unit 5 ap chem- emily dawson

what is the equation for a second order reaction?

(1/[A]t) - (1/[A]0) = kt

what are the 3 conditions of the collision theory?

1) particles must collide 2) particles must collide with sufficient energy 3) particles must collide with the proper orientation

what are the 2 criteria for plausible mechanism?

1. the sum of elementary steps must lead to the final overall reaction with the correct stoichiometry 2. the mechanism must agree with an experimentally determined rate law expression

a student conducts two trials of a reaction, but the reaction occurs at a much slower rate in the second trial. the student observes that the temperature in the lab is lower in the second trial. using particle-level reasoning, provide TWO explanations that help to account for the fact that the reaction rate was slower in the second trial

1. the temperature was lower in the second trial, so the average kinetic energy of the reactant particles was lower. therefore, there were fewer collisions between particles with sufficient energy to react 2. since the temperature was lower, the kinetic energy was lower and the average speed of the particles was lower. at lower speeds, the reactant particles collide less frequently

consider the reaction at which NO (g) reacts with oxygen to produce N2O3 (g) which destroys ozone in our atmosphere. how fast does the concentration of oxygen decrease when N2O3 forms at 8.0 x 10^-5 M/s? 4 NO + O2 --> 2 N2O3

4.0 x 10^-5 M/s

at 25 degrees Celsius, hydrogen iodide breaks down very slowly to hydrogen and iodine: rate= k[HI]^2. the rate constant at that temperature is 2.4 x 10^-21 (L/molxs). if 0.0100 mol of [HI] (g) is placed in a 1.0 L container, how long will it take for [HI] to reach 0.00900 mol/L (10% reacted)?

4.63 x 10^-21 s

the rate law for a reaction is rate= k[ClO2]^2[OH-] what is the reaction order in each reactant and the overall reaction order?

ClO2= second order OH-= first order overall reaction order= third order

step 1: ? SLOW step 2: NO3 (g) + CO (g) --> NO2 (g) + CO2 (g) FAST overall: NO2 (g) + CO (g) --> NO (g) + CO2 (g) a two-step reaction mechanism is proposed for a gas-phase reaction, as represented above. which of the following correctly identifies both the chemical equation for step 1 and the rate law for the overall reaction? a) the chemical equation for step 1 is 2 NO2 (g) --> NO (g) + NO3 (g), and the rate law is rate= [NO2]^2 b) the chemical equation for step 1 is NO2 (g) + CO (g) --> NO (g) + CO2 (g), and the rate law is rate= [NO2][CO] c) the chemical equation for step 1 is NO3 (g) + 2 CO (g) --> NO (g) + 2 CO2 (g), and the rate law is rate= [NO2][CO] d) the chemical equation for step 1 is NO3 (g) + NO2 (g) + 2 CO (g) --> NO (g) + NO2 (g) + 2 CO2 (g), and the rate law is rate= [NO3][NO2][CO]^2

a

step 1: H2 (g) + ICl (g) --> HI (g) + HCl (g) SLOW step 2: HI (g) + ICl (g) --> HCl (g) + I2 (g) FAST which of the following represents a rate law for the overall reaction that is consistent with the proposed mechanism? a) rate= k[H2][ICl] b) rate= k[H2][ICl]^2 c) rate= k[H2][HI][ICl]^2 d) rate= k[HI][ICl]

a

step 1: NO2 (g) + F2 (g) <--> NO2F2 (g) FAST step 2: NO2F2 (g) --> NO2F (g) + F (g) SLOW step 3: F (g) + NO2 (g) --> NO2F (g) FAST a proposed mechanism for the chemical reaction 2NO2(g)+F2(g)→2NO2F(g) is shown above. which of the following rate laws is consistent with this mechanism? a) rate= k[NO2][F2] b) rate= k[NO2]^2[F2] c) rate= k[NO2][F2]^2 d) rate= k[NO2]^2[F2]^2

