Unit 13
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Unit 13. Reaction Rates. Unit 13 Goals. Understand collision theory Understand how concentration, surface area & temperature, effect reaction rate. Understand potential energy diagrams Predict reaction rate change when concentration, temperature or surface area ∆’s

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Unit 13

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Unit 13

Unit 13

Reaction Rates


Unit 13 goals

Unit 13 Goals

  • Understand collision theory

  • Understand how concentration, surface area & temperature, effect reaction rate.

  • Understand potential energy diagrams

  • Predict reaction rate change when concentration, temperature or surface area ∆’s

  • Construct and interpret PE diagrams


Thinker

Thinker:

  • Make a list of RATES that you have heard about in your daily lives.

    • Rate of growth

    • Rate of pay

    • Rate of decay

    • Rate of aging

    • Rate of flow

    • Rate of power usage

    • Rate of speed


Rates of chemical change

Rates of Chemical Change

  • Rate is how fast something changes with time

    • Sprinters

      • Human 28.9 mph

      • Cheetah 62 mph

      • Quarter horse 55 mph

  • Reaction Rate – rate at which reactants are changed to products

  • Chemical Kinetics – study of Rx rates

http://en.wikipedia.org/wiki/Footspeed


Chemical kinetics

Karl Wenzel

German Scientist 1777

First to study Rx rates

Claude-Louis Berthollet

French – 1803

Published first book on Rx rates & concentration

Chemical Kinetics

http://en.wikipedia.org/wiki/Claude-Louis_Berthollet


Rate change over time

Rate = Change over time

  • CH3OCH3(g) CH4(g) + CO(g) + H2(g)

    • At 500 C, dimethyl ether decomposes

    • As Rx continues:

      • Conc. of dimethyl ether will decrease

      • Time will continue to change

  • Therefore:

  • Note: all reactants are -


Ch 3 och 3 g ch 4 g co g h 2 g

CH3OCH3(g) CH4(g) + CO(g) + H2(g)

  • We can talk about the rates of reactants OR products:

  • Note: all products are +


What about coefficients

What about coefficients?

  • Consider this Rx:

    • 2N2O5(s) 4NO2(g) + O2(g)

    • How do coefficients affect reactant/product?

    • Divide by each coefficient:

    • Think about how many per time unit


Measuring reactions

Measuring Reactions

  • Need to keep track of at least 1 concentration:

    • Color change

    • Pressure change

    • Temperature change

  • Rates is measured as is indicated by the rate you write.


Calculating rx rates

Calculate the rate of reaction between 0 and 20.0 s:

0.0000585 M/s

Calculate the rate of reaction between 20.0 and 40.0 s

0.0000528 M/s

Calculating Rx Rates:


Will reaction rate always be the same

Will Reaction Rate always be the same?

  • 2N2O5(s) 4NO2(g) + O2(g)

  • No!

  • Reactant is constantly decreasing

  • Product is constantly increasing

    slope of graph = rate of Rx


Factors affecting rate

Factors Affecting Rate

  • Concentration

  • Pressure

  • Temperature

  • Surface Area


Concentration

Concentration

  • Almost always:

    • Increase in concentration of reactants yields in increase of reaction rate

      NO2(g) + CO(g) NO(g) + CO2(g)

  • If [NO2] is doubled

    • More collisions with CO will occur

    • More collisions means more successful collisions

    • More successful collisions means more product

  • The inverse is also true


Non collision reactions

Non-collision reactions

  • Decomposition does not require collisions

    (CH2)3(g) CH2=CH-CH3(g)

  • The more cyclopropane you have:

    • The more is available to decompose

    • Thus a faster reaction rate


Pressure

Pressure

  • In liquid or solid, has almost no affect

  • In gasses:

    • Doubling pressure essentially doubles the concentration.

    •  by increasing pressure, we increase reaction rates in gasses.

    • Inverse is true


Temperature

Temperature

  • In General:

    • Increase in temperature will increase reaction rate.

    • Often increase of 10% temp yields in 2 times the reaction rate!

    • Our bodies work best around 37 C

      • Small changes cause great distress!

      • Snakes, lizards & other reptiles


Surface area

Surface Area

  • Increased surface area yields an increase in reaction rate.

    • Wood & fires

    • Grain Elevator


Questions

Questions?

