Question 9 Unsaved
What is the rate law for the following reaction, if the order of the reaction is m, an unknown?
H2O2(aq) → H2O(l) + ½O2(g)
a. k [H2O2]m
b.k [H2O]m [O2]1/2
c.k [H2O] m /[H2O][O2
d.k[H2O] m [O2]m
Thanks in advance. The k and m are meant to be exponents :)
k [H2O2]^m
Well, well, well, looks like we have ourselves a rate law problem here! Time to put on my chemistry wig and get wacky!
So, in order to determine the rate law for this reaction, we need to consider the stoichiometry of the reaction and the rate-determining step. In this case, the reaction involves one molecule of H2O2 producing one molecule of H2O and half a molecule of O2 (gotta love those sneaky coefficients).
Now, let's break it down. We have H2O2 on the reactant side and H2O and O2 on the product side. Based on the reaction equation, it seems like the rate of the reaction is directly proportional to the concentration of H2O2. So, that's a good start!
But wait, there's more! We also have to take into account that the order of the reaction is m, which means that the rate is proportional to the concentration of H2O2 raised to some exponent m. So our rate law should be k [H2O2]^m!
Now, let's dissect the answer choices. Option a looks good with k and [H2O2]^m, but the others throw in some distractions. Option b complicates matters by involving [H2O] and [O2], which we don't see in the rate-determining step. Option c further confuses the situation by introducing [H2O] and [O2] in the denominator, and option d is just a whole mix-up by bringing in [O2] with its own exponent.
So, after all that clowning around, the correct answer is a. k [H2O2]^m. Keep those laughter tears of joy flowing, my friend!
To determine the rate law for the given reaction, we need to examine the reaction's stoichiometry and the experimentally observed reaction rate.
Looking at the balanced chemical equation:
H2O2(aq) → H2O(l) + ½O2(g)
We can see that the concentration of H2O2 is decreasing while the concentrations of H2O and O2 are increasing at the same rate. Based on this, we can conclude that the rate of the reaction is directly proportional to the concentration of H2O2.
Therefore, the rate law for this reaction is:
Rate = k [H2O2]^m
Where:
- Rate is the reaction rate or the change in concentration of a reactant/product per unit of time,
- k is the rate constant, which depends on the temperature and catalyst, and
- m is the order of the reaction with respect to H2O2.
Looking at the provided options, the correct rate law is:
a. k [H2O2]^m
To determine the rate law for a reaction, you need to experimentally determine the dependence of the reaction rate on the concentrations of the reactants. In this case, the reaction is given by:
H2O2(aq) → H2O(l) + ½O2(g)
The rate law for this reaction can be expressed as:
Rate = k [H2O2]^m
where "k" is the rate constant and "[H2O2]" represents the concentration of hydrogen peroxide.
Since we are dealing with an unknown order of reaction, we need to investigate how changes in the concentration of hydrogen peroxide affect the rate of the reaction. This can be achieved by performing a series of experiments where the initial concentration of hydrogen peroxide is varied while keeping the concentrations of other reactants constant.
By comparing the initial rates of the reaction under different concentrations of hydrogen peroxide, you can determine the order of the reaction (m) by observing how the rate changes with the concentration.
For example, let's say you performed two experiments:
Experiment 1: [H2O2] = 0.2 M, Rate1 = k [H2O2]^m1
Experiment 2: [H2O2] = 0.1 M, Rate2 = k [H2O2]^m2
If the rate doubles when the concentration of hydrogen peroxide is halved (Rate2 = 2 * Rate1), then the order of the reaction with respect to hydrogen peroxide (m) is 1. This means the rate law for this reaction would be:
Rate = k [H2O2]
Based on this analysis, the correct choice for the rate law of the given reaction is:
a. k [H2O2]m