a

sucrose --> sucrase-sucrose complex + H2O --> glucose and fructose --> sucrase (the cycle repeats) the hydrolysis of sucrose is extremely slow in aqueous solution. however, when sucrase is added as shown in the reaction above, the rate of the reaction is about 6,000,000 times faster. based on this information, which of the following best explains the large increase in the rate of hydrolysis that occurs with the addition of sucrase? a) the reaction proceeds through a different reaction path with a lower activation energy in which the sucrase-sucrose complex is formed as an intermediate b) the addition of the sucrase reduces the number of effective collisions between the water and sucrose molecules c) the reaction proceeds through a different reaction path that is independent of temperature and concentration d) the addition of the sucrase decreases the number of sucrose molecules that have the minimum energy required to overcome the activation energy barrier

a

the proposed rate-determining step for a reaction is 2 NO2 (g) --> NO3 (g) + NO (g). the graph above shows the distribution of energies for NO2 (g) molecules at two temperatures. based on the graph, which of the following statements best explains why the rates of disappearance of NO2 (g) are different at temperature 2 and temperature 1? a) NO2(g) is consumed at a faster rate at temperature 2 because more molecules possess energies at or above the minimum energy required for a collision to lead to a reaction compared to temperature 1 b) NO2(g) is consumed at a faster rate at temperature 2 because the molecules have a wider range of energies allowing for a better orientation during a collision compared to temperature 1 c) fewer NO2(g) molecules have a relatively high energy at temperature 1, which favors collisions between molecules rather than between the molecules and the container, leading to a faster rate of disappearance compared to temperature 2 d) more NO2(g) molecules have a relatively low energy at temperature 1, which increases the number of effective collisions taking place and the rate of disappearance compared to temperature 2

a

when the chemical reaction 2 NO (g) + O2 (g) --> 2 NO2 (g) is carried out under certain conditions, the rate of disappearance of NO(g) is 5.0×10−5 Ms−1. what is the rate of disappearance of O2(g) under the same conditions? a) because two molecules of NO are consumed per molecule of O2, the rate of disappearance of O2(g) is 2.5×10−5 Ms−1 b) because two molecules of NO are consumed per molecule of O2, the rate of disappearance of O2(g) is 1.0×10−4 Ms−1 c) because the rate depends on [NO]^2 and [O2], the rate of disappearance of O2(g) is 2.5×10−9 Ms−1 d) because the O2(g) is under the same experimental conditions as NO, it is consumed at the same rate of 5.0×10−5 Ms−1

a

in a first-order decomposition reaction, 50.0% of a compound decomposes in 10.5 min a) what is the rate constant of the reaction? b) how long does it take for 75% of the compound to decompose?

a) 0.066 min-1 b) 21 min

the decomposition of hydrogen peroxide is a useful way to obtain oxygen. it is a first order reaction with a rate constant of 0.0410 min-1 2 H2O2 (l) --> 2 H2O (l) + O2 (g) a) if we start with a 3.00% hydrogen peroxide solution, what will it's concentration be after 30 min? b) how long will it take for its concentration to drop to 2.00%? c) how long will it take for one-half of the sample to decompose?

a) 0.87% hydrogen peroxide b) 9.76 min c) t1/2= 16.9 min

in a study of ammonia production, an industrial chemist discovers that the compound decomposes to its elements N2 and H2 in a first-order process. she collects the following data: time (s)- 0; [NH3] 4.00 time (s)- 1; [NH3] 3.986 time (s)- 2; [NH3] 3.974 a) use graphical methods to determine the rate constant b) what is the half-life for ammonia decomposition?

a) k= -0.003 s-1 b) 231 seconds

A --> B + C [A] (M)= 0.1; rate (M/min)= 0.2 [A] (M)= 0.2; rate (M/min)= 0.56 [A] (M)= 0.3; rate (M/min)= 1.04 [A] (M)= 0.4; rate (M/min)= 1.6 a) using this data, calculate the order of the reaction (m) b) write expression for reaction rate c) calculate k, choosing any set of data for concentration and rate d) what is the rate of reaction if [A] is 0.5 M

a) m= 1.5 b) rate= k[A]^1.5 c) k= 6.33 d) rate= 2.23 M/min

a 2-step mechanism is proposed for a reaction, as represented below step 1: NO (g) + Cl2 (g) --> NOCl2 (g) step 2: NOCl2 (g) + NO (g) --> 2 NOCl (g) a) write the rate law for elementary step 1 b) identify the intermediate and justify your choice c) what is the overall chemical equation?

