According to the collision theory all those factors which change the number of successful collisions per second affect the rates of chemical reactions. Some of the important factors are explained in the following paragraphs.
In order for the reaction to occur, the reactants must come in contact with each other. For this reason, the reactions are most often carried out in one phase, for example, in liquid solutions or in gas phase. When the reactants are present in the same phase, their particles are able to meet on the molecular level and thus are able to collide with each other easily. The higher the number of particles of the reactants per unit volume, higher will be chances of effective collisions and hence higher will be the rate of reaction. Hence an increase or decrease in concentrations of the reactants will result in increase or decrease in the rate of reaction respectively. For example, combustion of coal in air (21% oxygen) proceeds relatively slowly as compared to the reaction in pure oxygen. Similarly, limestone reacts with different concentrations of hydrochloric acid at different rates.
An experiment can be performed to show the increase in the rate of reaction with the increase in the concentration of hydrochloric acid. Take a flat-bottomed conical flask fitted with a delivery tube which is connected to a syringe as shown in Figure 17.3. The flask contains granules of limestone taken in excess and a fixed volume of one molar hydrochloric acid. After the start of reaction carbon dioxide gas is generated which is collected in the syringe. The scale present on the syringe directly measures the volume of carbon dioxide gas evolved. It is essential to keep all the variables constant in this experiment, for example, the volume of one molar hydrochloric acid, the temperature at which the experiment is being carried out, the mass and size of limestone granules, and the rate of stirring. The reaction is followed by measuring the volume of carbon dioxide gas evolved. The same experiment is then repeated taking the same volume of two molar hydrochloric acid.
Gas syringe
╭──────────╮
│ │← piston
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↑
delivery tube
↑
╭──────╮
│flask │ limestone + HCl
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(stand & clamp not shown)
Stand and clamp support the syringe and flask.
The change in the rate of reaction is then seen when a graph is drawn between the time in seconds or minutes against the volume of gas evolved as shown in Figure 17.4. The graph obtained is in the form of a curve. We can measure the rate of reaction during a time interval by drawing a tangent on this curve. The results show that the reaction starts with a very fast rate which then progressively slows down as the concentration of hydrochloric acid decreases. After a while the reaction stops because the whole of hydrochloric acid has been utilized.
The two curves A and B shown in Figure 17.5 represent two separate experiments done with one and two molar concentrations of HCl respectively. The greater the concentration, the steeper is the curve at the start which means faster rate. The more concentrated the acid, the more chances of effective collisions.
If reactants in a reaction are present in gas phase, their concentrations can be changed by changing their pressure. For example, the rate of reaction between hydrogen and chlorine can be doubled if the pressure of chlorine gas is doubled provided the other component hydrogen is present in excess.
Reactions involving solids take place on their surfaces and the rates depend on the area of surface contact. The larger the surface area, the more possibility of particles to come in contact and higher the rate. Finely divided solids react more rapidly than large pieces. For example, aluminum foil reacts slowly with warm sodium hydroxide but in finely divided state it reacts rapidly with even cold aqueous NaOH evolving hydrogen.
Similarly, a large piece of wood burns slowly while small pieces burn rapidly.
It is an everyday observation that rates increase with temperature. Food cooks faster at higher temperatures. Oxidation of iron is slow at room temperature but fast at high temperature. In many cases the rate nearly doubles per 10°C rise. Figure 17.6 explains this.
Rate increases sharply with temperature.
According to collision theory, as temperature increases, kinetic energy increases, thus frequency of successful collisions increases, enhancing rate.
A catalyst alters the rate without being consumed. It lowers activation energy. When activation energy is lowered, more particles form the high energy state and rate increases. For example, platinum catalyzes addition of hydrogen to ethene. Both adsorb on platinum surface, bonds weaken, helping reaction. Figure 17.7 shows reaction path with/without catalyst.
Catalyst lowers activation energy.
Enzymes are biological catalysts, proteins that speed up specific reactions in our body. They use active sites to bind reactant molecules, weakening bonds. Product then dissociates, freeing the enzyme. Active sites are very specific – each enzyme catalyzes a particular reaction.
Interesting: Kinetics helps increase yield and minimize waste. Understanding kinetics helps control reaction conditions.
📌 17.2 Quick Check! (from 17.4)
1. Why a catalyst is mostly used in finely divided form?
Answer: To increase surface area, providing more active sites for reactants to adsorb, thus increasing the rate of reaction.
2. How do you increase the number of successful collisions?
Answer: By increasing concentration, temperature, surface area, or using a catalyst; also by increasing pressure for gases.
3. How does the concentrations of the products change during reaction? Explain with the help of a graph.
Answer: Product concentration increases from zero to maximum as reaction proceeds. (Refer to Fig 17.2: products curve rises, reactants fall.)
🧠 Memorization tips – 17.4
✨ C.P.S.T.C.E. – Concentration, Pressure (gas), Surface area, Temperature, Catalyst, Enzymes.
✨ Higher concentration → more particles → more collisions → faster rate.
✨ Powdered solids react faster (more surface area).
✨ 10°C rule: rate doubles (approx).
✨ Catalyst lowers activation energy hill.
✨ Enzymes are specific: lock and key model.
📚 Guidelines – 17.4
• Understand each factor separately: concentration, surface area, temp, catalyst, pressure, enzymes.
• Study the activity 17.1 carefully – know the apparatus and variables kept constant.
• Interpret graphs: steeper slope = faster rate. Fig 17.5 shows concentration effect.
• Remember: catalyst lowers activation energy (Fig 17.7). Enzymes are specific.
• Use the Quick Check solutions to reinforce concepts.