General Chemistry Review

Ions, Protons, Electrons, and Neutrons

• Ions have an unequal number of protons and electrons.

• Atomic number (smaller number) = number of protons.

• Number of neutrons = mass number - atomic number.

• Number of electrons = atomic number - charge.

• Positive charge means more protons than electrons.

Naming Compounds

• Molecular compounds (nonmetal + nonmetal) use prefixes (mono, di, tri, etc.).

  • Example: N₂O₅ is dinitrogen pentoxide.

  • Do not use mono for the first element.

• Ionic compounds (metal + nonmetal) do not use prefixes.

  • Example: AlCl_3 is aluminum chloride.

• Review polyatomic ions.

Percent Composition

• Percent Composition = (Mass of element / Total mass) * 100\%.

Stoichiometry and Limiting Reactants

• Balance the chemical equation first.

• Grams to Grams Conversion:

  • Grams of A -> Moles of A -> Moles of B -> Grams of B

Molarity

• Molarity = Moles of solute / Liters of solution

Dilution Problems

• Use the equation: M₁V₁ = M₂V₂

• Ensure units of V₁ and V₂ are the same.

• When determining how much water to add, subtract the initial volume from the final volume.

Oxidation States

• The sum of oxidation states in a compound equals the net charge (usually zero).

• Sodium (Group 1) usually has a +1 charge.

• Oxygen usually has a -2 charge.

• Fluorine in a compound is always -1.

• Hydrogen is +1 with nonmetals, -1 with metals.

Titration

• Use M₁V₁ = M₂V₂ adjusting for molar ratios from the balanced equation.

• If H2SO4 is present and NaOH is the tritant M₁V₁(base) = 2 * M₂V₂(acid)

• Dimensional analysis requires a balanced reaction to apply molar ratios accurately.

Gas Laws

• Combined Gas Law:\frac{P_1V_1}{T_1}=\frac{P_2V_2}{T_2}

• Temperature must be in Kelvin.

Density of Gas at STP

• At STP (Standard Temperature and Pressure):

  • T = 0°C (273 K)

  • P = 1 atm.

• 1 mole of any gas occupies 22.4 L.

• Density = (Molar mass) / 22.4

• Ideal Gas Law: PV = nRT

Kinetic Molecular Theory

• Average kinetic energy of gas depends on temperature.

• Average velocity of gas particles depends on temperature.

• Pressure inside a container depends on the total number of moles.

• Real gases behave more ideally at high temperatures and low pressures.

Heat

• q = mc\Delta T where:

  • q = heat energy

  • m = mass in grams

  • c = specific heat capacity

  • \Delta T = change in temperature.

• When converting between phases use: q= n \Delta H_{fus}, where n is in moles.

Phase Changes

• Gas to solid - Deposition

• Solid to liquid - Melting (Endothermic)

• Liquid to gas - Vaporization (Endothermic)

• Solid to gas - Sublimation (Endothermic)

• Liquid to solid - Freezing (Exothermic)

• Gas to liquid - Condensation (Exothermic)

Enthalpy of Reaction

• \Delta H{rxn} = \Sigma \Delta H{products} - \Sigma \Delta H_{reactants}

Thermochemical Equations

• Grams of reactant can be converted to kilojoules using the balanced equation.

Energy of a Photon

• E = hf = \frac{hc}{\lambda}

  • h = Planck's Constant 6.626 \times 10^{-34} Js

  • c = speed of light 3 \times 10^{8} m/s

  • \lambda = wavelength (meters)

Electron Configuration

• Add up number of electrons from exponents. The sum should match the atomic number.

Quantum Numbers

• n = principal quantum number (energy level).

• l = azimuthal quantum number (sublevel shape: s=0, p=1, d=2, f=3).

• ml = magnetic quantum number (orbital orientation, ranges from -l to +l).

• ms = spin quantum number (+1/2 or -1/2).

• The number of orbitals = n^2

• l must be less than n. ml must be between -l and +l.

Electromagnetic Radiation

• From longest wavelength to shortest: Radio waves, Microwaves, Infrared, Visible, UV, X-rays, Gamma rays.

• (Left) Radio waves: longest wavelength.

• (Right) Gamma rays: most energy, highest frequency.

Periodic Trends

• Ionization energy increases up and to the right.

• Atomic radii increase down and to the left.

• Electronegativity increases toward fluorine.

• Metallic character increases down and to the left.

Molecular Geometry

• Trigonal Planar: 120° bond angle.

• Tetrahedral: 109.5° bond angle.

• Bent: less than 120° bond angle.

• Trigonal Pyramidal: 107° bond angle.

Hybridization

• sp3: 4 groups attached to the central atom (e.g., water).

• Know the number of groups for each hybridization.

Polarity

• Molecules with symmetrical arrangement of polar bonds can be nonpolar (e.g., CO_2).

Intermolecular Forces (IMFs)

• Hydrogen bonds > dipole interactions > London dispersion forces.

• H bonded to N, O, or F = hydrogen bonding.

• Nonpolar molecules have only London dispersion forces.

Molality

• Molality = Moles of solute / Kilograms of solvent

• To convert mass percent to molality, assume 100g solution.

Molarity of a Solution

• Molarity = Moles of solute / Liters of solution

Boiling Point Elevation

• \Delta T = K_b * m * i

  • K_b = constant

  • m = molality

  • i = van't Hoff factor (number of ions)

• Salt increases the boiling point.

• Salt decreases the freezing point.

Osmotic Pressure:

• \Pi = MRTi

  • \Pi = osmotic pressure

  • M = molarity

  • R = gas constant.

  • T = temperature in Kelvin

Vapor Pressure Lowering

• P{solution} = X{solvent} * P_{solvent}

  • P_{solution} = vapor pressure of the solution

  • X_{solvent} = mole fraction of the solvent

  • P_{solvent} = vapor pressure of the solvent

• Accounting for the van't Hoff factor is particularly important.

Average Atomic Mass:

• Average Atomic Mass = (mass1 * percent1) + (mass2 * percent2)

• Make sure the percentages are in decimal format.

Percent Yield

• Percent Yield = (Actual Yield / Theoretical Yield) * 100\%.

pH Calculations

• pH + pOH = 14

• pOH = -log[OH^-]

• For strong bases like barium hydroxide, account for the number of hydroxide ions in the formula.