ChemistryOnline is another engaging and interactive subject site from the ScholarNET Online Education stable of online learning resources. It covers the major Chemistry topics that are fundamental to every senior Chemistry course - Atomic Structure and Bonding, Aqueous Chemistry, Organic Chemistry, Redox Chemistry as well as Thermochemistry (Energy).
ChemistryOnline has been designed with both the teacher and student in mind, providing a wealth of interesting and interactive materials to help make learning Chemistry a more palatable experience.
Pure and simply, molecules have to physically touch each other to react. It is no good them being near each other. They have to be in direct contact to make the chemical change that turns them into new substances. And most reactants need to do it with a bump, to gain enough energy for the reaction to happen.
Activation energy is the name for the amount of energy it takes for two atoms or molecules to change from their original arrangement into a new one - to react together. For example, magnesium ribbon burns with an intense white light. It is reacting with oxygen molecules from the air. But it needs a great deal of activation energy to get the magnesium atoms in the piece of metal to react with the oxygen atoms from the air. So to make this reaction happen, you need to heat the piece of magnesium metal in a bunsen burner flame, to give the atoms sufficient activation energy to react with each other. Once they have, there is a change from Mg and O2 to a new substance, MgO.
Every reaction has its own activation energy - the amount of energy required for the two molecules to react. If the molecules collide with more than the required amount of activation energy, then the reaction goes ahead. We say that the molecules had sufficient energy to overcome the activation energy for that reaction. Remember, activation energy is a number. Reactions with a high activation energy are harder to get to work, happen very slowly, or happen rarely, because it isn't often that the molecules have enough energy to react that way. Reactions with a lower activation energy are easier to get to work. If the activation energy is low enough, then we say that the reaction happens spontaneously - because so little energy is needed that the molecules already have all the energy they need, and can just get on with it.
But it's not just about two molecules colliding with lots of energy - if they collide into the wrong part of each other, a reaction's still not going to happen. It's alright if we are talking about a tiny reactive species like an H+ ion - it is only a proton, so small that any direction you bump into it from, you can't help but hit it where it reacts. But what about a bigger molecule that only reacts at one end, like an alkene? Alkenes have a double bond, which is their reactive site. The rest of the alkene is an alkane straight chain which is not very reactive at all. For something to react with the alkene, it has to bang into the reactive, double-bond part. Otherwise nothing will happen. This is what people mean when they say that the molecules must collide in the correct orientation.
So to sum it up, a reaction takes place between two molecules when they collide in the correct orientation, and with sufficient energy to overcome the activation energy for that reaction.
Altering the rate of a reaction - how does it work?
Sometimes a reaction is working faster or slower than you'd like, and you want to be able to change that. The ways to affect reaction rate are listed at the top of this page - change concentration, surface area, temperature, or add (or remove) a catalyst if there is one.
Understanding how these things work is easy once you understand how reactions work in the first place - overcoming the activation energy and colliding the right way around. Let's look at how these four factors work to make an impact on the speed of a reaction.
Temperature is to do with heat, and heat is a form of energy. Particles are always moving (even solids are bouncing on the spot). The more energy they have, the more spread out they get and the faster they move. If you increase the temperature, you give the particles more energy, so they can move around faster and spread out further. And the faster they zing around, the more likely they are to bump into each other, and with enough energy to overcome the activation energy.
Conversely, if you want to slow down the reaction, cooling the particles down takes away some of their energy, so they move slower and travel less, making it less likely they'll collide with the right particles with enough oomph to react.
Concentration is about how much of your substance there is, floating around in the water (most of your reactions will be with molecules dissolved in water - the principle's the same for other solvents too). The more concentrated a solution is, the more of your stuff there is compared to the water. The weaker and less concentrated a solution is, the less of your stuff there is compared to water. Think of it as a glass of orange juice. A weak orange juice has not very much orange and lots of water. The orange juice is at a low concentration. A strong juice is mostly orange and hardly any water. This orange juice is at a high concentration.
In Chemistry, concentration is mostly talked about in the units moles per litre (molL-1). The higher the number, the more concentrated the solution is. For example, if you had two bottles of hydrochloric acid in front of you, one at 1 molL-1 and one at 0.1 molL-1, which one's the stronger concentration? The bigger number, 1 molL-1.
Right, back to the particles. If a solution is more concentrated, then there are more particles of your reactants floating around among the water molecules. A weaker concentration has less particles of your reactants floating around in it. And by the laws of chance, the less particles there are floating around, the less chance there is that they will bang in to each other, and do that reaction thing. And vice versa, the more particles there are in that amount of liquid, the better the chance that they will collide.
