Lab 1 (M) Flashcards
Kinetics
the study of reaction rate: how fast a reaction occurs.
factors that influence how fast a reaction occurs
temperature, the concentration of each of the reactants, and
the presence of catalysts.
differential rate equation.
rate = - 1/a d[A]/dt = - 1/b d[B]/dt = 1/c d[C]/dt = K[A]^m[B]^n
k is the overall rate constant
of the reaction and m and n are the reaction orders with respect to reactants A and B.
overall order of the reaction is found by summing the reaction orders with respect to the reactants, (m + n)
what differential rate equation is not/cant do
orders of the reaction with respect to each reactant and the overall order of reaction are not necessarily the same as the stoichiometric coefficients of the reaction; m
and n are not necessarily equal to a and b, respectively.
the individual reaction orders m & n and rate constant (k) can be predicted from the balanced chemical equation alone: these parameters must be determined experimentally.
Reaction rate
is simply the change in concentration of the products or reactants with respect to time, and may be described in terms of either the disappearance of the reactants or the formation of the products
The reaction rate must, by definition, always be a positive number: since reactant concentrations diminish as the reaction proceeds, the rate of change of each reactant will be a negative number and thus must be multiplied by a factor of -1 to furnish a positive value for reaction rate.
Integrated Rate Laws
a differential equation shows the relationship between the rate
of the reaction and the concentrations of the reactants.
Integrating the equation yields the
corresponding integrated rate law, which relates the concentration of reactants to time.
The form of the integrated rate law depends on the overall order of the reaction.
The rate of a zero-order reaction formula
rate = - d[A]/dt = k[A]^0 = k
The rate of a zero-order reaction, where m + n = 0, is not dependent on the concentration of the reactants.
Zero-Order Integrated Rate Law Formula
[A]t = −kt + [A]0
where [A]t is the concentration of reactant A at time t, [A]0 is the initial concentration of reactant A at the start of the reaction, and k is the overall rate constant.
equation forms a straight line, y = mx + b.
If the reaction is zero-order, plotting [A]t versus t will yield a
straight line with slope= –k and a y-intercept of [A]0.
The rate of a first-order reaction formula
rate = - d[A]/dt = k[A]
rate of reaction is directly proportional to the concentration of the reactants
First-Order Integrated Rate Law
For a first-order reaction, the order of the reaction with respect to the reactant is 1, which is often omitted from the rate law (if m or n do not appear in the rate law you can assume that they are 1). The overall reaction order is also 1.
First-Order Integrated Rate Law Formula
[A]t = [A]0e-kt
where [A]0 is the initial concentration, [A]t is the concentration at time t, and k is the overall rate constant.
Taking the natural logarithm of both sides of the first-order integrated rate law transforms the equation into an equation of the form y = mx + b:
in[A]t = -kt
+ in[A]0
From this equation it can be seen that a plot of in[A]t vs. t will yield a straight line, with slope = -k and y-intercept of in[A]0.
The rate of a second-order reaction: Single Reactant formula
Second-order reactions have one of two forms. The first occurs when a reaction has a single reactant and the rate is second order with respect to it.
A + A –> B
Assuming an elementary reaction, the overall rate would be proportional to the square of the
concentration of the reactant:
rate = - d[A]/dt = k[A]2
Notice that in this situation, the order of with respect to reactant A is 2 and the overall reaction
order is also 2.
Second-Order Integrated Rate Law: Single Reactant Formula
1/[A]t = -kt
+ 1/[A]0
From this equation it can be seen that a plot of 1/[A]t vs. t will yield a straight line, with slope = k and y-intercept of 1/[A]0
The rate of a second-order reaction: Two Reactants formula
Second-order reactions may also have two different reactants.
A + B –> C
overall rate of reaction would be:
rate = - d[A]/dt = k[A][B]
Notice that in this situation, the order of the reaction with respect to each reactant is 1, and the overall reaction order (m + n) is 2.
Second-Order Integrated Rate Law: Two Reactants Formula
in [B][A]0/[A][B]0 = k([B]0 - [A]0)t
Thus, a plot of [B][A]0/[A][B]0 versus t will yield a straight line, with slope = k([B]0 - [A]0)t
Pseudo-First Order Reactions
can greatly simplify the quantification of the reaction dynamics.
In order to transform a second or third order reaction into a more simple pseudo-order reaction, the concentrations of one (or more) of the reactants must stay constant through the course of the reaction.
Example of Pseudo-first Order Reaction
A + B → C
B is present at much higher concentrations (>100×) than
the other reactant A.
As the reaction progresses to completion, very little of reactant B will be used up (<1%) and the concentration of reactant B will remain approximately constant
throughout the reaction.
By assuming that [B] is constant, we can combine it with the overall rate constant (k; another constant), to form a new rate constant called the pseudorate constant, k’, where k’ = k[B].
rate = rate = - d[A]/dt = k’[A], where k’ = k[B]n
where k’ is the pseudo-first-order rate constant. It should be noted that because we assume that
the concentration of B is unchanging, at any point in time during the reaction
[B]t = [B]0.
Example of Pseudo-first Order Reaction continued
the order of the reaction with respect to A is, by definition, 1 (first order with respect to reactant A). However, the overall rate of reaction will still depend on [B], and our simplification makes no assumption of the reaction order with respect to B.
To find the effect [B] on the reaction rate (the order of reaction with respect to [B]), we must proceed systematically, that is, by changing one independent variable (in this case a concentration of B) while the other variables and parameters remain fixed.
By observing how the overall rate of reaction changes as we change [B], we can deduce what the reaction order is with respect to B.
overall rate law:
rate = k[A]1[B]1
hydroxide ion and crystal
violet equation and Rate law
OH– (aq) + CV+ (aq) ⟶ CVOH (aq)
Crystal violet (CV+) in water forms a bright purple solution. A colorless product, CVOH, is formed upon reaction with hydroxide.
rate = k[CV]m[OH-]n
The kinetics of hydroxide ion and crystal violet reaction
will be investigated by monitoring the disappearance of the purple colour using spectrophotometry.
collected data processed using spreadsheet software and reaction order with respect to [CV + ] (in other words, m ) will be determined by generating three plots: [CV + ] versus time, ln [CV + ] versus time, and 1/ [CV + ] versus time.
The pseudo rate constant (k’ ) will be determined from the appropriate plot, and the rate constant k will be calculated from k’.
The reaction order with respect to [NaOH] will be determined by varying its concentration amongst three runs.
