Study Notes for Chemical Reactions - Chapter 5

Chapter 5: Chemical Reactions

Chemical Equations

  • Definition: Chemical equations are a convenient way to represent chemical reactions.

  • Composition:

    • Written in terms of reactants (left side) and products (right side).

    • States of substances indicated by symbols in parentheses:

    • Gases: (g)

    • Liquids: (l)

    • Solids: (s)

    • Aqueous (substances dissolved in water): (aq)

  • Examples:

    • Solid sugar: $C{12}H{22}O_{11} (s)$

    • Liquid water: $H_{2}O (l)$

    • Sugar dissolved in water: $C{12}H{22}O_{11} (aq)$

Reactants and Products of a Chemical Equation

  • Reactants:

    • Defined as the substances written on the left side of the arrow pointing towards products.

    • When two or more reactants are present, they are separated by a plus sign (+).

    • Example: $2H{2} (g) + O{2} (g) → 2H_{2}O (l)$

  • Products:

    • Refers to the substances written on the right side of the arrow.

    • When two or more products exist, they are also separated by a plus sign (+).

Balanced Chemical Equations

  • Definition: A balanced chemical equation has the same number of atoms of each element on both sides of the equation.

  • Law Applied: Balancing is achieved via the law of conservation of matter.

  • Coefficients:

    • Written to the left of reactants or products to achieve balance.

    • Example: $2H{2}(g) + O{2}(g) → 2H_{2}O(l)$

Examples of Unbalanced and Balanced Equations

  • Example Reaction: Methane combustion.

    • Unbalanced Equation: $CH{4}(g) + O{2}(g) → CO{2}(g) + H{2}O(g)$

    • Balanced Equation: $CH{4}(g) + 2O{2}(g) → CO{2}(g) + 2H{2}O(g)$

    • Atom counting: 1 C, 4 H, 4 O on both sides.

Types of Chemical Reactions

  • Classification Scheme: Chemical reactions categorized based on characteristics.

  • Redox Reactions: Include both reduction and oxidation.

Redox Reactions

  • Definition: The term 'redox' combines reduction and oxidation processes.

  • Applicability: Oxidation numbers offer a straightforward approach for working with redox reactions.

Oxidation Numbers (Oxidation States)

  • Definition: Oxidation numbers are assigned to elements in chemical formulas based on specific rules, indicating positive or negative charges.

  • Rules:

    • Rule 1: The oxidation number of any uncombined element is 0. e.g., $Fe (0), Cl_{2} (0), Ca (0)$.

    • Rule 2: For a simple ion, oxidation number equals its charge. e.g., $Mg^{2+}(+2), O^{2-}(-2), Cl^{-}(-1)$.

    • Rule 3: Group IA and IIA elements always have +1 and +2 oxidation states, respectively. e.g., in $Na_{2}S$, $Na = +1$.

    • Rule 4: The oxidation number of hydrogen is always +1. e.g., in $HBr$, H = +1.

    • Rule 5: The oxidation number of oxygen is -2, except in peroxides where it is -1. e.g., in $MgO$, $O = -2$, in $H{2}O{2}$, $O = -1$.

    • Rule 6: The sum of oxidation numbers in a neutral compound equals zero.

    • Example: In $MgSO_{4}$: O.N. of Mg = +2; O.N. of S = +6; O.N. of O = -2. Total = +8 - 8 = 0.

    • Rule 7: The sum of oxidation numbers in a polyatomic ion equals the charge on the ion.

    • Example: For $HCO_{3}^{-}$: O.N. of H = +1; O.N. of C = +4; O.N. of O = -6 results in a total of -1 (charge).

Oxidizing and Reducing Agents

  • Oxidizing Agent: The substance that contains an element that is reduced during the reaction.

  • Reducing Agent: The substance that contains an element that is oxidized during the reaction.

Combination Reactions

  • Definition: Two or more substances react to form a single substance.

  • Types: May be classified as either redox or non-redox reactions.

