THE EFFECT OF CATALYSTS ON REACTION RATES

THE EFFECT OF CATALYSTS ON REACTION RATES This page describes and explains the way that adding a catalyst affects the rate of a reaction. It assumes that you are already familiar with basic ideas about the collision theory of reaction rates, and with the Maxwell-Boltzmann distribution of molecular energies in a gas.
	Note:  If you haven't already read the page about collision theory, you should do so before you go on. Use the BACK button on your browser to return to this page, or come back via the rates of reaction menu.
Note that this is only a preliminary look at catalysis as far as it affects rates of reaction. If you are looking for more detail, there is a separate section dealing with catalysts which you can access via a link at the bottom of the page. The facts What are catalysts? A catalyst is a substance which speeds up a reaction, but is chemically unchanged at the end of the reaction. When the reaction has finished, you would have exactly the same mass of catalyst as you had at the beginning. Some examples Some common examples which you may need for other parts of your syllabus include: reaction catalyst Decomposition of hydrogen peroxide manganese(IV) oxide, MnO2 Nitration of benzene concentrated sulphuric acid Manufacture of ammonia by the Haber Process iron Conversion of SO2 into SO3 during the Contact Process to make sulphuric acid vanadium(V) oxide, V2O5 Hydrogenation of a C=C double bond nickel
	Note:  You can find details of these and other catalytic reactions by exploring the menu for the main section on catalysis. You will find a link at the bottom of this page.
The explanation The key importance of activation energy Collisions only result in a reaction if the particles collide with a certain minimum energy called the activation energy for the reaction.
	Note:  What follows assumes you have a reasonable idea about activation energy and its relationship with the Maxwell-Boltzmann distribution. This is covered on the introductory page about collision theory. If you aren't confident about this, follow this link, and use the BACK button on your browser to return to this page.
You can mark the position of activation energy on a Maxwell-Boltzmann distribution to get a diagram like this: Only those particles represented by the area to the right of the activation energy will react when they collide. The great majority don't have enough energy, and will simply bounce apart. Catalysts and activation energy To increase the rate of a reaction you need to increase the number of successful collisions. One possible way of doing this is to provide an alternative way for the reaction to happen which has a lower activation energy. In other words, to move the activation energy on the graph like this: Adding a catalyst has exactly this effect on activation energy. A catalyst provides an alternative route for the reaction. That alternative route has a lower activation energy. Showing this on an energy profile: A word of caution! Be very careful if you are asked about this in an exam. The correct form of words is "A catalyst provides an alternative route for the reaction with a lower activation energy." It does not "lower the activation energy of the reaction". There is a subtle difference between the two statements that is easily illustrated with a simple analogy. Suppose you have a mountain between two valleys so that the only way for people to get from one valley to the other is over the mountain. Only the most active people will manage to get from one valley to the other. Now suppose a tunnel is cut through the mountain. Many more people will now manage to get from one valley to the other by this easier route. You could say that the tunnel route has a lower activation energy than going over the mountain. But you haven't lowered the mountain! The tunnel has provided an alternative route but hasn't lowered the original one. The original mountain is still there, and some people will still choose to climb it. In the chemistry case, if particles collide with enough energy they can still react in exactly the same way as if the catalyst wasn't there. It is simply that the majority of particles will react via the easier catalysed route.

