A chemical reaction can proceed only when reactant particles collide under the right conditions of energy and orientation. In macroscopic terms, the observable rate of reaction depends on several factors:
1. Concentration (or pressure) of reactants
2. Temperature of the system
3. Surface area of reactants (for heterogeneous reactions)
4. Physical state and nature of reactants
5. Presence of a catalyst
6. Solvent effects (for reactions in solution)

2.1 Concentration (or Pressure)

2.1.1 Definition
- In solution-phase reactions, concentration is measured in moles per liter (M).
- In gas-phase reactions, it is more convenient to use partial pressure (e.g., atmospheres or kilopascals) to represent how much of each gaseous reactant is present.

2.1.2 Collision Frequency and Concentration
According to collision theory, the frequency of collisions between species A and B in the gas phase is proportional to the product of their concentrations (or partial pressures). This means:
- If you double the concentration of A while keeping B constant, the number of A–B collisions per unit time doubles.
- If you double both the concentrations of A and B, the collision frequency quadruples.

2.1.3 Effect on Reaction Rate
For many simple (elementary) reactions, the rate law reflects this collision frequency. For instance, for the bimolecular reaction
A + B → products

the simplest possible rate law is
Rate = k × [A] × [B]
where k is the rate constant. In gas-phase notation, one can similarly write
Rate = k × PA × PB
if PA and PB are the partial pressures of A and B, respectively. Experimental verification is always required, since more complex mechanisms can lead to different dependencies on concentration.

2.2 Temperature

Temperature—measured in kelvins (K)—reflects the average kinetic energy of particles in a system. Raising the temperature has two main effects on reaction rate:
1. It increases the average speed of molecules, leading to more frequent collisions per unit time.
2. It increases the fraction of collisions whose energy exceeds the activation energy Ea (see Section 3).

Empirically, many reaction rates roughly double for every 10 K–20 K increase in temperature near room temperature. We will see quantitatively how the Arrhenius equation explains this behavior in Section 3.5.

2.3 Surface Area (for Heterogeneous Reactions)

When a reaction involves a solid reactant (or a solid catalyst) together with reactants in a different phase (gas or liquid), the exposed surface area of the solid controls the rate:
- A larger surface area provides more sites for reactant molecules to collide with the solid.
- Grinding a solid into a fine powder dramatically increases its surface area, causing the reaction to proceed more quickly than if the same mass were present as a single large piece.
For example, magnesium turnings (small, powdered pieces) react with hydrochloric acid much faster than a thick piece of magnesium ribbon, because the turnings expose far more metal surface to the acid.

2.4 Physical State and Nature of Reactants

2.4.1 Phase
- Gaseous reactants mix and diffuse rapidly, often corresponding to faster reaction rates (neglecting activation barriers).
- Liquid-phase reactions depend on diffusion rates in solution; stirring or agitation helps remove concentration gradients and speeds up mixing.
- Solid-phase reactions typically proceed by diffusion of reactants into or out of a solid lattice, which can be slower compared to gas or liquid processes.

2.4.2 Chemical Bonds and Molecular Structure
- Bond strength: Stronger bonds (for example, a carbon–carbon triple bond) require more energy to break, often slowing the reaction.
- Molecular complexity and steric hindrance: Large or bulky molecules may have a lower probability of colliding with the correct orientation, reducing the rate.
- Polarity and hydrogen bonding: Highly polar reactants or those that form extensive hydrogen bonds can be stabilized in solution, altering the effective activation energy.

2.5 Presence of a Catalyst

A catalyst is a substance that increases the rate of a chemical reaction by providing an alternate pathway with a lower activation energy Ea, yet it is not consumed in the overall process. Catalysts do not change the thermodynamic equilibrium; they accelerate both forward and reverse reactions equally.

2.5.1 Types of Catalysts
1. Homogeneous catalysts exist in the same phase as the reactants (for example, an acid or base dissolved in water, or a soluble transition-metal complex in organic solvent).
2. Heterogeneous catalysts are in a different phase (for example, a solid surface with gaseous or liquid reactants).
3. Enzymes are biological catalysts—proteins that operate under mild conditions with very high specificity and efficiency.

2.6 Solvent Effects (for Reactions in Solution)

For reactions carried out in solution, the choice of solvent can significantly influence the reaction rate by:
- Stabilizing reactants or transition states through solvation—via hydrogen bonding, dipole interactions, or dielectric screening—thus altering the effective activation energy.
- Changing diffusion rates of reactants (through viscosity), which affects how often molecules collide.
- Altering reaction pathways (for example, favoring ionic versus radical mechanisms).

For instance, in a nucleophilic substitution:
- An SN1 reaction (proceeding through a carbocation intermediate) often runs faster in a polar protic solvent (like water or an alcohol), because such solvents stabilize the carbocation.
- An SN2 reaction (which involves a charged nucleophile attacking a substrate) often runs faster in a polar aprotic solvent (like acetone or dimethyl sulfoxide), since these solvents do not strongly solvate the nucleophile, leaving it more reactive.