Factors Affecting Rates of Reactions

Nature of Reactants

There are several points to consider when we examine how the properties of the reactants affects reaction rate.

§  During chemical reactions, chemical bonds are broken and new bonds are formed. The nature (or type) of these chemical bonds - and how readily they are broken and formed - plays a critical role in the rate of a reaction. When the reaction involves primarily the exchange of electrons (as occurs in Redox reactions, the topic of our last unit of study), reactions tend to be very rapid.

For example, consider this very fast double displacement reaction that involves the formation of a yellow precipitate, barium chromate:

Ba(NO₃)₂(aq) + Na₂CrO₄(aq) → BaCrO₄(s) + 2 NaNO₃(aq)

We can write the net ionic equation for this reaction:

Ba²⁺(aq) + CrO₄^2- → BaCrO₄(s)

Reactions such as this that involve ions in solution tend to be very rapid.

§  As you might expect, however, if we had combined solid barium nitrate with solid sodium chromate, the reaction would be so slow that we would not be able to detect it. The phase of the reacting particles is important. Reactants in solution, liquids, and gases will react much faster than solids.

§  Closely related to phase is the amount of exposed surface area of the reacting particles. If you put a lit match to a large log you wouldn't get much of a fire started. However, if you took the same log and increased its surface area by chopping it into kindling, you'd have much more success in starting your fire.

§  Reactions involving covalently-bonded molecules tend to be slow unless highly exothermic. The decomposition of hydrogen peroxide into hydrogen gas and oxygen gas is a spontaneous reaction, but not a particularly fast one.

H₂O₂(l) → H₂(g) + O₂(g)

§  Finally, we will include stirring as another factor that increases the rate of a reaction. Stirring causes the particles to collide more frequently since they will come into contact with one another more frequently.

One other thing - notice we are discussing the nature of the reactants here. Characteristics of the products - whether they are solids or gases, ionic or covalently-bonded, etc., are important in determining reaction rate.

Try these practice questions before continuing to the next section.

Concentration & Pressure

Increasing the concentration of any one or more of the reactants will usually (but not always) increase the rate of a reaction if the reactants are all in the same phase (solid, liquid, gas, or aqueous).

Why? Concentration is a measure of how many particles are in a given volume (see note). By packing more particles in the same space, collisions will occur more often, thus increasing the rate of the reaction.

Remember our thought experiment with blindfolded students in a room? You would demonstrate how concentration increases reaction rate by packing more  blindfolded students into that room.

The concentration of gases is typically increased by decreasing the volume (making the container smaller). This forces the particles closer together, thus increasing their concentration.

Changing the concentration, or increasing the pressure, does not affect the amount of energy particles collide with. It simply increases the rate by increasing the frequency of collisions.

Changes in concentration also explain why the rate of a reaction slows down as a reaction proceeds. As the reactant particles collide with one another to form products, there are fewer and fewer reactant particles as the reaction continues. This lowers their concentration (fewer particles in a given volume), thus the rate declines.

Complete thePractice Questionsand checkyour answers. Then complete theassignment.

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Note:

Concentration (as was mentioned in section 1-3) is commonly measured as molarity, with the symbol M (the units for molarity are mol / L  or mol·L^-1).

A 6 M solution (containing 6 moles of a substance dissolved in 1 litre of solution) is more concentrated than a 1 M solution (1 mole of the substance dissolved in 1 litre of solution).

Temperature

The rate of almost all chemical reactions increases if the temperature of the system is increased. A general rule of thumb is that the rate will double with an increase of 10°C (but "rules of thumb" are not iron-clad rule; exceptions exist).

The collision theory helps us understand why this occurs. 

It is important to remember what this curve represents - all particles in a sample of matter are in constant motion because they possess kinetic energy, but not all particles have the same amount of KE - some have only a little energy and are slow moving, some have a lot of energy and move fast, but most are somewhere in the middle.

In this unit we have added the idea of threshold energy - only particles having at least as much energy as the threshold energy have sufficient energy for a successful collision leading to a reaction.

Why does increasing temperature increase reaction rate?

Students often predict that increasing temperature will cause the threshold energy to be lowered:

but that is NOT the case! The amount of energy required for successful collisions for a particular reaction is not changed when temperature increases.

Instead . . .

When temperature is increased, all of the particles in the sample of matter now have more kinetic energy, and hence are moving faster. There are still some that are slower than the rest, others that are really fast, but most still in the middle. But the average kinetic energy of the entire sample increases.

