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CHEM 112 - Week 2 Notes

by: Christopher Cooke

CHEM 112 - Week 2 Notes CHEM 112

Christopher Cooke
Penn State
GPA 3.5
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Class notes from the second week of CHEM 112, including unimolecular, bimolecular, and trimolecular elementary reactions, the Arrhenius equation, and reaction rates.
Chemical Principles II
Dr. Raymond Shaak
Class Notes




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This 3 page Class Notes was uploaded by Christopher Cooke on Tuesday September 6, 2016. The Class Notes belongs to CHEM 112 at Pennsylvania State University taught by Dr. Raymond Shaak in Fall 2016. Since its upload, it has received 146 views. For similar materials see Chemical Principles II in Chemistry at Pennsylvania State University.

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Date Created: 09/06/16
CHEM 112 Week 2 The Arrhenius Equation  Remember, ΔE and reaction rate are independent of each other.  Linearized Arrhenius equation: Y- X-axis Y-axis Slope intercept 1/T ln(k) -Ea/R ln(A)  In slope-intercept formula: y=mx+b  ln(k) = (-E /R)(a/T) + ln(A)  Deriving the formula if there is more than one k value (rate constant):  ln(k1) – ln(2 ) = (-a /R)(1/T) + ln(A) o ln(k1/k2) = (-Ea/R)(1/T2– 1/T 1  ln(k1/k2) = (Ea/R)(1/T2– 1/T 1  Remember, the Arrhenius Equation combines 3 factors: o Collision frequency o Orientation factor o Fraction of molecules above the energy threshold (E ≥ E ) a  The activation energy is typically about half as strong as an atomic bond (~100 kJ/mol vs ~200 kJ/mol) Reaction Rate  Reaction concentration varies as a function of time  Rate = Δ[X]/Δt (m/s) o Instantaneous, average, and initial rate  Can be expressed in terms of changes of concentration of each reactant and product as a function of time;  αA + ßB = γC + δ(D), where the reactants are A and B, the products are C and D, and α, ß, γ and δ represent their coefficients.  - (ΔA/αΔt) = -(ΔB/ßΔt) = (ΔC/γΔt) = (ΔD/δΔt); all reaction rates are positive Rate Law  Rate is proportional to the number of effective collisions divided by time  Rate = k[x] o Recall: αA + ßB = γC + δ(D) o Rate = -Δ[A]/Δt = k[A] [B] ; this is the generic rate law  Rate is always positive and measured in m/s  x: order with respect to A  y: order with respect to B o x and y are NOT the same as α and ß  x + y: overall order of the reaction  k: rate constant; varies with time and depends on the overall order (x+y) --------------------------------------------------------------------------------------------------------------------- --------------------- Unimolecular elementary reactions  A  products (e.g. O  O + O) 3 2  The relationship between the rate of a unimolecular reaction and the concentration of A is directly proportional by a factor of k (the rate constant)  The rate law for a unimolecular reaction is k[A] where x is always equal to 1; measured in M/s Bimolecular elementary reactions  A+B  products (NO + O  NO 3 O ) 2 2  The rate of product formation depends on the frequency of molecular collisions, which depends on the concentrations of A and B  Rate law for bimolecular elementary reactions x y o k[A] [B] ; x = y = 1; overall order = 2; measured in M/s  Keep in mind this is only for ELEMENTARY reactions -1 -1  k’s units are M s (1/Ms)  2A  products; rate = k[A] 2 Termolecular elementary reactions  A+B+C  products x y z o Rate: k[A] [B] [C] ; x = y = z = 1; overall order = 3  Again only for elementary reactions -2 -1 2  k’s units are M s (1/M s)  A+2B or 2A+B  products x 2 2 y o Rate: k[A] [B] or k[A] [B]  3A  products o Rate: k[A] 3 *** It is impossible to tell if a reaction is elementary from just looking at it. If a reaction is elementary, it will be specified. If otherwise, it will not be mentioned.*** Reaction mechanisms  Sequence of elementary steps that together will give the overall reaction  May contain 3 or more steps which are often unknown  For example, consider 2NO(g) + O (g)  2NO. The “steps” may be: o NO + NO  N O 2 2 o N 2 2 O  2NO 2  N 2 ,2then, is an intermediate, or a compound that forms and goes away within the steps of a reaction. o Intermediates are neither reactants or products, but are important to be aware of (for example, what if cyanide was an intermediate in the process of making certain medicines?)  Different reactions occur at different rates; whichever is the slowest will determine the rate of the entire reaction. o Higher E =aslower rate --------------------------------------------------------------------------------------------------------------------- --------------------- Mechanisms and rate-determining steps  Consider the reaction: 2NO  2N2 + O 2 o Mechanism could be:  NO 2 NO + O (slow)  O + NO  2 + N2 (fast) o Could also be:  2NO 2NO + N3 (slow)  NO 3 NO + O (fas2) o How would you determine which is correct?  For the first: Rate = k[NO ] 2 2  For the second: Rate = k[NO ] 2  Experimentally determine which rate matches – can be done  Consider the following: NO + CO2 NO + CO 2  Possible mechanism: NO + NO  NO + NO (slow); NO + CO  NO + CO 2 2 3 3 2 2 (fast) o Which rate law is consis2ent with this proposed mechanism?  Rate = k[NO ] 2 because the reaction is not elementary  Consider the following: 2NO + BR  2NOB2  Possible mechanism: NO + BR  NOBr 2fast); NO2r + NO  2NOBr2(slow) o Which rate law is consistent with this proposed mechanism?  Consider the following: Cl + 2HCl  HCl3+ CCl 4  Given the mechanism below, what should be the observed rate law? o Cl2 2Cl (fast); Cl + CHCl  H3l + CCl (slow)3 Cl + CCl  CCl (fa3t) 4 o k-1 k-2 k-3 o Reactants: Cl , 2HCl 3 o Intermediates: 2Cl, Cl, CCl 3 o Products: HCl, CCl 4 Autopilot – write out the “slow step”:  Rate = k[Cl][CHCl ] 3 cannot have an intermediate in the rate law expression.  Go back to the first step: Rate = k [Cl1] (2orward step); Rate = k [Cl] -1 2 (backwards step)  Set these rates equal to each other. k1 1/2 1/2  [Cl] = ( /k-1l ]2) = k obs[Cl2]  Combine, and this becomes acceptable.  Rate = k obsCl2] [CHCl ] 3


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