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Why most collisions fail, the two tests a collision must pass, and rate as collision frequency times the share that succeed.
Paper packet. Every task here also exists on screen, where it is checked automatically; answers written on paper are not assessed by Nydus. When you are back at a device, enter your answers there.
By the end of this lesson you will be able to explain a reaction rate using the collision model: particles must collide with at least the activation energy and a suitable orientation, so the rate depends on how often they collide and on the share of collisions that succeed. You will be able to calculate how many collisions succeed from the shares that pass each test, read the activation energy off an energy profile, and explain why a reaction slows down as it goes on.
You can measure a rate and read one off a curve, and you have seen that a rate falls as the reactants are used up. From Chemistry 1 you have the particle model — substances are made of particles in constant motion — and the energy profile, whose hump is the activation energy. This lesson joins the two: it explains a rate using particles and the hump.
A collision is two particles meeting. The activation energy, $E_a$, is the minimum energy a collision must carry for the particles to react. Orientation is which way round the particles are when they meet. A successful collision is one with at least the activation energy and a suitable orientation; it breaks bonds and makes products. Collision frequency is the number of collisions per second in a stated volume.
Particles in a gas or a solution are moving all the time and colliding all the time. In air at room temperature a single molecule is hit several thousand million times a second. If every collision caused a reaction, every mixture that could react would do so in a blink.
Particles collide far more often than they react. Only a collision with at least the activation energy, and with the particles meeting the right way round, turns reactants into products, so anything that changes the rate changes either how often particles meet or what share of the meetings succeed.
So a collision has to pass two tests.
The energy test. Reacting means breaking bonds, and breaking bonds takes energy. A collision that carries less than the activation energy simply bounces: the particles fly apart unchanged. On an energy profile the activation energy is the height of the hump above the reactants, and only collisions with at least that much energy get over it.
The orientation test. Particles are not points. For a hydrogen atom to be pulled off one molecule by another, the atoms that must bond have to meet; a collision that hits the wrong end of the molecule does nothing however hard it is.
Put together:
$$\text{rate} \propto (\text{collisions per second}) \times (\text{share with enough energy}) \times (\text{share correctly oriented})$$
That one line explains the rest of the unit. A change can speed a reaction up only by making particles meet more often or by making a bigger share of meetings succeed.
Another way: picture
Think of trying to post a letter through a letterbox by throwing it from across the room. Most throws miss (wrong orientation), and many that arrive are too weak to push the flap open (not enough energy). Throwing more letters a minute gets more through; throwing harder gets more through; a wider slot with a lighter flap gets more through. Those are concentration, temperature and a catalyst.
Another way: steps
To explain any change in rate with the collision model:
The numbers are worth seeing once, because they are what make the model convincing.
Suppose a small volume of a gas mixture has 8 000 000 collisions a second. If the activation energy is high enough that only 1 collision in 2000 carries it, then $8\,000\,000 \div 2000 = 4000$ collisions a second pass the energy test. If only 1 in 4 of those is lined up the right way, $4000 \div 4 = 1000$ succeed.
A thousand successful collisions out of eight million: one in eight thousand. In real reactions the share can be one in a million million or smaller. That is why reactions have measurable rates at all, and why a small change in the share that succeed — which is what temperature and catalysts do — can change a rate enormously.
Lesson one showed the rate falling from one interval to the next. The collision model explains it in three steps. As the reaction proceeds, reactant particles are turned into products, so fewer reactant particles are left in the same volume. Fewer particles means fewer collisions per second between them. The share that succeed has not changed, because the temperature and the activation energy have not, so fewer collisions per second means fewer successful collisions per second, and a lower rate. When one reactant runs out, its collisions stop altogether and the rate is zero.
Notice what the explanation does not say. Particles do not get tired, run out of energy or slow down because the reaction is old: their speeds are set by the temperature. If anything, an exothermic reaction warms its mixture as it goes, which pushes the other way.
Every collision makes products. Almost none do. Most collisions lack the activation energy or meet the wrong way round and bounce apart unchanged.
Particles collide only when something makes them. Particles in a gas or a solution collide constantly just by moving. What a spark or a flame provides is energy, not meetings.