  • General Form: $A + B → C$

  • Examples:

    • Redox combination: $S(s) + O{2}(g) → SO{2}(g)$

    • Non-redox combination: $N{2}O{5}(g) + H{2}O(l) → 2HNO{3}(aq)$

Single-Replacement Reactions

  • Definition: A single element reacts with a compound, displacing one element from the compound to form a new compound; always redox reactions.

  • General Form: $A + BX → B + AX$

  • Example: $Zn(s) + CuSO{4}(aq) → Cu(s) + ZnSO{4}(aq)$

Double-Replacement Reactions

  • Definition: Involve partner-swapping between two compounds; generally non-redox reactions.

  • General Form: $AX + BY → BX + AY$

  • Example: $Ba(NO{3}){2}(aq) + Na{2}S(aq) → BaS(s) + 2NaNO{3}(aq)$

Energy and Reactions

  • Energy Change: All chemical reactions accompany energy changes; can either absorb or release energy.

  • Exothermic Reactions: Release heat (e.g., combustion of logs).

  • Endothermic Processes: Absorb heat (e.g., melting of ice).

The Mole and Chemical Equations

  • Definition: Stoichiometry connects the mole concept to chemical equations to calculate mass relationships.

  • Balanced Equations: Indicate relationships that can be interpreted using the mole concept, aiding in factor-unit solutions to numerical problems.

Example of Mole Calculations

  • Balanced Reaction: $2H{2}S(g) + 3O{2}(g) → 2SO{2}(g) + 2H{2}O(l)$

  • Mole Statements:

    • In molecules: $2$ molecules $H{2}S$ + $3$ molecules $O{2}$ → $2$ molecules $SO{2}$ + $2$ molecules $H{2}O$

    • In moles: $2$ mol $H{2}S$ + $3$ mol $O{2}$ → $2$ mol $SO{2}$ + $2$ mol $H{2}O$

    • By molecular weight: $68.2$ g $H{2}S$ + $96.0$ g $O{2}$ → $128.2$ g $SO{2}$ + $36.0$ g $H{2}O$

Factor-Unit Method Steps

  1. Write down the known quantity (include numerical value and units).

  2. Set the known quantity equal to the unknown quantity.

  3. Multiply known by factors that cancel the units of the known quantity and yield the unknown.

  4. Perform arithmetic to arrive at the final answer.

Limiting Reactant

  • Definition: The limiting reactant determines the maximum product yield and runs out first.

  • Method: Calculate amounts of product that each reactant can produce; the one yielding the least is limiting.

Reaction Yields

  • Theoretical Yield: The calculated amount of product; can differ from actual lab yield.

  • Actual Yield: The amount of product produced in laboratory conditions.

  • Percentage Yield Calculation:

    • Formula: % Yield = (Actual Yield / Theoretical Yield) x 100%

Ionic Equations

  • Types:

    • Molecular Equations: Represent compounds by their formulas.

    • Total Ionic Equations: Show soluble ionic substances as ions; insoluble substances remain in formula form.

    • Net Ionic Equations: Show only unionized or insoluble materials and changing ions; removes spectator ions.

    • Example: Balanced chemical reaction: $NaCl(aq) + AgNO{3}(aq) → AgCl(s) + NaNO{3}(aq)$.

    • Total ionic: $Na^{+}(aq) + Cl^{-}(aq) + Ag^{+}(aq) + NO{3}^{-}(aq) → AgCl(s) + Na^{+}(aq) + NO{3}^{-}(aq)$.

    • Net ionic: $Cl^{-}(aq) + Ag^{+}(aq) → AgCl(s)$.

Summary of Reaction Types

  • Redox Reactions:

    • Combination: $A + B → AB$

    • Single replacement: $A + BC → AC + B$

    • Decomposition: $AB → A + B$

  • Non-Redox Reactions:

    • Combination: $AB + CD → ABCD$

    • Double replacement: $AB + CD → AD + CB$

    • Decomposition: $ABCD → AB + CD

Conclusion

  • Understanding the various types of chemical reactions and their representations through equations is essential in chemistry, as is the relation to energy changes and stoichiometry in calculating yields and reacting substances.