12.2 Factors Affecting Reaction Rates

Factors Influencing Reaction Rates

Reaction rates describe how fast reactants turn into products. Several factors control this speed: temperature, concentration, catalysts, surface area, and the chemical nature of the reactants themselves. Most of these come back to one core idea from collision theory: for a reaction to happen, molecules must collide with enough energy and the right orientation.

Temperature

Higher temperatures give molecules more kinetic energy. They move faster, collide more often, and those collisions are more forceful. The result is that more collisions exceed the activation energy threshold, so more of them actually lead to a reaction.

A common rule of thumb: raising the temperature by about 10°C roughly doubles the reaction rate. The precise relationship is captured by the Arrhenius equation:

$k = Ae^{-E_a/RT}$

• $k$ = rate constant
• $A$ = frequency factor (accounts for collision frequency and orientation)
• $E_a$ = activation energy
• $R$ = gas constant (8.314 J/mol·K)
• $T$ = absolute temperature in Kelvin

As $T$ increases, the exponent $-E_a/RT$ becomes less negative, so $k$ increases exponentially. That's why even a small temperature bump can have a big effect on rate.

Concentration

More molecules packed into the same volume means more frequent collisions. This is why increasing the concentration of reactants speeds up most reactions.

The rate law expresses this mathematically:

$\text{Rate} = k[A]^m[B]^n$

• $[A]$ and $[B]$ are molar concentrations of reactants (mol/L)
• $m$ and $n$ are reaction orders, determined experimentally (not from the balanced equation)
• $k$ is the rate constant

If a reactant has an order of 1 (first-order), doubling its concentration doubles the rate. If the order is 2, doubling the concentration quadruples the rate. If the order is 0, changing that reactant's concentration has no effect on rate at all.

Catalysts

A catalyst speeds up a reaction without being consumed. It works by providing an alternative reaction pathway that has a lower activation energy. The reactants and products stay the same, but the energy barrier to get there is smaller, so more collisions are successful at a given temperature.

Two main types:

• Homogeneous catalysts are in the same phase as the reactants (e.g., an acid dissolved in a reaction solution).
• Heterogeneous catalysts are in a different phase, typically a solid surface where gas or liquid reactants adsorb and react (e.g., the catalytic converter in a car).

Enzymes are biological catalysts. They're proteins with specific active sites that bind to particular substrates, making them highly selective for certain reactions.

Physical State and Surface Area

Physical state matters because it determines how easily reactant particles can mix and collide.

• Gas-phase reactions tend to be fast because gas molecules move freely and collide frequently.
• Reactions in solution are often faster than solid-state reactions because the solvent brings dissolved reactants into close contact.
• Solid reactants react more slowly because only the molecules at the surface are exposed. The interior is locked in a lattice structure and can't participate.

Surface area is the key variable for solid reactants. Grinding a solid into a fine powder dramatically increases the exposed surface, giving more sites where collisions can happen. This is why a sugar cube dissolves slowly in water, but powdered sugar dissolves almost instantly.

Heterogeneous catalysts take advantage of this same principle. Porous materials like zeolites and activated carbon have enormous internal surface areas, making them highly effective catalysts and adsorbents.

Chemical Nature of Reactants

Not all molecules are equally reactive. The identity of the reactants themselves affects how fast they react.

Bond strength plays a direct role. Breaking bonds requires energy, so reactants with weaker bonds generally react faster. For example, a C–C single bond is easier to break than a C=C double bond, which in turn is easier to break than a C≡C triple bond.

Polarity and solubility determine how well reactants mix. The "like dissolves like" principle applies here: polar reactants dissolve better in polar solvents (like water), and nonpolar reactants dissolve better in nonpolar solvents (like hexane). Better solubility means the reactants are more evenly distributed and collide more effectively.

Steric hindrance refers to bulky groups on a molecule physically blocking the reactive site. If large substituents surround the spot where a reaction needs to happen, it's harder for another molecule to get close enough to react. For instance, in $S_N2$ substitution reactions, primary substrates (less crowded) react much faster than tertiary substrates (more crowded).

Reaction Mechanisms and Energy Profiles

A reaction mechanism is the step-by-step sequence of elementary reactions that describes how reactants become products. Most reactions don't happen in a single collision. Instead, they proceed through intermediates (species that form and then get used up) and transition states (the highest-energy arrangement along each step).

A potential energy diagram plots energy on the y-axis against the reaction coordinate (progress of the reaction) on the x-axis. From this diagram you can read:

• The activation energy ($E_a$): the energy difference between the reactants and the highest transition state
• The overall energy change ($\Delta E$): the difference between reactant energy and product energy, which tells you if the reaction is exothermic or endothermic

One distinction worth keeping straight: thermodynamics tells you whether a reaction is energetically favorable (will it happen?), while kinetics tells you how fast it happens. A reaction can be thermodynamically favorable but kinetically slow if the activation energy is high. Diamonds converting to graphite is a classic example: it's thermodynamically favorable, but the rate is essentially zero at room temperature.