Reactants. Products. Reactants. Products. Chapter 9 Study Sheet. First Law: E universe = E system + E surroundings = 0 E system = - E surroundings

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1 First Law: E universe = E system + E surroundings = 0 E system = - E surroundings Chapter 9 Study Sheet E system = q + w system does work, w is negative work done on system, w is positive system loses heat, q is negative system gains heat, q is positive Enthalpy At constant pressure, q = H Reactants Products H H Products H < 0 H > 0 exothermic endothermic feels hot feels cold Ways to figure out enthalpy 1) How the reaction feels hot = exothermic = H <0 cold = endothermic = H >0 2) Bond Enthalpies Sum of bonds broken sum of bonds formed Be sure to draw Lewis Structures N 2 + 3H 2! 2NH 3 H ~+(N N) + 3(H-H) -6(N-H) H ~ + (941) + 3(436) -6(391)~ -97 kj 3) Phase changes Reactants Gas Gas Potential Energy E out E out Liquid Potential Energy Liquid E in E in Solid going down = exothermic = H <0 going up = endothermic = H >0 Solid

2 chapter 9 study sheet page 2 of 7 4) Bond strengths Weak bonds (low melting point) Entropy Strong bonds (high melting point) going down = exothermic = H <0 going up = endothermic = H >0 S = q / T or heat content / temperature in K. A measure of the concentration of energy. S universe = S system + S surroundings > 0 Whenever something happens, the universe is getting messier. Ways to figure out entropy 1) Gas in vs. Gas out 2 mol gas! 0 mol gas S < 0 0 mol gas! 1 mol gas S > 0 3 mol gas! 3 mol gas S ~ 0 2) Phase Changes Gas (dispersed energy, low order) liquid (more concentrated energy, medium order) solid (very concentrated energy, high order) going down = concentrating energy = S <0 going up = dispersing energy = S >0 3) Observation. If the reaction is bubbling, it s making gas and S >0. If the reaction requires some gas to work, S <0.

3 chapter 9 study sheet page 3 of 7 Free Energy From the second law of thermodynamics, the following equation is derived: G = H - T S Whenever G < 0 for a reaction in the forward direction, the forward reaction is spontaneously happening. Whenever G > 0 for a reaction in the forward direction, the forward reaction is not happening. However, the reaction is spontaneous in reverse. The sign of G is determined by the interaction of H and S. While you can memorize the chart, it might be easier to think of like this: At low temp, G is the same sign as H At high temp, G is the opposite sign of S. If S is near zero, at high temp G is the same sign as H. Free Energy at Standard State H S G is < always + - never + + At high temperature - - At low temperature The value of G changes with the concentration of the products and reactants. There is no chart out there with G values at every set of conditions. Chemists do have charts of G o, H o, and S o at standard state conditions. Standard state is the usual state of a material at 1 atmosphere and the defined temperature: Solids are pure. Liquids are pure. Solutes (aq) are at 1 mol/l. Gases are at 1 atm. A chemist can look up the values of G o, H o, and S o. Then a chemist can use the formula below to figure out what G is under the set of conditions the reaction is run. G = G o + RT ln ([products]/[reactants]) Whenever G o < 0 for a reaction in the forward direction, the reaction is extensive. Extensive = spontaneous at standard state conditions = products preferred at equilibrium. Whenever G o > 0 for a reaction in the forward direction, the reaction is non-extensive. However, the reaction is extensive in reverse. Non-extensive = non-spontaneous at standard state conditions = reactants preferred at equilibrium.

4 chapter 9 study sheet page 4 of 7 Reaction Composition Diagrams The relationship of G o vs G can be explained with reaction composition diagrams. Extensive Composition Diagram Imagine a ball starting at the far left of the figure shown (100% reactants). It will roll downhill until it reaches the lowest point. Rolling left to right would be the reaction in the forward direction. Rolling downhill means G <0 for the forward reaction at those compositions. G 100% R 0% P A(aq) + B(aq) 75% R 25% P C(aq) + D(aq) Rxn Composition 25% R 75% P Equilibrium Imagine a ball starting at the far right of the figure shown (100% products). It will roll downhill until it reaches the lowest point. Rolling right to left would be the reaction in the reverse direction. Rolling downhill means G <0 for the reverse reaction at those compositions. The low point in the curve is where G = 0. The reaction is at equilibrium at this point. Notice how this is an extensive reaction: G o <0 because the reactants are at higher energy than the products. The equilibrium concentration has more product. Notice also, that if you were to start the reaction at say 10% reactant and 90% product, the reaction would have G>0 for the forward reaction (uphill). Just because a reaction is extensive doesn t mean it will be spontaneous in the forward direction all the time, especially if you start out with more product than the equilibrium concentration allows. G o 0% R 100% P Non-extensive Composition Diagram The non-extensive reaction diagram is shown here. The reactants are at lower energy than the products, such that G o >0. Notice how a ball starting at the far left of the figure shown (100% reactants) will roll downhill a tiny bit, until equilibrium is reached at about 80% reactant, 20% product. G uphill G o Just because a reaction is non-extensive doesn t mean it won t produce any product. It will go until it reaches equilibrium, which is still mostly reactant with very little product.

