THE EFFECT OF CONCENTRATION ON REACTION RATES

THE EFFECT OF CONCENTRATION ON REACTION RATES This page describes and explains the way that changing the concentration of a solution affects the rate of a reaction. Be aware that this is an introductory page only. If you are interested in orders of reaction, you will find separate pages dealing with these. You can access these via the rates of reaction menu (link at the bottom of the page). The facts What happens? For many reactions involving liquids or gases, increasing the concentration of the reactants increases the rate of reaction. In a few cases, increasing the concentration of one of the reactants may have little noticeable effect of the rate. These cases are discussed and explained further down this page. Don't assume that if you double the concentration of one of the reactants that you will double the rate of the reaction. It may happen like that, but the relationship may well be more complicated.
	Note:  The mathematical relationship between concentration and rate of reaction is dealt with on the page about orders of reaction. If you are interested, you can use this link or read about it later via the rate of reaction menu (link at the bottom of the page).
Some examples The examples on this page all involve solutions. Changing the concentration of a gas is achieved by changing its pressure. This is covered on a separate page.
	Note:  If you want to explore the effect of changing pressure on the rate of a reaction, you could use this link. Alternatively, use the link to the rates of reaction menu at the bottom of this page.
Zinc and hydrochloric acid In the lab, zinc granules react fairly slowly with dilute hydrochloric acid, but much faster if the acid is concentrated. The catalytic decomposition of hydrogen peroxide Solid manganese(IV) oxide is often used as a catalyst in this reaction. Oxygen is given off much faster if the hydrogen peroxide is concentrated than if it is dilute. The reaction between sodium thiosulphate solution and hydrochloric acid This is a reaction which is often used to explore the relationship between concentration and rate of reaction in introductory courses (like GCSE). When a dilute acid is added to sodium thiosulphate solution, a pale yellow precipitate of sulphur is formed. As the sodium thiosulphate solution is diluted more and more, the precipitate takes longer and longer to form. The explanation Cases where changing the concentration affects the rate of the reaction This is the common case, and is easily explained. 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 solution, or whether one is in solution and the other a solid. If the concentration 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.
Suppose that at any one time 1 in a million particles have enough energy to equal or exceed the activation energy. If you had 100 million particles, 100 of them would react. If you had 200 million particles in the same volume, 200 of them would now react. The rate of reaction has doubled by doubling the concentration. Cases where changing the concentration doesn't affect the rate of the reaction At first glance this seems very surprising! Where a catalyst is already working as fast as it can Suppose you are using a small amount of a solid catalyst in a reaction, and a high enough concentration of reactant in solution so that the catalyst surface was totally cluttered up with reacting particles. Increasing the concentration of the solution even more can't have any effect because the catalyst is already working at its maximum capacity. In certain multi-step reactions This is the more important effect from an A' level point of view. Suppose you have a reaction which happens in a series of small steps. These steps are likely to have widely different rates - some fast, some slow. For example, suppose two reactants A and B react together in these two stages: The overall rate of the reaction is going to be governed by how fast A splits up to make X and Y. This is described as the rate determining step of the reaction. If you increase the concentration of A, you will increase the chances of this step happening for reasons we've looked at above. If you increase the concentration of B, that will undoubtedly speed up the second step, but that makes hardly any difference to the overall rate. You can picture the second step as happening so fast already that as soon as any X is formed, it is immediately pounced on by B. That second reaction is already "waiting around" for the first one to happen.
	Note:  The overall rate of reaction isn't entirely independent of the concentration of B. If you lowered its concentration enough, you will eventually reduce the rate of the second reaction to the point where it is similar to the rate of the first. Both concentrations will matter if the concentration of B is low enough. However, for ordinary concentrations, you can say that (to a good approximation) the overall rate of reaction is unaffected by the concentration of B.
The best specific examples of reactions of this type comes from organic chemistry. These involve the reaction between a tertiary halogenoalkane (alkyl halide) and a number of possible substances - including hydroxide ions. These are examples of nucleophilic substitution using a mechanism known as SN1.
	Note:  If you are interested in exploring nucleophilic substitution reactions further, you could follow this link. Otherwise, you can find more about how the relationship between concentration and rate of reaction is affected by reaction mechanisms by exploring the topics at the bottom of the rates of reaction menu (link below).

THE EFFECT OF SURFACE AREA ON REACTION RATES

THE EFFECT OF SURFACE AREA ON REACTION RATES This page describes and explains the effect of changing the surface area of a solid on the rate of a reaction it is involved in. This applies to reactions involving a solid and a gas, or a solid and a liquid. It includes cases where the solid is acting as a catalyst. The facts What happens? The more finely divided the solid is, the faster the reaction happens. A powdered solid will normally produce a faster reaction than if the same mass is present as a single lump. The powdered solid has a greater surface area than the single lump.
	Note:  Why normally? What exceptions can there be? Imagine a case of a very fine powder reacting with a gas. If the powder was in one big heap, the gas may not be able to penetrate it. That means that its effective surface area is much the same as (or even less than) it would be if it were present in a single lump. A small heap of fine magnesium powder tends to burn rather more slowly than a strip of magnesium ribbon, for example.
Some examples Calcium carbonate and hydrochloric acid In the lab, powdered calcium carbonate reacts much faster with dilute hydrochloric acid than if the same mass was present as lumps of marble or limestone. The catalytic decomposition of hydrogen peroxide This is another familiar lab reaction. Solid manganese(IV) oxide is often used as the catalyst. Oxygen is given off much faster if the catalyst is present as a powder than as the same mass of granules. Catalytic converters Catalytic converters use metals like platinum, palladium and rhodium to convert poisonous compounds in vehicle exhausts into less harmful things. For example, a reaction which removes both carbon monoxide and an oxide of nitrogen is: Because the exhaust gases are only in contact with the catalyst for a very short time, the reactions have to be very fast. The extremely expensive metals used as the catalyst are coated as a very thin layer onto a ceramic honeycomb structure to maximise the surface area. The explanation You are only going to get a reaction if the particles in the gas or liquid collide with the particles in the solid. Increasing the surface area of the solid increases the chances of collision taking place. Imagine a reaction between magnesium metal and a dilute acid like hydrochloric acid. The reaction involves collision between magnesium atoms and hydrogen ions. Increasing the number of collisions per second increases the rate of reaction.

