Chemical Bonding

• Chemical Bonding:
  • Atoms combine to achieve stable electronic configurations, primarily by filling their valence shells.
  • Types of chemical bonds include ionic, covalent, and metallic bonds.
  • Ionic bonds form between metals and nonmetals through the transfer of electrons. The resulting ions are held together by electrostatic attraction.
  • Covalent bonds form between nonmetals by sharing electrons. These bonds often result in molecules.
  • Metallic bonds occur in metals, involving a "sea" of delocalized electrons surrounding positively charged metal ions. This explains the conductivity and malleability of metals.
  • Bond polarity arises from differences in electronegativity. Electronegativity is the tendency of an atom to attract shared electrons.
  • Bond order describes the number of electron pairs shared between atoms in a covalent bond.
  • Bond length is the distance between the nuclei of two bonded atoms, and is related to bond strength.
  • Bond energy is the energy required to break a chemical bond.
  • VSEPR theory (Valence Shell Electron Pair Repulsion) predicts molecular shapes based on the arrangement of electron pairs around a central atom.

Types of Reactions

• Reactions of Acids and Bases:
  • Arrhenius concept defines acids as proton donors and bases as hydroxide ion donors.
  • Brønsted-Lowry concept defines acids as proton donors and bases as proton acceptors. This is more widely applicable.
  • Lewis concept defines acids as electron pair acceptors and bases as electron pair donors. This is the most broadly applicable definition, extending to many non-proton reactions.
  • Acid-base neutralization reactions produce water and a salt. The pH scale measures acidity: a lower pH indicates higher acidity.
  • Strong acids/bases completely dissociate in water. Weak acids/bases only partially dissociate.
• Redox Reactions:
  • Oxidation involves loss of electrons, and reduction involves gain of electrons. These occur simultaneously in a redox reaction.
  • Oxidizing agents cause oxidation; reducing agents cause reduction.
  • Oxidation states (numbers) track electron transfer.
  • Balancing redox reactions requires oxidation number methodology.

Chemical Equilibrium

• Equilibrium:
  • A dynamic state where forward and reverse reaction rates are equal. The concentrations of reactants and products remain constant.
  • Equilibrium constant (Kc) quantifies the extent of a reaction at equilibrium.
  • Le Chatelier's principle describes how changes in conditions (temperature, pressure, concentration) affect equilibrium. The system shifts to counteract the stress.

Thermochemistry

• Thermochemical Principles:
  • Enthalpy (ΔH) is the heat content of a system at constant pressure. Exothermic reactions release heat; endothermic reactions absorb heat.
  • Enthalpy changes (ΔH) accompany chemical reactions, and are tabulated for many reactions.
  • Hess's law allows calculation of enthalpy changes for complex reactions from simpler reactions.
  • Standard enthalpy of formation (ΔHf°) is the enthalpy change for the formation of one mole of a substance from its elements in their standard states. It is used to determine reaction enthalpy from tabulated values.

Energetics of Chemical Reactions

• Energy changes:
  • Activation energy (Ea) is the minimum energy required for a reaction to occur.
  • Reaction rates depend on temperature and activation energy.
  • Catalysts speed up reactions by lowering the activation energy, without being consumed by the reaction.

Solution Chemistry

• Solutions:
  • Solutions are homogeneous mixtures of solute and solvent.
  • Concentration of solutions can be expressed in several ways like molarity, molality, normality.
  • Colligative properties of solutions are affected by the number of solute particles, not their identity. Examples include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.

Electrochemistry

• Redox reactions in solution can be used to generate electricity or to drive non-spontaneous reactions:
  • Galvanic (voltaic) cells convert chemical energy into electrical energy via spontaneous redox reactions. Electrodes are involved, and a salt bridge.
  • Electrolytic cells use electrical energy to drive non-spontaneous redox reactions, and are used for electroplating.
  • Standard electrode potentials (E°) define the tendency of a half-reaction to occur.
  • Cell potential (Ecell) is the driving force for a redox reaction in a cell; related to standard free energy change (ΔG°) and equilibrium constant.

Nuclear Chemistry

• Radioactivity:
  • Radioactive decay involves spontaneous disintegration of unstable atomic nuclei.
  • Different types of radioactive decay exist (alpha, beta, gamma).
  • Radioisotopes are used for dating materials, medical imaging, and industrial applications.