CHEMISTRY 113 

Principles of Chemistry I

Students are responsible for reviewing these topics which they are expected to know from chem 11 and CHEM 12 or their equivalents.

• Accuracy.
• Precision.
• Uncertainties.
• Significant figures.
• Propagation of errors in math operations.

Nomenclature:

• Covalent and ionic compounds.
• Binary compounds.
• Hydroxides.
• Acids.
• Salts.
• Hydrates.

Stoichiometry:

• Percent composition.
• Empirical formula.
• Molecular formula.
• Reaction equations.
• Limiting reagent.
• Theoretical and actual yield, % conversion.

Solutions:

• Concentration.
• Mass %, molarity.
• Ionic solutions.
• Ionic equations.
• Ionic concentration.

Redox processes:

• Oxidation number.
• Oxidation.
• Reduction.
• Redox equations.
• Half-reactions.
• Balancing redox equations by the half-reaction method.

Electrochemistry:

• Galvanic cell.
• Electrode potential.
• Standard electrode.
• Standard electrode potential.
• Electrolytic cell.
• Electrolysis (active and inert electrodes).

Components of the Atom:

• Brief review of electron, proton, neutron, isotopes, atomic and molecular mass.

MATERIAL TO BE COVERED IN CLASS

Atom:

• Brief review.
• Bohr model.
• Relationship between frequency and wavelength.
• Quantization of energy.
• Energy of a photon.
• Electron transitions.
• Line spectra.
• Atomic orbitals.
• Principal quantum number n.
• Shapes of atomic orbitals. s-, p-, and d-orbitals.
• Orbital quantum number l.
• Relationship between l and n.
• Magnetic quantum number m.
• Relationship between m and l.
• Spin.
• Energy levels in the hydrogen atom.
• Energy levels in a multi-electron atom.
• Electron shells and subshells.
• Pauili principle.
• The aufbau.
• Hund's rule.
• Electron configurations.
• Paramagnetism. 

Periodic Properties:

• Core and valence electrons.
• Effective nuclear charge.
• Groups and periods.
• Position in the periodic table and ionic charges.
• Metals and nonmetals.
• Atomic and ionic radii.
• Ionization energy and electron affinity.
• Electronegativity.
• Chemical reactivity. 

Chemical Bond:

• Ionic bond.
• Covalent bond.
• Lewis diagrams.
• Single, double, and triple bonds.
• Lewis diagrams.
• Resonance structures.
• Formal charges.
• Co-ordinate covalent bonding.
• Exceptions to the octet rule: radicals, electron-deficient systems, d-electron systems.
• Relation between bond order, bond length, and strength of a bond.

Molecular Geometry:

• VSEPR model.
• Electron-pair configurations: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.
• Molecular shapes: linear, angular (bent), trigonal planar, trigonal pyramidal, T-shaped, tetrahedral, sawhorse, tetragonal (square) planar, trigonal bipyramidal, tetragonal (square) pyramidal, octahedral. Ideal bond angles - 90, 109.47, 120, and 180 deg, and deviations from them.
• Multiple bonds. 

MO AND VALENCE BOND THEORY

Diatomic molecules:

• Hydrogen molecule.
• Atomic and molecular orbitals.
• Bonding and antibonding orbitals.
  sigma- and pi- orbitals.
• Lone pairs and nonbonding orbitals.
• Core and valence electrons.
• Polyatomic molecules with a single central atom.
  sp3, sp2 and sp hybridization.
• Shapes of hybrid orbitals.
• Spatial arrangement of hybrid and remaining p-orbitals.
• Relationship between the number of valence-shell electrons and the type of hybridization.
• Polyatomic molecules with more than one central atom.
• Single, double, triple bonds.
• Ethane, ethene, ethyne, benzene. 

Intermolecular Interactions:

• Electric dipole.
• Dipole moment.
• Polar and nonpolar bonds.
• Polar and nonpolar molecules.
• Relationship between symmetry and polarity of molecules.
• Polarizability.
• Dipole-dipole, induction, and dispersion interactions.
• Hydrogen bond.
• Ion-dipole interaction.
• Solubility of polar and nonpolar compounds in polar and nonpolar solvents.
• Hydrophobicity. 

• Ideal gas.
• Equation of state.
• Partial pressure.
• Dalton's law.
• Molecular kinetic theory.
• Relation between temperature and Maxwell-Boltzmann distribution - general idea and shape of the curve.
• Types of solids: molecular, metallic, ionic, network. Examples. Liquid-vapor equilibrium. Boiling point and its dependence on pressure and the nature of liquid. 

Chemical Kinetics:

• Reaction rate.
• Concentration dependence of reaction rate.
• Rate equation.
• Rate constant.
• Order of reaction.
• Collision theory.
• Elementary processes.
• Molecularity.
• Uni-, bi- and trimolecular processes.
• Molecularity and order of reaction.
• Stoichiometric coefficients and partial orders.
• Complex kinetics.
• Multistep processes.
• Reaction mechanism.
• Intermediates.
• Limiting step of a multistep process.
• Chain reactions.
• Steady state approximation.
• Rate equations for multistep processes.
• Integrated rate equations for reactions of 1st order. Half-life.
• Relationship between the half-life and the rate constant.
• Temperature dependence of reaction rates.
• Arrhenius equation: exponential and logarithmic forms.
• Activation parameters: preexponential factor and activation energy.
• Reactive and elastic collisions.
• Steric factor and activation energy. Reaction coordinate and energy profile for a simple linear triatomic reaction A + BC. Reaction barrier and activation energy. Reaction energy. Reaction profile of a multistep process.
• Catalysis.
• Catalyst. Reaction profile for a catalyzed reaction. Activation energy and reaction energy of a catalyzed reaction.
• Enzymatic catalysis.
• Mass defect.
• Nuclear fusion.
• Nuclear fission.
• Kinetics of radioactive decay.