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Flashcards in Chemistry 12 Kinetics Deck (38):

Define a Chemical reaction

A Chemical Reaction: process in which chemical species react to form NEW SUBSTANCES
-different rates of chemical reactions


Define Reaction Rate

Reaction rate: change in concentration of a reactant or product of a chemical reaction per unit of time

Decrease in concentration of reactants per unit time
Increase in concentration of products per unit time

Reaction rate = reactant consumed/time interval (show a decrease over time)
Reaction rate = product formed/time interval (show an increase over time)


How to determine rates of reactions?

by monitoring the change in concentration of either the REACTANTS or PRODUCTS as a function of time


Average Reaction Rate

change in concentration of a reactant or product over a given time interval

-(measuring the concentration of a reactant or product at the beginning and the end of a time period)

R = change[C]/changet

*could measure other properties i.e. mass, colour, temperature, solids used up, gas produced, pH, conductivity (any property that differs between the reactants and products)

*change = y2 - y1/x2-x1


Calculating Average Reaction Rate

-given a table with time

If not given initial time, assume t1 = 0
If not given final concentration (reactant consumed), assumed c2 = 0


What does a negative rate mean?

negative sign obtained in a rate calculation means that the concentration is decreasing; reactant are being consumed

- reactant consumption
+ product formation


What happens to the average rate of reaction as the reaction proceeds (i.e. over time)?

the average rate decreases because as the reaction continues, there are FEWER COLLISIONS between reactant molecules due to a decrease in concentration

-when comparing averages, do not look at negative sign (absolute value)
-also, don't let 10 to the power scare you:
6.49x10^-7 is greater than 3.43x10^-7

*compare using same power


Reaction Rate Stoichiometry

rate of reactant consumption related to rate of product formation
-chemical reaction

**if see production and computation rates together = stoichiometry

units mol/L/s (comparing rates)


Reaction Rate Stoichiometry in Graphs

mole ratio in the chemical reaction is represented graphically (area under/above the curves)

Measured property vs. time

Decreasing over time = reactant (look at area above)
Increasing over time = product (look at area under)

If areas same, then same ratio
If one area is doubled another area, then ratio is 2:1


Instantaneous Rate of reaction

the rate of a chemical reaction at a single point in time (t = x)
-determined by finding the slope of a line tangent to the curve on a concentration-time graph

Instantaneous Rate can be found at several different points along the curve and shows how the reaction rate changes with time

-align ruler along the point
-draw tangent
-choose points (more convenient) on the line (typically broader range, equal both sides of point)
-calculate slope = rate

Slope = y2-y1/x2-x1 (rise over run)


Initial Rate of reaction

the instantaneous rate at time zero (t = 0), the beginning of the reaction
*initial rates greater than rates later on in the reaction due to high concentration = more collisions


Collision Theory

Chemical reactions involve collisions of reactant particles

Reaction = EFFECTIVE collisions (not all collisions lead to chemical reaction)

Successful collisions Requirements:
1. particles must collide with correct orientation/geometry
2. particles must collide with sufficient KINETIC energy (must meet minimum activation energy requirement)
-to break react bonds and form product bonds

Rate of reaction depends on the frequency of collisions and fractions of those collisions that are effective


Factors Affecting Reaction Rate (5)

1. Nature of the Reactants
2. Surface Area
3. Concentration
4. Temperature
5. Catalyst


Factor: Nature of Reactants

i) amount of bond-making/bond breaking
-reactions involving a lot of bond breaking (reactants) and new bond forming (products) are slow
-typically larger molecules have more bonds to break
Ex: if a molecule has a lot of bonds or strong bonds, then the reaction will be slower

ii) Physical States of reactants
-Solid + gas = SLOW ex: Cu + O2
-Liquid + gas = SLOW
-Gas + Gas = FAST (high speed of molecules moving)
-aq + aq = FAST (no bond breaking, only ions forming together)

*Solid and Liquid have slower molecules


Factor: Surface Area

more surface area exposed to or in contact with the reactants, the faster the reaction will take place -->Increases the number of collisions per second -->increases rate

-small particles, large surface area
-large partciles, small surface area

*only applies to heterogeneous reactions (ones involving DIFFERENT STATES)


Factor: Concentration

Increase in concentration--> an increase in reactant molecules in a given volume --> more effective collisions may occur --> increases rate

*increasing pressure in a gas


Rate Laws

mathematical measure of the average rate of a reaction and can only be determined from a series of EXPERIMENTS

R = k[A]^[B]^
where k is a rate constant (rate: concentration) determined using ANY experimental trial

-allows calculation of reaction rate as a function of reactant concentration (effect of a change in concentration of the reactants on the reaction rate)

