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Chem 102 Module 2: The Rates of Reactions

by: Lyna Nguyen

Chem 102 Module 2: The Rates of Reactions Chem 102

Marketplace > Texas A&M University > Chemistry > Chem 102 > Chem 102 Module 2 The Rates of Reactions
Lyna Nguyen
Texas A&M

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Module 3
General Chemistry 2
Dr. Bethel
Class Notes
Chemistry, Chem, chem 102, tamu
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This 4 page Class Notes was uploaded by Lyna Nguyen on Tuesday February 9, 2016. The Class Notes belongs to Chem 102 at Texas A&M University taught by Dr. Bethel in Spring 2016. Since its upload, it has received 16 views. For similar materials see General Chemistry 2 in Chemistry at Texas A&M University.

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Date Created: 02/09/16
Module 3: The Rates of Reactions (14.1 – 14.6) 14.1: Rates of Chemical Reactions  Reaction rate = change in concentration / change in time  Calculating a rate: o Concentrations:  Measuring the absorbance of light by a solution o Average rate is the change in concentration per unit time o Units: mol / (L * time) o Negative slope, but expressed as positive value o Instantaneous rate: single point in time  Slope of tangent line at specific point  Relative Rates and Stoichiometry o aA + bB -> cC + dD 14.2: Reaction Conditions and Rate  What effects reaction rate? o Reactant concentrations, temperature, and presences of catalyst  Also surface area if solid  Effects of concentration and temperature o Illustrated by “iodine clock reaction”  If concentration of a reactant is increase, the reaction rate will increase  Chemical reactions occur more rapidly at higher temperatures  Catalysts: substances that accelerate chemical reaction rates but are not consumed by reaction o Catalase: biological catalyst/enzyme o Increase reaction rate  Surface area (only solids) o Smaller particles -> more molecules on surface o More molecules, more reaction 14.3: Effect of Concentration on Reaction Rate  Rate Equation/ Rate Law o Relationship between reactant concentrations and reaction rate is expressed by an equation  Rate equations o Evaluating how the rate is affected when concentrations of the reactants are varied o K[Reactant]  The Order of Reaction o Order: exponent of its concentration term in the rate law expression  Not stoichiometric coefficients o Overall reaction order: sum of all exponents o Gives insight into how the reaction occurs  The Rate Constant o K: relates rate and concentration at a given temperature  Knowing k allows calculation at certain reactant concentrations  Units:  0 order: mol/L*time st  1 order: 1/time  2ndorder: L/mol*time  Wide range of values  Reaction dependent o Determining a rate equation  “method of initial rate”: the initial rate is the instantaneous reaction rate at the start of the reaction (rate @ time = 0) 14.4: Concentration-Time Relationships: Integrated Rate Laws  Integrated Rate Laws o Zero Order  [Ro] – [Rt] = kt o First Order  Ln([R t/[Ro]) = -kt o Second Order  1/[R t – 1/[Ro] = kt  Half Life o Zero Order  [Ao] / 2k o First Order  Ln2 / k o Second Order  1 / k[Ao] 14.5: A Microscopic View of Reaction Rates  Collision theory of reaction rates o 3 requirements must be met  reacting molecules must collide with one another  reacting molecules must collide with sufficient energy to initiate the process of breaking and forming bonds  molecules must collide in an orientation that can lead to rearrangement o the atoms and the formation of products  Collision Theory: Concentration and reaction Rate o Dependence of reaction rate on concentration because number of collisions between 2 reactant molecules is directly proportional to the concertation of each reactant  Collision Theory: Activation Energy o Activation energy: energy required to cross the carrier  If the barrier is low, energy is low, fast reaction  If the barrier is high, energy is high, slow reaction o Reaction Coordinate diagram  Horizontal axis: reaction progress Module 3 Page 2  Vertical axis: potential energy o Transition state: energy of the system reaches a maximum  Also known as activated complex  Collision theory: Activation Energy and temperature o As temperature increases, average energy of the molecules in sample increases  Collision theory: Effect of Molecular Orientation on Reaction Rate o Must come together in the correct orientation o Lower probability of achieving the proper alignment, smaller value of k, slower reaction o Bond rearrangement: both bond breaking and bond formation occurring  The Arrhenius equation o Summarized dependency on energy, frequency of collisions, temperature, and geometry o K = Ae^(-E /Ra)  R: 8.314 o Can be used to calculate E fram temperature dependence of the rate constant o Can be used to calculate the rate constant if E , a and A are known o Temperature increase, rate increase  Effects of catalyst on reaction rate o Speeds up the rate of a chemical reaction o Not consumed in a reaction o Function: provide a different pathway with a lower activation energy for the reaction o Reaction Intermediates: species formed in one step of the reaction and consumed in a later step 14.6: Reaction Mechanisms  Reaction mechanism: sequence of bond making and bond breaking steps that occurs during the conversion of reactants to products o 1 or more steps  Elementary step: each step in a multistep reaction sequence o Formation/rupture of chemical bond o Has its own activation energy and rate constant  Molecularity of Elementary Steps o Elementary steps: classified by number of reactant molecules that come together  Molecularity: number of molecules  1 mole: unimolecular  2 moles: bimolecular  3 moles: termolecular  Rate Equations for elementary steps o Rate equation for any elementary step is defined by its stoichiometry  Reaction Mechanisms and rate Equations o Products of a reaction can never be produced at a rate faster than the rate of the slowest step o Rate determining: slowest step, determines rate of reaction Module 3 Page 3 Module 3 Page 4


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