MCAT General Chemistry Syllabus

General Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for General Chemistry in MCAT, not a summary of it.

1. 1

   Atomic Structure and Properties

   5 topics

   • Atomic Theory
     • Dalton’s Atomic Theory
     • Thomson’s Discovery of Electrons
     • Rutherford’s Nuclear Model
     • Bohr’s Model of the Atom
     • Quantum Mechanical Model
   • Atomic Number and Mass Number
   • Isotopes
   • Electron Configuration
     • Aufbau Principle
     • Pauli Exclusion Principle
     • Hund’s Rule
   • Periodic Table Trends
     • Atomic and Ionic Radius
     • Ionization Energy
     • Electron Affinity
     • Electronegativity
2. 2

   Chemical Bonding and Molecular Structure

   5 topics

   • Types of Chemical Bonds
     • Ionic Bonds
     • Covalent Bonds (Polar and Nonpolar)
     • Metallic Bonds
   • Lewis Structures
     • Drawing Lewis Structures
     • Formal Charge Calculation
   • VSEPR Theory
     • Molecular Geometry
     • Bond Angles
   • Hybridization
     • sp
     • sp²
     • sp³
     • sp³d
     • sp³d²
   • Intermolecular Forces
     • London Dispersion Forces
     • Dipole-Dipole Interactions
     • Hydrogen Bonding
3. 3

   Stoichiometry

   4 topics

   • Mole Concept
     • Avogadro’s Number
     • Molar Mass
   • Balancing Chemical Equations
     • Steps to Balance Equations
   • Limiting Reactants
     • Identification of Limiting Reactant
     • Calculating Theoretical Yield
   • Percent Yield
     • Actual Yield vs. Theoretical Yield
4. 4

   Chemical Reactions

   3 topics

   • Types of Chemical Reactions
     • Synthesis
     • Decomposition
     • Single Replacement
     • Double Replacement
     • Combustion
   • Reaction Mechanisms
     • Elementary Steps
     • Rate-Determining Step
   • Oxidation-Reduction Reactions
     • Oxidation States
     • Balancing Redox Reactions (Half-Reaction Method)
5. 5

   Thermochemistry and Thermodynamics

   6 topics

   • Laws of Thermodynamics
     • First Law (Law of Energy Conservation)
     • Second Law (Entropy)
     • Third Law (Absolute Zero)
   • Enthalpy (ΔH)
     • Endothermic and Exothermic Reactions
     • Hess’s Law
     • Standard Enthalpy of Formation
   • Entropy (ΔS)
     • Predicting Entropy Changes
   • Gibbs Free Energy (ΔG)
     • ΔG = ΔH - TΔS
     • Spontaneity of Reactions
   • Calorimetry
     • Heat Capacity
     • Specific Heat
     • Coffee Cup and Bomb Calorimetry
   • Phase Changes
     • Enthalpy of Fusion
     • Enthalpy of Vaporization
6. 6

   Chemical Kinetics

   4 topics

   • Reaction Rates
     • Rate Laws
     • Rate Constants
   • Factors Affecting Reaction Rates
     • Concentration
     • Temperature
     • Catalysts
     • Surface Area
   • Reaction Mechanisms
     • Elementary Reactions
     • Reaction Intermediates
     • Rate-Determining Step
   • Activation Energy
     • Arrhenius Equation
7. 7

   Equilibrium

   4 topics

   • Dynamic Equilibrium
     • Characteristics of Equilibrium State
   • Equilibrium Constants (Kc, Kp)
     • Writing Equilibrium Expressions
     • Relationship Between Kc and Kp
   • Le Chatelier’s Principle
     • Effects of Concentration, Temperature, and Pressure Changes
   • Solubility Product (Ksp)
     • Solubility and Precipitation
     • Common Ion Effect
8. 8

   Acids and Bases

   5 topics

   • Acid-Base Theories
     • Arrhenius Definition
     • Brønsted-Lowry Definition
     • Lewis Definition
   • pH and pOH
     • Calculations Using pH, pOH, [H⁺], and [OH⁻]
     • pH of Strong and Weak Acids/Bases
   • Acid-Base Equilibria
     • Ka and Kb
     • Polyprotic Acids
   • Buffer Solutions
     • Henderson-Hasselbalch Equation
     • Buffer Capacity
   • Titrations
     • Strong Acid-Strong Base Titrations
     • Weak Acid-Strong Base Titrations
     • Indicators
9. 9

   Electrochemistry

   4 topics

   • Redox Reactions
     • Oxidation Numbers
     • Balancing Redox Equations
   • Galvanic (Voltaic) Cells
     • Cell Diagrams
     • Standard Electrode Potentials (E°)
     • Nernst Equation
   • Electrolytic Cells
     • Electrolysis
     • Faraday’s Laws of Electrolysis
   • Applications
     • Batteries
     • Corrosion
     • Electroplating

General Chemistry flashcards for MCAT

18 of 50 cards from the General Chemistry deck — real questions with worked answers.