Surface area means how much of the surface can be seen, reached, touched, or banged into at any particular time. Is now a good time to mention that this factor only counts if one of your reactants is a solid? Think about it - if something is dissolved in a solution, all its particles are floating around free already. If something is a solid, only the particles on the very outside can be bumped into by the other reactant molecules. So to speed things up, you want as much of your solid stuff exposed as possible. How to do this? Break it up into small chunks, the smaller the better. In fact, grinding it into a powder gives you many, many more exposed surfaces.
But wait! Is the ground up powder making everything react too fast? Is your mixture overflowing everywhere? Then you need to slow things down by using bigger chunks. Pellets, granules, marbles, pea-sized bits - these have a smaller surface area than a powder. That means that less of your reactant particles will be exposed to the other reactants, and so less particles will be available to be collided with. Phew!
A catalyst is another chemical that you can add to a reaction mixture to make a reaction happen faster. Sometimes the same reaction without a catalyst is so slow, it doesn't seem to be happening at all! The catalyst chemical isn't a reactant, or a product. It doesn't take part in the reaction. What it does do is help to set up the reactant molecules so that they collide the right way round more frequently, thus increasing the number of useful collisions that end up happening. Another way to describe it is that it provides an alternative reaction pathway that has the effect of lowering the level of activation energy required for that reaction.
Not all reactions have a matching catalyst, so this factor can't always be used to affect reaction rate. Those reactions that do use a catalyst will react very, very slowly if the catalyst isn't around. You could say that you can slow down the reaction by removing the catalyst, but it often seems to stop the reaction completely, not just slow it down.
You may hear that catalysts aren't ever used up. Well, that is true in the sense that the catalyst isn't a reactant, and so it doesn't run out or get used up, like a reactant does. So technically, there should be the same amount of catalyst at the end of a reaction as there was at the beginning. However, a lot of the reactions used by industry to mass produce important chemicals need a catalyst, and quite a few of these need to replace their catalysts over time. These reactions are being carried out on a large scale, many times bigger than a beaker and a stirring rod, like in our labs. Depending on the reaction and the catalyst, some catalysts eventually get affected by catalysing all those reactions. They can get worn out, and stop performing so well. Or a by-product of the reaction might interfere with the catalyst - this can be called "poisoning the catalyst". When this happens, the catalysts get replaced with new ones. Some companies replace their catalysts frequently, so that they are always working to full productivity.
Reactions happen when two molecules collide in the correct orientation and with sufficient energy to overcome the activation energy of that reaction.
Activation energy is the level of energy required before a reaction will occur between two molecules.
How fast a reaction occurs can be changed by altering the surface area, temperature or concentration of the reactants, or by using a catalyst if there is one for that reaction.
Speed up a reaction by increasing the temperature of the reaction mixture (heating it up), increasing the surface area of any solid reactants (grinding them up), increasing the concentration of any solutions (pick a stronger concentration - more molL-1), or adding a catalyst if there is one for that reaction.
Slow down a reaction by decreasing the temperature of the reaction mixture (putting it on ice), decreasing the surface area of any solid reactants (using bigger chunks), decreasing the concentration (pick a weaker concentration - less molL-1), or don't add that catalyst.
Increasing the temperature causes the particles to move with more energy, increasing the chance of a favourable collision in the correct orientation with sufficient energy to overcome the activation energy of the reaction.
Decreasing the temperature causes the particles to move with less energy, decreasing the chance of a favourable collision (you know, the one with both particles in the correct orientation and with sufficient energy to overcome the activation energy of the reaction).
Increasing the surface area exposes more particles to the other reactant particles, increasing the chances of a favourable collision (yep, that one with both particles in the correct orientation and with sufficient energy to overcome the activation energy of the reaction).
Decreasing surface area means that less particles will be exposed to the other reactant particles, increasing the chances of a favourable collision (you guessed it, that one with both particles in the correct orientation and with sufficient energy to overcome the activation energy of the reaction).
Increasing concentration of the reactants means that there are more particles of reactant present in the reaction mixture, increasing the chances of a favourable collision (do you think that would be the one with both particles in the correct orientation and with sufficient energy to overcome the activation energy of the reaction?).
Decreasing concentration of the reactants means that there are less particles of reactant present in the reaction mixture, decreasing the chances of a favourable collision (any chance of it being the one with both particles in the correct orientation and with sufficient energy to overcome the activation energy of the reaction?).
Adding a catalyst provides an alternative pathway for the reaction and in doing so lowers the activation energy of the reaction.
Catalysts aren't used up by a reaction, although in big scale reactions in industry they sometimes need replacing eventually.