 

T1 represents the energy distribution at the lower temperature; T2 is the distribution at the higher temperature. The threshold energy is unchanged, but for the new curve (T2) more particles will be above the threshold energy, and thus will have successful reactions.

This will increase the rate for two reasons: Because all particles are moving faster, there will be more frequent collisions, and thus an increase in rate. Remember our thought experiment, with blindfolded students moving around in a room? Increasing the temperature is like having the students start running around the room - obviously there will be more collisions. Since all of the particles have more energy, more particles will meet the minimum energy threshold (or activation energy) required for a successful collision. The graph below illustrates this:

Potential Energy CurveThe potential energy curve will be unchanged when we illustrate a reaction occurring at a higher temperature.

But at a higher temperature more particles will have enough energy to get over the activation energy "hump":

Surface Area

Increasing surface area has NO EFFECT on the energy of the particles. It does however increase the number of sites available for reactions to occur. So more  collisions can occur and therefore more "EFFECTIVE COLLISIONS" can happen.

Solids with a smaller particle size (e.g. powders or small chips) react more quickly than solids with a larger particle size (e.g large chips).

Diagrams of Particle Size and Collisions

Note the perimeter of the large cube has 12 sides for a possible reaction. The Small cubes combined have 36 (9*4) sides for a reaction to occur.  

Catalysts

When solid potassium chlorate is heated, potassium chloride and oxygen are produced. The reaction is not particularly fast.

2 KClO₃ (s) → 2 KCl (s) + 3 O₂(g) 

This reaction can take place much more rapidly, and at lower temperatures, if solid manganese dioxide (MnO₂) is added. After the reaction is complete, no more potassium chlorate remains, but all of the manganese dioxide remains.

Catalysts are very specific regarding which reactions they work with - the same catalyst will not necessarily work with many other (or any other) reactions. Finding a catalyst for a specific equation is one of the important jobs of an industrial chemist.

Catalysts most likely work by helping to promote a proper orientation between reacting particles. In doing so, it provides an alternate reaction pathway with lower activation energy.

In the potential energy diagram shown here, the regular, uncatalyzed pathway is shown as a solid line, and the catalyzed pathway is shown as the dotted line:

ΔH for the both the catalyzed and uncatalyzed reaction is -15 kJ.

Since ΔH is independent of the pathway, it is not changed by the presence of a catalyst.

Activation energy, E_a, does have different values for the catalyzed and uncatalyzed reactions.

E_a for the uncatalyzed reaction = +25 kJ

E_a for the catalyzed reaction = +10 kJ

Because more particles will possess the new 10 kJ energy minimum for a successful collision, the rate of the reaction will increase.

It is important to note that the original pathway is still present, and reacting particles will still follow that route.

Think of a catalyst as a new shortcut in your walk to school, or work, or whatever. The longer road still exists and can be used, but the shortcut makes for a faster trip.

Here's another example:

2 N₂O → 2 N₂ + O₂

E_a for the uncatalyzed reaction is 250 kJ, but when a gold metal catalyst is used, E_a is lowered to 120 kJ.

The N₂O is chemically adsorbed on the metal surface. A bond is formed between the O of the N₂O and an Au atom. This weakens the bond joining the O to the N, thus making it easier for the molecule to break apart.

Inhibitors

Some substances, known as inhibitors, slow down chemical reactions. These work by tying the reactants up in "side reactions".

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CHEMISTRY TUTORIAL: FACTORS AFFECTING REACTION RATE

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The Theoretical Model of the Reaction Process

     In our lives, there is no harm on trying things; likewise to the Chemists who proposed a theoretical model known as collision theory which states that molecules must collide in order to react. The kinetic energy of the molecules can be used to break bonds leading to chemical reactions upon collision. If molecules are moving slowly, they have little kinetic energy that they merely bounce off without any reaction It is because not all collisions are effective. If it leads to a formation of products, then it is a effective collision. If every collision leads to a reaction, chemical reactions would be thousands of times faster than they really are.

     There are two factors that determine the effectiveness of a collision: the energy of the colliding particles and the orientation of the molecules at the time of collision Effective collisions depend on the colliding particles having sufficient kinetic energy to achieve the threshold energy or minimum energy required for a chemical reaction to occur This minimum energy is called activation energy. As the temperature is increased, the frequency collision is increase and more reactant particles acquire kinetic energies that exceed the activation energy.