The reaction H2 SeO3(a q)+6 I^-(a q)+ 4 H^+(a q) ⟶ Se(s)+2 I3^-(a q)+3 H2 O(l) was stu

The reaction H2 SeO3(a q)+6 I^-(a q)+ 4 H^+(a q) ⟶ ...

The reaction $$\begin{aligned} \mathrm{H}_{2} \mathrm{SeO}_{3}(a q)+6 \mathrm{I}^{-}(a q)+& 4 \mathrm{H}^{+}(a q) \\ \longrightarrow & \mathrm{Se}(s)+2 \mathrm{I}_{3}^{-}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l) \end{aligned}$$ was studied at $0^{\circ} \mathrm{C},$ and the following data were obtained: These relationships hold only if there is a very small amount of $\mathrm{I}_{3}^{-}$ present. What is the rate law and the value of the rate constant? $$\text { (Assume that rate }=-\frac{\Delta\left[\mathrm{H}_{2} \mathrm{SeO}_{3}\right]}{\Delta t}.)$$

The reaction $$\begin{aligned}
\mathrm{H}_{2} \mathrm{SeO}_{3}(a q)+6 \mathrm{I}^{-}(a q)+& 4 \mathrm{H}^{+}(a q) \\
\longrightarrow & \mathrm{Se}(s)+2 \mathrm{I}_{3}^{-}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l)
\end{aligned}$$ was studied at $0^{\circ} \mathrm{C},$ and the following data were obtained: These relationships hold only if there is a very small amount of $\mathrm{I}_{3}^{-}$ present. What is the rate law and the value of the rate constant? $$\text { (Assume that rate }=-\frac{\Delta\left[\mathrm{H}_{2} \mathrm{SeO}_{3}\right]}{\Delta t}.)$$

Key Concepts

Rate Law

The rate law is an expression that relates the rate of a chemical reaction to the concentrations of its reactants. In experimental kinetics, it is determined by measuring how changes in reactant concentration affect the reaction rate. It typically takes the form rate = k[A]^m[B]^n, where k is the rate constant and m and n are the orders with respect to each reactant.

Reaction Order

The reaction order refers to the exponent on a reactant’s concentration in the rate law, indicating how sensitively the rate depends on a change in that concentration. The overall order is the sum of the individual orders and determines the units of the rate constant as well as the kinetics (for example, a first?order reaction behaves differently than a second?order one).

Rate Constant

The rate constant, k, is a proportionality factor in the rate law that is specific to a given reaction at a given temperature. Its value and units depend on the overall reaction order. The constant essentially scales the concentration terms in the rate law so that the product equals the rate of the reaction, and it can be determined by measuring initial reaction rates under controlled conditions.

Initial Rate Method

The initial rate method involves measuring the rate of reaction at the very start (initial time) when the reactant concentrations are known and when product concentrations (which might interfere with the reaction) are negligible. This method simplifies analysis by ensuring that reverse reactions and side effects (such as product inhibition) do not affect the measurements, as noted in the condition of a very small amount of I?? present.

Negligible Product Interference

Many kinetics experiments assume that the concentration of products is initially very low or negligible. This assumption helps in accurately determining the rate law, because the product does not affect the measurement of the initial rate. It is particularly important when products might participate in secondary reactions or feedback effects that would otherwise complicate the kinetic analysis.

Transcript

00:01 So we have the data varying the concentrations of three reactants and looking at the effects on the initial rate.

00:15 And if we look at the first two experiments, the hydronium ion h -plus is at 2 times 10 to the minus 2, and the iodide is at 2 times 10 to the minus 2.

00:30 So those two are held constant in the first two experiments.

00:33 And we can look at the effect of doubling the h2 -s -e -o -3, i guess that's selenic acid.

00:41 And so i've set it up to compare the ratio of the rates in the first two experiments to that concentration in the first two experiments.

00:57 And so i plugged in the rates for the first two experiments here.

01:03 And notice that when we determined, the rate law, we're working on the value of that exponent, which is the exponent is on the concentration.

01:13 So you want to be sure to show that on the concentration in each case.

01:16 Basically, the rate doubled when the concentration doubled.

01:19 And so the exponent on the h2 seo3 is a 1.

01:25 Now, and you can see also when you go from the first to the third, that again, the h plus and the i minus remain constant.

01:35 And there was a tripling of the rate.

01:38 And so that one is worked out.

01:40 Now, looking at the last four experiments, the h -2 -s -e -o -3 was constant, and the last and not the one above it, but the second one above it, also holds the h -plus constant.

01:56 So we can look at the iodide there.

01:58 Let's see, that's two, four, six.

02:01 Look at the seventh experiment over the fifth experiment.

02:12 You have to be careful that everything remains constant except the one that you're changing.

02:18 So again, the h2 -sco -3 is 1 times 10 to the minus 4 in both of those experiments.

02:25 The h -plus is 1 times 10 to the minus 2 in both of those experiments.

02:29 And in the seventh one, okay, so this is going to be related to the iodide in trial number 7 and the iodide in trial number 5.

02:43 And so the rate in experiment 7 was 3 .36 times 10 to the minus 7...