Chapter 1: Chemical Equilibrium 1.1 The Equilibrium Constant ...

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

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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 < Keq: reaction must proceed in forward direction  (not enough product) Q > Keq: reaction must proceed in reverse direction  (not enough reactant) Q = Keq: equilibrium 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

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Van’t Hoff equation determines K values at different temperatures (

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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  Ksp is 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 > Ksp, solid 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 (Ksp doesn’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 H3O+ ions 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)  pK and K share the same relationship as pH and [H+]

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If %ionization < 5%, using approximation that c – x ~ c

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More dilute the solution of weak acid / base, the greater the percent ionization Initial concentration is decreased, but in order for Keq to remain constant the % ionization increases (more products relative to reactants at a lower initial concentration) A- is the conjuage base of the weak acid HA and behaves like a base in solution, HA has one more H and is more positive BH+ is the conjugate acid of the weak base B and behaves like an acid in solution Stronger the weak acid, weaker the weak conjugate base will be Spectator ions do not affect pH at all When salt dissolves in water, its cation and anion separate from each other Amphiprotic anion can act as both acid or base Strength of weak acid depends on relative thermodynamic stabilities of their conjugate bases More stable conjugate = stronger the weak will be Delocalized charges make the species thermodynamically stable and thus, more favorable Resonance delocalizes charge over two or more atoms and increase stability of the ion (cumulative) More electronegative atom is better able to bear a negative charge, only works for atoms of same size Distance effect caused by an electronegative atom stabilizing a negative charge somewhere else on an ion Larger the area to delocalize, more stable the ion will be Common ion effect also comes into play with ionization of weak acids and bases Presence of the common ion = even less weak species will ionize Amount suppressed by the common ion is pretty significant Polyprotic: contain more than one ionizable hydrogen atom K1 >> K2; much more H+ is produced in the first step than in the second step pH of the solution is due almost entirely to the H+ produced in the first step

1.4 Buffer Solutions  Solution able to withstand changes in pH  Must contain a weak acid that will react with any OH- ions and a weak base that will react with any H3O+ ions, and must not react with each other  Amounts of species and conjugate must be roughly equal  Three ways to make an acid buffer: start with both weak acid and conjugate base, start with only HA and use strong base to convert, start with only A- and use strong acid to convert  Analogous combinations can be used to form base buffers

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Identifying buffer solutions: identify all species, will there be a reaction, what species are present after pH of a solution must remain nearly constant, %ionization is fairly small Value of x is always negligible when compared to the concentration of the parent species Preferable to use mole amounts pH of a solution is dependent on the ration of parent to conjugate species Diluting a buffer solution does not change its pH Once a buffer solution is formed, it is treated as a equilibrium containing a common ion Adding a small amount of strong acid / base to a buffer solution results in a small change in pH Titrations are used to determine amounts of acid or base in a solution Equivalence point: moles H+ = moles OHNeutralization has occurred but pH is not always 7 Any reaction involving a strong species will go to completion, unidirectional arrow () is used If one species is weak acid and one is strong, the conjugate of the weak species is produced in the reaction Acid-base indicators are used to detect equivalence point Point at which the indicator changes colour is called the endpoint pK for indicator = pH at equivalence point Titration curve shows pH resulting from all the reactions occurring Strong Acid-Strong Base, no buffering action occurs and pH change is very large at the equivalence point (~7) Weak Acid-Strong Base, magnitude and slope of the increase near the equivalence point is less and contains a buffer (>7) Weak Base-Strong Acid, equivalence point (