Unit 3: Rates and Chemical Equilibrium

Reaction Rate How fast a reactant is used up or how fast a product is formed Decreases over time Given the reaction: A  B Rate of reaction can be expressed as the rate of disappearance of A or the rate of appearance of B

Average Rate of Reaction Given as a concentration per unit time Looking at general equation: A  B 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑟𝑎𝑡𝑒 𝑜𝑓 𝑎𝑝𝑝𝑒𝑎𝑟𝑎𝑛𝑐𝑒 𝑜𝑓 𝐵= ∆[𝐵] ∆𝑡 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑟𝑎𝑡𝑒 𝑜𝑓 𝑑𝑖𝑠𝑎𝑝𝑝𝑒𝑎𝑟𝑎𝑛𝑐𝑒 𝑜𝑓 𝐴=− ∆[𝐴] ∆𝑡

Average Rate of Reaction The rate of disappearance of HCl was measured for the following reaction: CH3OH(aq) + HCl(aq)  CH3Cl(aq) + H2O(l) The following data was collected: Time (min) [HCl] (M) 0.0 1.85 54.0 1.58 107.0 1.36 215.0 1.02 430.0 0.580

Average Rate of Reaction What is the average rate of disappearance of HCl in M/s in the first 107.0 minutes? 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑟𝑎𝑡𝑒 𝑜𝑓 𝑑𝑖𝑠𝑎𝑝𝑝𝑒𝑎𝑟𝑎𝑛𝑐𝑒 𝑜𝑓 𝐴=− ∆ 𝐴 ∆𝑡 ∆ 𝐻𝐶𝑙 =1.36𝑀−1.85𝑀=−0.49𝑀 ∆𝑡=6420𝑠−0.0𝑠=6420𝑠 A𝑣. 𝑟𝑎𝑡𝑒 𝑟𝑒𝑎𝑐𝑡𝑖𝑜𝑛 𝐻𝐶𝑙=− −0.49𝑀 6420𝑠 =−7.63× 10 −5 𝑀/𝑠

Speeding up Reaction Rate Solids: Higher subdivision Higher concentration of reactants Gases: Higher pressure

Concrete Formation Five ingredients: Coarse aggregates Fine aggregates Water Mixtures Portland cement Concrete formation is an exothermic reaction

Aggregates Serve as reinforcement to add strength to the concrete Coarse (gravel, crushed stone) Fine (sand)

Mixtures Accelerating (accelerators) Retarding (retarders) Added to concrete to reduce setting time Accelerate early strength (calcium chloride in non steel applications) Retarding (retarders) Used in hot weather to delay setting time Many act as water reducers (sugar)

Mixtures Microscopic air bubbles (air entraining) Fly ash (by-product of coal burning plants) Can replace 15-30% of concrete in mixture Improves workability and reduces heat generated by concrete Microscopic air bubbles (air entraining) Used when exposed to freezing/thawing and deicing salts When hardened concrete freezes the ice inside will expand into the air bubbles instead of cracking concrete

Temperature and Concrete Cold weather Concrete sets more slowly Water expands when frozen, causes cracks if not set before exposure to cold Hot weather Top layer solidifies faster than bottom Concrete shrinks as it solidifies and will cause an uneven product

Concrete Formation

Collision Theory When particles of reactants collide there are only a certain percentage of collisions that are successful In successful collision particles must collide at proper orientation and with enough energy to break existing bonds

Collision Theory Increasing the temperature and concentration of reactants increasing the frequency of collisions as well as the energy in each collision Increases the likelihood of a successful collision

Activation Energy Difference between the initial energy of reactants and the transition state Is the level of energy that must be met in order for a collision to be successful Transition state is the energy maximum at which the activated complex is formed

Energy Diagrams

Catalyst Catalyst is a substance that influences the rate of reaction Remains unchanged at the end of the reaction Lower the activation energy of the reaction Examples: Enzymes Platinum/Rhodium

Catalysis

Rate Law (Rate Equation) Represents the dependence of the rate of reaction on the concentrations of the reactants Rate constant (k) represents the effect of temperature on the reaction

Rate Law (Rate Equation) 𝑎𝐴+𝑏𝐵→𝑐𝐶+𝑑𝐷 𝑅𝑎𝑡𝑒=𝑘 [𝐴] 𝑚 [𝐵] 𝑛 Exponents are usually small whole numbers and are determined experimentally Must be given experimental data or told the values

