General Chemistry: Key Concepts and Practice Questions

Intermolecular Forces and Molecular Structure

Types of Intermolecular Forces

Intermolecular forces are the forces of attraction between molecules, which influence physical properties such as boiling and melting points.

• London Dispersion Forces: Weak, temporary attractions due to momentary dipoles in all molecules, especially nonpolar ones.

• Hydrogen Bonding: A strong type of dipole-dipole interaction occurring when hydrogen is bonded to highly electronegative atoms (N, O, or F). For N-H bonds, the attraction between dipoles is a hydrogen bond.

• Ionic Bonds: Electrostatic attraction between oppositely charged ions.

• Covalent Bonds: Sharing of electron pairs between atoms within a molecule (not an intermolecular force).

• Example: Water (H2O) exhibits hydrogen bonding due to O-H bonds.

Hybridization and Molecular Orbitals

Types of Hybrid Orbitals

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• sp3 Hybridization: Four electron domains (e.g., methane, CH4).

• sp2 Hybridization: Three electron domains (e.g., ethene, C2H4).

• sp Hybridization: Two electron domains (e.g., acetylene, C2H2).

• Example: In ethane (CH3CH3), all carbons are sp3 hybridized.

Physical Properties: Boiling Points

Boiling Point Trends in Alkanes

Boiling points increase with molecular size and mass due to stronger London dispersion forces.

• Order: C2H6 < C3H8 < C4H10 < C5H12

• Example: Pentane (C5H12) boils at a higher temperature than ethane (C2H6).

Atomic Structure and Periodic Trends

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom. It increases across a period and decreases down a group.

• Order (smallest to largest): Cs < Kr < Ar < Ne

• Example: Neon (Ne) has the highest first ionization energy among the listed elements.

Isotopes and Atomic Mass

Atomic mass is the weighted average of the masses of an element's isotopes.

• Calculation: $\text{Atomic mass} = (\%\ ^{12}X \times 12) + (\%\ ^{13}X \times 13)$

• Example: For 20% $^{12}X$ and 80% $^{13}X$, atomic mass ≈ 12.8

Electron Configuration

Valence electron configuration determines chemical reactivity and group placement.

• Example: Period 3, Group 7A element (Chlorine): $3s^2 3p^5$

Maximum Electrons in Orbitals

• Each orbital (including f): Maximum of 2 electrons (Pauli exclusion principle).

Rutherford's Gold Foil Experiment

• Conclusion: Atoms are mostly empty space; most alpha particles passed through, some deflected at large angles.

Redox Chemistry

Oxidation Numbers and Agents

• Oxidation Number: The charge an atom would have if electrons were assigned according to certain rules.

• Chromium Example: In CrO42-, Cr is +6.

• Oxidizing Agent: Substance that causes oxidation by accepting electrons.

• Example: In the reaction of Fe2+ to Fe3+, an oxidizing agent is required.

Paramagnetism

• Paramagnetic Species: Have unpaired electrons (e.g., Fe2+).

Molecular Geometry and Bonding

VSEPR Theory

• ML4 with 2 lone pairs: Square planar geometry (e.g., XeF4).

Thermochemistry

First Law of Thermodynamics

• Equation: $\Delta E = q + w$

• Example: System absorbs 500 kJ heat, does 20 kJ work: $\Delta E = 500 - 20 = 480$ kJ

Enthalpy Calculations

• Combustion and Formation: Use standard enthalpies of formation to calculate reaction enthalpy.

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

Stoichiometry and Chemical Quantities

Empirical Formula Determination

• Steps: Convert grams to moles for each element, divide by smallest number of moles, write formula.

Limiting Reactant and Yield Calculations

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Example: Calculating grams of AgCl formed from given volumes and concentrations of AgNO3 and NH4Cl.

Electrolytes

• Strongest Electrolyte: Compound that dissociates completely and produces the most ions per formula unit (e.g., AlCl3).

Periodic Table and Atomic Structure

Protons, Neutrons, and Electrons

• Notation: $^{A}_{Z}X$ where A = mass number, Z = atomic number.

• Example: $^{131}_{53}I$ has 53 protons, 78 neutrons, and (for I-) 54 electrons.

Species Identification

• Given: p+ = 25, n = 27, e- = 23 → Mn2+

Nomenclature and Formula Writing

Compound Naming

• TiCl4: Titanium (IV) chloride

• Ammonium sulfate: (NH4)2SO4

Precipitation and Solution Chemistry

Precipitate Formation

• Solubility Rules: Most hydroxides are insoluble except those of alkali metals and Ba2+.

• Example: FeCl2 + NaOH forms Fe(OH)2 precipitate.

Stoichiometry in Solutions

• Example: Calculating moles of Iodine in ZnI2: 4.50 mol ZnI2 × 2 = 9.00 mol I

Tables

Periodic Table (Partial)

1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
H												B	C	N	O	F	Ne
Li	Be											Na	Mg			Al	Si