a) rate= k[NO][Cl2] b) NOCl2 is the intermediate because it is produced in step 1 and consumed in step 2 c) 2 NO (g) + Cl2 (g) --> 2 NOCl (g)

a reaction occurs according to the proposed 3 step mechanism below step 1: A + A --> C step 2: C+ B --> D + F step 3: D+ B --> E + F a) write the rate law for each elementary step b) what is the overall chemical equation? c) it is unlikely that this overall reaction occurs in a single elementary step, explain why

a) step 1 rate= k[A]^2; step 2 rate= k[C][B]; step 3 rate= k[D][B] b) 2 A + 2 B --> E + 2 F (C and D are intermediates) c) unlikely they would simultaneously collide with proper orientation and enough energy to overcome the activation energy barrier

CO(NH2)2 <--> NH4+ + OCN- a student studying the decomposition reaction of urea (CO(NH2)2) runs the reaction at 90 degrees Celsius. the student collects data on the concentration of urea as a function of time, as shown by the table 0 hours- 0.1 M 5 hours- 0.0707 M 10 hours- 0.05 M 15 hours- 0.0354 M 20 hours- 0.0250 M 25 hours- 0.0177 M 30 hours- 0.0125 M the student proposes that the rate law is rate= k[CO(NH2)2] a) explain how the data support the student's proposed rate law b) determine k (include units)

a) the data show a constant half-life of 10 hours, which indicates the reaction is a first order reaction b) k= 0.069 hours-1

a mechanism has been proposed for the following reaction: Cl2 (g) + CHCl3 (g) --> HCl (g) + CCl4 (g) step 1: Cl2 (g) <--> 2 Cl (g) FAST step 2: Cl (g) + CHCl3 (g) --> CCl3 (g) + HCl (g) SLOW step 3: Cl (g) + CCl3 (g) --> CCl4 (g) FAST a) does the overall balanced equation from the mechanism agree with the given stoichiometry? b) identify intermediate(s) and catalyst(s) in this mechanism c) determine the rate law expression predicted from this reaction d) write an expression for the observed rate constant in terms of rate constants from the proposed mechanism

a) yes b) Cl (g) and CCl3 (g) are intermediates, there are no catalysts c) rate= k2 (k1/k-1)^1/2 [Cl2]^1/2[CHCl3] d) kobs= k2 (k1/k-1)^1/2

energy required to break a bond

activation energy

(FLIP FLASHCARD FOR MODELS) the particle models shown above represent a proposed two-step mechanism for the destruction of ozone (O3) in the upper atmosphere. based on the proposed mechanism, which of the following best describes the concentration of the species represented above as the reaction occurs? a) the concentration is larger than that of reactant or product b) the concentration is very low for the duration of the reaction c) the concentration increases steadily as the reaction proceeds d) the concentration fluctuates between high and low during the reaction

b

2 HBr (g) + O2 (g) --> H2O2 (g) + Br2 (g) based on a kinetics study of the reaction represented by the equation above, the following mechanism for the reaction is proposed. step 1: HBr (g) + O2 (g) --> HO2Br (g) SLOW step 2: HO2Br (g) + HBr (g) --> 2 HOBr (g) FAST step 3: 2 HOBr (g) --> H2O2 (g) + Br2 (g) FAST which of the following rate laws is consistent with the proposed mechanism? a) rate= k[HBr]^2[O2] b) rate= k[HBr][O2] c) rate= k[HO2Br][HBr] d) rate= k[HOBr]^2

b

Br2 (g) + 2 NO (g) --> 2 NOBr (g) the reaction represented by the equation above has the following proposed mechanism step 1: NO (g) + Br2 (g) <--> NOBr2 (g) FAST EQUILIBRIUM step 2: NOBr2 (g) + NO (g) --> 2 NOBr (g) SLOW based on the information, which of the following is the initial rate law for the reaction? a) rate= k[Br2][NO] b) rate= k[Br2][NO]^2 c) rate= k[Br2]^2[NO] d) rate= k[NOBr2][NO]