1. What are the four main postulates of Dalton's atomic theory?

   1) All matter is composed of indivisible atoms. 2) Atoms of a given element are identical in mass and properties. 3) Compounds form from fixed whole-number ratios of atoms. 4) Chemical reactions rearrange atoms but do not create or destroy them.

2. Define atomic number ($Z$) and mass number ($A$).

   $Z$ = number of protons in the nucleus (defines the element). $A$ = total number of protons plus neutrons: $A = Z + N$, where $N$ is the neutron number.

3. How do you determine the number of neutrons in an atom from $A$ and $Z$?

   $N = A - Z$ (mass number minus atomic number).

4. What are isotopes?

   Atoms of the same element (same $Z$, same number of protons) that differ in the number of neutrons, and therefore have different mass numbers $A$.

5. How is the average atomic mass of an element calculated from its isotopes?

   It is the weighted average of isotopic masses: $$\bar{m} = \sum_i f_i m_i$$ where $f_i$ is the fractional natural abundance and $m_i$ the mass of each isotope.

6. State the maximum number of electrons in a shell of principal quantum number $n$.

   $2n^{2}$ electrons (e.g., $n=1$ holds 2, $n=2$ holds 8, $n=3$ holds 18).

7. How many electrons can each subshell ($s$, $p$, $d$, $f$) hold?

   $s$ holds 2, $p$ holds 6, $d$ holds 10, and $f$ holds 14 electrons (each orbital holds 2).

8. State the Aufbau principle, Pauli exclusion principle, and Hund's rule.

   Aufbau: fill lowest-energy orbitals first. Pauli: no two electrons share all four quantum numbers (max 2 per orbital, opposite spins). Hund: fill degenerate orbitals singly with parallel spins before pairing.

9. Write the ground-state electron configuration of iron ($\ce{Fe}$, $Z=26$).

   $1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,4s^{2}\,3d^{6}$, or $[\ce{Ar}]\,4s^{2}\,3d^{6}$.

10. Why are chromium and copper exceptions to the normal filling order?

    A half-filled or fully-filled $d$ subshell is extra stable, so $\ce{Cr}$ is $[\ce{Ar}]4s^{1}3d^{5}$ and $\ce{Cu}$ is $[\ce{Ar}]4s^{1}3d^{10}$ rather than $4s^{2}3d^{4}$/$4s^{2}3d^{9}$.

11. How do atomic radius and ionization energy trend across a period (left to right)?

    Atomic radius decreases (greater effective nuclear charge pulls electrons in); ionization energy increases (electrons held more tightly).

12. How do atomic radius and ionization energy trend down a group?

    Atomic radius increases (more electron shells); ionization energy decreases (valence electrons farther from and more shielded from the nucleus).

13. Which element in the periodic table has the highest electronegativity, and what is the general electronegativity trend?

    Fluorine is the most electronegative. Electronegativity increases up a group and left-to-right across a period (toward the upper right, excluding noble gases).

14. Contrast ionic, covalent, and metallic bonding.

    Ionic: electrostatic attraction between cations and anions (electron transfer, large electronegativity difference). Covalent: sharing of electron pairs (small difference). Metallic: positive metal ions in a delocalized 'sea' of shared electrons.

15. What electronegativity differences roughly distinguish nonpolar covalent, polar covalent, and ionic bonds?

    Nonpolar covalent: $\Delta\text{EN} < 0.5$; polar covalent: $0.5 \leq \Delta\text{EN} < 1.7$; ionic: $\Delta\text{EN} \geq 1.7$ (approximate cutoffs).

16. State the octet rule and name common exceptions.

    Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons. Exceptions: incomplete octets ($\ce{H}$, $\ce{Be}$, $\ce{B}$), odd-electron species ($\ce{NO}$), and expanded octets (period 3+ elements like $\ce{S}$, $\ce{P}$).

17. What are the steps to draw a Lewis structure?

    1) Count total valence electrons. 2) Place least electronegative atom in the center. 3) Connect atoms with single bonds. 4) Distribute remaining electrons to complete octets on outer atoms then the center. 5) Form multiple bonds if the central atom lacks an octet.

18. How is formal charge calculated?

    $$\text{FC} = (\text{valence } e^{-}) - (\text{nonbonding } e^{-}) - \tfrac{1}{2}(\text{bonding } e^{-})$$ The best Lewis structure minimizes formal charges.

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