THE EFFECT OF TEMPERATURE ON REACTION RATES

THE EFFECT OF TEMPERATURE ON REACTION RATES This page describes and explains the way that changing the temperature affects the rate of a reaction. It assumes that you are already familiar with basic ideas about the collision theory, and with the Maxwell-Boltzmann distribution of molecular energies in a gas.
	Note:  If you haven't already read the page about collision theory, you should do so before you go on. Use the BACK button on your browser to return to this page, or come back via the rates of reaction menu.
The facts What happens? As you increase the temperature the rate of reaction increases. As a rough approximation, for many reactions happening at around room temperature, the rate of reaction doubles for every 10°C rise in temperature. You have to be careful not to take this too literally. It doesn't apply to all reactions. Even where it is approximately true, it may be that the rate doubles every 9°C or 11°C or whatever. The number of degrees needed to double the rate will also change gradually as the temperature increases.
	Note:  You will find the effect of temperature on rate explored in a slightly more mathematical way on a separate page.
Examples Some reactions are virtually instantaneous - for example, a precipitation reaction involving the coming together of ions in solution to make an insoluble solid, or the reaction between hydrogen ions from an acid and hydroxide ions from an alkali in solution. So heating one of these won't make any noticeable difference to the rate of the reaction. Almost any other reaction you care to name will happen faster if you heat it - either in the lab, or in industry. The explanation Increasing the collision frequency Particles can only react when they collide. If you heat a substance, the particles move faster and so collide more frequently. That will speed up the rate of reaction. That seems a fairly straightforward explanation until you look at the numbers! It turns out that the frequency of two-particle collisions in gases is proportional to the square root of the kelvin temperature. If you increase the temperature from 293 K to 303 K (20°C to 30°C), you will increase the collision frequency by a factor of: That's an increase of 1.7% for a 10° rise. The rate of reaction will probably have doubled for that increase in temperature - in other words, an increase of about 100%. The effect of increasing collision frequency on the rate of the reaction is very minor. The important effect is quite different . . . The key importance of activation energy Collisions only result in a reaction if the particles collide with enough energy to get the reaction started. This minimum energy required is called the activation energy for the reaction.
	Note:  What follows assumes you have a reasonable idea about activation energy and its relationship with the Maxwell-Boltzmann distribution. This is covered on the introductory page about collision theory. If you aren't confident about this, follow this link, and use the BACK button on your browser to return to this page.
You can mark the position of activation energy on a Maxwell-Boltzmann distribution to get a diagram like this: Only those particles represented by the area to the right of the activation energy will react when they collide. The great majority don't have enough energy, and will simply bounce apart. To speed up the reaction, you need to increase the number of the very energetic particles - those with energies equal to or greater than the activation energy. Increasing the temperature has exactly that effect - it changes the shape of the graph. In the next diagram, the graph labelled T is at the original temperature. The graph labelled T+t is at a higher temperature. If you now mark the position of the activation energy, you can see that although the curve hasn't moved very much overall, there has been such a large increase in the number of the very energetic particles that many more now collide with enough energy to react. Remember that the area under a curve gives a count of the number of particles. On the last diagram, the area under the higher temperature curve to the right of the activation energy looks to have at least doubled - therefore at least doubling the rate of the reaction.

THE EFFECT OF PRESSURE ON REACTION RATES

THE EFFECT OF PRESSURE ON REACTION RATES This page describes and explains the way that changing the pressure of a gas changes the rate of a reaction. The facts What happens? Increasing the pressure on a reaction involving reacting gases increases the rate of reaction. Changing the pressure on a reaction which involves only solids or liquids has no effect on the rate. An example In the manufacture of ammonia by the Haber Process, the rate of reaction between the hydrogen and the nitrogen is increased by the use of very high pressures. In fact, the main reason for using high pressures is to improve the percentage of ammonia in the equilibrium mixture, but there is a useful effect on rate of reaction as well.
	Note:  If you want to explore equilibria you will find the topic covered in a separate section of the site.
The explanation The relationship between pressure and concentration Increasing the pressure of a gas is exactly the same as increasing its concentration. If you have a given mass of gas, the way you increase its pressure is to squeeze it into a smaller volume. If you have the same mass in a smaller volume, then its concentration is higher. You can also show this relationship mathematically if you have come across the ideal gas equation: Rearranging this gives: Because "RT" is constant as long as the temperature is constant, this shows that the pressure is directly proportional to the concentration. If you double one, you will also double the other.
	Note:  If you should be able to do calculations involving the ideal gas equation, but aren't very happy about them, you might be interested in my chemistry calculations book.
The effect of increasing the pressure on the rate of reaction Collisions involving two particles The same argument applies whether the reaction involves collision between two different particles or two of the same particle. In order for any reaction to happen, those particles must first collide. This is true whether both particles are in the gas state, or whether one is a gas and the other a solid. If the pressure is higher, the chances of collision are greater. Reactions involving only one particle If a reaction only involves a single particle splitting up in some way, then the number of collisions is irrelevant. What matters now is how many of the particles have enough energy to react at any one time.
	Note:  If you aren't sure about this, then read the page about collision theory and activation energy before you go on. Use the BACK button on your browser to return to this page.