b

N2 (g) + 3 H2 (g) --> 2 NH3 (g) the reaction above occurs on the surface of Ru (s). the rate of this reaction is determined by the amount of energy required to break the bond in N2. this bond is weakened when N2 is absorbed on Ru (s). based on this information, which of the following provides the best reason for the use of Ru (s) for the synthesis of NH3? a) it promotes the proper orientation of N and H atoms to form new N-H bonds b) it provides a reaction path with a lower activation energy c) it forms a bond between Ru and N that is stronger than the bond in N2 d) it decreases the frequency of collisions between N2 (g) and H2 (g)

b

NO (g) + NO3 (g) --> 2 NO2 (g) rate= k[NO][NO3] the reaction represented above occurs in a single step that involves the collision between a particle of NO and a particle of NO3. A scientist correctly calculates the rate of collisions between NO and NO3 that have sufficient energy to overcome the activation energy. The observed reaction rate is only a small fraction of the calculated collision rate. which of the following best explains the discrepancy? a) the energy of collisions between two reactant particles is frequently absorbed by collision with a third particle b) the two reactant particles must collide with a particular orientation in order to react c) the activation energy for a reaction is dependent on the concentration of the reactant particles d) the activation energy for a reaction is dependent on the temperature

b

NO (g) and O2 (g) react to form NO2 (g). the rate law of the reaction is rate= k[NO]^2[O2]. if the reaction occurs in a single elementary step that is a three-body molecular collision, then which of the following is the equation for the elementary step? a) N + O + O --> NO2 b) 2 NO + O2 --> 2 NO2 c) N + O2 --> NO2 d) N2 + 2 O2 --> 2 NO2

b

S2O8^2- (aq) + 3 I- (aq) --> 2 SO4^2- (aq) + I3- (aq) in aqueous solution, the reaction represented by the balanced equation shown above has the experimentally determined rate law: rate= k[S2O8^2-][I-] if the concentration of [S2O8^2-] is doubled while keeping [I-] constant, which of the following experimental results is predicted based on the rate law, and why? a) the rate of reaction will remain the same, because k will decrease by half b) the rate of reaction will double, because the rate is directly proportional to [S2O8^2-] c) the rate of reaction will increase by a factor of four, because 2 moles of SO4^2- are produced for each mole of S2O8^2- consumed d) the rate of reaction will increase by a factor of 4, because the reaction is second order overall

b

step 1: H2 (g) + ICl (g) --> HI (g) + HCl (g) SLOW step 2: HI (g) + ICl (g) --> HCl + I2 (g) FAST the reaction is carried out at constant temperature inside a rigid container. based on this mechanism, which of the following is the most likely reason for the different rates of step 1 and step 2? a) the only factor determining the rate of step 2 is the orientation of the HI and ICl polar molecules during a collision, but it has a negligible effect when H2 and ICl molecules collide b) the amount of energy required for a successful collision between H2 and ICl is greater than the amount of energy required for a successful collision between HI and ICl c) the fraction of molecules with enough energy to overcome the activation energy barrier is lower for HI and ICl than for H2 and ICl d) the frequency of collisions between H2 and ICl is greater than the frequency of collisions between HI and ICl

b

trial 1: [A2]= 0.10 [B]= 0.50 initial rate (Ms-1)= 2.5 x 10^-4 trial 2: [A2]= 0.2 [B]= 0.50 initial rate (Ms -1)= 5.0 x 10^-4 trial 3: [A2]= 0.30 [B]= 0.05 initial rate (Ms -1)= 5.0 x 10^-5 trial 4: [A2]= 0.30 [B]= 0.10 initial rate (Ms -1)= 1.0 x 10^-4 an experiment was conducted to determine the rate law for the reaction A2 (g) + B (g) --> A2B (g). the table above shows the data collected. based on the data in the table, which statement is correct? a) since the rate law can be expressed as rate= k[A2]2, tripling the concentration of A2 will cause a 9-fold increase in the rate of the reaction b) since the rate law can be expressed as rate= k[A2][B], doubling the concentrations of A2 and B will quadruple the rate of the reaction c) since the rate law can be expressed as rate= k[A2]2[B], tripling the concentration of A2 while keeping the concentration of B constant will triple the rate of the reaction d) since the rate law can be expressed as rate= k[A2][B]2, doubling the concentration of B while keeping the concentration of A2 constant will double the rate of the reaction

