Acids and Bases: Properties, Definitions, and Calculations – GOB Chemistry Study Guide

Acids and Bases

Introduction to Acids

Acids are covalent compounds that contain a hydrogen ion (H+) connected to a nonmetal or a polyatomic ion. The hydrogen is typically found at the beginning of the formula, except for acetic acid.

• Common Acids: HCl, H2S, HCN, H3PO4, CH3COOH

• Structure: Acids generally have hydrogen attached to a nonmetal or polyatomic group.

• Example: HCl (hydrochloric acid) is a strong acid commonly found in laboratory and industrial settings.

Characteristics of Acids

Acids exhibit distinct properties in aqueous solutions due to their ability to donate protons (H+).

• Dissociation: Acids dissociate in water to produce H+ ions.

• Taste and Feel: Acids typically taste sour and can be corrosive.

• Reactivity: Acids react with metals to form hydrogen gas.

• Litmus Paper: Acids turn blue litmus paper red.

Dissociation	Taste and Feel	Reactivity	Litmus Paper
Acids dissociate in H2O (e.g., HCl → H+ + Cl-)	Sour taste, corrosive	React with metals to form H2	Blue litmus turns red

Introduction to Bases

Bases are ionic compounds containing a metal cation connected to a basic anion, or hydrogen-containing covalent compounds called amines.

• Common Bases: NaOH, KOH, Sr(OH)2, NH3

• Structure: Bases often contain the hydroxide ion (OH-).

• Example: NaOH (sodium hydroxide) is a strong base used in cleaning agents.

Characteristics of Bases

Bases share common properties in aqueous solutions due to their ability to accept protons or release hydroxide ions.

• Dissociation: Bases dissociate in water to produce OH- ions.

• Taste and Feel: Bases taste bitter and feel slippery.

• Litmus Paper: Bases turn red litmus paper blue.

Dissociation	Taste and Feel	Litmus Paper
Bases ionize in H2O (e.g., NaOH → Na+ + OH-)	Bitter taste, slippery feel	Red litmus turns blue

Definitions of Acids and Bases

Arrhenius Definition

Developed by Svante Arrhenius, this definition states:

• Acid: Increases the concentration of H+ ions when dissolved in water.

• Base: Increases the concentration of OH- ions when dissolved in water.

Brønsted-Lowry Definition

Introduced by Johannes Brønsted and Thomas Lowry in 1923, this definition expands the concept:

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

• Brønsted-Lowry acids and bases always occur in pairs called conjugate acid-base pairs.

Conjugate Acids and Bases

When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

• Conjugate Acid: Species formed when a base gains a proton.

• Conjugate Base: Species formed when an acid loses a proton.

• Example: HF + H2O ⇌ F- + H3O+

Strength of Acids and Bases

Strong vs. Weak Acids and Bases

Acids and bases are classified based on their degree of ionization in water.

• Strong Acid: Completely dissociates in water (e.g., HCl, HNO3).

• Weak Acid: Partially dissociates in water (e.g., CH3COOH).

• Strong Base: Completely dissociates in water (e.g., NaOH, KOH).

• Weak Base: Partially dissociates in water (e.g., NH3).

Strong Acids	Strong Bases
HCl, HBr, HI, HNO3, H2SO4, HClO4	LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Strength of Conjugate Acids and Bases

There is an inverse relationship between the strength of an acid and the strength of its conjugate base.

• Strong acid → weak conjugate base

• Weak acid → strong conjugate base

• Example: HCl (strong acid) forms Cl- (weak conjugate base)

Equilibrium and Ionization Constants

Acid Dissociation Constant ($K_a$)

The strength of a weak acid is measured by its acid dissociation constant, $K_a$.

• Expression: $K_a = \frac{\text{Products}}{\text{Reactants}}$

• Example: $HA + H_2O \rightleftharpoons A^- + H_3O^+$

Base Dissociation Constant ($K_b$)

For weak bases, the base dissociation constant, $K_b$, is used.

• Relationship: $K_w = K_a \times K_b$

• Where: $K_w$ is the ion product constant for water.

pH and pOH Calculations

pH and pOH Definitions

pH and pOH are logarithmic measures of the concentration of hydrogen and hydroxide ions in solution.

• pH: $pH = -\log[H^+]$

• pOH: $pOH = -\log[OH^-]$

• Relationship: $pH + pOH = 14$

Classification by pH

• pH > 7: Basic solution, [H+] < [OH-]

• pH < 7: Acidic solution, [H+] > [OH-]

• pH = 7: Neutral solution, [H+] = [OH-]

Practice and Application

Sample Dissociation Equations

• $Sr(OH)_2 (s) \rightarrow Sr^{2+} (aq) + 2OH^- (aq)$

• $H_2SO_4 (l) \rightarrow 2H^+ (aq) + SO_4^{2-} (aq)$

• $HBO_3 (s) \rightarrow 3H^+ (aq) + BO_3^{3-} (aq)$

Conjugate Acid-Base Pairs

• Example: HF + H2O ⇌ F- + H3O+

• Acid: HF, Base: H2O, Conjugate Acid: H3O+, Conjugate Base: F-

Calculating Ion Concentrations

• Hydroxide Ion: $[OH^-]$

• Hydronium Ion: $[H_3O^+]$

• Relationship: $[H^+][OH^-] = K_w$

Normality (N)

Normality is a measure of concentration equivalent to molarity multiplied by the number of equivalents per mole.

• Formula: Normality = Equivalents / Liter

Summary Table: Strong Acids and Bases

Strong Acids	Strong Bases
HCl, HBr, HI, HNO3, H2SO4, HClO4	LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Key Equations

• $K_a = \frac{[A^-][H_3O^+]}{[HA]}$

• $K_b = \frac{[BH^+][OH^-]}{[B]}$

• $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ (at 25°C)

• $pH = -\log[H^+]$

• $pOH = -\log[OH^-]$

• $pH + pOH = 14$

Additional info:

• Some context and examples were inferred to ensure completeness and clarity for GOB Chemistry students.

• Tables and equations were reconstructed for clarity and academic accuracy.