Graduate Medical School Admissions Test (GAMSAT) Section III: Chemistry Syllabus

Section III: Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Section III: Chemistry in Graduate Medical School Admissions Test (GAMSAT), not a summary of it.

1. 1

   General and Physical Chemistry

   4 topics

   • Atomic structure and periodicity
     • Electron configuration and orbitals
     • Periodic trends
   • Chemical bonding
     • Ionic, covalent and metallic bonding
     • Molecular geometry and VSEPR
     • Intermolecular forces
   • Stoichiometry and the mole
     • Balancing equations
     • Mole calculations and concentration
   • States of matter and gas laws
2. 2

   Thermodynamics and Kinetics

   4 topics

   • Energetics of reactions
     • Enthalpy and Hess's law
     • Entropy and Gibbs free energy
   • Reaction rates
     • Factors affecting rate
     • Rate laws and order of reaction
   • Collision theory and activation energy
   • Catalysis
3. 3

   Equilibrium and Solutions

   4 topics

   • Chemical equilibrium
     • Equilibrium constant and Le Chatelier's principle
     • Position of equilibrium
   • Acids, bases and pH
     • Strong and weak acids and bases
     • pH, pKa and buffers
     • Titration curves
   • Solubility and precipitation
     • Solubility product
   • Electrochemistry
     • Redox reactions and oxidation states
     • Electrochemical cells and electrode potentials
4. 4

   Organic Chemistry

   5 topics

   • Structure and nomenclature
     • Functional groups
     • IUPAC naming
     • Drawing structural formulae
   • Isomerism and stereochemistry
     • Structural isomers
     • Chirality and optical isomerism
     • Cis-trans isomerism
   • Reaction mechanisms
     • Substitution, addition and elimination
     • Nucleophiles and electrophiles
   • Functional group chemistry
     • Alcohols, aldehydes, ketones and acids
     • Amines and amides
   • Biomolecules
     • Carbohydrates, lipids and proteins
     • Amino acids and peptide bonds

Section III: Chemistry flashcards for Graduate Medical School Admissions Test (GAMSAT)

21 of 61 cards from the Section III: Chemistry deck — real questions with worked answers.

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

   The atomic number $Z$ is the number of protons in an atom's nucleus (defines the element). The mass number $A$ is the total number of protons plus neutrons: $A = Z + N$.

2. What are isotopes, and how do they differ from one another?

   Isotopes are atoms of the same element (same number of protons, same $Z$) that have different numbers of neutrons, and therefore different mass numbers $A$. They share chemical properties but differ in mass.

3. State the relationship between the energy of light and its frequency and wavelength.

   $E = h\nu = \dfrac{hc}{\lambda}$, where $h$ is Planck's constant, $\nu$ is frequency, $c$ is the speed of light, and $\lambda$ is wavelength. Energy is directly proportional to frequency and inversely proportional to wavelength.

4. List the maximum number of electrons each of the s, p, d and f subshells can hold.

   s holds 2, p holds 6, d holds 10, and f holds 14 electrons. A subshell with angular momentum quantum number $l$ holds $2(2l+1)$ electrons.

5. State the order of orbital filling described by the Aufbau principle (first few subshells).

   Electrons fill the lowest-energy orbitals first: $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, \ldots$ (lowest $n+l$ first; for equal $n+l$, lower $n$ first).

6. State Hund's rule and the Pauli exclusion principle.

   Hund's rule: electrons fill degenerate orbitals singly with parallel spins before pairing. Pauli exclusion: no two electrons in an atom can have the same set of all four quantum numbers (an orbital holds at most 2 electrons, with opposite spins).

7. How do atomic radius and ionization energy trend across a period and down a group?

   Across a period (left to right): atomic radius decreases, ionization energy increases (greater nuclear charge, same shell). Down a group: atomic radius increases, ionization energy decreases (more shells, increased shielding).

8. Define electronegativity and state how it trends in the periodic table.

   Electronegativity is the tendency of an atom to attract a shared bonding electron pair. It increases across a period (left to right) and decreases down a group; fluorine is the most electronegative element.

9. Distinguish ionic, covalent and metallic bonding.

   Ionic: electrostatic attraction between oppositely charged ions (metal + non-metal, electron transfer). Covalent: shared electron pairs between non-metals. Metallic: positive metal ions in a 'sea' of delocalised electrons.

10. Distinguish a polar covalent bond from a non-polar covalent bond.

    A non-polar covalent bond has equal sharing of electrons (zero or negligible electronegativity difference). A polar covalent bond has unequal sharing due to an electronegativity difference, producing partial charges ($\delta+$ and $\delta-$).

11. According to VSEPR theory, what shapes correspond to 2, 3 and 4 electron domains around a central atom?

    2 domains: linear ($180^{\circ}$); 3 domains: trigonal planar ($120^{\circ}$); 4 domains: tetrahedral ($109.5^{\circ}$). Lone pairs reduce these ideal bond angles.

12. List the three types of intermolecular force in order of increasing strength.

    London dispersion forces (weakest) < dipole-dipole interactions < hydrogen bonding (strongest). All are weaker than covalent or ionic bonds.

13. What conditions are required for hydrogen bonding?

    A hydrogen atom must be covalently bonded to a highly electronegative atom (N, O, or F), and there must be a lone pair on an N, O, or F atom on a neighbouring molecule to accept it.

14. Define a mole and state the value of Avogadro's number.

    A mole is the amount of substance containing as many entities as there are atoms in 12 g of carbon-12. Avogadro's number is $N_A \approx 6.022 \times 10^{23}\ \text{mol}^{-1}$.

15. State the formula relating number of moles to mass and molar mass.

    $n = \dfrac{m}{M}$, where $n$ is moles, $m$ is mass (g), and $M$ is molar mass ($\text{g mol}^{-1}$).

16. Define concentration (molarity) and give its formula.

    Molarity is moles of solute per litre of solution: $c = \dfrac{n}{V}$, with $c$ in $\text{mol L}^{-1}$ (M), $n$ in mol, and $V$ in litres.

17. Distinguish empirical formula from molecular formula.

    The empirical formula is the simplest whole-number ratio of atoms (e.g. $\ce{CH2O}$). The molecular formula gives the actual number of each atom in a molecule (e.g. $\ce{C6H12O6}$), an integer multiple of the empirical formula.

18. What is the limiting reagent in a reaction?

    The limiting reagent is the reactant that is completely consumed first, limiting the amount of product formed. It is identified by comparing the mole ratios of reactants to the stoichiometric ratio.

19. State the ideal gas law and identify each variable.

    $PV = nRT$, where $P$ is pressure, $V$ is volume, $n$ is moles, $R$ is the gas constant ($8.314\ \text{J K}^{-1}\text{mol}^{-1}$), and $T$ is absolute temperature (K).

20. State Boyle's law and Charles's law.

    Boyle's law: at constant $T$ and $n$, $P \propto \dfrac{1}{V}$ (so $P_1V_1 = P_2V_2$). Charles's law: at constant $P$ and $n$, $V \propto T$ (so $\dfrac{V_1}{T_1} = \dfrac{V_2}{T_2}$).

21. State the postulates of the kinetic theory of ideal gases.

    Gas particles have negligible volume, are in constant random motion, undergo perfectly elastic collisions, have no intermolecular forces, and their average kinetic energy is proportional to absolute temperature.

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