b

which of the following statements best explains why an increase in temperature of 5-10 degrees Celsius can substantially increase the rate of a chemical reaction? a) the activation energy for the reaction is lowered b) the number of effective collisions between reactant particles is increased c) the rate of the reverse reaction is increased d) delta H for reaction is lowered e) delta G for the reaction becomes more positive

b

when two molecules are involved

bimolecular

(FLIP FLASHCARD FOR MODEL) the particle models shown above represent a proposed two-step mechanism for the destruction of ozone (O3) in the upper atmosphere. based on the models, which of the following represents a species that acts as a catalyst for the reaction? a) O3 b) O2 c) Cl d) ClO

c

2 H2 (g) + 2 NO (g) --> N2 (g) + 2 H2O (g) for the chemical reaction represented above, the following mechanism is proposed step 1: 2 NO (g) <--> N2O2 (g) FAST EQUILIBRIUM step 2: N2O2 (g) + H2 (g) --> N2O (g) + H2O (g) SLOW step 3: N2O (g) + H2 (g) --> N2 (g) + H2O (g) FAST which of the following initial rate law expressions is consistent with this proposed mechanism? a) rate= k[H2]^2[NO]^2 b) rate= k[NO]^2 c) rate= k[H2][NO]^2 d) rate= k[H2][N2O]^2

c

2 HO2 (g) --> H2O2 (g) + O2 (g) the reaction represented by the chemical equation shown above occurs in Earth's atmosphere. in an experiment, [HO2] was monitored over time and the data was plotted in a time vs. 1/[HO2] graph (the graph was linear). based on the information, which of the following is the rate law expression for the reaction? a) rate= k b) rate= k[HO2] c) rate= k[HO2]2 d) rate= k[H2O2][O2]

c

H2O2 --> H2O + 1/2O2 hydrogen peroxide decomposes to produce water and oxygen according to the equation above. the decomposition of hydrogen peroxide is first-order. which of the following best identifies the rate constant k for the reaction based on the order? a) k= ln[H2O2] at t= 0 s b) k= ln[H2O2] at t= 500 s c) k= - (slope of plot) d) k= - 1/(slope of plot)

c

the gas-phase reaction A2 (g) + B2 (g) --> 2 AB (g) is assumed to occur in a single step. two experiments were done at the same temperature inside rigid containers. the initial partial pressures of A2 and B2 used in experiment 1 were twice the initial pressures used in experiment 2. which statement provides the best comparison of the initial rate of formation of AB in experiments 1 and 2? a) the initial rate of formation of AB is the same in both experiments because they were done at the same temperature and the frequency and energy of the collisions between A2 and B2 would have been about the same b) the initial rate of formation of AB is slower in experiment 1 than in experiment 2 because at the same temperature, a higher pressure would reduce the volume available for A2 and B2 molecules to achieve the proper orientation for a successful collision c) the initial rate of formation of AB is faster in experiment 1 than in experiment 2 because at a higher pressure the collisions between A2 and B2 molecules would have been more frequent, increasing the probability of a successful collision d) the initial rate of formation of AB is faster in experiment 1 than in experiment 2 because at a higher pressure a larger fraction of the A2 and B2 molecules would have the minimum energy required to overcome the activation energy barrier

c

identify catalyst(s) and intermediate(s) in the mechanism show below (overall balanced equation: O3 (g) + O (g) --> 2 O2 (g) elementary step 1: O3 (g) + Cl (g) --> ClO (g) + O2 (g) elementary step 2: ClO (g) + O (g) --> Cl (g) + O2 (g)

