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# Purdue CHM 11600 - Exam 1 Study Guide

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CHM 116 1st EditionExam # 1 Study Guide Lectures: 1 - 8Lecture 1 (January 12)Units of Concentration- Molarity M = moles solute/volume solution (L)- Molality m = moles solute/mass solvent (kg)- % by mass = mass solute/total mass solution x 100- Mole fraction: Xsolute = moles solute/total # moles in solutionLearning Objectives- Calculate molarity, molality, mass percent and mole fraction in solution- Interconvert among the various terms for expression concentration- Calculate how to prepare a dilute solution for a concentrated oneo Cwant x Vwant = Chave x VhaveLecture 2 (January 14) Chemical Kinetics- Rate = change in moles of A/ change in time where A is any reactant or producto Rate = d[A]/dt at constant volume- Rate Law: Rate = k[A]x[B]y…o Exponents are the orders of each reactant and rely on experimental datao The variables in the brackets represent the molar concentrations of each reactantLearning Objectives- Express rates of reaction with regard to reactant or product concentration changes- Express rate laws- Describe reaction ordersLecture 3 (January 21)Method of Instantaneous Rates- Given a table of experimental concentrations and rates, determine the order of the reaction for the given reaction by calculating how much the rate changes as the concentration changes from one experiment to the nexto For example, if the concentration was doubled from one reaction to the next andit was observed that the rate doubled as a result of this, then it can be concludedthat the order the reaction is 1 because both rate and concentration changed to the same degree- Note that there is no correlation between the stoichiometric coefficient of the reaction and the order of the reaction- K is the only thing that remains constant when comparing reactions- The method of initial rates is where several experiments are performed with different concentrations of reactants and the initial rate is measured for eachLearning Objectives- Derive the rate equation, rate constant, and reaction order from experimental data- Determine rate constant and reaction order from experimental rate dataLecture 4 (January 26)Method When Rate Depends on 2+ Reactants- Write the general from of the rate law- Evaluate relationship between rate and concentration and determine the order given experimental datao The overall reaction order is the sum of all orders appearing in the rate law- Solve for ko Generally, calculate k for all reactions and take an averageMethod of Integrated Rate Laws- Express reactant concentration as a function of time- Will only need to be able to integrate reactions with respect to one reactantCollision Model of Chemical Kinetics- Rate is directly proportional to the # of collisions between reaction species over a given period of timeo Actual rates are much slower than calculated rates predict because there are many collisions that don’t react- Particles must have a certain minimum kinetic energy at impact that is about the activation energy of the reaction in order to produce a chemical reactionLearning Objectives- Describe differential (rate as a function of concentration) and integrated (concentration as a function of time) rate laws for zero, first, and second order reactionsLecture 5 (January 28)Activation Energy of Reactions- Dictates the speed of reactiono Larger activation energy  slowero Smaller  faster because more particles are able to obtain a kinetic energy above the activation energy threshold- On a reaction energy graph, the transition states are the peaks of the grapho This is the point where a bond is partially broken and another is partially formed- Fraction of collisions with energy greater than or equal to EA = e-Ea/RT - Arrhenius Equation: K = Ae-Ea/RTo A is the frequency factor (A= z (collision frequency) x p (steric factor))- Use the equation ln(k) = -Ea/R (1/T) +ln(A) to determine the activation energy of a reaction- Use the equation ln(K2/K1) = -Ea/R (1/T2 – 1/T1) to determine k at various temperatures- Catalysts speed up reactions without being consumedLearning Objectives- Be able to interpret the effect of temperature and concentration on reaction rate in terms of the number of collisions and whether each collision has sufficient energy to achieve the transition state- Be able to find the activation energy from a reaction energy diagramo The change in energy from the reactants to the top of the first peak on the graph- Use the Arrhenius equation to find Ea from experimental data- Be able to interpret a reaction energy diagramLecture 6 (February 2)Elementary Steps- A reaction that occurs exactly as written and can’t be broken down into simpler steps- Intermediate: Species formed and consumed in mechanismo Does not appear in overall equation- How to tell if a mechanism is an acceptable possibility for a reaction:o Do elementary steps sum to produce observed (overall) reaction?o Does mechanism agree with experimentally determined rate law? Identify the rate determining step (slowest) and write the rate law for thisstepMolecularity- The number of reactant particles colliding in an elementary step- One reactant particle  uni-molecular, first order overall- Two reactant particles  bi-molecular, second order overallLearning Objectives- Understand that an elementary step represents a single molecular event and its molecularity equals the number of colliding particles- Describe relationships between reaction mechanism and elementary steps- Describe rate-limiting steps and kinetics of overall reactions- Be able to describe the role of catalysts in affecting forward and reverse reactionsLecture 7 (February 4)Learning Objectives- Be able to describe the role of catalysts in affecting forward and reverse reactionsLecture 8 (February 9)Chemical Equilibrium- A state where amounts of reactants and products remain constant with time- The rates of the forward and reverse reactions are equal- Law of mass actiono At given T, chemical system reaches state where particular ratio of reactant and product concentration is constant- Kp = Kc(RT)dno Dn = the sum of product coefficients – sum of reactant coefficients- Small k  mostly reactants at equilibrium- Large k  mostly products at equilibrium- Magnitude of k and time required to reach equilibrium are not related- Pure liquids and solids do not appear in k expressionLearning Objectives- Describe a system at chemical equilibrium- Describe the law of mass action- Describe the “position” at equilibrium-

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