5 chapter 9 study sheet page 5 of 7 Kinetics While lots of reactions can happen, not all of them will happen quickly. Kinetics is the study of how fast a reaction occurs. The speed of a reaction is dictated by the concentrations of the species involved, and the energy of the transition state. There are 3 possible scenarios: 6 kj/mol 7 kj/mol 2 kj/mol X-Y, Z X, Y-Z N-O, P N, O-P G (forward reaction) < 0 G = 3 6 = -3kJ/mol If S small, H (forward reaction) < 0 E act (forward) < E act (reverse) E act (forward) = + (9-6) E act (reverse) = +6 kj/mol (9-3) k forward > k reverse K > 1 At equilibrium, [reactant] < [product] G (forward reaction) = 0 G = 2 2 = 0 kj/mol If S small, H (forward reaction) = 0 E act (forward) = E act (reverse) E act (forward) = +5 kj/mol (7-2) E act (reverse) = +5 kj/mol (7-2) k forward = k reverse K = 1 At equilibrium, [reactant] = [product] 18 kj/mol A, B-C A-B, C G > 0 G (forward reaction) = 9 3 = +6kJ/mol If S small, H (forward reaction) > 0 E act (forward) > E act (reverse) E act (forward) = +15 kj/mol (18-3) E act (reverse) = + (18-9) k forward < k reverse K < 1 At equilibrium, [reactant] > [product] Rate of Reaction Forward Rate = k forward * [reactants] Reverse Rate = k reverse * [products] Remember that only (aq) and (g) states are included in the rate expressions, and you need to raise them to the power of their coefficient in the transition state. A(aq) + C(aq)! 2M(aq) Forward rate = k forward * [A] 1 * [C] 1 Reverse rate = k reverse * [M] 2

6 chapter 9 study sheet page 6 of 7 Equilibrium Equilibrium occurs when G = 0 at which point Rate forward = Rate reverse. This means when G forward reaction + G reverse reaction = 0. K = k forward = [products] k reverse [reactants] For a reaction in which reactants are at higher energy than products, [reactants]<[products] and K >1 at equilibrium. X-Y, Z 6 kj/mol X, Y-Z For a reaction in which reactants are at lower energy than products, [reactants]>[products] and K < 1 at equilibrium. 18 kj/mol A, B-C A-B, C Remember that only (aq) and (g) states are included in K. 2A(aq) + B(aq) C(aq) + 2D(aq) K = [C][D] 2 / [A] 2 [B] A(s) + 3B(aq) 2C(aq) + D(g) K = [C] 2 [D] / [B] 3 LeChatelier s Principle When a system at equilibrium is placed under a stress, it will respond so as to appear to relieve the stress. Remember that at equilibrium, only changes in (aq) and (g) species or heat impact the concentrations of other (aq) and (g) species. I know many of you think that if more solid is added, you can impact the equilibrium. This is not true if there is solid already present in the equilibrium mixture. Imagine a really sweet pitcher of Kool-Aid with sugar on the bottom. You pour your drink of Kool- Aid from the top, without disturbing the sugar on the bottom too much. The Kool-Aid is sweet because the sugar is at the maximum concentration in water. If you add more sugar to the pitcher, can you make the Kool-Aid any sweeter? No. The extra sugar just sits on the bottom with the sugar that was already there. Once the sugar has reached its maximum concentration, it s not going to go over that amount. Therefore, adding more

7 chapter 9 study sheet page 7 of 7 sugar to a solution already saturated in sugar won t make the drink any sweeter. It s the same with chemical reactions. If the solid is already there and dissolved to it s maximum amount (which it will be at equilibrium), adding more solid won t increase the concentration. If you can t change the concentration of something, adding more won t affect the equilibrium. You can think of LeChatlier s as either pushing or pulling the reaction in a particular direction. You can push by adding something. You can pull by taking something away. If a reaction is endothermic, think of heat as a reactant. If a reaction is exothermic, think of heat as a product. Al(OH) 3 (s) Al 3+ (aq) + 3OH 1- (aq) H > 0 Result Stress Visualization Reaction Al 3+ d OH 1- adding Al 3+ Pushing from the right shifts left decreases decreases removing OH 1- Pulling from right shifts right increases increases adding Al(OH) same same adding heat (endothermic so heat is reactant) adding water Pushing from the left shifts right increases increases More Al(OH) 3 dissolves, but as long as solid is present, concentration of ions the same. - same same