catalyst= Cl intermediate= ClO

present at beginning and end of an experiment (consumed and reformed); cancels out in final balanced equation

catalysts

(FLIP FLASHCARD FOR MODEL) the particle models shown above represent a proposed two-step mechanism for the destruction of ozone (O3) in the upper atmosphere. based on the proposed mechanism, what is the balanced chemical equation for the overall reaction? a) O3 (g) + Cl (g) --> O2 (g) + ClO (g) b) 2 O3 (g) + 2 Cl (g) --> 2 O2 (g) + 2 ClO (g) c) O3 (g) + Cl (g) + ClO (g) --> 2 O2 (g) + Cl2 (g) d) 2 O3 (g) --> 3 O2 (g)

d

(FLIP FLASHCARD FOR MODEL) the particle models shown above represent a proposed two-step mechanism for the destruction of ozone (O3) in the upper atmosphere. based on the proposed mechanism, which of the following is the rate-law expression for the destruction of O3? a) rate= k[O3]^2 b) rate= k[Cl]^2 c) rate= k[ClO][O3] d) rate= k[O3][Cl]

d

2 NO (g) + Cl2 (g) --> 2 NOCl (g) Experiment 1: [NO]= 0.0250 [Cl2]= 0.0510 rate of NOCl formation (M s-1)= 4.55 x 10^-5 Experiment 2: [NO]= 0.0250 [Cl2]= 0.1020 rate of NOCl formation (M s-1)= 9.10 x 10^-5 Experiment 3: [NO]= 0.0500 [Cl2]= 0.0510 rate of NOCl formation (M s-1)= 1.82 x 10^-4 The initial rates of the reaction represented by the equation shown above were measured for different initial concentrations of NO (g) and Cl2 (g). Based on the data given in the table above, which of the following is the rate law expression for the reaction, and why? a) rate= k[NO]2 because the initial rate quadrupled when [NO] was doubled but remained constant when [Cl2] was doubled b) rate= k[NO][Cl2] because the initial rate doubled when either [NO] or [Cl2] was doubled c) rate= k[NO][Cl2]2 because the initial rate doubled when [NO] was doubled and quadrupled when [Cl2] was doubled d) rate= k[NO]2[Cl2] because the initial rate quadrupled when [NO] was doubled and doubled when [Cl2] was doubled

d

step 1: H2 (g) + ICl (g) --> HI (g) + HCl (g) SLOW step 2: HI (g) + ICl (g) --> HCl (g) + I2 (g) FAST which of the following represents the overall chemical equation for the reaction and the rate law for elementary step 2? a) the overall reaction is H2 (g) + 2 ICl (g) --> HCl (g) + I2 (g); the rate law for step 2 is rate= k[H2][HI][ICl] b) the overall reaction is H2 (g) + 2 ICl (g) --> 2 HCl (g) + I2 (g); the rate law for step 2 is rate= k[H2][ICl]^2 c) the overall reaction is H2 (g) + ICl (g) --> HCl (g) + I2 (g); the rate law for step 2 is rate= k[H2][ICl] d) the overall reaction is H2 (g) + 2 ICl (g) --> 2 HCl (g) + I2 (g); the rate law for step 2 is rate= k[HI][ICl]

d

which of the following best helps explain why an increase in temperature increases the rate of reaction? a) at increased temperatures, reactions have decreased activation energy b) at increased temperatures, reactions have increased activation energy c) at increased temperatures, every collision leads to formation of product d) at increased temperatures, high-energy collisions happen more often

d

Maxwell-Boltzmann distribution curve

describes the distribution of particle energies; can be used to gain a qualitative estimate of the fraction of collisions with enough energy to lead to a reaction, and how it is temperature dependent

the following mechanisms are proposed for the gas phase decomposition of ozone, O3. a student claims that the rate of the first mechanism is faster than the rate of the second mechanism because the first mechanism has fewer steps. do you agree or disagree? justify your claim mechanism 1- step 1: O3 --> O2 + O step 2: O3 + O --> 2 O2 overall: 2 O3 --> 3 O2 mechanism 2- step 1: NO + O3 --> NO2 + O2 step 2: O3 --> O2 + O step 3: NO2 + O --> NO + O2 overall: 2 O3 --> 3 O2

disagree; mechanism 1 has an intermediate (O) and leads to the correct balanced equation. mechanism 2 has intermediates (NO2 & O) AND a catalyst (NO). catalysts increase the rate of a reaction, therefore mechanism 2 has a lower activation energy and is likely the faster reaction

