D13.1 Factors Affecting Reaction Rates
Thus far, we have considered the kinetics of a chemical reaction on the molecular scale: how some collisions between atoms and molecules can lead to chemical reactions, how the activation energy (and hence temperature) influences the rate at which a reaction occurs, how a chemical reaction may actually be sequential elementary reactions that transform reactant molecule(s) to product molecule(s) in a step-wise manner. A reaction coordinate diagram is a visual model that can help us when thinking about these various aspects of a chemical reaction.
But, how do we know if a particular multi-step reaction actually occurs in the way the reaction mechanism describes? In other words, how do we know if a particular reaction energy diagram correctly describes a reaction? Experimentally, it is nearly impossible to follow a single molecule as it undergoes the various transformations from reactant to product. So we have to take a step back, and consider what we can actually observe on the laboratory scale, and connect the macroscopic observations to our understanding of a chemical reaction. It turns out that measuring the effects of varying concentrations of reactants, products, and catalysts provides information that can help determine the sequence of elementary reaction steps by which most reactions occur.
We will explore how we can experimentally measure reaction rates and figure out plausible mechanisms in the upcoming sections. But first, let us briefly and qualitatively consider the various factors that affect the rate of a chemical reaction, some of which (e.g. temperature) we have already discussed.
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:
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.
Temperature
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!
Concentrations
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. Oftentimes, 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.
Surface Area
If a reaction occurs on 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.
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