MCAT General Chemistry Review

Electronic Structure and Periodic Table

• Electronic structure
  • Orbital structure of hydrogen atom, principal quantum number n, number of electrons per orbital
  • Ground state, excited states
  • Absorption and emission spectra
  • Quantum numbers l, m, s, and number of quantum states (electrons) per orbital
  • Common names and geometric shapes for orbitals s, p, d
  • Conventional notation for electronic structure
  • Bohr atom
  • Effective nuclear charge
• The periodic table: classification of elements into groups by electronic structure; physical and chemical properties of elements
  • Alkali metals
  • Alkaline earth metals
  • Halogens
  • Noble gases
  • Transition metals
  • Representative elements
  • Metals and nonmetals
  • Oxygen group
• The periodic table: variations of chemical properties with group and row
  • Electronic structure
    • the representative elements
    • the noble gases
    • transition metals
  • Valence electrons
  • First and second ionization energies
    • definition
    • prediction from electronic structure for elements in different groups or rows
  • Electron affinity
    • definition
    • variation with group and row
  • Electronegativity
    • definition
    • comparative values for some representative elements and important groups
  • Electron shells and the sizes of atoms

Bonding

• The ionic bond (electrostatic forces between ions)
  • Electrostatic energy α q₁q₂/r
  • Electrostatic energy α lattice energy
  • Electrostatic force α q₁q₂/r²
• The covalent bond
  • sigma and pi bonds
    • hybrid orbitals (sp³, sp², sp and respective geometries)
    • valence shell electron pair repulsion (VSEPR) theory, predictions of shapes of molecules (e.g., NH₃, H₂O, CO₂)
  • Lewis electron dot formulas
    • resonance structures
    • formal charge
    • Lewis acids and bases
  • Partial ionic character
    • role of electronegativity in determining charge distribution
    • dipole moment

Phases and Phase Equilibria

• Gas phase
  • Absolute temperature, K scale
  • Pressure, simple mercury barometer
  • Molar volume at 0 degrees Celcius and 1 atm = 22.4 L/mol
  • Ideal gas
    • definition
    • ideal gas law PV=nRT
      • Boyle's law
      • Charles' law
      • Avogadro's law
  • Kinetic molecular theory of gases
  • Deviation of real-gas behavior from ideal gas law
    • qualitative
    • quantitative (Van der Waals' equation)
  • Partial pressure, mole fraction
  • Dalton's law relating partial pressure to composition
• Liquid phase: intermolecular forces
  • Hydrogen bonding
  • Dipole interactions
  • Van der Waals' forces (London dispersion forces)
• Phase equilibria
  • Phase changes and phase diagrams
  • Freezing point, melting point, boiling point, condensation point
  • Molality
  • Colligative properties
    • vapor pressure lowering (Raoult's law)
    • boiling point elevation (ΔT_b = K_b*m)
    • freezing point depression (ΔT_f = -K_f*m)
    • osmotic pressure
  • Colloids
  • Henry's Law

Stoichiometry

• Molecular weight
• Empirical formula versus molecular formula
• Metric units commonly used in the context of chemistry
• Description of composition by % mass
• Mole concept; Avagadro's number
• Definition of density
• Oxidation number
  • common oxidizing and reducing agents
  • disproportionation reactions
  • redox titration
• Description of reactions by chemical equations
  • conventions for writing chemical equations
  • balancing equations, including oxidation-reduction equations
  • limiting reactants
  • theoretical yields

Thermodynamics and Thermochemistry

• Energy changes in chemical reactions- thermochemistry
  • Thermodynamic system, state function
  • Conservation of energy
  • Endothermic/exothermic reactions
    • enthalpy H and standard heats of reaction and formation
    • Hess' law of heat summation
  • Bond dissociation energy as related to heats of formation
  • Measurement of heat changes (calorimetry), heat capacity, specific heat capacity (specific heat of water = 4.184 J/g·k)
  • Entropy as a measure of "disorder"; relative entropy for gas, liquid, and crystal states
  • Free energy G
  • Spontaneous reactions and ΔG°
• Thermodynamics
  • Zeroth law (concept of temperature)
  • First law (ΔE = q + w, conservation of energy)
  • Equivalence of mechanical, chemical, electrical and thermal energy units
  • Second law (concept of entropy)
  • Temperature scales, conversion
  • Heat transfer (conduction, convection, radiation)
  • Heat of fusion, heat of vaporization
  • PV diagram (work done = area under or enclosed by curve)
  • Calorimetry

Rate Processes in Chemical Reactions - Kinetics and Equlibrium

• Reaction rates
• Dependence of reaction rate upon concentration of reactants; rate law
  • rate constant
  • reaction order
• Rate determining step
• Dependence of reaction rate on temperature
  • activation energy
    • activated complex or transition state
    • interpretation of energy profiles showing energies of reactants and products, activation energy, ΔH for the reaction
  • Arrhenius equation
• Kinetic control versus thermodynamic control of a reaction
• Catalysts; the special case of enzyme catalysis
• Equilibrium in reversible chemical reactions
  • Law of Mass Action
  • the equilibrium constant
  • application of LeChatelier's principle
• Relationship of the equilibrium constant and standard free energy change

Solution Chemistry

• Ions in solution
  • Anion, cation (common names, formulas and charges for familiar ions; e.g., NH₄⁺, ammonium; PO₄^3-, phosphate; SO₄^2-, sulfate)
  • Hydration, the hydronium ion
• Solubility
  • Units of concentration (e.g., molarity)
  • Solubility product constant, the equilibrium expression
  • Common-ion effect, its use in laboratory separations
  • Complex ion formation
  • Complex ions and solubility
  • Solubility and pH

Acids/Bases

• Acid / base equilibria
  • Bronsted definition of acid, base
  • Ionization of water
    • Kw, its approximate value (Kw = [H+][OH-] = 1*10^-14 at 25°C)
    • definition of pH; pH of pure water
  • Conjugate acids and bases (e.g., amino acids)
  • Strong acids and bases (common examples, e.g., nitric, sulfuric)
  • Weak acids and bases (common examples, e.g. acetic, benzoic)
    • dissociation of weak acids and bases with or without added salt
    • hydrolysis of salts of weak acids or bases
    • calculation of pH of solutions of salts of weak acids or bases
  • Equilibrium constants Ka and Kb: pKa, pKb
  • Buffers
    • definition and concepts (common buffer systems)
    • influence on titration curves
• Titration
  • Indicators
  • Neutralization
  • Interpretation of titration curves
  • Redox titration

Electrochemistry

• Electrolytic cell
  • electrolysis
  • anode, cathode
  • electrolyte
  • Faraday's law relating amount of elements deposited (or gas liberated) at an electrode to current
  • electron flow; oxidation, and reduction at the electrodes
• Galvanic or voltaic cell
  • half reactions
  • reduction potentials; cell potential
  • direction of electron flow

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Old topics

The topics below are outdated. They have been either modified or replaced by the most recent aamc publication.

• E = kQ1Q2/d
• E = lattice energy
• Force attraction = R(n+e)(n-e)/d^2

• Measurement of heat changes (calorimetry); heat capacity; specific heat (specific heat of water = 1 cal per degrees Celcius)
• First law: ΔE = Q - W (conservation of energy)

Source: Official MCAT general chemistry topics were obtained directly from the AAMC.