THE COLLISION THEORY OF REACTION RATES

THE COLLISION THEORY OF REACTION RATES This page describes the collision theory of reaction rates. It concentrates on the key things which decide whether a particular collision will result in a reaction - in particular, the energy of the collision, and whether or not the molecules hit each other the right way around (the orientation of the collision). The individual factors which affect the rate of a reaction (temperature, concentration, and so on) are discussed on separate pages. You can get at these via the rates of reaction menu - there is a link at the bottom of the page. We are going to look in detail at reactions which involve a collision between two species.
	Species:  This is a useful term which covers any sort of particle you like - molecule, ion, or free radical.
Reactions where a single species falls apart in some way are slightly simpler because you won't be involved in worrying about the orientation of collisions. Reactions involving collisions between more than two species are going to be extremely uncommon (see below). Reactions involving collisions between two species It is pretty obvious that if you have a situation involving two species they can only react together if they come into contact with each other. They first have to collide, and then they may react. Why "may react"? It isn't enough for the two species to collide - they have to collide the right way around, and they have to collide with enough energy for bonds to break. (The chances of all this happening if your reaction needed a collision involving more than 2 particles are remote. All three (or more) particles would have to arrive at exactly the same point in space at the same time, with everything lined up exactly right, and having enough energy to react. That's not likely to happen very often!) The orientation of collision Consider a simple reaction involving a collision between two molecules - ethene, CH2=CH2, and hydrogen chloride, HCl, for example. These react to give chloroethane. As a result of the collision between the two molecules, the double bond between the two carbons is converted into a single bond. A hydrogen atom gets attached to one of the carbons and a chlorine atom to the other.
	Note:  The mechanism for this reaction is dealt with on a separate page. This might help you to understand why the orientation of the two molecules is so important. If you want to read a bit more about this, follow this link and use the BACK button on your browser to return to this page.
The reaction can only happen if the hydrogen end of the H-Cl bond approaches the carbon-carbon double bond. Any other collision between the two molecules doesn't work. The two simply bounce off each other. Of the collisions shown in the diagram, only collision 1 may possibly lead on to a reaction. If you haven't read the page about the mechanism of the reaction, you may wonder why collision 2 won't work as well. The double bond has a high concentration of negative charge around it due to the electrons in the bonds. The approaching chlorine atom is also slightly negative because it is more electronegative than hydrogen. The repulsion simply causes the molecules to bounce off each other.
	Note:  If you aren't sure about electronegativity , you might like to follow this link. Use the BACK button on your browser to return to this page.
In any collision involving unsymmetrical species, you would expect that the way they hit each other will be important in deciding whether or not a reaction happens. The energy of the collision Activation Energy Even if the species are orientated properly, you still won't get a reaction unless the particles collide with a certain minimum energy called the activation energy of the reaction. Activation energy is the minimum energy required before a reaction can occur. You can show this on an energy profile for the reaction. For a simple over-all exothermic reaction, the energy profile looks like this:
	Note:  The only difference if the reaction was endothermic would be the relative positions of the reactants and products lines. For an endothermic change, the products would have a higher energy than the reactants, and so the green arrow would be pointing upwards. It makes no difference to the discussion about the activation energy.
If the particles collide with less energy than the activation energy, nothing important happens. They bounce apart. You can think of the activation energy as a barrier to the reaction. Only those collisions which have energies equal to or greater than the activation energy result in a reaction. Any chemical reaction results in the breaking of some bonds (needing energy) and the making of new ones (releasing energy). Obviously some bonds have to be broken before new ones can be made. Activation energy is involved in breaking some of the original bonds. Where collisions are relatively gentle, there isn't enough energy available to start the bond-breaking process, and so the particles don't react. The Maxwell-Boltzmann Distribution Because of the key role of activation energy in deciding whether a collision will result in a reaction, it would obviously be useful to know what sort of proportion of the particles present have high enough energies to react when they collide. In any system, the particles present will have a very wide range of energies. For gases, this can be shown on a graph called the Maxwell-Boltzmann Distribution which is a plot of the number of particles having each particular energy.
	Note:  The graph only applies to gases, but the conclusions that we can draw from it can also be applied to reactions involving liquids.
The area under the curve is a measure of the total number of particles present.
	Note:  The reason for this lies in some maths beyond the scope of an A'level chemistry course. It is important that you remember that the area under the curve gives a count of the number of particles even if you don't understand why!
The Maxwell-Boltzmann Distribution and activation energy Remember that for a reaction to happen, particles must collide with energies equal to or greater than the activation energy for the reaction. We can mark the activation energy on the Maxwell-Boltzmann distribution: Notice that the large majority of the particles don't have enough energy to react when they collide. To enable them to react we either have to change the shape of the curve, or move the activation energy further to the left. This is described on other pages.
	Note:  You can change the shape of the curve by changing the temperature of the reaction. You can change the position of the activation energy by adding a catalyst to the reaction. You could either go straight to these pages if you are interested, or access them later via the rates of reaction menu (link at the bottom of the page).