Full AAMC Science Content Outline

Covalent Bond (GC)

• Lewis Electron Dot formulas

  • Resonance structures

  • Formal charge

  • Lewis acids and bases

• Partial ionic character

  • Role of electronegativity in determining charge distribution

  • Dipole Moment

• σ and π bonds

  • Hybrid orbitals: sp3, sp2, sp and respective geometries

  • Valence shell electron pair repulsion and the prediction of shapes of molecules (e.g., NH3, H2O, CO2)

  • Structural formulas for molecules involving H, C, N, O, F, S, P, Si, Cl

  • Delocalized electrons and resonance in ions and molecules

• Multiple bonding

  • Effect on bond length and bond energies

  • Rigidity in molecular structure

Stoichiometry (GC)

• Molecular weight

• Empirical versus molecular formula

• Metric units commonly used in the context of chemistry

• Description of composition by percent mass

• Mole concept, Avogadro’s number NA

• Definition of density

• Oxidation number

  • Common oxidizing and reducing agents

  • Disproportionation reactions

• Description of reactions by chemical equations

  • Conventions for writing chemical equations

  • Balancing equations, including redox equations

  • Limiting reactants

  • Theoretical yields

Solubility (GC)

• Units of concentration (e.g., molarity)

• Solubility product constant; the equilibrium expression Ksp

• Common-ion effect, its use in laboratory separations

  • Complex ion formation

  • Complex ions and solubility

  • Solubility and pH

Electrochemistry (GC)

• Concentration cell: direction of electron flow, Nernst equation

• 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 cells

  • Half-reactions

  • Reduction potentials; cell potential

  • Direction of electron flow

• Concentration cell

• Batteries

  • Electromotive force, Voltage

  • Lead-storage batteries

  • Nickel-cadmium batteries

Gas Phase (GC, PHY)

• Absolute temperature, (K) Kelvin Scale

• Pressure, simple mercury barometer

• Molar volume at 0°C and 1 atm = 22.4 L/mol

• Ideal gas

  • Definition

  • Ideal Gas Law: PV = nRT

  • Boyle’s Law: PV = constant

  • Charles’ Law: V/T = constant

  • Avogadro’s Law: V/n = constant

• Kinetic Molecular Theory of Gases

  • Heat capacity at constant volume and at constant pressure (PHY)

  • Boltzmann’s Constant (PHY)

• 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

Energy Changes in Chemical Reactions – Thermochemistry, Thermodynamics (GC, PHYC)

• Thermodynamic system – state function

• Zeroth Law – concept of temperature

• First Law - conservation of energy in thermodynamic processes

• PV diagram: work done = area under or enclosed by curve (PHY)

• Second Law – concept of entropy

  • Entropy as a measure of “disorder”

  • Relative entropy for gas, liquid, and crystal states

• Measurement of heat changes (calorimetry), heat capacity, specific heat

• Heat transfer – conduction, convection, radiation (PHY)

• Endothermic/exothermic reactions (GC)

  • Enthalpy, H, and standard heats of reaction and formation

  • Hess’ Law of Heat Summation

• Bond dissociation energy as related to heats of formation (GC)

• Free energy: G (GC)

• Spontaneous reactions and ΔG° (GC)

• Coefficient of expansion (PHY)

• Heat of fusion, heat of vaporization

• Phase diagram: pressure and temperature

Rate Processes in Chemical Reactions - Kinetics and Equilibrium (GC)

• Reaction rate

• Dependence of reaction rate on concentration of reactants

  • Rate law, rate constant

  • Reaction order

• Rate-determining step

• Dependence of reaction rate upon temperature

  • Activation energy

    • Activated complex or transition state

    • Interpretation of energy profiles showing energies of reactants, products, activation energy, and ΔH for the reaction

  • Use of the Arrhenius Equation

• Kinetic control versus thermodynamic control of a reaction

• Catalysts

• Equilibrium in reversible chemical reactions

  • Law of Mass Action

  • Equilibrium Constant

  • Application of Le Châtelier’s Principle

• Relationship of the equilibrium constant and ΔG°

Principles of Bioenergetics (BC, GC)

• Bioenergetics/thermodynamics

  • Free energy/Keq

    • Equilibrium constant

    • Relationship of the equilibrium constant and ΔG°

  • Concentration

    • Le Châtelier’s Principle

  • Endothermic/exothermic reactions

  • Free energy: G

  • Spontaneous reactions and ΔG°

Acid/Base Equilibria (GC, BC)

• Brønsted–Lowry definition of acid, base

• Ionization of water

  • Kw, its approximate value (Kw = [H+][OH−] = 10−14 at 25°C, 1 atm)

  • Definition of pH: pH of pure water

• Conjugate acids and bases (e.g., NH4+ and NH3)

• Strong acids and bases (e.g., nitric, sulfuric)

• Weak acids and bases (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

Ions in Solutions (GC, BC)

• Anion, cation: common names, formulas and charges for familiar ions (e.g., NH4+ ammonium, PO43− phosphate, SO42− sulfate)

• Hydration, the hydronium ion

Liquid Phase - Intermolecular Forces (GC)

• Hydrogen bonding

• Dipole Interactions

• Van der Waals’ Forces (London dispersion forces)