CHEM 1160 Exam 3 Study Guide (copy)

🧪 CHEM 1160 Exam 3 Study Guide

🧾 Covered Topics:

• Chapter 7: Acids & Bases, Acid Strength, pH, Redox Reactions, Enthalpy

• Chapter 8: Rates, Equilibria, Gibbs Free Energy, Rate Laws

🧱 Chapter 7: Acids & Bases, Redox, Enthalpy

Acid Strength Trends

• Across a Period (→): Acid strength increases with electronegativity
  Example: HF > H₂O > NH₃ > CH₄

• Down a Group (↓): Acid strength increases with size of atom
  Example: HI > HBr > HCl > HF

• Why?

  • ΔH (Enthalpy): Larger atoms = weaker H–X bonds = easier to break.

  • ΔS (Entropy): Larger ions disrupt fewer water molecules = increase in disorder = favorable.

Resonance & Stability

• More resonance in conjugate base = more stable = stronger acid.

• More oxygen atoms (e.g., HClO₄ > HClO₃ > HClO₂) = more resonance = stronger acid.

Strong Acids (Know These!)

• Strong acids fully dissociate:

  • HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄

  • [H₃O⁺] = [acid].

Ka, pKa Relationships

• ext{Ka} = \frac{[\text{products}]}{[\text{reactants}]}

• ext{pKa} = -\log(\text{Ka})

• Large Ka → Strong acid

• Small pKa → Strong acid.

pH and Concentrations

• \text{pH} = -\log[H₃O⁺]

• [H₃O⁺] = 10^{-\text{pH}}

• [H₃O⁺] \cdot [OH⁻] = K_w = 1 \times 10^{-14}

• Percent Ionization: \% = \frac{[H₃O⁺]}{[\text{acid initial}]} \times 100

Strong vs Weak Acids/Bases

• Strong Acids/Bases: Complete ionization (no ICE tables!).

• Weak: Partial ionization → requires ICE tables.

Redox Basics

• Oxidation: Loss of electrons

• Reduction: Gain of electrons.

• Assign oxidation states:

  • Free elements: 0

  • Ions = their charge

  • Oxygen = –2 (except peroxides)

  • Hydrogen = +1 (with nonmetals), –1 (with metals).

How to ID Redox Reactions

• Look for changes in oxidation numbers

• Combustion: redox + O₂ as reactant → produces CO₂ + H₂O.

Bond Enthalpies

• \Delta H = \sum \text{bonds broken} - \sum \text{bonds formed}

• Triple > Double > Single bonds in strength.

• Shorter = Stronger, Longer = Weaker.

⚙ Chapter 8: Rates, Equilibria, Gibbs

Reaction Rates

• Average Rate: \Delta[\text{concentration}] / \Delta\text{time}

• Instantaneous Rate: Slope of tangent to curve at a point.

Rate Laws

• General form: \text{Rate} = k[A]^m[B]^n

• Order = exponent.

• Use initial rate method:

  • Hold one reactant constant

  • Compare how rate changes with other.

Units of k

• Depends on overall reaction order.
  Example: For 3rd order → units: \frac{1}{M^2 \cdot s}

Graphing Method (To Find Order)

• [A] vs time → zero order

• \ln[A] vs time → first order

• \frac{1}{[A]} vs time → second order.

Energy Diagrams

• Activation Energy (Ea): Energy needed to start reaction.

• Catalyst: Lowers Ea via alternate path.

• Exothermic: \Delta H < 0

• Endothermic: \Delta H > 0.

Gibbs Free Energy & Equilibrium

• \Delta G = \Delta H - T\Delta S

• \Delta G^\circ = -RT \ln K

• At equilibrium:

  • \Delta G = 0

  • Q = K.

ICE Tables & Equilibrium Constants

• Use to find concentrations when they change:

• K = \frac{[products]^{coeff}}{[reactants]^{coeff}}

  • K > 1: favors products

  • K < 1: favors reactants.

Le Châtelier’s Principle

• Add reactant: shifts right

• Remove product: shifts right

• Increase Pressure: shifts to side with fewer gas moles

• Increase Temp (exo): shift left

• Catalyst: No shift, just faster equilibrium.

🔥 Exam Tips

• Got it — here's a clean, exam-ready summary of those points, formatted for memorization or quick review. You can paste this into your notes or flashcards:

🔥 CHEM 1160 Exam Strategy Checklist

🧪 Strong Acids & Bases

• Memorize:

  • Strong Acids: HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄

  • Strong Bases: Group 1 and heavy Group 2 hydroxides (e.g., NaOH, Ca(OH)₂)

  • Properties: Fully ionize in water

  • Do not use ICE tables (use direct stoichiometry).

📊 ICE Tables + Equilibrium

• Practice solving for:

  • [H₃O⁺], pH, percent ionization

  • Using KaKa and extpKaextpKa: pKa=−log⁡KapKa=−logKa

  • Know how to simplify when x \ll [\text{initial}].

⚡ Redox Reactions

• Oxidation: Loss of electrons, Reduction: Gain.

• Assign oxidation states using these quick rules:

  • Free element: 0

  • Monoatomic ion: its charge

  • Oxygen: –2 (except peroxides)

  • Hydrogen: +1 (except in metal hydrides: –1).

• Balance redox with half-reactions in acidic or basic solution.

🧠 Rate Laws

• Understand: \text{Rate} = k[A]^m[B]^n

  • Using Experimental Data: Collect data by conducting experiments at different initial concentrations of reactants while keeping others constant. The rate of reaction (D) can be expressed as: ext{Rate} = k[A]^m[B]^n Where:

    • k = rate constant

    • [A] and [B] = concentrations of reactants

    • m and n = reaction orders with respect to each reactant.

  • Finding Reaction Order from Graphs: The reaction order can be identified through plots of concentration against time. Each order has a characteristic graph:

    • Zero Order: A plot of [A] versus time yields a straight line, indicating a constant rate independent of concentration.

    • First Order: Plotting ext{ln}[A] versus time produces a straight line; this indicates that the rate is directly proportional to the concentration of one reactant.

    • Second Order: A graph of rac{1}{[A]} versus time will yield a straight line, signifying that the rate depends on the square of the concentration of the reactant.

  • Determining Units of k: The units of the rate constant k vary depending on the overall reaction order:

    • For a Zero Order reaction: k has units of M imes s^{-1} (concentration per time).

    • For a First Order reaction: k has units of s^{-1} (inverse time).

    • For a Second Order reaction: k has units of M^{-1} imes s^{-1} (inverse concentration times inverse time).

📈 Graphs in Kinetics

• [A] vs. time → zero order

• \ln[A] vs. time → first order

• \frac{1}{[A]} vs. time → second order

• Use the one that gives a straight line to find order.

⚖ Le Châtelier’s Principle

• Predict shifts when:

  • Concentration changes: ↑ Reactant → shift right

  • Pressure increases: shift to side with fewer gas moles

  • Temperature increases:

  • Exothermic → shift left

  • Endothermic → shift right

  • Catalysts do NOT affect equilibrium — only speed.

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