Chem1120 1st Edition Lecture 7 Outline of Last Lecture I. Method of Initial RatesA. Deriving the Rate Law II. Concentration and TimeA. Quick Logarithm ReviewB. Reaction Orders III. Half LifeOutline of Current Lecture I. How to Control the Rate of a Chemical ReactionA. Collision Model II. The Arrhenius Equation III. Reaction Mechanisms IV. Rate Laws for Multistep Mechanisms V. CatalysisCurrent LectureI. To answer the question of how to control the rate of a chemical reaction, we must first understand what factors modify the speed of a chemical reaction. Four main factors can control the reaction: 1. Concentration- molecules must collide to react2. Physical State- molecules must mix to collide 3. Temperature- molecules must collide with enough energy to react (rates generally increase with increasing temperature)These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.4. The use of a catalyst A. Collision Model (i.e. Collision theory)- reaction rates depend oncollisions, which in turn will likely depend on at least 3 factors:1. Collision Frequency = number of collisions per second per literhigher concentrations ——> more frequent collisionshigher temperatures ——> more frequent collisions2. Collision Energy = fraction of the collisions that are sufficiently forcefulpowerful collisions ——> reactiongentle collisions —— > no reaction3. Collision Orientation = fraction of the collisions with correctly oriented moleculescorrect alignment ——> reaction incorrect alignment ——> no reaction II. Activation Energy, Ea = minimum collision energy required for molecules to reactExample: A + BC ——> AB + CThe lower the activation energy, the faster the reaction!Typically, only a fraction of the molecules in a sample possess the rightamount of energy to react. The higher the temperature, the higher this fraction. The Arrhenius Equation explains that if the reaction rate varies with temperature, so must the rate constant. k = rate constant A = frequency factor = related to collision frequency and collision orientation)Ea = activation energyR = gas constant = 8.314 J/(mol*K)T = Temperature in KelvinHigher T —> larger k —> increased rate Example: 2N2O5 ——> 4NO2 + O2 (g)-measure k at different temperatures, plot k vs temperature-then, using the Arrhenius equation, plot ln(k) vs 1/T-if it yields a straight line, then we know the slope = -Ea / R III.reaction mechanism = the sequence of elementary steps in an overall chemical reactionelementary step = a simple reaction having no intermediates and only having one transition state (A.K.A elementary reaction or elementary process)Example: Consider the reaction:2NO2 (g) + F2 (g) ——> 2NO2F (g) For which the accepted mechanism is: You may be wondering what an intermediate would be. If step 1 were to proceed as:Then NO2F2 would be an intermediate and steps 1a and 1b would be elementary steps rather than the original step 1. The temporary existence of an intermediate implies that there are two transition states in step 1.In order for a mechanism to be judged acceptable, it must satisfy 2 requirements:1. The sum of the elementary steps must result in the overall reaction2. The mechanism must be consistent with the experimentally determine rate lawSince elementary steps are simple collision processes, their rate laws are determined by their molecularity: IV. Rate determining step = a mechanism’s slowest stepThe rate determining step determines the overall reaction rate (analagous toan assembly line’s rate determining step). The rate overall is equal to the rate of the rate determining step, or the slowest step. Example: V. Catalyst = a substance that increases the rate of a chemical reaction without undergoing permanent chemical changeCatalysts make the activation energy smaller or easier to surmount. It is much easier to surmount two energy barriers (humps on the energy graph)than to surmount one giant energy barrier.Homogeneous catalysis = catalyst and reactants in the same phaseHeterogeneous catalysis = catalyst and reactants in different phasesCatalysts are extremely important in industry in the production of NH3, gasoline, plastics, etc…Enzymes are biological catalysts that have a region where the reactants attach (active site). The reactants are referred to as substrates. Most enzymes are large proteins with molar masses from 10^4 to 10^6