Factors affecting rate

All six factors reduce to the same two collision-theory levers: how often reactant particles meet, and what fraction of those meetings are energetic enough to succeed (i.e., meet or exceed the activation energy $E_a$).

Surface area matters only when a solid reactant is involved. Breaking a lump into powder does not change how much solid there is — it changes how much of it is exposed. More exposed surface means more particles available for collision with the other reactant at any instant, so the collision frequency rises and the rate rises with it. This is why powdered marble reacts with dilute HCl far faster than a single chip of identical mass.

Nature of reactants is about intrinsic chemistry, not amount or temperature. Reactions between ions in solution are typically near-instantaneous because no covalent bond has to be broken — the species are already primed to combine. Reactions that require breaking a strong covalent bond (as in most organic reactions) are intrinsically slower, because a larger energy barrier separates reactants from products regardless of how concentrated or hot the mixture is.

Concentration raises the rate because more reactant particles packed into the same volume means they collide more often — a purely frequency effect, not an energy effect. The exact mathematical form (rate law and order) belongs to the Rate Law topic; here the qualitative direction is what matters: more concentrated reactants, on average, a faster reaction.

Temperature is the most powerful of the six because it acts on both levers at once — faster-moving particles collide more often, but far more significantly, a larger fraction of them now carry energy $\ge E_a$. A convenient rule of thumb, accurate for many reactions over a modest temperature range, is that rate roughly doubles for every $10\,^{\circ}\text{C}$ rise. Writing $r_1$ for the rate at temperature $T_1$ and $r_2$ for the rate at a higher temperature $T_2$, this rule of thumb is

where: $r_1, r_2$ are the reaction rates at temperatures $T_1$ and $T_2$ (in $^{\circ}\text{C}$), and the exponent counts how many complete $10\,^{\circ}\text{C}$ steps separate the two temperatures. Equation (1) is a rough empirical approximation, not a law derived from first principles — the rigorous temperature dependence is the Arrhenius equation, met later in the chapter.

Catalyst raises the rate by opening an alternative reaction pathway with a lower activation energy. Lowering $E_a$ means a larger fraction of molecules already clear the barrier at the same temperature — the catalyst does not add energy or heat the mixture, it simply makes the barrier easier to cross. A true catalyst is chemically unchanged in mass and composition at the end of the reaction, even though it may take part in intermediate steps of the mechanism.

Radiation accelerates photochemical reactions by supplying energy directly to reactant molecules, either helping them cross the activation barrier or generating reactive intermediates such as free radicals. $\text{H}_2$ and $\text{Cl}_2$ are negligibly slow to react in the dark but combine explosively in sunlight — the absorbed light energy substitutes for what a temperature rise would otherwise have to supply.

Question. A reaction proceeds at a rate of $1.0\times10^{-4}\ \text{mol dm}^{-3}\text{s}^{-1}$ at $20\,^{\circ}\text{C}$. Using the rule that rate approximately doubles for every $10\,^{\circ}\text{C}$ rise, estimate the rate at $60\,^{\circ}\text{C}$.

Step 1. Find the number of $10\,^{\circ}\text{C}$ steps. The temperature rise is $T_2 - T_1 = 60 - 20 = 40\,^{\circ}\text{C}$, so $n = 40/10 = 4$ steps.

Step 2. Apply equation (1): $r_2 = r_1 \times 2^{n} = (1.0\times10^{-4}) \times 2^{4} = (1.0\times10^{-4}) \times 16$.

Step 3. Compute the result: $r_2 = 1.6\times10^{-3}\ \text{mol dm}^{-3}\text{s}^{-1}$.

The estimated rate at $60\,^{\circ}\text{C}$ is $1.6\times10^{-3}\ \text{mol dm}^{-3}\text{s}^{-1}$ — sixteen times the rate at $20\,^{\circ}\text{C}$.

Competitive stretch

If the same reaction mixture were also given a freshly powdered (rather than lump) solid reactant and a $\text{MnO}_2$-type catalyst at the same $60\,^{\circ}\text{C}$, the temperature contribution above ($16\times$) would combine with independent boosts from surface area (more collision frequency) and catalyst (lower $E_a$, larger effective fraction) — the true observed rate would exceed the $16\times$ temperature-only estimate, since real reactions rarely have only one factor changing at a time.

• Treating "nature of reactants" as vague filler — it is a distinct, examinable factor referring to intrinsic bond type and physical state, not to be confused with concentration or temperature.
• Believing a catalyst is consumed or permanently altered — it is chemically unchanged at the end, even though it may participate in intermediate steps.
• Applying "doubles every 10°C" additively instead of multiplicatively — a $30\,^{\circ}\text{C}$ rise gives $\times 2^3 = \times 8$, not $\times 2 \times 3 = \times 6$.
• Assuming surface area matters for reactions between two liquids or gases already homogeneously mixed — it is relevant only when a solid or heterogeneous interface is present.

Ask the tutor: ask your AI tutor to give you a new everyday scenario and have you identify which of the six factors — and which lever, frequency or effective fraction — explains the rate change.