Collisions between particles

This is asking for the exact meaning of activation energy ($E_{\text{a}}$) as used in collision theory (rates of reaction). It’s about what energy condition must be met for a collision to turn into a reaction.

In IB Chemistry kinetics (collision theory), particles only react when they collide (1) with the right orientation and (2) with enough energy to break bonds and form new ones.

That required “energy barrier” is defined as the minimum energy needed for a successful collision.

Here’s why the other words are not the definition:

• Total energy: means all energy in the system added up, not the barrier a single collision must overcome.
• Maximum energy: reactions don’t have an upper energy cutoff; having more than enough energy doesn’t prevent reaction.
• Average energy: many particles sit near the average kinetic energy at a given temperature, but only those above $E_{\text{a}}$ can react—so $E_{\text{a}}$ is not an average.

Using the textbook/markscheme wording, activation energy is the minimum energy needed for a collision to lead to reaction.

Collision theory describes reactions as needing collisions with correct orientation and enough energy. That energy requirement behaves like a barrier that must be overcome.

So we need the word that best describes a “threshold” energy.

If $E_a$ were a maximum, then collisions with more energy than that would fail to react.

But higher-energy collisions are generally more likely to be successful, not less. There isn’t an upper cap for a collision to be allowed to react, so “maximum” can’t be right.

The average energy of particles depends on temperature (as shown by a Maxwell–Boltzmann distribution).

Activation energy ($E_a$) is a property of the reaction pathway itself—changing the average particle energy doesn’t redefine $E_a$. Therefore it isn’t an “average.”

Total refers to the sum of energies a particle/system has.

Even though particles have a total energy, the definition of $E_a$ is about the required amount to start bond breaking/forming in a collision. “Total” describes a state, not a minimum requirement, so it doesn’t fit.

That leaves minimum, which matches the idea of a barrier: collisions must have at least $E_a$ to react.

As stated in markschemes: $E_a$ is the minimum energy colliding particles must have for a reaction to occur.

Think of an energy profile diagram showing potential energy from reactants to products.

You can imagine the reactants starting in a valley and needing to reach the top of a hill (the transition state) before products can form.

The Activation Energy ($E_{\text{a}}$) is the vertical energy difference from the reactants’ level up to the peak of the curve.

In collision terms, this is the barrier: particles must collide with enough kinetic energy to get over that peak.

If particles collide with energy higher than the hill height, the reaction still proceeds—so there is no maximum requirement.

Also, the average particle energy is often below the barrier, which explains why only a fraction of collisions lead to reaction at a given temperature.

Although particles have a total energy, the diagram highlights the required barrier height for reaction, not the sum of all energy present.

So “total” doesn’t capture what activation energy is meant to represent.

Any collision with energy below the barrier won’t reach the transition state and will not react—it just results in particles separating again.

So $E_{\text{a}}$ is the minimum energy needed for a collision to cause a reaction.

The sentence says that raising temperature "lowers the activation energy ($E_{\text{a}}$)." We must judge whether that matches what activation energy means in Topic 3: Collision theory and how temperature affects reacting particles.

The activation energy ($E_{\text{a}}$) is the smallest energy that two colliding particles need for a reaction to happen.

Most importantly, $E_{\text{a}}$ belongs to the reaction itself (its mechanism/pathway). It is like the size of the energy “hump” reactants must get over. For a particular reaction, this value does not change just because you warm up or cool down the mixture.

Heating the mixture increases the average kinetic energy of the particles.

From the Maxwell–Boltzmann distribution, increasing temperature increases the fraction of particles with energy at or above the activation energy ($E\ge E_{\text{a}}$).

A helpful picture is a jump challenge:

• Temperature affects how much energy the jumpers (particles) typically have.
• Activation energy is the fixed height of the hurdle.

Warming up helps more jumpers clear the hurdle, but it does not lower the hurdle itself.

We can cross off the options that conflict with collision theory:

• Option A is wrong because particle kinetic energy increasing is not the same thing as the reaction barrier changing. The barrier is about the reaction pathway; kinetic energy is about the particles.
• Option C is wrong because $E_{\text{a}}$ does not have an inverse relationship with temperature; it does not depend on temperature.
• Option D is wrong because temperature changes the energy distribution (the Maxwell–Boltzmann curve shifts), not only the collision rate.

The statement is false. Increasing temperature changes particle energies, not the activation energy value. The only common way to reduce $E_{\text{a}}$ is to add a catalyst, which offers a different pathway with a lower barrier.

The claim says that raising temperature reduces the activation energy ($E_{\text{a}}$). To check this, we must keep two ideas apart: the kinetic energy particles have versus the minimum barrier energy required for reaction. This distinction is central to Collision theory and Chemical reaction kinetics.

A Maxwell–Boltzmann distribution describes how particle kinetic energies are spread out at a given temperature.

• The x-axis shows Kinetic Energy ($E$).
• The y-axis shows the Number of Particles.

Even at one temperature, particles do not all have the same energy. The particles able to react are those with energies to the right of the vertical line at the activation energy ($E_{\text{a}}$), i.e., the area under the curve where energy is high enough.

