Acids and Bases: Bronsted-Lowry Theory, pH, and Buffer Solutions (Chapter 10 Study Notes)

Acids and Bases: Bronsted-Lowry Theory, pH, and Buffer Solutions

Bronsted-Lowry Acids and Bases

The Bronsted-Lowry theory defines acids and bases based on their ability to donate or accept protons (H+ ions).

• Bronsted-Lowry Acid: A substance that donates a proton (H+) to another substance.

• Bronsted-Lowry Base: A substance that accepts a proton (H+) from another substance.

• Example: In the reaction NH3 + H2O → NH4+ + OH-, water acts as a Bronsted-Lowry acid (donates H+), and ammonia acts as a Bronsted-Lowry base (accepts H+).

Amphoteric Substances

Some substances, such as amino acids, can act as either acids or bases depending on the reaction conditions. These are called amphoteric substances.

• Example: The amino acid alanine can donate a proton (act as an acid) or accept a proton (act as a base).

Conjugate Acid-Base Pairs

When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid. These pairs differ by one proton.

• Example: In the reaction H2O + NH3 → OH- + NH4+:

  • H2O (acid) and OH- (conjugate base)

  • NH3 (base) and NH4+ (conjugate acid)

Acid and Base Strength

The strength of an acid or base is determined by its tendency to donate or accept protons. This is quantified by the acid dissociation constant (Ka) and base dissociation constant (Kb).

• Strong acids and strong bases dissociate completely in water.

• Weak acids and weak bases only partially dissociate.

• pKa is the negative logarithm of Ka: $pK_a = -\log K_a$

• Lower pKa means a stronger acid.

pH and pOH

pH is a measure of the hydrogen ion concentration in a solution. It is calculated as:

• $pH = -\log[H^+]$

• $pOH = -\log[OH^-]$

• At 25°C, $pH + pOH = 14$

Buffer Solutions

A buffer solution resists changes in pH when small amounts of acid or base are added. Buffers are typically made from a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Example: A mixture of acetic acid (CH3COOH) and sodium acetate (CH3COONa) forms a buffer.

• Buffer capacity depends on the concentrations of the acid and its conjugate base.

Calculating Concentrations in Acid-Base Reactions

When mixing acids and bases, use stoichiometry to determine the final concentrations of ions in solution.

• Example: If 25 mL of 0.24 M HCl is added to a flask, the number of moles of HCl is:

• $\text{Moles HCl} = 0.024 \text{ L} \times 0.24 \text{ mol/L} = 0.006 \text{ mol}$

• To find the final concentration, divide by the total volume after mixing.

Identifying Acidic and Basic Salts

Salts can be acidic, basic, or neutral depending on the strengths of the acids and bases from which they are derived.

• Acidic salt: Derived from a strong acid and a weak base.

• Basic salt: Derived from a strong base and a weak acid.

• Neutral salt: Derived from a strong acid and a strong base.

Sample Table: Acid Strength and pKa Values

Acid	pKa	Relative Strength
Acetic acid	4.76	Weak
Formic acid	3.74	Stronger (than acetic acid)
Propanoic acid	4.87	Weaker (than acetic acid)
Benzoic acid	4.20	Stronger (than acetic acid)

Key Equations

• $K_a = \frac{[H^+][A^-]}{[HA]}$

• $pK_a = -\log K_a$

• $pH = -\log[H^+]$

• $pOH = -\log[OH^-]$

• $pH + pOH = 14$

Additional info:

• Understanding the periodic table is essential for predicting acid-base properties, as electronegativity and atomic size influence acid strength.

• Buffer solutions are crucial in biological systems to maintain stable pH.