General Chemistry

Elements of General Chemistry necessary for the the Department of Biology students to know. In the Annexes of the course modules from the Viticulture and Food Chemistry of the Department of Chemistry are included.

Objectives

At the end of this course the student should be able to: 1. Formulate net ionic equations, classify acids and bases as strong or weak, assign oxidation numbers, balance simple oxidation – reduction reactions, calculate and use molarity. 2. Write and handle thermochemical equations, calculate the heat of reaction from stoichiometry, apply the Hess’ law, calculate enthalpy of a reaction using standard enthalpies of formation. 3. Use Lewis symbols to represent ionic bond formation and write electron configurations of ions, obtain relative bond polarities, write Lewis formulas using formal charges, relate bond order and bond length. 4. Predict molecular geometries, relate dipole moment and molecular geometry, apply valence bond theory, describe molecular orbital configurations. 5. Calculate solution concentration, find mole fractions, calculate the vapor pressure lowering, the boiling-point elevation, the freezing-point depression and using them calculate the molecular weight of the solute. Calculate osmotic pressure and determine colligative prorerties of ionic solutions. 6. Use the Arrhenius equation, write overall chemical equation from a mechanism, determine the moleculalarity of an elementary reaction and write its rate equation. Determine the rate law from a mechanism with an initial slow step. 7. Apply stoichiometry to an equilibrium mixture, write equilibrium-constant expressions and obtain them from reaction composition. Use the reaction quotient, solve equilibrium problems and apply Le Chatelier’s principle changing the reaction conditions. 8. Identify acid and base species according to the Brønsted-Lowry and Lewis concepts, decide whether reactants or products are favoured in an acid-base reaction, calculate concentrations of Η3Ο+ and ΟΗ– in solutions of a strong acid or base. 9. Determine Ka and Kb from the solution pH and vice versa, calculate concentrations of species in solutions of a weak acid or base, calculate the pH of a buffer solution. 10. Calculate the entropy change for a phase transition, calculate ΔGο from ΔΗο and ΔSο, calculate Κ from the standard free-energy change and ΔGο and Κ at various temperatures.

Prerequisites

None

Syllabus

1. Calculations with Chemical Formulas and Equations. Molecular weight and formula weight. The mole concept. Mass percentages from the formula. Elemental analysis: Percentages of carbon, hydrogen and oxygen. Determining formulas. Molar interpretation of a chemical equation. Amounts of substances in a chemical reaction. Limiting reactant: Theoretical and percentage yields. 2. Chemical Reactions: Introduction. Ionic theory of solutions. Molecular and ionic equations. Precipitation reactions. Acid – base reactions. Oxidation – reduction reactions. Balancing simple oxidation – reduction reactions. Molar concentration. Diluting solutions. Gravimetric analysis. Volumetric analysis. 3. Thermochemistry. Energy and its units. Heat of reaction. Enthalpy and enthalpy change. Thermochemical equations. Applying stoichiometry to heat of reaction. Measuring heat of reaction. Hess’s law. Standard enthalpies of formation. Fuels-foods, commercial fuels and rocket fuels. 4. Quantum Theory of the Atom. The wave nature of light. Quantum effects and photons. The Bohr theory of the hydrogen atom. Quantum mechanics. Quantum numbers and atomic orbitals. 5. Electron Configurations and Periodicity. Electron spin and the Pauli exclusion principle. Building-up principle and the periodic table. Writing electron configurations using the periodic table. Orbital diagrams of atoms – Hund’s rule. Mendeleev’s predictions from the periodic table. Periodic properties (atomic radii, ionization energies, electron affinities). Periodicity in the main-group elements. 6. Ionic and Covalent Bond. Describing ionic bonds. Electron configuration of ions. Ionic radii. Describing covalent bonds. Polar covalent bonds. Electronegativity. Writing Lewis electron-dot formulas. Delocalized bonding – Resonance. Formal charge and Lewis formulas. Bond length and bond order. Bond energy. 7. Molecular Geometry and Chemical Bonding Theory. The VSEPR model. Dipole moment and molecular geometry. Valence bond theory. Description of multiple bonding. Principles of molecular orbital theory. Electron configurations of diatomic molecules of the second-period elements. 8. Solutions. Types of solutions. Solubility and the solution process. Effect of tempetrature and pressure on solubility. Ways of expressing concentration. Vapor pressure of a solution. Boiling-Point elevation and Freezing-point depression. Osmosis. Colligative properties of ionic solutions. Coloids. 9. Rates of reaction. Definition of reaction rate. Experimental determination of rate. Dependence of rate on concentration. Change of concentration with time. Temperature and rate; Collision and transition-state theories. Arrhenius equation. Elementary reactions. The rate law and the mechanism. Catalysis. 10. Chemical Equilibrium. Chemical Equilibrium-A dynamic equilibrium. The equilibrium constant. Heterogeneous equilibria. Solvents in homogenius equilibria. Qualitatively interpreting the equilibrium contant. Predicting the direction of reaction. Calculating equilibrium concentrations. Removing products or adding reactants. Changing the pressure and temperature. Effect of a catalyst. 11. Acids and Bases. Arrhenius concept of acids and bases. Brønsted–Lowry concept of acids and bases. Lewis concept of acids and bases. Relative strengths of acids and bases. Molecular structure and acid strength. Self ionization of water. Solutions of a strong acid or base. The pH of a solution. 12. Acid-Base Equilibria. Acid-ionization equilibria. Polyprotic acid. Base-ionization equilibria. Acid-base properties of salt solutions. Common-ion effect. Buffers. Acid-base titration curves 13. Solubility and complex ion balance. The constant product of solubility. Solubility and common ion effect. Subsidence calculations. Effect of pH on the solubility. Complex ion formation. Complexes ions and solubility. Qualitative analysis of metal ions. 14. Thermodynamics and Equilibrium. First Law of Thermodynamics. Enthalpy. Entropy and the second law of thermodynamics. Standard entropies and the third law of thermodynamics. Free energy and spontaneity. Interpretation of free energy. Relating ΔGο to the equilibrium constant. Change of free energy with temperature. 15. Laboratory. Preparation of solutions, dilution and pH measurement. Qualitative analysis. Volumetric determinations. Exercises of the previous sections.