Rate Law (Rate Equation) The exponents (m and n) in the rate law are called reaction orders Exponent = 1, rate is first order in that reactant Exponent = 2, rate is second order in that reactant The overall reaction order is the sum of all orders in the rate law Sum = 1, reaction is first order overall Sum = 2, reaction is second order overall “Zero order” means the disappearance of the reactant is independent of the reaction rate

Rate Law (Rate Equation) 2NO(g) + 2H2(g)  N2(g) + 2H2O(g) Rate = k[NO]2[H2] Reaction is second order in NO Reaction is first order in H2 Reaction is third order overall

Rate Law (Rate Equation) Calculate the rate of reaction when the reaction is third order in reactant A and zero order in reactant B. The rate constant for this reaction is 2.7 x 10-4M/s and the concentrations of A and B are 0.01M and 0.2M respectively Rate = k [A]m[B]n Rate = 0.00027[0.01M]3[0.2]0 Rate = 2.7 x 10-10M/s

Half Life (t1/2) Time required for the concentration of a reactant to reach half its initial value Fast reactions have a short half life Time that it takes for [A]0 to become 1 2 [A]0

Chemical Equilibrium Reversible reaction: Any reaction that can proceed forward (left to right) or reverse (right to left) direction Represented by a double arrow (⇋) Equilibrium is the dynamic balance between these two processes Catalysts DO NOT effect equilibrium compositions

Equilibrium Constant (Kc) At equilibrium there is a constant changing of products and reactants but the overall concentrations remain constant Equilibrium constant (Kc) relates the concentrations of reactants and products that are in equilibrium

Equilibrium Constant (Kc) For the general equation: aA +bB ⇋ cC + dD 𝐾 𝑐 = [𝐶] 𝑐 [𝐷] 𝑑 [𝐴] 𝑎 [𝐵] 𝑏 Kc >1 = forward reaction favoured Kc <1 = reverse reaction favoured Kc ≠ 1 Concentration is in mol/L

Equilibrium Constant (Kc) Concentration of pure liquids (l) and solids (s) doesn’t change so they do not appear in equilibrium equations Heterogeneous = involving different states Homogeneous = involving one state Given: H2O(g) + C(s) ⇋ H2(g) + CO(g) 𝐾 𝑐 = [𝐻2] 1 [𝐶𝑂] 1 [𝐻2𝑂] 1

Equilibrium Constant (Kc) Given the following information calculate Kc: N2(g) + 3H2(g) ⇋ 2NH3(g) [N2] = 0.184M [H2] = 0.352M [NH3] = 1.032M 𝐾 𝑐 = [𝑁𝐻3] 2 [𝑁2] 1 [𝐻2] 3 = [1.032] 2 [0.184] 1 [0.352] 3 = 46.71 Kcreverse = 1 46.71 =0.021

Equilibrium Constant (Kp) With gases the pressure (atm) of the reactants and products can be compared Kc becomes Kp Kc and Kp vary with temperature

La Châtelier’s Principle When a system at equilibrium is disturbed it will compensate for the disturbance Examples of altering: Addition of Catalyst (Trick! No change!) Concentrations Pressures Temperature

La Châtelier’s Principle Concentration Adding substance, system responds by using it up Removing substance, system responds to make more N2(g) + 3H2(g) ⇋ 2NH3(g) Adding N2, favours right side Removing H2, favours left side Adding NH3, favours left side

La Châtelier’s Principle Pressure Increasing pressure (decreasing volume of container) favours side with fewer mols Decreasing pressure (increasing volume of container) favours side with more mols N2(g) + 3H2(g) ⇋ 2NH3(g) Increasing pressure, favours right side Increasing volume, favours left side

La Châtelier’s Principle Temperature Heat acts as a product or reactant Exothermic reactions, heat = product Endothermic reactions, heat = reactant When temperature of system is increased Adding reactant to endothermic reaction Adding product to exothermic reaction When temperature of system is decreased Adding product to endothermic reaction Adding reactant to exothermic reaction

La Châtelier’s Principle C(s) + H2O(g) ⇋ CO(g) + H2(g) ΔH = +131.3kJ Endothermic reaction (ΔH = +) Heat is a reactant ΔH + C(s) + H2O(g) ⇋ CO(g) + H2(g) Adding heat, favours right side Removing heat, favours left side

La Châtelier’s Principle Because heat can be seen as a product of reactant it will effect our K just like any other product or reactant 𝐾= 𝑃𝑟𝑜𝑑𝑢𝑐𝑡𝑠 𝑅𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 Endothermic: ↑T = ↑K Heat is a reactant, adding more favours the products Exothermic: ↑T = ↓K Heat is a product, adding more favours the reactants

La Châtelier’s Principle