If temperature increases from $T_1$ to $T_2$ ($T_2>T_1$):

1. The most common energy shifts rightward (higher average kinetic energy).
2. The curve spreads out and becomes lower and wider so the total area (total particles) stays the same.
3. The key point: the activation energy ($E_{\text{a}}$) is tied to the reaction route, so the vertical $E_{\text{a}}$ line on the graph stays in the same place.

Activation energy ($E_{\text{a}}$) is the minimum energy for an effective collision, set by which bonds must be broken and formed.

Heating increases particle motion and energy, so a greater proportion of particles can reach or exceed $E_{\text{a}}$, but the barrier height (the $E_{\text{a}}$ value) does not change. A catalyst is what can reduce $E_{\text{a}}$ by offering a different pathway.

Based on the distribution model:

• Option A is incorrect: the barrier is a property of the reaction, not something that drops just because particles move faster.
• Option C is incorrect: there is no rule that $E_{\text{a}}$ is inversely linked to $T$; $E_{\text{a}}$ does not depend on temperature.
• Option D is incorrect: temperature changes the distribution of energies (the curve shifts), not only how often collisions occur.

The statement is false. Temperature increases average kinetic energy, but activation energy stays fixed. On a Maxwell–Boltzmann diagram, the curve moves while the $E_{\text{a}}$ line does not.

The claim says that a higher temperature makes the activation energy ($E_{\text{a}}$) smaller. To judge this, we use the collision theory definition.

Activation energy is the minimum energy needed for a collision to lead to reaction, and it depends on the bonds involved. In IB, we treat $E_{\text{a}}$ as fixed for a given reaction unless a catalyst changes the pathway.

Option A says the claim is true because more kinetic energy lowers the barrier.

That mixes up two different things. Heating does raise the particles’ average kinetic energy ($KE$), but it does not reduce the energy barrier ($E_{\text{a}}$). Like a high-jump event: giving the jumper more speed doesn’t lower the bar; it only improves the chance they clear it.

Option C claims $E_{\text{a}}$ must drop as temperature rises (an inverse proportion).

There is no evidence for that. In the Arrhenius model taught later, $E_{\text{a}}$ is not controlled by $T$; it is a characteristic of the reaction pathway, so it stays the same when temperature changes.

Option D correctly says the original statement is false, but the reasoning is wrong. It suggests temperature only increases collision frequency and does not change the energy spread.

However, the Maxwell–Boltzmann distribution shows that heating shifts and flattens the curve, meaning a larger fraction of particles now have $E\ge E_{\text{a}}$.

With the other choices removed, the correct explanation is that increasing temperature increases average kinetic energy, leading to more successful collisions (those with $E\ge E_{\text{a}}$), while $E_{\text{a}}$ itself remains unchanged.

We are asked to describe the particle-level effect of increasing pressure by reducing the volume. This links directly to collision theory, which says reactions depend on particles colliding (with enough energy and the right orientation).

When a gas is compressed into a smaller space, the main thing that changes is how much room the particles have to move around in, not what the reaction “requires” energetically.

If the volume ($V$) is decreased while the number of particles ($n$) stays the same, the gas becomes more concentrated. Concentration is:

$C=\frac{n}{V}$

With a smaller $V$, the fraction becomes larger, so $C$ increases. In particle terms, there are more particles in each unit of space, meaning the particles are effectively pushed closer together.

Now we can dismiss the options that describe changes in energy rather than spacing:

• Option A: Activation energy ($E_a$) depends on the reaction itself (bond energies and pathway). Changing container volume doesn’t alter $E_a$; a catalyst is what can lower it.
• Options B and C: Particle speed and kinetic energy ($E_k$) depend on absolute temperature ($T$). Because no temperature change is mentioned, we assume $T$ stays constant, so average speed and individual kinetic energy do not increase.

With particles packed into a smaller volume, the particle density is higher. Because the particles are closer together, they collide more frequently.

In collision theory terms, this means the collision frequency increases. More collisions each second increases the chance of successful collisions, so the reaction rate goes up.

After removing answers about changing activation energy or particle speed/energy, the remaining correct idea is that compression makes particles closer together, leading to more collisions.

The situation describes pressure increasing because the volume of the container is reduced.

In collision theory, changing volume mainly changes how “crowded” the gas particles are. We can solve this by eliminating options that depend on variables that are not being changed (like activation energy or temperature).

Activation energy ($E_a$) is the minimum energy needed for a collision to lead to reaction.

Importantly, the activation energy barrier is set by the reaction pathway. It can be lowered by a catalyst that provides an alternative pathway, but not by compressing the gas. Decreasing volume does not reduce the energy needed to break and form bonds.

Options B and C describe particles moving faster or having more kinetic energy. From kinetic molecular theory, average speed and average kinetic energy are tied to the absolute temperature ($T$) in Kelvin.

Since the question does not say temperature changes, we treat $T$ as constant. Therefore, particles do not gain extra kinetic energy and do not have a higher average speed, so both options can be eliminated.

Eliminating A, B, and C leaves Option D. When volume is reduced, the same number of gas particles occupy a smaller space.

That raises the concentration (particles per unit volume). With particles closer together, they collide more often, so collision frequency increases. This is the correct molecular-level explanation for increasing pressure by decreasing volume.