Kinetics and Pseudo-Order Conditions As you have seen in class and in your textbook (Silberberg sections 14-2 & 14.3), the rate of reaction can be related to the reactant concentrations through an expression called the rate law for a particular reaction. For the iodine clock reaction studied in this experiment the general form of the rate is: Rate = k[H*]*[1] [H₂O₂)² Where: kis the rate constant for the reaction (temperature dependent) x is the order with respect to H* y is the order with respect to iodide z is the order with respect to hydrogen peroxide We assume that x, y, and z are whole numbers (most commonly 0, 1, or 2). In this experiment you will determine the values for k, y and z. Solving a system with four unknown variables is possible but is much easier if the experiments are designed so that only one variable changes at a time. Since k is already a constant value, we will control the reactant concentrations in order to manipulate the variables in the equation. You will approximate a constant [H*] by using a large excess of this reagent. Since the [H*] is so much larger than the [1] and the [H₂O₂] it remains effectively constant, while the relative changes in [1] and the [H₂O₂] are very large. If we make the (reasonable) assumption that [H*] is constant under the conditions described above, we can rewrite Equation [4] as: Run-1 Run-2 Run-3 Run-4 k' Where: k' = k[H*]* The constant k' is a pseudo-order constant - it behaves and can be treated as if it were a rate constant for a reaction of orders of y and z, but it is not a true rate constant because its value depends on the concentration of a reactant (in this case, [H*]). Quantity [Na₂S₂O3] stock [KI] stock [H₂O₂] stock Value 210.6 74.3 314 82 0.005M 0.2M 0.1M Rate = k'[-]³[H₂0₂1¹ 0.0386 Average time(s) [H₂O₂](M) [I-](M) 0.04 0.04 0.0125 0.025 0.04 0.08 0.1 0.1 [S₂03²-] (M) [1] 0.001 0.001 0.00125 0.00125 [2] Average rate of reaction (M/s) 2.37 x 10-6 M/s 6.73 x 10-6 M/s 1.99 x 10-6 M/s 7.62 x 10-6 M/s
Kinetics and Pseudo-Order Conditions As you have seen in class and in your textbook (Silberberg sections 14-2 & 14.3), the rate of reaction can be related to the reactant concentrations through an expression called the rate law for a particular reaction. For the iodine clock reaction studied in this experiment the general form of the rate is: Rate = k[H*]*[1] [H₂O₂)² Where: kis the rate constant for the reaction (temperature dependent) x is the order with respect to H* y is the order with respect to iodide z is the order with respect to hydrogen peroxide We assume that x, y, and z are whole numbers (most commonly 0, 1, or 2). In this experiment you will determine the values for k, y and z. Solving a system with four unknown variables is possible but is much easier if the experiments are designed so that only one variable changes at a time. Since k is already a constant value, we will control the reactant concentrations in order to manipulate the variables in the equation. You will approximate a constant [H*] by using a large excess of this reagent. Since the [H*] is so much larger than the [1] and the [H₂O₂] it remains effectively constant, while the relative changes in [1] and the [H₂O₂] are very large. If we make the (reasonable) assumption that [H*] is constant under the conditions described above, we can rewrite Equation [4] as: Run-1 Run-2 Run-3 Run-4 k' Where: k' = k[H*]* The constant k' is a pseudo-order constant - it behaves and can be treated as if it were a rate constant for a reaction of orders of y and z, but it is not a true rate constant because its value depends on the concentration of a reactant (in this case, [H*]). Quantity [Na₂S₂O3] stock [KI] stock [H₂O₂] stock Value 210.6 74.3 314 82 0.005M 0.2M 0.1M Rate = k'[-]³[H₂0₂1¹ 0.0386 Average time(s) [H₂O₂](M) [I-](M) 0.04 0.04 0.0125 0.025 0.04 0.08 0.1 0.1 [S₂03²-] (M) [1] 0.001 0.001 0.00125 0.00125 [2] Average rate of reaction (M/s) 2.37 x 10-6 M/s 6.73 x 10-6 M/s 1.99 x 10-6 M/s 7.62 x 10-6 M/s
Chapter12: Chemical Kinetics
Section: Chapter Questions
Problem 3RQ: One experimental procedure that can be used to determine the rate law of a reaction is the method of...
