the initial concentration of the acid. Similarly, for strong bases, the hydroxide concentration $OH⁻$ equals the initial concentration multiplied by the number of hydroxide ions per formula unit.

For example, in a 0.057 M HBr solution:

HBr(aq) + H₂O(l) → H₃O⁺(aq) + Br⁻(aq)

$H₃O⁺$ = 0.057 M, giving a pH of 1.24

Important: Remember that only a few acids and bases are strong - most are weak and follow equilibrium principles!

Weak Acids and Acid Ionization Constants

Weak acids only partially ionize in water, establishing an equilibrium between the acid and its ions. The extent of this ionization depends on both the acid's concentration and its acid ionization constant, Ka.

For a weak acid HA:

HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)

The acid ionization constant is:

Ka = [H₃O⁺][A⁻]/[HA]

Common weak acids include acetic acid $Ka = 1.8 × 10⁻⁵$, formic acid $Ka = 1.7 × 10⁻⁴$, and hydrofluoric acid $Ka = 7.1 × 10⁻⁴$.

To calculate the pH of a weak acid solution, we set up an equilibrium table with:

• Initial concentrations
• Changes in concentration (often designated as x)
• Equilibrium concentrations

If Ka is small compared to the initial concentration, we can use the approximation:

Ka ≈ x²/Ci

This approximation is valid if x < 5% of the initial acid concentration.

Problem-Solving Tip: When calculating weak acid pH, start by assuming x is small. If your answer shows x > 5% of the initial concentration, you'll need to solve the quadratic equation without the approximation!

Weak Bases and Base Ionization Constants

Similar to weak acids, weak bases only partially ionize in water. Their ionization is characterized by the base ionization constant, Kb.

For a generic weak base B:

B(aq) + H₂O(l) ⇌ BH⁺(aq) + OH⁻(aq)

The base ionization constant is:

Kb = [BH⁺][OH⁻]/[B]

Common weak bases include ammonia $Kb = 1.8 × 10⁻⁵$, methylamine $Kb = 4.4 × 10⁻⁴$, and pyridine $Kb = 1.7 × 10⁻⁹$.

The approach for solving weak base problems mirrors that for weak acids, except that we solve for $OH⁻$ rather than $H₃O⁺$. Once we know $OH⁻$, we can calculate pOH and then pH.

For example, if a 0.50 M weak base solution has pH 9.59, we can determine that:

• pOH = 4.41
• $OH⁻$ = 3.89 × 10⁻⁵ M
• Kb = 3.0 × 10⁻⁹

Real-Life Connection: Many medications and household products are weak bases. Understanding their chemistry helps pharmacists formulate drugs and chemists develop effective cleaning products!

Conjugate Acid-Base Pairs and Their Relationship

There's an inverse relationship between the strength of an acid or base and its conjugate. A strong acid will have a weak conjugate base since it readily gives up its proton. Conversely, a weak acid will have a strong conjugate base that's eager to reclaim a proton.

For example:

• HCl (strong acid) forms Cl⁻ (weak conjugate base)
• HCN (weak acid) forms CN⁻ (strong conjugate base)
• OH⁻ (strong base) forms H₂O (weak conjugate acid)
• NH₃ (weak base) forms NH₄⁺ (strong conjugate acid)

This relationship can be quantified. For any conjugate acid-base pair, the product of their ionization constants equals the water ionization constant:

Ka × Kb = Kw = 1 × 10⁻¹⁴

This means if we know the Ka of an acid, we can calculate the Kb of its conjugate base:

Kb = Kw/Ka

For example, if benzoic acid has Ka = 6.5 × 10⁻⁵, then its conjugate base (benzoate ion) has Kb = 1.5 × 10⁻¹⁰.

Study Strategy: When you learn the strength of an acid, you automatically know something about its conjugate base. This pattern helps you predict chemical behavior without memorizing every value!