Exam II Study Guide: Reactions in Aqueous Solution, Thermochemistry, and Electronic Structure

Chapter 4: Reactions in Aqueous Solution

Types of Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water due to the presence of ions.

• Strong Electrolytes: Completely dissociate into ions (e.g., NaCl, HCl).

• Weak Electrolytes: Partially dissociate into ions (e.g., CH3COOH).

• Nonelectrolytes: Do not produce ions in solution (e.g., sugar, ethanol).

Example: NaCl in water dissociates into Na+ and Cl- ions, making it a strong electrolyte.

Reactions in Water: Predicting Products

• Metathesis (Double Displacement) Reactions: Exchange of ions between two compounds. Often forms a precipitate, gas, or weak electrolyte.

• Total Ionic Equation: Shows all soluble ionic substances as ions.

• Net Ionic Equation: Shows only the species that change during the reaction.

Example: Mixing AgNO3 and NaCl in water:

• Molecular: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

• Total Ionic: Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(s) + Na+(aq) + NO3-(aq)

• Net Ionic: Ag+(aq) + Cl-(aq) → AgCl(s)

Table 4.1: (Additional info: Table 4.1 typically lists solubility rules for common ionic compounds.)

Acids and Bases: Neutralization

• Acid: Proton (H+) donor.

• Base: Proton acceptor or OH- donor.

• Neutralization: Acid reacts with base to form water and a salt.

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

Acid Ionization and Polyprotic Acids

• Strong Acids: Completely ionize in water. The 7 strong acids are HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4.

• Polyprotic Acids: Can donate more than one proton (e.g., H2SO4).

Example: H2SO4 ionizes in two steps:

• First: H2SO4 → H+ + HSO4-

• Second: HSO4- ⇌ H+ + SO42-

Metal Carbonates and Acids: Neutralization

• Metal carbonates react with acids to produce a salt, water, and carbon dioxide gas.

Example: CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)

Redox Reactions: Oxidation Numbers

• Oxidation Number: A value assigned to an atom to indicate its degree of oxidation or reduction.

• Rules: (Additional info: Assign based on element, compound, or ion; O is usually -2, H is +1, etc.)

• Oxidized: Species that loses electrons (increase in oxidation number).

• Reduced: Species that gains electrons (decrease in oxidation number).

• Mnemonics: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

Activity Series and Single Replacement Reactions

• Activity Series: Ranks metals by their ability to displace other metals from solution.

• A more active metal will replace a less active metal in a compound.

Example: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) (Zinc is above copper in the activity series.)

Molarity and Dilution

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution: $M_1V_1 = M_2V_2$

Example: To dilute 2.0 M HCl to 1.0 M, mix 50 mL of 2.0 M with 50 mL water.

Solution Stoichiometry

• Use molarity and volume to calculate moles, then use stoichiometry to relate reactants and products.

Example: How many moles of NaCl are in 250 mL of 0.20 M solution? $0.20 \times 0.250 = 0.050$ mol

Titration Calculations

• Titration: Process of determining concentration by reacting a known volume with a standard solution.

• At equivalence point: $M_1V_1 = M_2V_2$ (for 1:1 stoichiometry)

Chapter 5: Thermochemistry

Energy: Kinetic and Potential

• Kinetic Energy (KE): Energy of motion. $KE = \frac{1}{2}mv^2$

• Potential Energy (PE): Energy due to position or composition.

First Law of Thermodynamics

• Law: Energy cannot be created or destroyed, only transferred or converted.

• Formula: $\Delta E = q + w$

• q: Heat exchanged; w: Work done.

• Sign conventions: q > 0 (system gains heat), w > 0 (work done on system).

Energy and Stoichiometry

• Relate energy changes to chemical equations using enthalpy ($\Delta H$).

Example: If $\Delta H = -285.8$ kJ for H2 + 1/2 O2 → H2O, then forming 2 mol H2O releases $2 \times -285.8 = -571.6$ kJ.

State Functions

• Properties that depend only on the initial and final states, not the path (e.g., energy, enthalpy, pressure, temperature).

Enthalpy: Exothermic and Endothermic Reactions

• Enthalpy ($\Delta H$): Heat change at constant pressure.

• Exothermic: Releases heat ($\Delta H < 0$).

• Endothermic: Absorbs heat ($\Delta H > 0$).

Diagram: (Additional info: Exothermic reactions have products lower in energy than reactants; endothermic are higher.)

Calorimetry

• Specific Heat (c): Amount of heat to raise 1 g by 1°C.

• Molar Heat Capacity: Amount of heat to raise 1 mol by 1°C.

• Formula: $q = mc\Delta T$

• Coffee Cup Calorimetry: Measures heat at constant pressure (e.g., dissolving salt or heating metal).

Hess’s Law

• The enthalpy change for a reaction is the same, no matter how many steps it takes.

• Formula: $\Delta H_{\text{overall}} = \sum \Delta H_{\text{steps}}$

Heats of Formation

• Standard Heat of Formation ($\Delta H_f^\circ$): Enthalpy change for forming 1 mol of a compound from its elements in standard states.

• Formula: $\Delta H_{\text{rxn}}^\circ = \sum n\Delta H_f^\circ(\text{products}) - \sum n\Delta H_f^\circ(\text{reactants})$

Bond Enthalpies

• Energy required to break one mole of a bond in the gas phase.

• Formula: $\Delta H = \sum \text{Bond Energies (broken)} - \sum \text{Bond Energies (formed)}$

Chapter 6: Electronic Structure of Atoms

Wave Properties: Speed of Light

• Formula: $c = \lambda \nu$

• c: Speed of light ($3.00 \times 10^8$ m/s)

• \lambda: Wavelength (m)

• \nu: Frequency (Hz)

Particle Properties: Energy of Photons

• Formula: $E = h\nu$

• h: Planck's constant ($6.626 \times 10^{-34}$ J·s)

Bohr Model: Hydrogen Atom and Rydberg Calculations

• Energy Levels: $E_n = -2.18 \times 10^{-18} \frac{1}{n^2}$ J (for H atom)

• Rydberg Equation: $\frac{1}{\lambda} = R \left( \frac{1}{n_1^2} - \frac{1}{n_2^2} \right)$, where $R = 1.097 \times 10^7$ m-1

de Broglie Wavelength Calculations

• Formula: $\lambda = \frac{h}{mv}$

• Shows that particles have wave-like properties.

Quantum Numbers

• n: Principal quantum number (energy level)

• l: Angular momentum quantum number (shape; s=0, p=1, d=2, f=3)

• ml: Magnetic quantum number (orientation)

• ms: Spin quantum number (+1/2 or -1/2)

Orbital Shapes

• s orbital: Spherical shape

• p orbital: Dumbbell shape

• d orbital: Cloverleaf shape (recognize, not draw)

Electron Configurations

• Shorthand: Uses noble gas core (e.g., [Ne]3s23p4)

• Condensed: Lists all occupied orbitals

• Orbital Diagram: Shows electrons as arrows in boxes representing orbitals

Anomalies in Electron Configurations

• Some elements (Cr, Mo, Cu, Ag) have electron configurations that differ from the expected order due to stability of half-filled or filled subshells.

• Example: Cr: [Ar]4s13d5 (not [Ar]4s23d4)

Previous Material

• Review all previously memorized material as needed for Exam I topics.