  • HW: Read 16.1

  • P. 585 Section Review:

    • 1-3, 6, 8, 11 – 13


Rate law

Rate Law

  • Describes how reactant concentration affect the reaction rate.

  • Provides the ability to form a reaction mechanism.

    • Model of how the reaction occurs

    • Discusses # of steps necessary


Determining a general rate law equation

Determining a General Rate Law Equation

  • When there is a single reactant:

    • Rate of reaction is proportional to the concentration raised to some power (n)

    • The power is the order of the Rx.

      Rate = k[reactant]n

    • k is the rate constant

      • Proportionality that varies with temperature

    • Order is usually a whole number (1, 2)

      • Can be fractions

    • Occasionally is 0

      • Reaction rate is independent of concentration


Determining order example

Determining Order (example)

  • Experiments were preformed to measure the initial rate of the reaction 2HI(g)  H2(g) + I2(g). Conditions were identical except that the HI concentrations were varied.

  • rate = k[HI]n

  • n=?

  • Find ratio of the reactant concentrations b/t experiments.

  • Find ratio of the reaction rates.

  • Plug into rate = k[HI]n

  • Verify the Results (compare 3 and 1)


Other questions

Other questions

  • Given the previous reaction:

    • What will happen if [HI] is increased from 4 M to 8 M?

    • What will happen if [HI] is decreased from 6 M to 2 M?

    • What will happen if [HI] is quadrupled?


Questions1

Questions?

  • Homework:

    • Worksheet


Rate laws for several reactants

Rate Laws for Several Reactants

  • When there are several reactants the rate law applies to each.

  •  an order applies to each reactant

    NO(g) + O3(g) NO2(g) + O2(g)

  • rate = k[NO3]n1[O3]n2

  • It turns out:

    • n1 = n2 = 1

  • Since each order is equal, it suggests that this reaction has a 1 step mechanism.

  • If they weren’t 1, then the mechanism could be longer than 1 step.


Rate determining step

Rate Determining Step

  • Despite being written as 1 step, many Rx, aren’t quite that simple.

  • The mechanism is more complex.

  • Example:

    2Br-(aq) + H2O2(aq) + 2H3O+(aq) Br2(aq) + 4H2O(l)

  • Has 4 separate steps (each reactants order 1)

    • Br-(aq) + H3O+(aq) HBr(aq) + H2O(l)

    • HBr(aq) + H2O2(aq) HOBr(aq) + H2O(l)

    • Br-(aq) + HOBr(aq)  Br2(aq) + OH-(aq)

    • OH-(aq) + H3O+(aq)  2H2O(l)

  • Each step produces intermediates.

    • Species that are produced, but consumed

  • If one of the steps is slower, it is known as the rate determining step

    • In this case it is step 2


Collision theory

Collision Theory:

  • Reactions occur because:

    • Substances (atoms, ions, molecules) must collide

      • As particles get closer, they repel, so they must have enough KE to actually collide

    • Collisions must occur in correct orientation

    • Collisions must have sufficient energy to form the activated complex

      • This energy is known as activation energy (Ea)


Potential energy pe diagram

A + B  C + D

Imagine trying to bowl uphill

This is what molecules need to do

Activation energyis required to formactivated complexat top of curve.

# peaks = # of steps

Potential Energy (PE) Diagram

http://commons.wikimedia.org/wiki/Image:Coordenada_reaccion.GIF


Exothermic vs endothermic

Exothermic vs. Endothermic


Collision theory1

Collision Theory

  • Reacting substances (atoms, ions, or molecules) must collide

  • Collisions must occur with the correct orientation

  • Collisions must have sufficient energy to form the activated complex.

  • Activated complex is a temporary unstable arrangement of atoms that may form products or break apart to re-form the reactants


Unsuccessful collision not enough energy

Unsuccessful Collision (not enough energy)


Unsuccessful collision wrong orientation

Unsuccessful Collision (wrong orientation)


Successful collision

Successful Collision


Catalysts potential energy

Catalyst – chemical in a reaction that lowers the activation energy, but never is used.

Save energy

Usually costly

Biological catalysts are called enzymes

Catalysts & Potential Energy

http://en.wikipedia.org/wiki/Image:Activation_energy.svg


Homework

Homework:

  • PE Diagram Worksheet


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