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CHY-103 General Chemistry I

Chapter 4: Chemical Reactions and Stoichiometry

Types of Chemical Reactions

  • Acid-base reactions

  • Oxidation Reduction reactions (Redox)

  • Note: Students are not responsible for balancing redox equations on pages 121 - 124.

Arrhenius Definitions of Acids and Bases

Acid-Base Reactions

  1. Acid: Any substance that dissociates (or ionizes) in water to generate the hydrogen cation H^+ (or hydronium cation H_3O^+).

    • Example: HNO3(aq) ightarrow H^+(aq) + NO3^-(aq)

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

    • Example: NaOH(aq)
      ightarrow Na^+(aq) + OH^-(aq)

    • Reaction: H^+(aq) + OH^-(aq)
      ightarrow H_2O(l)

    • Additional Example: NH3(aq) + H2O(l)
      ightleftharpoons NH_4^+(aq) + OH^-(aq)

Features of Acids and Bases

  • Strong Acids and Bases: Completely ionize in solution, thus are strong electrolytes.

    • Example Reaction for Strong Acid: HCl(aq) + H2O(l) ightarrow H3O^+(aq) + Cl^-(aq)

  • Weak Acids and Bases: Partially ionize in solution, thus are weak electrolytes.

    • Example Reaction for Weak Acid: HF(aq) + H2O(l) ightleftharpoons H3O^+(aq) + F^-(aq)

Common Strong Acids and Bases

Strong Acids

  • Hydrobromic Acid (HBr)

  • Hydrochloric Acid (HCl)

  • Hydroiodic Acid (HI)

  • Nitric Acid (HNO3)

  • Perchloric Acid (HClO4)

  • Sulfuric Acid (H2SO4) (for loss of first H only)

Strong Bases

  • Group 1 Metal Hydroxides (e.g., LiOH)

  • Heavy Group 2 Metal Hydroxides (e.g., Ca(OH){2}, Sr(OH){2}, Ba(OH)_{2})

Exam Alert

  • Students must know and be able to identify these acids and bases.

Acid-Base Neutralization Reactions

  • Definition: Reaction of an acid and a base to yield an ionic compound (or salt) and possibly water.

    • Typical Reaction Example:

      • HCl(aq) + NaOH(aq)
        ightarrow H_2O(l) + NaCl(aq)

      • Where:

        • HCl = acid

        • NaOH = base

        • H_2O = water

        • NaCl = salt

    • Another Example:

      • KOH(aq) + HCN(aq)
        ightarrow KCN(aq) + H_2O(l)

    • Net Ionic Equation for the Reaction: Part where students are prompted to determine the net ionic equation for the above reaction.

Acid-Base Reactions Involving Gas Formation

  • Table of Types of Compounds That Undergo Gas-Evolution Reactions:

Reactant Type

Intermediate

Product

Gas Evolved

Sulfides

None

H_2S

H_2S

Carbonates and Bicarbonates

H2CO3

CO_2

CO_2

Sulfites and Bisulfites

H2SO3

SO_2

SO_2

Ammonium

NH_4Cl

NH_3

NH_3

Example Reactions for Gas Formation

  • 2 HCl(aq) + K2S (aq) ightarrow H2S(g) + 2 KCl(aq)

  • 2 HCl(aq) + K2CO3(aq)
    ightarrow H2O(l) + CO2(g) + 2 KCl(aq)

  • 2 HCl(aq) + K2SO3(aq)
    ightarrow H2O(l) + SO2(g) + 2 KCl(aq)

  • NH4Cl(aq) + KOH(aq) ightarrow H2O(l) + NH_3(g) + KCl(aq)

Oxidation and Reduction (Redox) Reactions

  • Observations: Atoms and molecules can possess various degrees of charge. E.g., Iron can exist as Fe, Fe^{2+}, Fe^{3+}.

  • The oxidation state or oxidation number is used to track electrons in redox reactions, assuming the more electronegative atom will take all electrons in a covalent bond.

Definitions

  1. Oxidation: Loss of electrons by an atom, molecule, ion, etc.

    • Example: Ca
      ightarrow Ca^{2+} + 2e^-

    • Here, calcium is oxidized.

  2. Reduction: Gain of electrons by an atom, molecule, ion, etc.

    • Example: Cl_2 + 2e^-
      ightarrow 2Cl

    • Here, chlorine is reduced.

Memory Aid

  • LEO GER: Lose Electrons = Oxidation; Gain Electrons = Reduction

Example Reactions

  1. 2 H2(g) + O2(g)
    ightarrow 2 H_2O(g) : Hydrogen oxidized and oxygen reduced.

  2. 2 Na(s) + Cl_2(g)
    ightarrow 2 NaCl(s) : Sodium is oxidized, chlorine is reduced.

Rules for Assigning Oxidation Numbers (O.N.)

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

  2. The oxidation state of an atom in a monoatomic ion equals the ion's charge.

  3. The sum of oxidation states:

    • a. In a neutral molecule, the sum is always zero.

    • b. In a polyatomic ion, it equals the charge of the ion.

  4. In compounds, metals have positive oxidation states:

    • a. Group 1 metals: +1

    • b. Group 2 metals: +2

  5. The common oxidation state of hydrogen in compounds is +1; in metal hydrides, it is -1.

  6. Nonmetals typically have negative oxidation states:

    • a. Fluorine: -1

    • b. Group 17 elements: typically -1

    • c. Oxygen: usually -2

    • d. Group 16 elements: usually -2

    • e. Group 15 elements: usually -3

Example Assignments

  • Assign Oxidation Numbers to the Following:

    • CO_2

    • PbCl_2

    • Tough Example: Fe3O4

    • NH_3

    • Concept Check: Determines oxidation number of different compounds and ions.

Redox Reaction Composition

  • A redox reaction involves both the gain and loss of electrons.

    • Consider oxidation and reduction as two half-reactions:

    • Oxidation Half-Reaction: 2 Na(s)
      ightarrow 2 Na^+ + 2 e^-

    • Reduction Half-Reaction: Cl_2(g) + 2 e^-
      ightarrow 2 Cl

  • Overall Reaction: 2 Na(s) + Cl_2(g)
    ightarrow 2 NaCl(s)

More Definitions

  1. Oxidizing Agent: Causes another species to be oxidized by accepting/taking electrons from it.

  2. Reducing Agent: Causes another species to be reduced by donating electrons to it.

  • Key Statement: "The oxidizing agent is reduced, and the reducing agent is oxidized."

  • Oxidation increases oxidation state (loss of electrons).

  • Reduction decreases oxidation state (gain of electrons).

Concept Check Questions

  • Identify the species being oxidized and reduced in examples.

  • Additional questions to reinforce understanding may be posed to students as part of the learning assessment.