Anything that will increase the number of effective collisions in a period of time will increase the reaction rate. There are five factors that are commonly known to affect reaction rates. We will investigate each one in terms of collisions.
Factors Affecting Reaction Rates
- the chemical nature of the reactants
- the ability of the reactants to collide with each other
- the concentration of the reactants
- the temperature
- the presence of a catalyst
Chemical Nature of Reactants
The tendency of atoms, molecules, and ions to break and form bonds is the most important factor affecting reaction rates. When a fresh surface of sodium metal is exposed to air, it tarnishes almost immediately. Yet, when iron is exposed to air, it tarnishes (or rusts) much more slowly. Sodium atoms are unstable and tend to lose electrons easily, so when they collide with oxygen atoms, the collision easily produces a product. But the transition element iron is not as unstable and collisions between iron and oxygen atoms are not as effective in producing products.
Ability of the Reactants to Collide with Each Other
Reactions are classified by the states of the reactants into two categories: homogeneous reactions and heterogeneous reactions. An acid and base neutralization reaction, where each substance was aqueous, is an example of a homogeneous reaction. Each reactant is in the liquid state.
The formation of rust from solid iron and oxygen found either in liquid water or gaseous water vapor is an example of a heterogeneous reaction. In these reactions, the reactants are only able to interact at the interface between the phases. For this reason, the area of contact between the phases affects reaction rate. The smaller the particles, the faster the reaction is. Crushing a solid will increase surface are and produce more collisions per second.
Dust explosions are an example of the how increased surface area affects reaction rate. The explosions occur when tiny dust particles of a combustible material such as grain, sugar, or coal are ignited. The increased surface area of the dust particles present an increase in the effective collisions, causing rapid explosions that are a real danger to those that work in these fields.
Watch the video of the aftermath of a grain silo explosion available in the sidebar. When very small grain dust particles accumulate and mix with air in a silo, an explosion can occur.
Concentration of the Reactants
When reactants are more concentrated, more of the molecules, atoms, or ions are present. With that increase comes more chances for collisions to take place. In the image at the right, the green ball must come into contact with a purple ball for a reaction to occur.
By increasing the number of purple balls, effectively increasing their concentration, the probability of an effective collision is larger.
Nearly all chemical reactions speed up when the temperature in increased. Remember that temperature is a measure of the kinetic energy of the particles. With increased kinetic energy, the molecules move faster, which increases the chances of a successful collision. A general rule of thumb is that a reaction rate will double for each 10°C that the temperature is raised. This does not hold true for all reactions.
Presence of a Catalyst
Catalysts increase the rate of a chemical reaction without being used up in the chemical reaction themselves. Enzymes are common catalysts that are continually at work in our bodies. A reaction that is influenced by a catalyst is said to be undergoing catalysis. Important classes in catalysis include acid-base catalysis, surface catalysis, and enzyme catalysis. In order to understand how catalysts work, we need to discuss reaction mechanisms.
In the potential energy diagrams below, the blue curves show the addition of a catalyst. Notice that the catalyst does not change the heat of reaction because it does not change the PE of the reactants or the products. It only changes the activation energy and PE of the activated complex.
The synthesis of ammonia from hydrogen and nitrogen, 3H2(g) + N2(g) → 2NH3(g), is a common example of a catalyzed reaction. The reaction takes place on the surface of an iron catalyst that contains traces of aluminum and potassium oxides. The hydrogen and nitrogen bind to the catalyst and dissociate into atoms. This allows for more effective collisions easier, lowering the activation energy. This process is known as the Haber process.