# Chemistry - Has IUPAC been inaccurate in their 1994 definition of kinetic stationary state?

## Solution 1:

IUPAC should say they are defining the "steady state approximation" rather than the steady state.

There are two approximations in this technique.

First, that final reaction step is not an equilibrium (reverse reaction rate is zero).

Second, that the intermediate is so unstable that its concentration (and therefore the rate of change of the concentration) is approximately zero.

Compare the IUPAC definition to the following 1954 U.S. National Research Council explanation:

STEADY STATE APPROXIMATION

The most generally applicable of these simple methods is the so-called "steady-state approximation." A steady state may be defined as a condition in which the rates of change of the concentrations of the several intermediates are very small compared to the rates of change of the concentrations of the reactants and products. This condition is realizable whenever the ratio of the concentrations of the intermediates to the concentrations of the reactants is very much less than unity. When this condition is not attained, the method is not applicable; however, it should not then be necessary since the (larger) concentrations of the intermediates could be measured by experimental means. ... The

steady-state approximation consists in setting the rates of change of each of the intermediates equal to zeroand in solving simultaneously the resulting algebraic equations.

Overall, if the situation was truly, not just approximately, steady state, yes the concentration of the intermediate would be constant, but the point is that the "stead state approximation" is useful even when the concentration of the intermediate, though always near zero in absolute terms, decreases to say half of its original concentration over the course of an experiment (as the starting material, A, decreases to half its original concentration).

Additionally, a very small $\ce{[X]}$ does not purely-mathematically imply a small rate of change of $\ce{[X]}$, because $\ce{[X]}$ could rapidly oscillate within a small, near-zero, range, resulting in brief instances of large rate of change. Instead, $\ce{[X]}$ being small over a given period of time really just places a limit on how long of a time the rate of change can exceed a certain value. For example, if $\ce{[X]} < 0.001\ \mathrm{M}$,

$|\mathrm{d}[\ce{X}]/\mathrm{d}t|$ must not exceed $1\ \mathrm{M/s}$ for more than $0.001\ \mathrm{s}$.

## Solution 2:

Ok, this isn't an answer, but below is the text in question from Muller's *Glossary of Terms Used in Physical Organic Chemistry (IUPAC Recommendations 1994)*. The definition contains two parts.

I agree with DavePHD that the *first definition*, which is for a batch reaction, would be better called what is known as the **STEADY STATE APPROXIMATION**

The *second definition* is for a flow reactor which is at a true **STEADY STATE**. In this case all the reactants, intermediates, and products would have constant concentrations. That is to say the relative fluctuations (%) in those concentrations would be on the order of the precision with which the flows could be controlled.

**Text is below**

steady state (or stationary state)

(1)In a kinetic analysis of a complex reaction involving unstable intermediates in low concentration, the rate of change of each such intermediate is set equal to zero, so that the rate equation can be expressed as a function of the concentrations of chemical species present in macroscopic amounts. For example, assume that $\ce{X}$ is an unstable intermediate in the reaction sequence:$$\ce{A <-->[$k_1$][k_{-1}] X}$$

$$\ce{X + C ->[$k_2$] D}$$

Conservation of mass requires that:

$$\ce{[A] + [X] + [D] = [A]_0}$$

which, since $\ce{[A]_0}$ is constant, implies:

$$\dfrac{-d\ce{[X]}}{dt} = \dfrac{d\ce{[A]}}{dt} + \dfrac{d\ce{[D]}}{dt}$$

Since $\ce{[X]}$ is negligibly small, the rate of formation of $\ce{D}$ is essentially equal to the rate of disappearance of $\ce{A}$, and the rate of change of $\ce{[X]}$ can be set equal to zero. Applying the steady state approximation ($d\ce{[X]}/dt = 0$) allows the elimination of $\ce{[X]}$ from the kinetic equations, whereupon the rate of reaction is expressed

$$ \dfrac{d\ce{[D]}}{dt} = -\dfrac{d\ce{[A]}}{dt} = \dfrac{k_1k_2\ce{[A][C]}}{k_{-1} + k_2\ce{[C]}}$$

Note: The steady-state approximation does not imply that $\ce{[X]}$ is even approximately constant, only that its absolute rate of change is very much smaller than that of $\ce{[A]}$ and $\ce{[D]}$. Since according to the reaction scheme $\ce{d[D]/dt = k_2[X][C]}$, the assumption that $\ce{[X]}$ is constant would lead, for the case in which $\ce{C}$ is in large excess, to the absurd conclusion that formation of the product $\ce{D}$ will continue at a constant rate even after the reactant $\ce{A}$ has been consumed.

(2)In a stirred flow reactor a steady state implies a regime so that all concentrations are independent of time.

I think definitions in

A GLOSSARY OF TERMS USED IN CHEMICAL KINETICS, INCLUDING REACTION DYNAMICS(IUPAC Recommendations 1996) edited by KEITH J. LAIDLER, Pure & Appl. Chem., Vol. 68, No. 1, pp. 149-192, 1996.

are much better.

Pre-Equilibrium or Prior Equilibrium

The mechanism of a reaction may involve two or more consecutive reactions: If any step except the first is rate-controlling,those steps that precede it are essentially at

$\ce{A <=> B <=> C <=> ... <=> X -> Y ->Z}$

equilibrium, and there is said to be a pre-equilibrium, or a prior equilibrium; for example

$\quad\quad\quad\quad\quad\quad\quad\quad\quad\text{rate-controlling step}$

$\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\blacktriangledown$

$\ce{A <=> B <=> C <=> ... <=> X -> Y ->Z}$$\text{|}\blacktriangleleft\text{---- pre-equilibrium ------} \blacktriangleright\text{|}$

Steady State (or Stationary State)If during the course of a chemical reaction the concentration of an intermediate remains constant, the intermediate is said to be in a steady state.

In a static system a reaction intermediate reaches a steady state if the processes leading to its formation and those removing it are approximately in balance. The steady-state hypothesis leads to a great simplification in reaching an expression for the overall rate of a composite reaction in terms of the rate constants for the individual elementary steps. Care must be taken to apply the steady-state hypothesis only to appropriate reaction intermediates. An intermediate such as an atom or a free radical, present at low concentrations, can usually be taken to obey the hypothesis during the main course of the reaction.

In a flow system a steady state may be established even for intermediates present at relatively high concentrations.