D + 3 E --> 2 F when the chemical reaction is carried out under certain conditions, the rate of disappearance of D is 2.5 x 10^-2 Ms-1. what is the rate of disappearance of E and rate of appearance of F under these same conditions?

disappearance of E: 7.5 x 10^-2 Ms-1 appearance of F: 5.0 x 10^-2 Ms-1

a process in a chemical reaction that occurs in a single event or step

elementary reaction

energy is absorbed

endothermic

if Ep > Er is the reaction exothermic or endothermic?

endothermic

the rate of the forward reaction is equal to the rate of the reverse reaction

equilibrium

energy is released

exothermic

if Ep < Er is the reaction exothermic or endothermic?

exothermic

true or false: intermediates should be in rate law expressions

false

true or false: the value of the constant (k) is NOT dependent on temperature

false

true or false: we CAN use the stoichiometry coefficients from a balanced, overall equation to determine the orders of a reaction

false

true or false: the units for a reaction rate have to include seconds

false (any unit of time can be used)

true or false: changing the temperature changes the activation energy

false (catalysts are one of the only things that change activation energy)

true or false: if a reaction is first order with respect to a reactant being monitored, the graph of ln[A] vs. time will be EXPONENTIAL

false (it will be linear)

true or false: as activation energy increases, reaction rate also INCREASES

false (reaction rate DECREASES)

true or false: the rate of a reaction ISN'T influenced by reactant concentrations, temperature, surface area, catalysts, and other environmental factors

false (the rate of the reaction IS influenced by those factors)

true or false: elementary reactions involving the simultaneous collision of three or more particles are COMMON

false (very rare)

half-life is a critical parameter for _________ order reactions because the half-life is constant and related to the rate constant

first

the amount of time it takes for a radioactive sample to decay to 1/2 it's original amount

half-life

reaction is a DIFFERENT phase than the catalyst

heterogeneous catalysis

in a reaction energy profile does the slow step have a higher or lower activation energy than the fast step?

higher slow= high activation energy; fast= low activation energy

reaction is the SAME phase as the catalyst (ex: both aqueous)

homogenous catalysis

how can the reaction rate be increased?

increase temperature (more forceful collisions), increase concentration (more collisions), decrease activation energy, improve orientation of particle

substances that are neither reactants or products but are formed in one elementary reaction and consumed in the next (these aren't included in the overall chemical equation or rate law expression)

intermediates

rate constant

k

what is the effect on the k value and rate if temperature is increased, activation energy is decreased, and orientation is improved

k and rate both increase

how can the slope of a ln[A] vs. time graph for a first order reaction be used to find k?

k= - slope

how can the slope of a concentration vs. time graph for a zero order reaction be used to find k?

k= -slope

iodide-123 is used to study thyroid gland function. this radioactive isotope breaks down in a first order process with a half-life of 13.1 hours. what is the rate constant for this process?

k= 0.053 h-1

how can the slope of a 1/[A] vs. time graph for a second order reaction be used to find k?

k= slope

the study of the rates of a chemical reaction

kinetics

what is the equation for a first order reaction?

ln[A]t - ln[A]0 = -kt

what is the equation for determining the units on k?

m^1-n x t^-1 n= # of overall order zero order= m^1t^-1 first order= m^0t^-1 (t^-1) second order= m^-1t^-1

what does collision theory state?

molecules must collide, molecules must collide with sufficient energy, molecules must collide with correct orientation

reaction coordinate

monitors progress of reaction over time in a reaction energy profile

what happens to the activation energy when temperature is increased?

