DAT (Dental Admission Test) 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 DAT (Dental Admission Test), not a summary of it.

1. 1

   Atomic and Molecular Structure

   4 topics

   • Atomic Theory and Structure
     • Subatomic particles and isotopes
     • Quantum numbers and orbitals
     • Electron configuration and Aufbau principle
   • Periodic Table and Trends
     • Atomic radius and ionization energy
     • Electronegativity and electron affinity
   • Chemical Bonding
     • Ionic and covalent bonds
     • Lewis structures and formal charge
     • VSEPR theory and molecular geometry
     • Hybridization and bond polarity
   • Intermolecular Forces
     • Dipole-dipole, hydrogen bonding, London dispersion
2. 2

   Stoichiometry and Reactions

   4 topics

   • Moles and Molar Mass
     • Avogadro's number and mole conversions
     • Percent composition and empirical formulas
   • Chemical Equations
     • Balancing equations
     • Limiting reagents and percent yield
   • Reaction Types
     • Synthesis, decomposition, combustion
     • Single and double displacement
   • Oxidation-Reduction
     • Assigning oxidation numbers
     • Balancing redox reactions
3. 3

   States of Matter and Solutions

   4 topics

   • Gas Laws
     • Ideal gas law
     • Boyle's, Charles's, and Dalton's laws
     • Kinetic molecular theory
   • Liquids and Solids
     • Phase diagrams and phase changes
     • Vapor pressure
   • Solutions and Concentration
     • Molarity, molality, and dilution
     • Solubility rules
   • Colligative Properties
     • Boiling point elevation and freezing point depression
     • Osmotic pressure
4. 4

   Thermodynamics and Kinetics

   4 topics

   • Thermochemistry
     • Enthalpy and Hess's law
     • Calorimetry and specific heat
   • Laws of Thermodynamics
     • Entropy and Gibbs free energy
     • Spontaneity of reactions
   • Chemical Kinetics
     • Rate laws and reaction order
     • Activation energy and catalysts
   • Reaction Mechanisms
     • Rate-determining step
5. 5

   Equilibrium and Acid-Base Chemistry

   4 topics

   • Chemical Equilibrium
     • Equilibrium constant (Keq)
     • Le Chatelier's principle
   • Acids and Bases
     • Arrhenius, Bronsted-Lowry, and Lewis definitions
     • pH, pOH, and Kw
   • Buffers and Titrations
     • Henderson-Hasselbalch equation
     • Titration curves and equivalence points
   • Solubility Equilibria
     • Ksp and common ion effect
6. 6

   Electrochemistry and Nuclear Chemistry

   3 topics

   • Galvanic Cells
     • Cell potential and standard reduction potentials
     • Nernst equation
   • Electrolytic Cells
     • Electrolysis and Faraday's laws
   • Nuclear Chemistry
     • Radioactive decay types
     • Half-life calculations

General Chemistry flashcards for DAT (Dental Admission Test)

18 of 77 cards from the General Chemistry deck — real questions with worked answers.

1. What are the three fundamental subatomic particles, and what is the relative charge and approximate mass (in amu) of each?

   Proton (+1 charge, ~1 amu, in nucleus), neutron (0 charge, ~1 amu, in nucleus), and electron (-1 charge, ~1/1836 amu, in orbitals).

2. How do you determine the number of protons, neutrons, and electrons in a neutral atom from its atomic number (Z) and mass number (A)?

   Protons = Z; electrons = Z (neutral atom); neutrons = A - Z.

3. What are isotopes?

   Atoms of the same element (same number of protons) that have different numbers of neutrons, and therefore different mass numbers.

4. State Hund's rule and the Pauli exclusion principle for filling electron orbitals.

   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 four quantum numbers (max 2 electrons per orbital, opposite spins).

5. What do the four quantum numbers (n, l, m_l, m_s) describe?

   n = principal (energy level/size); l = azimuthal (subshell shape: s,p,d,f); m_l = magnetic (orbital orientation); m_s = spin (+1/2 or -1/2).

6. How does atomic radius trend across a period and down a group?

   Atomic radius decreases left-to-right across a period (increasing nuclear charge) and increases down a group (added shells).

7. How do ionization energy and electronegativity trend on the periodic table?

   Both increase across a period (left to right) and decrease down a group. They are highest toward the upper right (excluding noble gases).

8. How does electron affinity generally trend across a period, and which family has the most negative (most exothermic) values?

   Electron affinity becomes more negative (more exothermic) across a period toward the right; halogens (Group 17) have the most negative electron affinities.

9. Why is a cation smaller and an anion larger than its parent neutral atom?

   A cation loses electrons (often a whole shell) and has higher effective nuclear charge per electron, shrinking it. An anion gains electrons, increasing electron-electron repulsion and size.

10. Distinguish ionic, covalent, and metallic bonding.

    Ionic: transfer of electrons between metal and nonmetal (electrostatic attraction of ions). Covalent: sharing of electron pairs between nonmetals. Metallic: positive metal ions in a delocalized 'sea' of electrons.

11. According to VSEPR theory, what are the electron-pair geometries for 2, 3, 4, 5, and 6 electron domains?

    2 = linear (180°); 3 = trigonal planar (120°); 4 = tetrahedral (109.5°); 5 = trigonal bipyramidal (90°/120°); 6 = octahedral (90°).

12. What is the difference between a sigma (σ) bond and a pi (π) bond?

    A sigma bond forms by head-on (axial) orbital overlap and allows rotation; a pi bond forms by side-to-side (parallel p) overlap above and below the bond axis. A single bond is 1 σ; a double bond is 1 σ + 1 π; a triple bond is 1 σ + 2 π.

13. How does bond order relate to bond length and bond strength?

    Higher bond order means shorter bond length and greater bond strength (energy). Triple > double > single bonds in strength; single > double > triple in length.

14. List the common types of intermolecular forces from weakest to strongest.

    London dispersion forces < dipole-dipole forces < hydrogen bonding < ion-dipole forces.

15. What conditions are required for hydrogen bonding to occur?

    An H atom must be covalently bonded to a highly electronegative atom (N, O, or F) and interact with a lone pair on a nearby N, O, or F atom.

16. How are intermolecular forces related to boiling point, melting point, and vapor pressure?

    Stronger intermolecular forces raise boiling and melting points and lower vapor pressure (less volatile).

17. What is Avogadro's number, and what does one mole represent?

    Avogadro's number is 6.022 × 10^23. One mole is that number of particles (atoms, molecules, ions), and its mass in grams equals the substance's molar mass.

18. How do you convert grams of a substance to number of molecules?

    Divide grams by molar mass (g/mol) to get moles, then multiply moles by 6.022 × 10^23 molecules/mol.

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