A reaction slows because the particles get tired. It slows because there are fewer reactant particles to meet. The particles' speeds depend on the temperature.
Faster particles always mean a faster reaction, whatever else changes. A faster reaction needs more successful collisions per second. Speed matters because it changes the share that carry the activation energy — the subject of the lesson on temperature.
A mixture has 6 000 000 collisions a second; 1 in 3000 has the activation energy, and 1 in 5 of those is correctly oriented.
Two tests, each keeping a share of what reaches it.
Energy test: $6\,000\,000 \div 3000 = 2000$ a second.
Apply the energy share to all the collisions.
Orientation test: $2000 \div 5 = 400$ successful collisions a second.
Apply the orientation share only to the collisions that already passed the energy test.
Two identical pieces of magnesium are put in hydrochloric acid, one in acid twice as concentrated as the other.
Name the change first: more acid particles in the same volume.
More acid particles in the same volume hit the magnesium surface more often, so collisions per second go up. The temperature and activation energy are the same, so the share that succeed does not change.
Decide which factor the change moves, and say that the other one stays put.
So there are more successful collisions per second and the more concentrated acid reacts faster.
The rate follows the number of successful collisions per second.
Energy test: $9\,000\,000 \div 1500 = 6000$ a second.
The energy share applies to every collision.
Orientation test: $6000 \div 2 = \ldots$
…$3000$ successful collisions a second, out of nine million.
In a small volume of a gas mixture there are $9\,000\,000$ collisions every second. Only one collision in $5000$ has at least the activation energy, and only one in $4$ of those has the particles lined up the right way. How many successful collisions happen each second?
Answer: per second
A mixture of methane and air above a gas hob can sit for years without reacting, although its molecules collide thousands of millions of times a second. A single spark sets the whole mixture off. What does the collision model say is going on?
A reaction has an activation energy of $90$ kJ/mol. Five collisions are described by their energy (per mole of such collisions) and by whether the particles meet the right way round. Decide whether each one reacts.
| collision energy, in kJ/mol | lined up the right way? | does it react? | |
|---|---|---|---|
| collision 1 | 110 | yes | |
| collision 2 | 80 | yes | |
| collision 3 | 120 | no | |
| collision 4 | 90 | yes | |
| collision 5 | 65 | no |
The rate of a reaction is how often particles collide multiplied by the share of collisions that succeed. Match each change to what it does in that model.
| more collisions per second in the solution | more particles of the solid exposed to be hit | a larger share of collisions have enough energy | a lower energy barrier for the same collisions | |
|---|---|---|---|---|
| dissolving more acid in the same volume of water | ||||
| grinding a lump of solid into a powder | ||||
| warming the mixture | ||||
| adding a catalyst |
A learner explains why magnesium ribbon reacts more and more slowly in acid. Select every sentence that is wrong.
This task has no paper form; do it on a device.
Put the events of one successful reaction between a nitrogen monoxide molecule and an ozone molecule in order.
Number the steps in order (write the number in the box):
An atmospheric chemist estimates that in one cubic centimetre of air near a busy road, nitrogen monoxide and ozone molecules collide about $3\,000\,000$ times a second, while $12$ of those collisions actually produce nitrogen dioxide. One collision in how many succeeds?
Answer: collisions
Lesson test: one question per skill, one attempt each, no hints. Your answers are checked when you submit.
An energy profile for a reaction shows the reactants at $65$ kJ and the products at $99$ kJ. The activation energy is $63$ kJ. Give the energy at the top of the hump, and the activation energy for the reverse reaction.
An energy profile: reactants on the left, one hump, products on the right.
Energy at the top of the hump, in kJ:
Activation energy of the reverse reaction, in kJ:
You can say what a collision needs in order to react and use that to explain a rate. Say out loud why a mixture of hydrogen and oxygen can sit unreacted for years, and what a spark changes. Next: the first of the four ways to change a rate — concentration.
10. Your turn: a gas mixture has 9 000 000 collisions a second. One in 1500 carries the activation energy and one in 2 of those is correctly oriented. How many succeed each second?, step 3