Chapter_16_Lecture_acid_base_equillibria1

Chapter 16: Acid-Base Equilibria

Introduction

• Chemistry as the Central Science

  • Fundamental understanding of acids and bases.

Definitions of Acids and Bases

Arrhenius Definition

• Acid: Substance that increases the concentration of hydrogen ions (H+) in water.

• Base: Substance that increases the concentration of hydroxide ions (OH-) in water.

Brønsted–Lowry Definition

• Acid: Proton donor.

• Base: Proton acceptor.

Properties of Acids and Bases

Characteristics of Brønsted–Lowry Acids and Bases

• Brønsted–Lowry Acid: Must possess a removable (acidic) proton.

• Brønsted–Lowry Base: Must have a pair of non-bonding electrons.

Amphiprotic Compounds

• Amphiprotic: Substances that can act as both acids and bases. Examples include HCO3-, HSO4-, and H2O.

Acid Dissolution in Water

• When an acid dissolves in water,

  • Water acts as a Brønsted–Lowry base and abstracts a proton from the acid.

  • Results in the formation of the conjugate base and a hydronium ion (H3O+).

Conjugate Acid-Base Pairs

• Definition: An acid and base pair that differs only by the presence or absence of a proton.

  • Example: HA (acid) and A- (conjugate base).

Identifying Conjugate Acids and Bases

• Conjugate Bases of several acids:

  • HClO4 → ClO4-

  • H2S → HS-

  • PH4+ → PH3

  • HCO3- → CO3 2-

• Conjugate Acids of several bases:

  • CN- → HCN

  • SO4 2- → HSO4-

  • H2O → H3O+

  • HCO3- → H2CO3

Proton-Transfer Reactions

Example of Equations for HSO3- with Water:

• As Acid:

  • HSO3- + H2O → SO3 2- + H3O+

• As Base:

  • HSO3- + H2O → H2SO3 + OH-

Acid and Base Strength

• Strong Acids: Completely dissociate in water; conjugate bases are weak.

• Weak Acids: Partially dissociate; conjugate bases are weak bases.

• Substances with negligible acidity do not dissociate in water.

Equilibrium in Acid-Base Reactions

• Equilibrium favors the reaction that moves the proton to the stronger base.

  • Example: HCl + H2O → H3O+ + Cl- (Equilibrium lies to the right)

  • Example: CH3CO2H + H2O ⇌ H3O+ + CH3CO2- (Equilibrium favors left)

Autoionization of Water

• Water is amphoteric and can act both as an acid and a base, resulting in autoionization:

  • 2H2O ⇌ H3O+ + OH-

Ion Product Constant

• Ion Product Constant for Water (Kw):

  • Kw = [H3O+][OH-]

  • At 25°C, Kw = 1.0 x 10^-14.

pH Scale

• Definition: pH = -log[H3O+].

• In pure water, pH = 7, hence:

  • Acidic solutions: pH < 7

  • Basic solutions: pH > 7.

Measuring pH

• Less accurate: Litmus paper or indicators.

• More accurate: pH meter.

Strong Acids and Bases

• Strong Acids: Seven strong acids include HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.

• Strong Bases: Soluble hydroxides of alkali and alkaline earth metals.

Weak Acids and Bases

• Dissociation Constants (Ka, Kb):

  • Ka for acid dissociation, Kb for base dissociation.

  • Relationship: Ka * Kb = Kw.

Percent Ionization

• Defined as:

  • % Ionization = ([H+] ionized / [HA] initial) * 100.

Calculating pH from Ka

• Equilibrium calculations involve setting up concentration tables and using the equilibrium expression for dissociation.

Polyprotic Acids

• Can donate more than one proton. The first dissociation is often the most significant.

Hydrolysis of Ions

• Cations and anions react with water producing H+ or OH- ions, respectively.

  • Cations: Lower pH, making solutions acidic (e.g., NH4+).

  • Anions: Act as bases, depending on acid strength from which they are derived.

Factors Affecting Acid Strength

• Bond polarity and strength, stability of the conjugate base, and number of electronegative atoms influence acidity.

Lewis Acids and Bases

• Lewis Acids: Electron-pair acceptors.

• Lewis Bases: Electron-pair donors.

Summary

• Understanding acid-base equilibria is crucial in predicting behavior, strength, and reactions of acidic and basic substances in chemical reactions.