Chapter 10: Acids and Bases

10.1 Acids and Bases: Definitions

This section introduces the fundamental definitions of acids and bases, including Arrhenius and Brønsted-Lowry concepts, and the idea of conjugate acid-base pairs.

• Arrhenius Acid: A substance that produces hydronium ions (H3O+) when dissolved in water.

• Arrhenius Base: A substance that produces hydroxide ions (OH-) when dissolved in water.

• Brønsted-Lowry Acid: A substance that donates a proton (H+) to another molecule or ion; acids are proton donors.

• Brønsted-Lowry Base: A substance that accepts a proton; bases must have a lone pair of electrons to accept H+.

• Monoprotic, Diprotic, Triprotic Acids: Acids that can donate one, two, or three protons, respectively (e.g., HCl is monoprotic, H2SO4 is diprotic, H3PO4 is triprotic).

• Conjugate Acid-Base Pair: Two species that differ by one proton; found on opposite sides of a reaction.

Example: The cyanide ion (CN-) is a Brønsted-Lowry base; its conjugate acid is HCN.

10.2 Acid and Base Strength

Acids and bases vary in strength, which affects their behavior in solution and the direction of proton transfer reactions.

• Strong Acids: Completely dissociate in water; examples include HCl, HNO3, H2SO4.

• Weak Acids: Partially dissociate; examples include acetic acid (CH3COOH).

• Strong Bases: Completely dissociate; examples include NaOH, KOH.

• Weak Bases: Partially accept protons; examples include ammonia (NH3).

• Polyprotic Acids: Undergo stepwise dissociation; each successive proton is harder to remove.

• Inverse Relationship: The stronger the acid, the weaker its conjugate base, and vice versa.

Example: The equilibrium of a proton-transfer reaction favors formation of the weaker acid and base.

10.3 Acid Dissociation Constants

The strength of a weak acid is quantified by its acid dissociation constant (Ka), which is derived from the equilibrium expression.

• Acid Dissociation Constant (Ka):

• Strong Acids: Have Ka values much greater than 1.

• Weak Acids: Have Ka values much less than 1.

• Polyprotic Acids: Each successive Ka is lower.

Example: Most organic acids (containing –COOH) have Ka values near 10-5.

10.4 Water as Both an Acid and a Base

Water is amphoteric, meaning it can act as both an acid and a base. Its self-ionization is described by the ion-product constant.

• Amphoteric: A substance that can act as either an acid or a base.

• Self-Ionization of Water:

• Ion-Product Constant (Kw): at 25°C

• Acidic, Neutral, Basic Solutions: Determined by relative concentrations of H+ and OH-.

Example: Milk is slightly acidic if its [H+] is greater than 1.0 × 10-7 M.

10.5 Measuring Acidity in an Aqueous Solution: The pH Scale

The pH scale is a logarithmic measure of the acidity or basicity of a solution, ranging typically from 0 to 14.

• pH Definition:

• pH Range: 0 (very acidic) to 14 (very basic); 7 is neutral.

• Logarithmic Nature: Each unit change in pH represents a tenfold change in [H+].

• Acid-Base Indicators: Dyes that change color depending on pH.

Example: Coffee with [H+] = 1.0 × 10-5 M has pH = 5.0.

10.6 Working with pH

Converting between pH and [H+] involves logarithms and antilogarithms, with attention to significant figures.

• From pH to [H+]:

• From [H+] to pH:

• Significant Figures: The number of decimal places in pH equals the number of significant figures in [H+].

Example: A soft drink with pH 3.1 has [H+] = 7.9 × 10-4 M.

10.7 Acid and Base Equivalents

Acid and base equivalents are used to quantify neutralization reactions, and normality is a related concentration unit.

• Equivalent of Acid: Amount containing 1 mole of H+ ions.

• Equivalent of Base: Amount containing 1 mole of OH- ions.

• Normality (N): Number of equivalents per liter; number of H+ or OH- ions per formula unit.

• Monoprotic Acids: Normality equals molarity; for polyprotic acids, normality is higher.

Example: 3.1 g of a diprotic acid (molar mass 34.0 g) contains 0.182 equivalents.

10.8 Some Common Acid-Base Reactions

Acids and bases undergo characteristic reactions, including neutralization and reactions with carbonates and ammonia.

• Acid + Hydroxide: Neutralization yields water and a salt.

• Acid + Bicarbonate/Carbonate: Yields carbonic acid, which decomposes to CO2 and water.

• Acid + Ammonia: Yields ammonium salts.

• Amines: Organic compounds with nitrogen react similarly to ammonia.

Example: Marble (CaCO3) reacts with HCl to release CO2 bubbles.

10.9 Acidity and Basicity of Salt Solutions

The pH of a salt solution depends on the acid and base from which the salt is formed.

• Salt from Strong Acid + Strong Base: Neutral solution (e.g., NaCl).

• Salt from Strong Acid + Weak Base: Acidic solution (e.g., NH4Cl).

• Salt from Weak Acid + Strong Base: Basic solution (e.g., NaCN).

• Both Ions React: The ion that reacts more strongly with water determines pH.

Example: NaCN forms a basic solution because CN- is a weak base.

10.10 Buffer Solutions

Buffers are solutions that resist changes in pH, typically composed of a weak acid and its conjugate base.

• Buffer: Combination of substances that prevent drastic pH changes.

• Henderson-Hasselbalch Equation:

• Effective Buffer: Ka close to desired pH; ratio of [A-] to [HA] near 1; buffer components much greater than added acid/base.

Example: A buffer with 0.100 M HF and 0.120 M NaF can be used to calculate pH using the Henderson-Hasselbalch equation.

10.11 Titration

Titration is a technique for determining the concentration of an acid or base by reacting it with a solution of known concentration.

• Titration: Procedure for determining total acid or base concentration.

• End Point: The point at which neutralization is complete, indicated by a color change.

• Calculation: For 1:1 reactions, ; for non-1:1, use equivalents and normality.

• Normality Equation:

Example: 5.00 mL of vinegar titrated with 44.5 mL of 0.100 M NaOH yields a molarity of acetic acid, which can be converted to equivalents and milliequivalents per liter.

Additional info:

• Acid-base reactions are central to many biological and environmental processes.

• Buffer systems are critical in maintaining physiological pH, such as in blood.

• Understanding titration is essential for laboratory analysis of unknown solutions.