Delaware Technical Kinetics Determination of Rate Law Calculations Experiment 17 Calculations to be peformed in Excel using rate laws. Please help. Picture

Delaware Technical Kinetics Determination of Rate Law Calculations Experiment 17 Calculations to be peformed in Excel using rate laws. Please help. Picture of questions attached Experiment 17
Kinetics: Determination of a Rate Law
Purpose
studied? What techniques are being introduced or reviewed?
What chemical principles are being examined? What physical or chemical properties are being
Introduction
E
The speed at which a reactant disappears, or at which a product forms, is known as the rate of
products during a period of time.
reaction. Reaction rate is measured by monitoring the change in concentration of the reactants or
The factors affecting the rate of a reaction are:
Chemical Nature of the Reactants: The ease with which the reactants break existing
bonds and form new bonds.
Phase of the reactants: most reactions occur in the gas or liquid phase as molecules are
moving freely and are able to contact one another.
Concentration of reactants: increasing concentration of reactants increases the rate.
Temperature of the System: almost all chemical reactions are faster at higher
temperatures than they are at lower temperatures.
Presence of Catalyst: catalysts lower the activation energy of a reaction, causing the
reaction rate to increase.
The rate will vary for different reactants because the change in concentration depends on the
reaction stoichiometry. For instance, if a reaction consumes two molecules of A for every one
molecule of B, the rate of disappearance of A will be twice the rate of disappearance of B. For
the general equation:
TOT
a A + b B ? C+ dD (1)
30 V
the relationship can be expressed as:
1 A[A]
a ?t
1 A[B] _ 1 A[C] _1 A[D]
b At C ?t d ?t
(2)
1 A[A]
?t
1 A[B] 1 A[C] 1 A[D]
b At
?t d ?t
a
(2)
?
In equation (2) A[X] refers to the chance in concentration of X ([X]2 -[X]?) that was measured in
the time period At (t2 – t?). [X] indicates concentration is in units of molarity (mol/L). Time is
most frequently measured in seconds.
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Rate Law
Equation (2) allows us to determine the rate relative to any reactant or product but does not
indicate how changes in concentration affect the reaction rate. The relationship between the
reaction rate and reactant concentration is given in equation (3):
Rate= k[A][B]”
(3)
k= the rate constant
[A] = molar concentration of A
m = the order of the reaction with respect to A
[B] = molar concentration of B
n = the order of the reaction with respect to B
The orders and the rate constant cannot be predicted, they must be determined experimentally.
If the concentration of one of the reactants is held constant, it can be folded into to k to simplify
the rate law:
Rate= kobs[B]”
(4)
Reaction under these conditions is known as the initial rate method. Measurement of the rate at
different concentrations of B makes it possible to determine the order with respect to B. In
theory, doubling the concentration should double the rate for a first order reaction; quadruple the
rate for a second order reaction, and so on. In practice results may be ambiguous so a ratio of
results at two concentrations is used to mathematically determine the order n:
rate2
=
nd
[B]2
(5)
o tom odio
ratei [B]1) Domodom
nonnupost
The experiment is performed at several different concentration ratios, and the results are
averaged. Once n is determined, kobs can be calculated from equation (4).
The Arrhenius equation relates the rate constant to temperature. The following equation, derived
from the Arrhenius equation, is used to calculate the activation energy (Ea) by measuring the rate
constant at two Kelvin temperatures:
ki
In
in G, T)
(6) A
di olio di bor
com) som to give the
Orion
40
The reaction monitored in this experiment is:
IO3 + 3 SO32- + I + 3 SO42-
The corresponding rate law is:
Rate = k[IO3-]P[S032-19
The order of the reaction with respect to [103] is measured by holding the initial concentration of
H2SO3 constant and varying the concentration of HIO3. Under these conditions the rate law is!
rate = kobs [103-]P (7)
The reaction occurs in 3 steps:
103 + 3 SO32- + I + 3 SO42-
5 1 + 6 H+ +IO3 ?3 12 + 3 H2O
3 12 + 3 H2O + 3 SO32- ? 61 + 6 H+ +3 SO42-
Completion of the reaction is signaled by the formation of 12, which is detected using starch
indicator. Iz produces a deep blue color when it complexes with starch. Because the third step of
the reaction consumes 12, the solution turns blue only once the SO32- is fully consumed. The time
elapsed from the mixing the reactants to the color change is measured. This produces the two
data points used to calculate the rate; concentration of H2SO3 at the beginning of the reaction
(time = 0) and time when the blue color appears (concentration of H2SO3 is zero):
A[s03] ([SO]”]initial – 0)
(0-t)
b) [SOTinitial
rate=
units = mol L’s-1
?t
(8)
texn,
Note that the rate is negative, as expected for disappearance of a reactant. As the ratio of rates is
calculated (equation 9), the sign may be neglected. The calculation also permits use of the rate
with respect to sulfite as the ratio of rates is the same for all reactants and products.
Once obtained, the reaction rates at different concentrations of IO3 are then substituted into
equation (9) to allow calculation of the order, p.
rate_ ([103],
rate 1 [IO3]1
(9)
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