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Department
Chemistry
Course
Chemistry 1027A/B
Professor
Felix Lee
Semester
Fall

Description
Chemistry 1200B Final 1 Chapter 1: Chemical Equilibrium 1.1: The Equilibrium Constant Equilibrium: amounts remain constant, overall process reaches equilibrium (dynamic, not static situation) For any equilibrium, there is a mathematical relationship between the pressures or concentrations of the reactants and products K, the equilibrium constant Equilibrium constants are temperature-dependent If all the components of an equilibrium system are in the same phase, the system is homogeneous (if they are in different phases, it is heterogeneous) The concentration of any solid is considered to be constant, and is therefore omitted from the equilibrium constant expression How k is affected by: o Reversing a reaction: k is inverted (1/k) o Multiplying or dividing: k is put to the power of whatever you multiply by o Combining equilibria: add together to give overall equation The magnitude of the equilibrium constant is an indication of the relative amount of product or reactant present at equilibrium o If k is greater than 1, the equilibrium lies towards the right (more products than reactants) o If k is less than one, it lies to the left (more reactants than products) If the amounts of products and reactants can be determined, a quantity known as the reaction quotient, Q, can be calculated this number indicates whether or not the system is at equilibrium, and the direction in which the reaction must proceed to attain equilibrium o Q expression is the same as K expression, except that the concentrations used to calculate Q are not necessarily the equilibrium concentrations used as comparison o Q can be found at some point, and is compared to K o A system is at equilibrium only when Q=K o If Q is less than K, the reaction must proceed to the right to obtain equilibrium. If Q is greater than K, then the reaction must proceed to the left to obtain equilibrium Once a system has reached equilibrium, the concentrations or pressures of the reactants and products remain constant (system hasnt stopped reacting, but forward and reverse rates are equal) o Changing the system by changing the temperature, pressure, adding reactants will disturb the equilibrium Le Chateliers principle: If a chemical system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of a participant in the equilibrium, the equilibrium will shift in such a way as to minimize the disturbance One way to disturb a gaseous equilibrium is to change the total pressure of the system o Can be achieved by changing the volume of a system o If an inert gas is added, the result will be an increase in the total pressure, but there will not be any changes to the partial pressures of the reactant and products (so no change in position of equilibrium) If the temperature of the system is changed, the value of K will change o In an exothermic reaction (heat is lost to the surroundings): increasing the temperature has the same effect as increasing the amount of one of the products (shift of equilibrium to the left, decreasing K) Chemistry 1200B Final 2 o Adding heat to an endothermic reaction will shift to the right, increasing K Gibbs Free Energy change can be used to determine whether a reaction proceeds spontaneously or not o G = spontaneous: equilibrium lies to the right and k is large o +G = nonspontaneous: equilibrium lies to the left, k is small Delta G = Standard Delta G + RTlnQ Vant Hoff equation: used to determine K values at given temperatures (given) 1.2: Solubility of Ionic Compounds Solubility: the amount of that substance that will dissolve in a certain volume of a solvent o Soluble: 10g or more can dissolve in a litre of solvent o Slightly soluble: 0.1g-10g o Insoluble: less than 0.1g can dissolve Soluble: nitrates, alkali metal4 NH , halides (Cl, Br, I), sulfates Insoluble: sulfides, carbonates, phosphates, hydroxides Solvent: liquid or a gas that dissolves another solid, liquid, or gas (solutes dissolve in a solvent) Saturated: no more additional solute can dissolve in the solution (at equilibrium) Salts that are very soluble have a very larsp K value (usually go to completion) Solubility product constant is an equilibrium constant, and is the same for a given solid at a given temperature vs. solubility, which can vary at a given temperature due to other conditions The reaction quotient: o If QK , solid will formed precipitation, solution is supersaturated sp Common Ion Effect: the solution may already contain an ion in common with the dissolving salt can significantly reduce the solubility of the solid relative to that in pure water o Common ion shifts equilibrium 1.3: Weak Acids and Bases Acids: increase hydrogen ion concentration Bases: increase hydroxide ion concentration o Strong acids and bases completely ionize in solution, whereas weak acids and bases do not ionize completely o Weak acid and base solutions proceed until an equilibrium is achieveda(K andbK Arrhenius theory of acids an+ bases: o An acid produces H 3 in water o A base produces OH in water Bronsted-Lowry theory o An acid is a proton donor o A base is a proton acceptor Lewis Acids and Bases o A Lewis acid can accept a pair of electrons from another atom, forming a coordinate covalent bond o A Lewis base donates a pair of electrons o The species formed is called an acid-base adduct Ammonia and amines behave as Bronsted-Lowry bases o The nitrogen atom of NH 3arries a nonbonding pair of electrons, which attracts the partially positive hydrogen of water hydrogen is transferred to the amine Chemistry 1200B Final 3 o The amine donates the nonbonding pair of electrons to form a new N-H bond, resulting in NH 4 Metal cations act as Lewis acids: form complex ions when they react with Lewis bases such as water, NH 3r OH sometimes called coordination complexes o Oxides of non-metals also behave as Lewis acids when they react with hydroxide + - Weak acid equilibrium: K a H A / HA Weak base equilibrium: K b BH OH/B K andaK valubs are much less than one % ionization = x/c times 100% o If the % ionization is less than 5%, the approximation that x is much smaller than c is used (making c-x around c) The more dilute the solution of a weak acid or base, the greater the percent ionization In the equation of the ionization of a weak acid in solution: o HA + H 2 H O3+ A - - o HA is the weak acid, and A is the conjugate base of the weak acid o The conjugate base has one less hydrogen, more negatively charged, will behave like a weak base when in solution -14 The equilibrium constant for the ionization of water iswK = 1x10 o Ka x Kb = Kw The stronger the weak acid is, the weaker its conjugate base will be Strong acids have extremely weak conjugate bases their bases are so weak, that they are spectator ions and do not affect pH at all (like Cl as the spectator ion for HCl A salt is an ionic solid containing cations and anions when a salt dissolves in water, its cation and anion separate from each other o If the cation and anion can affect the pH, then if Ka> Kb, the salt is acidic The relative strengths of weak acids are primarily determined by the relative thermodynamic stabilities of the conjugate bases What affects the stability of conjugate bases? o Resonance: increases the stability of the ion (cumulative effects also) o Electronegativity of the atom bearing the negative charge: more stable = more electronegative o Inductive effect: distance effect, further stabilizes by pulling electrons away Polyprotic: they contain more than one ionisable hydrogen atom these acids ionize in stages, each with a different a value 1.4: Buffer Solutions Buffer: both the weak acid and its conjugate base are added to the same solution maintains pH. Requirements: o Must contain a weak acid that will react with any OH o Contain a weak base that will react with any 3 O o The acid and the base in the buffer must not react with each other (weak acid or base and its own conjugate, in equilibrium) Acid buffer: weak acid and its conjugate base start with acid, and add a salt that contains the conjugate base, or react the weak acid wi
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