Chemical Energetics and Dynamics

Upon successful completion of this course, students will be able to:

• apply the first law of thermodynamics to calculate changes in internal energy, heat transferred, and work done for physical and chemical processes;
• solve heat transfer problems using heat capacity relationships and calorimetry data;
• calculate pressure-volume work and relate gas expansion or compression to thermodynamic energy changes;
• distinguish between internal energy and enthalpy;
• determine thermodynamic quantities under constant-volume and constant-pressure conditions;
• determine reaction enthalpies with Hess's Law;
• calculate standard enthalpy changes for reactions using standard enthalpies of formation;
• explain how intermolecular forces influence the properties and phase behaviour of liquids and solids;
• solve thermochemical and phase-equilibrium problems involving heating curves, phase changes, and the Clausius-Clapeyron equation;
• predict phase changes and stable states using phase diagrams;
• relate molecular structure and kinetic energy distributions to fusion, freezing, vapourization, condensation, sublimation, and deposition;
• classify crystalline solids by bonding type;
• explain how a solid's structure influences its physical properties;
• calculate and interconvert solution concentration units including molarity, molality, mole fraction, percent by mass, and parts per million;
• predict solution formation and solubility using intermolecular forces and entropy of mixing;
• analyze the enthalpy changes associated with solution formation, hydration, and ionic dissolution;
• interpret solubility curves to determine saturation behaviour and precipitation under changing temperature conditions;
• apply Henry’s Law to quantitatively determine gas solubility in solutions;
• apply Raoult’s Law to quantitatively determine solution vapour pressure;
• distinguish between ideal solution behaviour and non-ideal solution behaviour;
• define dynamic equilibrium in terms of forward and reverse reaction rates;
• write equilibrium constant expressions (K, K_c, and K_p) for homogeneous chemical reactions using balanced chemical equations;
• calculate equilibrium constants, equilibrium concentrations, and equilibrium pressures using algebraic methods and ICE tables;
• calculate the reaction quotient, Q, and predict the direction a chemical system will shift to reach equilibrium by comparing Q with the thermodynamic equilibrium constant, K, value;
• predict the effects of changes in chemical systems (e.g. concentration, pressure, volume, temperature, catalysts) on chemical equilibria using Le Châtelier’s principle;
• relate the magnitude of an equilibrium constant to the relative amounts of reactants and products present at equilibrium;
• combine equilibria equations and determine the equilibrium constant for the net reaction;
• distinguish between Arrhenius and Brønsted-Lowry acids and bases;
• identify acids, bases, and conjugate acid-base pairs in chemical reactions;
• write acid and base ionization equilibria and calculate equilibrium constants (K_a, K_b, K_w);
• determine pH, pOH, hydronium concentration, and hydroxide ion concentration in an aqueous solution using stoichiometry, equilibrium expressions, ICE tables, and appropriate approximations;
• predict how acid strength, base strength, concentration, and molecular structure influence ionization, conjugate base stability, and solution pH;
• classify ionic compounds and dissolved ions as acidic, basic, or neutral;
• analyze acid-base equilibria involving acid mixtures and polyprotic acids to determine the dominant contributors to solution pH;
• explain the composition, action, effectiveness, and capacity of buffer solutions using weak acid-base equilibria and Le Châtelier’s principle;
• calculate the pH of a buffer solution before addition of strong acid or base using stoichiometry and the Henderson-Hasselbalch equation;
• calculate the pH of a buffer solution after addition of strong acid or base using stoichiometry and the Henderson-Hasselbalch equation;
• sketch acid-base titration curves;
• calculate titration solution pH in the buffer region, at the equivalence point, and after the equivalence point for strong and weak acid-base systems;
• select an appropriate acid-base indicator for a titration by comparing an indicator pK_a value to the expected equivalence point pH range;
• calculate the solubility product constant (K_sp) and molar solubility for ionic compounds using equilibrium expressions and stoichiometric relationships;
• predict the effects of system changes (e.g. common ions, pH) on solubility equilibria and molar solubility of an ionic compound;
• distinguish between thermodynamic spontaneity and reaction kinetics using activation energy and Gibbs energy;
• predict the sign and relative magnitude of entropy changes for physical and chemical processes using molecular motion, phase, temperature, and gas-particle considerations;
• calculate entropy changes for phase changes and chemical reactions using thermodynamic equations and tabulated standard molar entropy data;
• determine the spontaneity of chemical processes using Gibbs energy and entropy of the universe;
• determine the standard Gibbs energy of a chemical reaction from Gibbs energy of formation values;
• calculate Gibbs energy changes under standard and non-standard conditions;
• calculate changes in the thermodynamic equilibrium constant, K, with temperature using the van't Hoff equation;
• predict equilibrium behaviour using Gibbs energy;
• assign oxidation states to atoms in compounds;
• identify the species oxidized, the species reduced, the oxidizing agent, and the reducing agent in an oxidation-reduction reaction;
• sketch galvanic and electrolytic cells and identify half-reactions, direction of electron flow, the electrode functions, and the salt bridge;
• describe galvanic cells with electrochemical cell notation;
• calculate standard cell potentials from standard reduction potentials;
• predict the spontaneity of electrochemical reactions under standard conditions;
• predict the spontaneity of electrochemical reactions under non-standard conditions with the Nernst equation;
• relate Gibbs energy, the equilibrium constant, and cell potential using thermodynamic and electrochemical equations;
• predict the effects of concentration changes on electrochemical cell behaviour;
• solve quantitative electrolysis problems involving charge, current, time, moles of electrons, and mass of substances produced or consumed;
• safely perform demonstrated laboratory techniques;
• collect and record quantitative data to the correct number of significant digits;
• collect, record, and analyze experimental data using appropriate mathematical, graphical, and statistical methods to obtain quantitative results;
• report quantitative results to the correct number of significant digits;
• evaluate experimental results by assessing random and systematic error sources;
• design and apply appropriate experimental procedures and techniques to investigate chemical problems;
• communicate experimental methods, results, and conclusions in a clear scientific format using appropriate chemical theory and evidence.