CHEM1200 - Chemical Equilibrium Study Note.docx

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Department
Chemistry
Course Code
Chemistry 1027A/B
Professor
Mark Workentin

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Chapter 1: Chemical Equilibrium 1.1 The Equilibrium Constant  Reversible chemical reactions – at equilibrium, system contains both reactants and products (dynamic)  Equilibrium achieved when rate of forward reaction = rate of reverse  At the molecular level, both forward and reverse processes are still occurring  Equilibrium constant (K) is the mathematical relationship between pressures / concentrations of the reactants and products  Concentration of pure solids and liquids are not included in K  Numerical value of K only changes if the temperature of the system changes  Same phase = homogenous / different phase = heterogeneous  Summary of manipulation rules: √  Magnitude of the equilibrium constant is an indication of the relative amount of product / reactant  If K is large – more products are present, equilibrium lies to the right  If K is small – more reactants are present, equilibrium lies to the left  Reaction Quotient (Q) indicates whether or not the system is at equilibrium and the direction in which the reaction must proceed  Q < K eqreaction must proceed in forward direction  (not enough product) Q > K eqreaction must proceed in reverse direction  (not enough reactant) Q = K : equilibrium eq  Le Châtelier’s Principle: “If chemical system at equilibrium is disturbed, equilibrium will shift in such a way as to minimize the disturbance”  Concentration changes – system will attempt to use additional / make more  Pressure changes – Increase pressure = decrease volume, it will shift to right towards smaller number of moles / decrease pressure = increase volume, it will shift to left towards larger number of moles  If equal number of moles of gaseous reactants and products, position of equilibrium is unaffected by pressure changes  Numerical value of K is not affected by concentration or pressure changes  For exothermic reaction, heat is treated like a product (K decrease) For endothermic reaction, heat is like a reactant (K increase)  ΔG < 0 – reaction will proceed spontaneously  If ΔG is +, equilibrium lies to the left and if ΔG is –, equilibrium lies to the right ( )  Van’t Hoff equation determines K values at different temperatures ( ) 1.2 Solubility of Ionic Compounds  Solubility: amount of substance that will dissolve in a certain volume of a specific solvent  Ionic compounds become more soluble as the temperature increases, completely dissociate into ions once solid dissolves  Solutes dissolve in solvent  Equilibrium is established between undissolved solid and the free ions once no more will dissolve  K sp the equilibrium constant called the solubility product  Excess solid does not affect the position of the equilibrium  Saturated implies that there is an equilibrium between solid and ions, if quantity is less than required it is unsaturated  When Q > K ,spolid will form from ions in solution (supersaturated)  Solution may already contain an ion in common with the dissolving salt – decreases solubility of an ionic compound (K dspsn’t change) 1.3 Weak Acid and Bases  Strong acids and bases completely ionize in solution whereas weak acids and bases do not ionize completely  Arrhenius: Acids produce H O 3ons water while bases produce OH in water  Weak acids and bases proceed until an equilibrium is achieved  Brønsted-Lowry: Acid is proton donor, base is proton acceptor  When a weak acid (HA) ionizes, it donates a hydrogen ion to water  When a weak base (B) ionizes, it accepts a hydrogen ion from water  Lewis Acid: accept a pair of electrons, results in formation of a coordinate covalent bond  Lewis Base: donates a pair of electrons to another atom, forms acid-base adduct  Ammonia and amines behave as bases, metal cations act as acids, and oxides of non-metals behave as acids when reacted with hydroxide  If K > 1, equilibrium lies to the right and K < 1, to the left  Larger the K value, the greater the ionization (stronger = ionize more)
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