Chemistry Exam Info

Atomic theory and atomic structure

• Evidence for the atomic theory
• Atomic masses; determination by chemical and physical means
• Atomic number and mass number; isotopes and mass spectroscopy
• Electron energy levels: atomic spectra and atomic orbitals
• Periodic relationships, including, for example, atomic radii, ionization energies, electron affinities, and oxidation states
• Nuclear chemistry: nuclear equations, half-lives, and radioactivity; chemical applications

Chemical bonding

• Binding forces
  • Types: covalent, ionic, metallic, macromolecular (or network), dispersion, and hydrogen bonding
  • Relationships to structure and to properties
  • Polarity of bonds and electronegativities
• VSEPR theory and Lewis electron-dot diagrams
  • Hybridization of orbitals
  • Geometry of molecules, ions, and coordination complexes
  • Structural isomerism
  • Resonance
  • Sigma and pi bonds
  • Dipole moments of molecules
  • Relation of properties to structure

Gases

• Laws of ideal gases; equations of state for an ideal gas
• The mole concept; Avogadro's number
• Kinetic-molecular theory
  • Interpretation of ideal gas laws on the basis of this theory
  • Dependence of kinetic energy of molecules on temperature: Boltzmann distribution
  • Deviations from ideal gas laws

Liquids and solids

• Liquids and solids from the kineticmolecular viewpoint
• Phase diagrams of one-component systems
• Changes of state and critical phenomena

Solutions

• Types of solutions and factors affecting solubility
• Methods of expressing concentration
• Colligative properties; for example, Raoult's law
• Effect of interionic attraction on colligative properties and solubility

Acid-base reactions; concepts of Arrhenius, Brønsted-Lowry, and Lewis; amphoterism

Reactions involving coordination complexes

Precipitation reactions

Oxidation-reduction reactions

• Oxidation number
• The role of the electron in oxidation-reduction
• Electrochemistry; electrolytic cells, standard half-cell potentials, prediction of the direction of redox reactions, and the effect of concentration changes

Ionic and molecular species present in chemical systems; net-ionic equations

Stoichiometry: mass and volume relations with emphasis on the mole concept

Balancing of equations, including those for redox reactions

Concept of dynamic equilibrium of physical and chemical; LeChâtelier's principle; equilibrium constants

Quantitative treatment

• Equilibrium constants for gaseous reactions in terms of both molar concentrations and partial pressure (K_c, K_p)
• Equilibrium constants for reactions in solutions
• Constants for acids and bases; pK; pH
• Solubility product constants and their application to precipitation and the dissolution of slightly soluble compounds
• Constants for complex ions
• Common ion effect; buffers

Concept of rate of reaction

• Order of reaction and rate constant
• Determination of order of reaction and rate constant from experimental data
• Effect of temperature change on rates

Activation energy and the role of catalysts

The relationship between the rate-determining step and reaction mechanism

State functions

First law:

• Heat of formation
• Heat of reaction, change in enthalpy, and Hess's law
• Heat capacity; heats of vaporization and fusion

Second law:

• Free energy of formation
• Free energy of reaction
• Dependence of change in free energy on enthalpy and entropy changes

Relationship of change in free energy to equilibrium constants and electrode potentials

The accumulation of certain specific facts of chemistry is essential to enable students to:

• comprehend the development of principles and concepts
• demonstrate applications of principles
• relate fact to theory and properties to structure
• develop an understanding of systematic nomenclature that facilitates communication

The following areas are normally included on the exam:

• Chemical reactivity and products of chemical reactions
• Relationships in the periodic table: horizontal, vertical, and diagonal
• Chemistry of the main groups and transition elements, including typical examples of each
• Organic chemistry, including such topics as functional groups and isomerism (may be treated as a separate unit or as exemplary material in other areas, such as bonding)

Some questions are based on laboratory experiments widely performed in general chemistry and ask about the equipment used, observations made, calculations performed, and interpretation of the results. The questions are designed to provide a measure of understanding of the basic tools of chemistry and their applications to simple chemical systems.