-given experimental data with multiple trials -->([reactant 1], [reactant 2], [reactant 3]...), (reaction rate mol/L/min)
-in each trial, changing one reactant concentration, seeing effect on overall reaction rate
-if changing concentration (while other reactant concentrations constant) causes reaction rate to be directly proportional (x2, x2), then exponent is 1
-If rate change appears to be directly proportional to reactant change squared (x2, x4 or x3, x9), then exponent 2
-if changing concentration has no effect on reaction rate, then exponent 0 (do not include in rate law)


Tips when calculating rate laws and constant (k)

to solve for k, rearrange rate law so k is isolated
-choose any trial for values
-don't forget to apply exponents if in rate law
ex: R = k[A]^2
then k = R/([A]^2)
-also solve for units (M/min divide by M...)
-show all work for communication marks


Reaction Mechanisms

For reactions involving more than 3 collisions, most reactions occur through a series of steps called a reaction mechanism

A Reaction Mechanism is a sequence of simple 2 or 3 particle reactions by which a complex reaction occurs

-essentially breaking up the net reaction into steps with smaller particle collisions


Characteristics of a Reaction Mechanism

1. Each step in the mechanism is usually BIMOLECULAR (2 particle) and never more than termolecular (3 particle)

2. The sum of the steps in the reaction mechanism must give the balanced equation for the chemical reaction (intermediates are consumed and does not appear in net reaction)

3. the slowest step in a series of steps determines the rate of the overall reactions. This step is known as the rate determining step or R.D.S.

4. To increase the rate of the reaction, speed up the slowest step (any of the reactants in the slow step)

5. If the reactant is not in the rate determining step (slow step), then changing its concentration will have no effect on the reaction rate

*Rate law = slow step
1) only reactants from the rate determining step may appear in the Rate Law
2) the coefficient from the Rate Determining Step (slow step) becomes the power in the rate law
Ex: HBr + O2 --> HOOBr (Slow)
R = k [HBr][O2]


Reaction Intermediate and Catalysts

Intermediate: neither a reactant nor a product; FORMED then CONSUMED (not part of the net chemical reaction)
-every step needs an intermediate (product side) except for the final step
-first appear on the PRODUCT side
*as long as used in a subsequent step

Catalyst: a substance that enters a reaction and can be recovered at the end of a reaction, unchanged
-not part of the net reaction
-opposite of intermediate
-"Enters on REACTANT side", "Exists on product side"


Reaction Order

The sum of the exponents in the rate law (overall reaction order)

Can be for the overall reaction or for each reactant

Exponents in the rate law represent the number of molecules reacting in the Rate determining Step. The rate law rule states that the sum of the exponents in the rate law can never exceed 3, because it is highly unlikely/never have more than a 3 particle collision.

Example of overall reaction order:
1st Order Reaction: R = k[A]
2nd Order Reaction: R = k[A][B] or [A]^2
3rd Order Reaction: R = k[A][B][C] (not to the power of three)

Example of Order of each reactant:
Zero order [A]^0
1st order [A]
2nd order [A]^2

*Differentiate between reactant and reaction order


What does Reactant Order look like on a graph? (Rate vs. Concentration Graph)

Zero Order: horizontal line --> no change in rate when concentration changes

1st Order: linear line -->directly proportionate

2nd Order: Exponential -->square


To consider when designing Mechanisms

-no more than 3 molecules reacting
-rate law is the slow step (label)
-form an intermediate in each step, except final step
-make sure intermediates cancel
-all steps add up to net reaction


Given experimental data with several trials, concentration of reactants, and rate of reaction...

For a multi-step reaction...

1. Write the rate law R = k[A]^[B]^...
-change one reactant's concentration, keep other concentration constant, see effect on overall rate
0 no change
1 directly proportionate
2 square

2. Solve for k
-rearrange rate law, use info from trial
-correct units

3. Afterwards, one can find information given the rate law and constant (no need for experimental data)
a) given mols of reactants find rate of reaction
-convert mols to concentration using C = n/v
-sub concentration of each reactant and value of k into rate law, solve for R

b) change volume of container
-convert mols to concentration using C = n/v (new volume)
-sub concentration of each reactant and value of k into rate law, solve for R
Ex: from 5 L to 1 L, means that rate is 5x faster

c) calculate how much the rate has changed
New rate/old rate
sub in the change (i.e. divide by 5 to get to new volume) into rate law and solve
Ex: R = k[A][B]^2
=125 x faster
*if fraction, means slower by denominator

d) find what happens to rate of reaction according to situations...
-reactant quadrupled (sub 4 in place of reactant)
-reactant tripled (sub 3 in place of reactant)
-reactant is quartered (sub 1/4 in place of reactant)
-container is halved (means concentration increased, sub 2 into all reactants, multiplication). If volume is increasing, then fraction (because lower concentration and rate)

