Chemical Nature of the Reacting Substances
Some substances react faster than others. For example, potassium and calcium, which are next to each other in the fourth row of the periodic table, both react with water to form H2 gas and a basic solution:
2K(s) + 2H2O(ℓ) ⟶ 2KOH(aq) + H2(g)
2Ca(s) + 2H2O(ℓ) ⟶ 2CaOH(aq) + H2(g)
As the video below shows, calcium reacts at a moderate rate, whereas potassium reacts so rapidly that the reaction is almost explosive. One factor affecting these different rates is that the reactions involve loss of electrons from potassium or calcium atoms, and potassium has a smaller first ionization energy, making loss of an electron easier. In other words, the smaller IE1 of potassium makes the activation energy of the potassium reaction lower than that of the calcium reaction.
Chemical reactions typically occur faster at higher temperatures. At higher temperatures, the rate constant is larger, as shown by the Arrhenius equation. Therefore, assuming the concentrations of reactants are the same, a larger rate constant means a faster reaction. For example, methane (CH4) does not react rapidly with air at room temperature, but strike a match and POP!
Reaction rates usually increase when the concentration of one or more of the reactants increases. For example, calcium carbonate (CaCO3) deteriorates as a result of its reaction with the pollutant sulfur dioxide (SO2). Specifically, sulfur dioxide reacts with water vapor to produce sulfurous acid:
Sulfurous acid then reacts with calcium carbonate:
The rate of the overall reaction depends on the concentration of sulfur dioxide in the air. In a more polluted atmosphere where the concentration of sulfur dioxide is higher, calcium carbonate deteriorates more rapidly.
In another example, a cigarette burns slowly in air, which contains about 21% oxygen by volume. Put it in pure oxygen and the rate of the reaction accelerates, as shown in the video below.
Presence and Concentration of a Catalyst
A catalyst is a substance that increases the rate of a chemical reaction by providing an alternative reaction pathway but is not consumed by the reaction. The greater the concentration of a catalyst the more the catalyst can speed up a reaction. How catalysts work will be discussed in detail later on. Watch the video below to see how a catalyst can speed up the decomposition of hydrogen peroxide to form oxygen and water.
If a reaction occurs at a surface, an increase in the surface area of the intersection of two phases (such as the surface of a solid in contact with a gas) can increase the rate. A finely divided solid (like a powder) has more surface area available for reaction than one large solid piece of the same substance. For example, large pieces of wood smolder, smaller pieces burn rapidly, and sawdust burns explosively. The video below shows how large pieces of iron can be held in a burner flame for a long time and hardly react, whereas iron powder blown into the flame sparkles as the tiny particles burn.