Reaction Kinetics (1) An Introduction

What is Reaction Kinetics? 

Suppose we are watching the reaction between magnesium and dilute hydrochloric acid.

What do we notice as the reaction proceeds?

The hydrogen gas is evolved faster at the start of the reaction than it is at the end of the reaction.  You can watch the reaction here.

This kind of an observation begs some questions:

1.   How can we measure the speed of this chemical change?

2.   What characteristic/dimension/property of the reaction is the best to measure the change in speed?

3.   Do all reactions under different conditions between magnesium and hydrochloric acid react in the same way?

4.   Is the way the reaction between magnesium and hydrochloric acid behaves also true for other otherwise unrelated chemical changes?  In other words are we able to see a general or universal pattern to the way reactions go?

Choosing a property of the reaction that changes to measure the kinetics.

In our reaction above involving magnesium and hydrochloric acid, we noticed that a gas was evolved. 

We can use this property of the reaction in order to measure the reaction’s kinetics.

The set up could look like these:

We can measure the volume of the gas evolved over time. 

We have two continuous variables and can plot this data to see how the volume of gas and concentration of hydrochloric acid changes over time.

Here is a typical set up with results.

First the raw data sets showing the volume of hydrogen collected (blue) and the corresponding concentration of hydrochloric acid (red)

First, notice that information about reaction kinetics is only accessible from experimental data.

Second, notice that the volume of hydrogen evolved is proportional to the concentration of hydrochloric acid remaining in the reaction mixture.

Third it is the concentration of hydrochloric acid that is actually changing in the reaction mixture to slow down the reaction till it stops.

Fourth, no more gas is produced when either

1.   the reaction runs out of magnesium ribbon meaning excess acid was used.

2.   the reaction runs out of hydrochloric acid meaning excess magnesium was used.

In this case it is the hydrochloric acid that has run out because at 5 minutes its concentration is zero.

Fifth, if a tangent is drawn at certain points to the hydrochloric acid (red) curve the rate of the reaction at that specific concentration can be calculated.

Here’s an example of how to draw and use the tangent method to get at the rate of reaction at a specific time and volume of gas.

You can see that I have drawn two tangents to the curve for the change in hydrochloric acid concentration.

One is the initial concentration and the other close to the final concentration.

You would need to draw probably five in total to produce data fit to draw a plot of rate of reaction against concentration at say concentrations 5, 10, 15, 25 and 35.

(You should have realised that the concentration units are arbitrary yet proportional to the volume of hydrogen evolved.)

You can try drawing these graphs for yourself as good practice.

Here is my attempt at drawing tangents to the curve and then plotting the reaction rate vs concentration of hydrochloric acid graph:

This approach is called the graphical method to determining the reaction kinetics of the reaction between magnesium and hydrochloric acid. 

What you will notice about this second graph is that it is a linear plot.

The conclusion seems to be that the rate of this reaction is proportional to the concentration of hydrochloric acid remaining in the reaction mixture. 

So we can possibly write what is called a rate equation

Rate of reaction ∝ [HCl]  

Where [  ]  is the symbol for the concentration of the acid and ∝ represents proportionality.

But we can go s step further and say that

Rate of reaction =  k [HCl]¹  

Where k is called the rate constant and the power to which the concentration is raised is called the order of the reaction. 

In this case the reaction is first order with respect to the hydrochloric acid concentration. 

Thus we have attempted to quantify the relationship between reaction rate and concentration in this one specific reaction. 

We can attempt to do a similar thing for other reactions that we decide to study and when we do we find that this pattern of first order reactions is repeated in other different chemical contexts. 

But it is not the only pattern found. 

In my next blog I’m going to describe other mathematical models of reaction kinetics.

And I’m going to demonstrate a different experimental method of getting at the mathematical models . 

What you have read about here isn’t the only experimental way to model the kinetics of a reaction.

We will consider the reaction between hydrogen peroxide (H₂O₂) and potassium iodide (KI).

This reaction can be set up to give a sudden change in colour, the colour of iodine in starch.

We call it the iodine clock reaction because the blue-black colour appears instantaneously in a flash.

You can watch the reaction here and listen to the brilliant Professor Martin Poliakoff explain what’s going on here.

In this reaction we measure the time taken for the blue-black colour to appear. 

The time taken for the blue black colour to appear depends on several factors associated with the reaction.

The time for the blue black colour to appear depends on the concentration of potassium iodide solution, the concentration of the hydrogen peroxide solution, and indeed the temperature of the reaction mixture. 

So it is possible to set up experiments that investigate the effect of changing these factors on the time the blue black colour appears. 

Again using several experiments perhaps repeated for reliability reasons, we can access considerable information about the way the concentrations of the two reactants, hydrogen peroxide (H₂O₂) and potassium iodide (KI) affect the reaction rate. 

But more of this in the next blog…..