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![Kinetics and Pseudo-Order Conditions
As you have seen in class and in your textbook (Silberberg sections 14-2 & 14.3), the rate of reaction can be
related to the reactant concentrations through an expression called the rate law for a particular reaction. For the
iodine clock reaction studied in this experiment the general form of the rate is:
Rate = k[H*]*[1] [H₂0₂]²
Where: kis the rate constant for the reaction (temperature dependent)
x is the order with respect to H*
y is the order with respect to iodide
z is the order with respect to hydrogen peroxide
We assume that x, y, and z are whole numbers (most commonly 0, 1, or 2).
In this experiment you will determine the values for k, y and z.
Solving a system with four unknown variables is possible but is much easier if the experiments are designed so
that only one variable changes at a time. Since k is already a constant value, we will control the reactant
concentrations in order to manipulate the variables in the equation. You will approximate a constant [H+] by
using a large excess of this reagent. Since the [H+] is so much larger than the [1] and the [H₂O₂] it remains
effectively constant, while the relative changes in [I] and the [H₂O₂] are very large.
If we make the (reasonable) assumption that [H*] is constant under the conditions described above, we can
rewrite Equation [4] as:
Run-1
Run-2
Run-3
Run-4
k'
Where: k' = k[H*]*
The constant k'is a pseudo-order constant - it behaves and can be treated as if it were a rate constant for a
reaction of orders of y and z, but it is not a true rate constant because its value depends on the concentration of a
reactant (in this case, [H*]).
Quantity
[Na₂S₂O3] stock
[KI] stock
[H₂O₂] stock
Value
210.6
74.3
314
82
0.005M
0.2M
0.1M
Rate = k'[-]³[H₂0₂]²
0.0386
Average time(s) [H₂O₂](M) [I](M)
0.04
0.04
0.0125
0.025
0.04
0.08
0.1
0.1
[S₂03²-] (M)
[1]
0.001
0.001
0.00125
0.00125
[2]
Average rate of
reaction (M/s)
2.37 x 10-6 M/s
6.73 x 10-6 M/s
1.99 x 10-6 M/s
7.62 x 10-6 M/s](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fe31c5e57-f927-44e9-a19f-7b1a4337091e%2F3eef22df-638b-467d-97de-3e25e94ea8a6%2F9mjxfop.png&w=3840&q=75)
Transcribed Image Text:Kinetics and Pseudo-Order Conditions
As you have seen in class and in your textbook (Silberberg sections 14-2 & 14.3), the rate of reaction can be
related to the reactant concentrations through an expression called the rate law for a particular reaction. For the
iodine clock reaction studied in this experiment the general form of the rate is:
Rate = k[H*]*[1] [H₂0₂]²
Where: kis the rate constant for the reaction (temperature dependent)
x is the order with respect to H*
y is the order with respect to iodide
z is the order with respect to hydrogen peroxide
We assume that x, y, and z are whole numbers (most commonly 0, 1, or 2).
In this experiment you will determine the values for k, y and z.
Solving a system with four unknown variables is possible but is much easier if the experiments are designed so
that only one variable changes at a time. Since k is already a constant value, we will control the reactant
concentrations in order to manipulate the variables in the equation. You will approximate a constant [H+] by
using a large excess of this reagent. Since the [H+] is so much larger than the [1] and the [H₂O₂] it remains
effectively constant, while the relative changes in [I] and the [H₂O₂] are very large.
If we make the (reasonable) assumption that [H*] is constant under the conditions described above, we can
rewrite Equation [4] as:
Run-1
Run-2
Run-3
Run-4
k'
Where: k' = k[H*]*
The constant k'is a pseudo-order constant - it behaves and can be treated as if it were a rate constant for a
reaction of orders of y and z, but it is not a true rate constant because its value depends on the concentration of a
reactant (in this case, [H*]).
Quantity
[Na₂S₂O3] stock
[KI] stock
[H₂O₂] stock
Value
210.6
74.3
314
82
0.005M
0.2M
0.1M
Rate = k'[-]³[H₂0₂]²
0.0386
Average time(s) [H₂O₂](M) [I](M)
0.04
0.04
0.0125
0.025
0.04
0.08
0.1
0.1
[S₂03²-] (M)
[1]
0.001
0.001
0.00125
0.00125
[2]
Average rate of
reaction (M/s)
2.37 x 10-6 M/s
6.73 x 10-6 M/s
1.99 x 10-6 M/s
7.62 x 10-6 M/s
Transcribed Image Text:SAMPLE TABLE (PSEUDO RATE): Find the missing information in this table
Pseudo-rate constants from experimental data
Run 1
Run 2
Run 3
Run 4
Average:
k' (indicate units)
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