no change

how much the concentration impacts the rate

order (zero order- concentration doesn't matter; first order- when concentration increases 4x the rate also increases 4x; second order- when concentration increases 4x the rate increases 16x)

the sum of the powers of the reactant concentrations in the rate law is the __________ order of the reaction

overall

determine the overall balanced equation, identify catalyst(s), and identify intermediate(s) in the mechanism shown below step 1: CH3COCH3 + H3O+ <--> CH3COHCH3+ + H2O step 2: CH3COHCH3+ + H2O <--> CH3COHCH2 + H3O+ step 3: CH3COHCH2 + Br2 --> CH3COCH2Br + HBr

overall balanced equation- CH3COCH3 + Br2 --> CH3COCH2Br + HBr intermediates- CH3COHCH3+, H2O, & CH3COHCH2 catalyst- H3O+

consists of one or more elementary reactions or steps

overall chemical reaction

determine the rate law expression predicted from the following proposed mechanism step 1: O3 (g) <--> O2 (g) + O (g) FAST step 2: O3 (g) + O --> 2 O2 SLOW overall: 2 O3 (g) --> 3 O2 (g)

rate= k' ([O3]^2/[O2])

the rate law expression reaction of H2O2 and I- in an acidic solutions is found to be first order with respect to H2O2 and first order with respect to I-. does the proposed mechanism below support the rate law expression? step 1: H2O2 (aq) + I- (aq) --> H2O (l) + OI- (aq) SLOW step 2: H+ (aq) + OI- (aq) --> HOI (aq) FAST step 3: HOI (aq) + H+ (aq) + I- (aq) --> I2 (aq) + H2O (l) FAST overall: H2O2 (aq) + 2 I- (aq) + 2 H+ (aq) --> I2 (aq) + 2 H2O (l)

rate= k[H2O2][I-] yes, the slow step is first order in H2O2 and I-

what is the rate law equation?

rate=k[A]^m[B]^n [A] and [B] are concentrations of reactants m and n are reaction orders with respect to each reactant

the process by which a reaction occurs

reaction mechanism

when three molecules are involved

termolecular

how does INCREASING the aqueous concentration impact collisions and reaction rate?

the number of collisions and reaction rate also increase

how does INCREASING the surface area impact collisions and reaction rate?

the number of collisions and reaction rate also increase (as surface area increases, reactant molecules collide more frequently, leading to increased rates)

how does INCREASING the temperature impact collisions and reaction rate?

the number of collisions and reaction rate also increase (atoms move faster, colliding with more energy and increased frequency)

how does DECREASING the volume of gases impact collisions and reaction rate?

the number of collisions and reaction rate increase (reactants are closer together)

A --> B + 2 C what is true about the rate of disappearance of B and C relative to the rate of appearance of A?

the rate of disappearance of A is equal to the rate of appearance of B, and 1/2 the rate of appearance of C rate of disappearance of A= -[A]/time rate of appearance of B= [B]/time rate of appearance of C= [C]/time

true or false: a catalyst increases reaction rate by providing an alternate pathway with lower potential energy for the activated complex (catalyst decreases activation energy)

true

true or false: activated complex is the transition state from reactant to product (bonds are partially broken and formed- committed to forming a product)

true

true or false: catalysts are one of the only things that change the activation energy

true

true or false: if a reaction is second order with respect to a reactant being monitored, the graph of 1/[A] vs. time will be LINEAR

true

true or false: if a reaction is zero order with respect to a reactant being monitored, the graph of [A] vs. time will be LINEAR

true

true or false: in most reactions, only a small fraction of collisions lead to a reaction

true

true or false: it is common for catalysts to show up in a rate law expression

true

true or false: the chemical equations for the elementary reactions in a multistep mechanism must always add to give the chemical equation for the overall process

true

true or false: the kinetics of a chemical reaction can be defined as the rate of change of a reactant or product concentrations per unit of time

true

true or false: the units of k reflect the OVERALL reaction order

true

true or false: we CAN use the stoichiometry coefficients from an elementary reaction to determine the orders of a reaction

true

true or false: the rate of a reaction can be no faster than the slowest step

true (the slowest step is the rate determining step)

when only 1 molecule is involved

unimolecular

what is the enthalpy equation?

∆Hrxn = Hproducts - Hreactants