CH 4-2: Chemical Reactions and Stoichiometry: Acid-Base and Redox Reactions

Chapter 4: Chemical Reactions and Stoichiometry

Types of Chemical Reactions

This chapter focuses on two major types of chemical reactions: acid-base reactions and oxidation-reduction (redox) reactions. Understanding these reactions is fundamental to general chemistry, as they are common in both laboratory and real-world contexts.

Acid-Base Reactions

Arrhenius Definitions

• Acid: Any substance that dissociates (or ionizes) in water to generate the hydrogen cation H+ (or H3O+).

• Base: Any substance that reacts with H+ ions formed by acids OR any substance that produces hydroxide ion (OH-) in aqueous solution.

Examples:

• $\mathrm{HNO_3 (aq) \rightarrow H^+ (aq) + NO_3^- (aq)}$

• $\mathrm{NaOH (aq) \rightarrow Na^+ (aq) + OH^- (aq)}$

• $\mathrm{NH_3 (aq) + H_2O (l) \rightarrow NH_4^+ (aq) + OH^- (aq)}$

Features of Acids and Bases

• Strong acids and bases completely ionize in solution and are strong electrolytes.

• Weak acids and bases partially ionize in solution and are weak electrolytes.

Example:

• Beaker A: $\mathrm{HCl (aq) + H_2O (l) \rightarrow H_3O^+ (aq) + Cl^- (aq)}$ (complete ionization)

• Beaker B: $\mathrm{HF (aq) + H_2O (l) \rightleftharpoons H_3O^+ (aq) + F^- (aq)}$ (partial ionization)

Common Strong Acids and Bases

• Acids: HBr, HCl, HI, HNO3, HClO4, H2SO4 (first H only)

• Bases: Group 1 metal hydroxides (e.g., LiOH), Group 2 metal hydroxides (e.g., Ca(OH)2, Sr(OH)2, Ba(OH)2)

Note: You must be able to identify these strong acids and bases.

Acid-Base Neutralization Reactions

Definition: The reaction of an acid and a base to yield an ionic compound (salt) and possibly water.

• General equation: $\mathrm{Acid + Base \rightarrow Salt + Water}$

• Example: $\mathrm{HCl (aq) + NaOH (aq) \rightarrow H_2O (l) + NaCl (aq)}$

• Another example: $\mathrm{KOH (aq) + HCN (aq) \rightarrow KCN (aq) + H_2O (l)}$

The salt product consists of the cation from the base and the anion from the acid.

Net Ionic Equations for Acid-Base Reactions

• Remove spectator ions to show only the species that change during the reaction.

• Example: $\mathrm{OH^- (aq) + NH_4^+ (aq) \rightarrow NH_3 (g) + H_2O (l)}$

Acid-Base Reactions Involving Gas Formation

Certain acid-base reactions produce a gas as a product. These are called gas-evolution reactions.

Compound Type	Example	Gas Formed
Sulfides	$\mathrm{Na_2S + H_2SO_4}$	$\mathrm{H_2S}$
Carbonates/Bicarbonates	$\mathrm{Na_2CO_3 + HCl}$	$\mathrm{CO_2}$
Sulfites/Bisulfites	$\mathrm{Na_2SO_3 + HCl}$	$\mathrm{SO_2}$
Ammonium salts	$\mathrm{NH_4Cl + NaOH}$	$\mathrm{NH_3}$

Practice Example:

• $\mathrm{(NH_4)_2SO_4 (aq) + Ca(OH)_2 (aq) \rightarrow 2 NH_3 (g) + 2 H_2O (l) + CaSO_4 (s)}$

• $\mathrm{HCl (aq) + KHCO_3 (aq) \rightarrow H_2O (l) + CO_2 (g) + KCl (aq)}$

Oxidation-Reduction (Redox) Reactions

Oxidation States and Oxidation Numbers

Atoms and molecules can possess various degrees of charge, known as oxidation states. The oxidation state (or oxidation number) is a formal way to keep track of electrons in redox reactions.

• The more electronegative atom in a covalent bond is assigned all the shared electrons.

Rules for Assigning Oxidation Numbers

1. The oxidation state of each atom in an element is 0.

2. The oxidation state of a monatomic ion is equal to the ion's charge.

3. The sum of the oxidation states of all atoms in:

   • a neutral molecule is 0

   • a polyatomic ion equals the charge of the ion

4. Group 1 metals: +1; Group 2 metals: +2

5. Hydrogen: +1 (except in metal hydrides, where it is -1)

6. In compounds:

   • Fluorine: -1

   • Other group 17 elements: usually -1

   • Oxygen: usually -2

   • Group 16 elements: usually -2

Definitions: Oxidation and Reduction

• Oxidation: Loss of electrons by an atom, molecule, or ion.

  • Example: $\mathrm{Ca \rightarrow Ca^{2+} + 2e^-}$ (Ca was oxidized)

• Reduction: Gain of electrons by an atom, molecule, or ion.

  • Example: $\mathrm{Cl_2 + 2e^- \rightarrow 2Cl^-}$ (Cl2 was reduced)

Memory Aid: OIL RIG (Oxidation Is Loss, Reduction Is Gain); LEO GER (Lose Electrons Oxidation, Gain Electrons Reduction)

Redox Reactions

A redox reaction involves both a gain and loss of electrons, i.e., both an oxidation and a reduction occur. It is convenient to think of the overall redox reaction as two individual half-reactions: an oxidation half-reaction and a reduction half-reaction.

• Oxidation half-reaction: $\mathrm{2Na (s) \rightarrow 2Na^+ (aq) + 2e^-}$

• Reduction half-reaction: $\mathrm{Cl_2 (g) + 2e^- \rightarrow 2Cl^- (aq)}$

• Overall reaction: $\mathrm{2Na (s) + Cl_2 (g) \rightarrow 2NaCl (s)}$

Oxidizing and Reducing Agents

• Oxidizing Agent: Causes another species to be oxidized by accepting/taking electrons from it. The oxidizing agent is reduced.

• Reducing Agent: Causes another species to be reduced by donating electrons to it. The reducing agent is oxidized.

Oxidation results in an increase in oxidation state; reduction results in a decrease in oxidation state.

Identifying Oxidized and Reduced Species

To identify which species are oxidized and reduced in a reaction, assign oxidation numbers to all elements and track their changes.

• Example: $\mathrm{8H^+ + MnO_4^- + 5Fe^{2+} \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O}$

  • Fe2+ is oxidized to Fe3+ (oxidation)

  • MnO4- is reduced to Mn2+ (reduction)

• Another example: $\mathrm{N_2H_4 + O_2 \rightarrow 2NO_2 + 2H_2O}$

  • N2H4 is oxidized; O2 is reduced.

Summary Table: Key Features of Acid-Base and Redox Reactions

Reaction Type	Key Feature	Example
Acid-Base	Transfer of H+ or OH-	$\mathrm{HCl + NaOH \rightarrow NaCl + H_2O}$
Redox	Transfer of electrons (change in oxidation state)	$\mathrm{2Na + Cl_2 \rightarrow 2NaCl}$

Additional info: For more practice, refer to Concept Checks 4.4, 4.5, and 4.6 as indicated in the original notes.