Chemical Kinetics and Thermodynamics
Chemical Kinetics- concerned with:
1. Rates of Chemical Reactions# of moles of reactant used up or product formed
Unit time
Or
2. Reaction Mechanisms-
Rate of Reaction and the Collision Theory
Reactions take place at different rates
Collision Theory
· Used to explain why reactions take place at different rates
·
Effective Collision
Factors that Affect the Rate of a Chemical Reaction
1.
· What bonds are being broken and formed?
Fast = slight rearrangement of electrons
AgNO3 + NaCl à AgCl ppct. + NaNO3 ---FAST!
In reactions in which ionic bonds are broken and formed occur quickly at room
temperature
Slow= many covalent bonds broken
2H2O2 à 2H2O + O2 ---Slow!
Reactions that include covalent bonds being broken and formed occur very slowly at
room temperature
2.
· An increase in concentration of any one of the reactants usually, but not always
increases the rate of the reaction if the reaction is homogeneous
i. Homogeneous reactions- all the reactants are in the same phase.
ii. Heterogeneous reactions- reactants are in more than one phase.
I.e.- rusting
· How can we increase the concentration of reactants in the following:
i. Gasesii. Liquids-
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Collision Theory-
Heterogeneous systems- solid and liquid- increase……
Surface Area
Grinding, pulverizing, smashing, stirring, pureeing, blending, etc, etc, etc
Reactions that incorporate two gases
Partial pressure increase in one of the gases results in an increase in reaction rate:
· As a result of decreased volume or
· Increase in the number of molecules of gas
3. Temperature
For many reactions- 10oC increase in temperature = 2X rate
Collision TheoryTemperature increase =
·
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Energy Distribution Curve
4.
· Catalyst·
·
Pt metal
Enzymes
2H2O2 à 2H2O + O2 ---Slow!
2H2O2 à 2H2O + O2 ---Fast of MnO2 present as catalyst
5.
· Only for systems in which the reactants are in more than one phase
Solid and a liquid
Reaction Mechanisms
2C2H2(g) + 5O2(g) à 4CO2(g) + 2H2O(l)
What is the likelihood that 2 acetylene molecules will come into contact with 5 oxygen
molecules with the correct orientation and energy?
Most reactions proceed through a series of simple steps
· Each step- collision of two particles
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Reaction Mechanism·
Consider:
2A + B à A2B
It is unlikely that this reaction occurs in one step. One proposed mechanism may be-
OR
Often, chemists don‛t know the mechanism. They may only know the end result.
Why might it be difficult?
It is difficult to determine the reaction mechanisms because the intermediate
products have short lives
Intermediate products have structures unlike the structure of either the products or
the reactants
·
Activated Complex:
·
Reaction Mechanisms and Rates of Reactions
Rate determining stepWe stated earlier:
An increase in concentration not always increases the rate of a homogeneous reaction
WHY?
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Potential Energy Diagrams
·
A pictorial way to describe the energy involved in a reaction
Activation Energy and Temperature, Concentration, and Catalysts
Activated Complex:
· Short lived, unstable particle that will temporarily exist when reacting molecules collide
at the proper angle with the proper amount of energy
Rate Laws
·
A rate law is an equation that can be used to calculate the reaction rate for any given
concentration of reactants.
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·
The rate law can be determined by keeping the concentrations of all but one reactant
constant while measuring the reaction rate for various concentrations of that reactant.
The process is repeated for each reactant.
·
·
[A] and [B] represent the molar concentrations of reactants A and B in moles / liter (M)
exponents x and y are the powers of the concentrations of the reactants (determined
experimentally)
k = proportionality constant, rate constant, has a fixed value for a reaction at a
particular temperature (determined experimentally)
The actual form of the rate law varies from reaction to reaction
·
·
NO2(g) + O3(g) à NO3(g) + O2(g)
·
·
·
Rate = k[NO2][O3]
In this reaction, the rate is directly proportional to the concentration of both NO2 and
O3
Suppose the initial concentration of each reactant is 1.00M. If the concentration of
either reactant is doubled to 2.00M, the rate increases by a factor of 2.
Also, if the concentration of either reactant is multiplied by 4 to 4.00M, the rate is
increased by a factor of 4.
2 NO(g) + O2(g) à 2NO2(g)
·
·
·
·
Rate = k[NO]2[O2]
Notice the squaring of the NO concentration. If the concentration of NO is multiplied
by 6, the rate increases by a factor of 36 (6)2.
If the concentration of O2 increases by a factor of 6, the rate only increases by a
factor of 6.
Question: By how much would the rate increase if the concentration of both reactants
is multiplied by 2?
Important: the exponent in the rate law has no relationship to the coefficient in the
balanced chemical equation. The coefficients are determined experimentally!
2N2O5 (g) à 4 NO2 (g) + O2(g)
Rate = k[N2O5]
·
Important: not necessarily all reactants appear in the rate law. If changing the
concentration of a particular reactant does not change the rate, that reactant does not
appear in the rate law.
