In this article we will discuss about:- 1. Nature of Chemical Reactions 2. Rates of Chemical Reactions 3. Types.
Nature of Chemical Reactions:
Let us look at a very simple example of a chemical reaction and consider what happens when two hydrogen atoms approach one another. For our purposes, a hydrogen atom is, simply, a positively charged proton and a negatively charged electron. When two hydrogen atoms approach one another, a variety of forces are set in motion.
First, as they come closer, the negative charges of the electrons repel the atoms from one another. As they come closer still, the two atoms encounter one another’s positive charges and are repelled. But there is a third force at work- The nucleus of the first atom attracts the electron of the second, and vice versa.
When the two nuclei are a certain distance (about 7.4 nanometers or 0.74 × 10a centimeter-in other words, a little less than a hundred millionth of a centimeter) apart from one another, these forces of repulsion and attraction are equalised. At this distance, the two electrons are shared between the two nuclei, forming a strong covalent bond.
This simple example illustrates two universal facts about chemical reactions:
1. All involve a change in the electron configurations of the atoms or molecules involved.
2. All require energy to make the transition between one stable electron configuration and another. This energy is known as the energy of activation.
Rates of Chemical Reactions:
According to present concepts, atoms or molecules react with one another only when they collide with sufficient force to overcome the initial forces of repulsion.
The force required varies with the nature of the atoms or molecules; the more stable their initial state, the more forceful the collision must be. In any given group of atoms, it is likely that some proportion is moving with sufficient energy to cause a reaction to occur, but often this proportion is so small that the reaction, for all practical purposes, does not take place.
Reaction rates can be increased by increasing the likelihood of forceful collisions. One way to do this is to raise the temperature, thereby increasing the average velocity at which the atoms or molecules move and so increasing their likelihood of colliding with sufficient force.
Sometimes, as in the case of natural gas, a spark is all that is needed. Once the reaction begins, it liberates heat that is transferred to the other CH4 molecules until all are moving rapidly enough to react almost simultaneously with explosive force. Driving a reaction by heat is a method commonly used in both chemical laboratories (for example, the familiar Bunsen burner) and industry.
A second way to increase the rate of collisions is to increase the number of reacting molecules (or to decrease the volume in which they move). Laboratories and chemical factories usually work with pure chemicals in high concentrations.
A third way is to use catalysts. Catalysts lower the energy of activation of a reaction. Metals, such as iron, nickel, and platinum, are commonly used as catalysts in industrial laboratories. Certain molecules apparently tend to cluster on the surfaces of such metals and so the likelihood of close encounters of the necessary kinds between the reactants is increased.
The metals, although they participate in the reactions, are not used up by them, and so they can be used over and over again.
Types of Chemical Reactions:
Chemical reactions can be classified into a few general types.
One type can be represented by the expression:
A + B → AB
An example of this sort of reaction is the combination of hydrogen gas with oxygen gas to produce water-
2H2 + O2 → 2H2O
A reaction may also take the form of dissociation-
AB → A + B
For example, the equation above showing the formation of water can be reversed-
2H2 O → 2H2 + O2
This means that water molecules yield hydrogen and oxygen gases. A reaction may also involve an exchange, taking the form-
AB + CD → AD + CB
An example of such a reaction is the combination of hydrochloric acid (HCl) and sodium hydroxide (NaOH) to make table salt and water-
HCl + NaOH → NaCl + H2O
The passing of an electron from one molecule to another is known as an oxidation-reduction (or redox) reaction. The loss of an electron is known as oxidation, and the compound that loses the electron is said to be oxidised. The reason electron loss is called oxidation is that oxygen, with its high electronegativity, is often the electron acceptor.
Reduction is, conversely, the gain of an electron. Oxidation and reduction take place simultaneously because an electron that is lost by the oxidised atom is accepted by another atom, which is then reduced. (The fact that reduction means a gain of an electron seems paradoxical and may make it hard to remember the distinction between the two terms. It may help to recall that gain of an electron reduces the charge of an atom or molecule.)
Redox reactions may involve only a solitary electron, as for instance when sodium loses an electron and becomes oxidised to Na+, and chlorine gains an electron, thereby reducing its charge (CI–). Often the electron travels with a proton, so oxidation involves removal of hydrogen atoms and reduction the gain of hydrogen atoms.
When glucose is burned (oxidised), hydrogen atoms are lost by the glucose molecule and gained by oxygen:
C6H12O6 + 6O2 → 6CO2 + 6H2O
Chemical reactions can go in either direction. When net change ceases, the reaction is said to be at equilibrium. In the reaction the point of equilibrium is reached when as many molecules of C and D are being converted to molecules of A and B as molecules of A and B are being converted to molecules of C and D.
The concentration of reactants does not have to equal the concentration of products in order for equilibrium to be established only the rates of the forward and reverse reactions must be the same. Consider the reaction shown at the bottom. The different lengths of the arrows indicate that at equilibrium there is more C + D present than A + B.
If only A and B molecules are present initially, the reaction occurs at first to the right, with A and B molecules converting into C and D molecules. Figure 4.3 shows the relative changes in concentration as the reaction continues. As C and D accumulate, the rate of the reverse reaction increases, and at the same time, the rate of the forward reaction decreases because of the decreasing concentrations of A and B.
At about minute 6, the rates of the forward and reverse reactions equalise and no further changes in concentration take place. The proportions of A + B and C + D will remain the same, but there will always be more C and D molecules in the system.