-volume decreases by ____ (divide), then concentration increases by ____ (multiple)
-same temperature means constant k


Factor: Temperature

temperature increase always increases the rate of a chemical reaction

-increase in temperature causes molecules to move faster with higher kinetic energy--> 1) more successful collisions between reacting particles (increase number of collisions) 2) collisions that are more forceful/violent (more likely to collide with proper orientation)
-increase temperature -->increases the number of particles with the necessary activation energy (Ea) -->increases rate


Define Activation Energy (Ea)

Ea is the minimum energy required for a successful collision, assuming optimum collision geometry

-the more particles that have energy greater or equal to Ea, the faster the reaction
-Ea differs for each reaction


Heat Content (H) and Heat of Reaction (changeH)

Heat content (H): amount of (potential) energy contained within a substance and is measured in KJ

Heat of Reaction (changeH) -->enthalphy= Hproducts - Hreactants
-->net energy added or lost in a chemical reaction


Determining endothermic and exothermic reactions

-heat enters on reactant side
-Hp > Hr
-reactants more stable
-positive heat of reaction
-feels cold

-heat exits on product side
-products more stable
-negative heat of reaction


Graph: Kinetic Energy (Maxwell-Boltzmann) Distribution Curve:


Y axis = # of particles (%)
X axis = Kinetic energy (left to right: low, medium, high)

Bell shape distribution
-Ea typically right hand side (only small number of reacting molecules under curve to the right of the Ea vertical line) -->stable (not readily react)
-If Ea line is more towards left, then reaction is not as stable

How to graph:
-bell shape distribution for T1 low
-for increase in temperature, Ea DOES not move (activation energy does NOT CHANGE). The graph shifts to the right because more molecules have the sufficient activation energy to react


Factor: Catalyst

A catalyst is a substance that increases the reaction rate without being consumed in the chemical reaction
-does not appear in net reaction

Catalysts lower the Ea so that a greater percentage of molecules will have sufficient energy to react (E greater or equal to Ea)

on a Kinetic Energy Distribution Curve with a catalyst:
-the Ea line shifts left, causing more reacting molecules

Catalyst lowers Ea for both forward and reverse reactions

• Provide alternative pathway for a reaction, which has a lower activation energy, end up with same products→larger fraction of reactant will have kinetic energy equal to or greater than the lower activation energy
• Catalysts do not affect the energy difference between products and reactants (enthalpy change remains the same between regular reaction and catalyzed reaction)


Explaining catalysts (how do reactions take place)

1. reactant particles approach each other
-slow down as particles approach because of electron repulsive forces.
-Kinetic energy (motion) is decreasing and Potential energy is increasing

2. Collide
-Forms activated complex (transition state)
-if the molecules have enough activation energy, they will react on collision

3. Once product molecules form they will move apart from one another because of repulsive forces. As they separate, kinetic energy will increase and potential energy will decrease


Define Activated Complex

An unstable cluster of molecules (high potential energy) formed when reactant particles collide
-very top of a Potential energy diagram


Potential energy diagram

Y axis: potential energy (KJ/mol)
X axis: Reaction Coordinate

R --> P

If Products energy is lower than reactant energy, then exothermic reaction

Ea(forward): the energy required by the reactants to react.
-->difference between activated complex and reactant energy

Ea(reverse): the energy that is lost by the products as they separate -->difference between activated complex and product energy

Net energy: Hp-Hr (INCLUDE SIGN) -->doesn't change for catalyst
+ endothermic
- exothermic


Reversing chemical reactions

-many are reversible
-products can combine to produce reactants
-reaction pathway remains the same but for the reverse reaction the Ea is different due to the direction of the pathway

-Endothermic becomes exothermic

Net energy for the reverse reaction:
-retrace path
-Products of original become reactants of reverse
-Net energy = Hp - Hr
=Old reactants - old products


How to draw potential energy diagrams

Given enthalpy change, Ea(f) or Ea(r).
Asked to determine Ea(f) or Ea(r) and show effect of catalyst if given Ea(f) cat or Ea(r).

1. locate reactants, products, and activated complex
-look at enthalpy change to determine if endothermic or exothermic
2. complete curve
3. label (find Ea(f), Ea(r), catalyst)


What is the rate law for a one step reaction?

one step reaction = R.D.S

Therefore rate law = k [reactant 1]^coefficent...


What is the rate law rule?

The sum of the exponents in the rate law can never exceed 3.

-->exponents in the rate law represent the number of molecules reacting in the slow step.
-->can never have more than 3 particle collision (highly unlikely)