NO2(g) + CO(g) à NO (g) + CO2(g)
·
Rate = k[NO2]2
Because changes in the concentration of carbon monoxide do not affect the reaction
rate, [CO] does not appear in the rate law for this reaction.
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Questions:
1.
The rate law for the reaction in which nitrogen monoxide reacts with oxygen to produce
nitrogen dioxide was shown to be
Rate = k[NO]2[O2]
Suppose you measure the rate of the reaction with the concentration of each reactant at
1.00M. What will happen to the reaction rate if the concentration of NO is doubled? What
will happen to the reaction rate if the concentration of O2 is doubled instead?
2. The chemical equation and the rate law for the decomposition of hydrogen iodide are
shown. What will the effect on the reaction rate if the concentration of HI is raised
from 1.00M to 4.00M?
2HI (g) à H2(g) + I2 (g)
Rate = k[HI]2
Thermodynamics
·
Enthalpy =
·
1st Law of Thermodynamics – heat energy of a system is constant as long as no energy
enters or leaves the system
·
In all chemical reactions, there is a change in enthalpy = heat of reaction
Heat of Formation
·
When one mole of a compound is formed from its elements
Dependent on:
· Temperature of reaction
· Pressure of the reaction
· Phase of the product
Standard Heat of Formation - DHf
Heat of formation at 25oC and 760 mmHg
· Phases of the reactant and the products must be stated
H2(g) + 1/2O2(g) à H2O(l) + 286 kJ
·
When one mole of hydrogen gas reacts with ½ mole of oxygen gas, liquid water is formed
and heat is given off.
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Exothermic reaction
Energy is regarded as a product
H2O(l) + 286 kJ à H2(g) + ½ O2(g)
Heat must be absorbed by the water in order for one mole of liquid water decomposes
into its elements
___________________ reaction
Energy is regarded as a ______________
**Notice that the standard heat of formation is written so that one mole of product is
formed
·
Stability of Compounds
Unstable compounds- tend to break down into simpler substances or elements
Which of the following compounds is most stable?
Compound
Carbon Monoxide
Nitric Oxide
·
Standard Heat of Formation DHf
-110 kJ/mol
+90.4 kJ/mol
Compounds that give off large amounts of heat during their formation are considered
stable
Compounds that give off low amounts of heat during their formation or require heat to be
formed are unstable
· Require little or no net input of energy to cause them to decompose
Explosives!!!
Hess‛ Law of Constant Heat Summation
When a reaction can be expressed as the algebraic sum of two or more other reactions,
then the heat of the reaction is the algebraic sum of the heats of these other reactions
CuO(s) + H2(g) à Cu(s) + H2O(g) DH = ?kJ
Express as the sum of a series of other reactions
CuO(s) à Cu(s) + ½ O2(g) DH=155kJ
H2(g) + ½ O2(g) à H2O(g) DH = -242kJ
CuO(s) + H2(g) à Cu(s) + H2O(g) DH=-87kJ
Heat of Formation =
Spontaneous Chemical Reactions
·
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Consider:
A waterfall
Skiing
Cycling
·
·
The bond strength of the products is greater than the bond strength of the
reactants—heat is give off in the reactions
Energetically favorable, should occur spontaneously
But…
H2O(l) à H2O(s) +6.03 kJ
This reaction should be spontaneous, however, we know that this occurs only at 0oC
Entropy
·
·
A system has a large entropy if its is in a great state of disorder
DS = Sf -Si
·
Both the final and the initial entropies will be positive numbers, but the change between
final and initial can be negative
+DS = ________________ in disorder
-DS = ________________ in disorder
Substances in solid phase have fixed particle = low entropy
·
As a solid changes to a liquid, entropy increases- particles that make up the liquid have
greater freedom of movement
·
Particles making up a gas have a great deal of random motion- highest entropy of all the
phases
·
·
What changes in entropy occur when substances react?
Formation of a compound- ________________________________ The bonding
of elements to form a compound creates a more orderly state for the atoms
involved.
Compound decomposition- order and organization breaks down_______________________
What effect does temperature have on Entropy?
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Entropy and Spontaneous Change
·
·
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An increase in entropy (+DS) favors spontaneous reactions
Nature prefers a change from a more orderly system to a less orderly system
When a system is observed to have a special order, you can assume that there is
some reason for it. Some restraint on the system must prevent it from assuming a
more random arrangement.
In Summary:
Two factors that determine if a reaction will occur spontaneously are
1.
2.
Gibbs Free Energy Equation
·
·
·
·
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DH =
T=
DS =
DG =
Spontaneous reactions DG –
*Notice – temperature is a factor in this equation
Pb(s) + ½ O2(g) à PbO(s)
PbO(s) DHf = -215kJ
DSf = -0.092 kJ/K
Calculate DG at room temperature
But… What about when T = 3000K?
½ N2(g) + ½ O2(g) à NO(g)
DHf = + 90kJ
DHf = + 0.012kJ/K
At room temp, T = 298K, will this reaction be spontaneous?
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