1.4.4 Thermodynamics

Unit 1 reaction equation as it can in the case of an elementary reaction, the rate law is a direct result of the sequence of elementary steps that constitute the reaction mechanism. As such, it provides our best tool for determining an unknown mechanism. Once we know the sequence of elementary steps that constitute the reaction mechanism, we can quite quickly deduce the rate law. Conversely, if we do not know the reaction mechanism, we can carry out experiments to determine the orders with respect to each reactant and then try out various ‘trial’ reaction mechanisms to see which one fits best with the experimental data. At this point it should be emphasised again that for multi-step reactions, the rate law, rate constant, and order are determined by experiment, and the orders are not generally the same as the stoichiometric coefficients in the reaction equation. A final important point about rate laws is that overall rate laws for a reaction may contain reactant, product and catalyst concentrations, but must not contain concentrations of reactive intermediates A rate law is a differential equation that describes the rate of change of a reactant (or product) concentration with time. If we integrate the rate law then we obtain an expression for the concentration as a function of time, which is generally the type of data obtained in an experiment. In many simple cases, the rate law may be integrated analytically. Otherwise, numerical (computer-based) techniques may be used. Four of the simplest rate laws are given below in both their differential and integrated form

In the above [A]0 and [B]0 represent the initial concentrations of A and B i.e. their concentrations at the start of the reaction.

Half lives

The half life, t1/2, of a substance is defined as the time it takes for the concentration of the substance to fall to half of its initial value. Note that it only makes sense to define a half life for a substance not present in excess at the start of the reaction. We can obtain equations for the half-lives for reactions of various orders by substituting the values t = t 1/2 and [A] = ½ [A]0 into the integrated rate laws. We obtain

1.4.4 Thermodynamics

Thermodynamics is a branch of science, which deals with energy associated with different atoms, molecules and chemical bonds. Thermodynamics helps us to understand the reaction kinetics and driving force for the reactions. Let us examine the basic terminologies of the Thermodynamics, which are required from CSIR point of view.

1. System

 Any solution/solid part in which molecules are there and some chemical reactions are happening  E.g.: Solution in a beaker, air filled in balloon, you are a system.

Unit 1

2. Surrounding

 Whatever is present surrounding a system  Eg: Air present above the solution in beaker

3. Enthalpy ( H)

 Greek: enthalpein, to warm in  In simple terms, it is the heat content of the system.  What does it means? It is the energy stored in every chemical bond in the molecules of the system.  As we all know, to form a bond we need to add energy and to break a bond we need to supply energy, when a bond is broken energy is again given out.  If there is a system, called A, whose H = 20, is was heated or some reaction happened, its H changes to 30.  Now what is the change in Enthalpy? = H2 – H1  We write it as Change in enthalpy = ∆H = H2 – H1

 Here U is energy  q represents the heat absorbed by the system from the surrounding  W is the work done on the system by the surroundings.  Heat is due to random motion of atoms and molecules  Work, which is defined as force times the distance moved under its influence, is associated with organized motion.  Force can be of different forms Gravitational force exerted by one mass on another,  The expansion force exerted by a gas,  the tensional force exerted by a spring or muscle fibre,  the electrical force of one charge on another  Forces of friction and viscosity.

Figure 1.32 You are a system…..you are surrounded by air , we call it as Surrounding

Many chemical reactions release energy in the form of heat, light, or sound. These are exothermic reactions. Exothermic reactions may occur spontaneously and result in higher randomness or entropy (ΔS > 0) of the system. They are denoted by a negative heat flow (heat is lost Figure 1.33 An exothermic to the surroundings) and decrease in enthalpy (ΔH < 0). In the lab, thermite reaction using iron(III) exothermic reactions produce heat or may even be explosive. oxide. The sparks flying outwards are globules of molten There are other chemical reactions that must absorb energy in order to iron trailing smoke in their wake. proceed. These are endothermic reactions. Endothermic reactions cannot occur spontaneously. Work must be done in order to get these reactions to occur. When endothermic reactions absorb energy, a temperature drop is measured during the reaction. Endothermic reactions are characterized by positive heat flow (into the reaction) and an increase in enthalpy (+ΔH).

Examples of exothermic processes are:

 Condensation of rain from water vapour  Combustion of fuels such as wood, coal and oil petroleum  Mixing water and strong acids  Mixing alkalis and acids  The setting of cement and concrete  Some polymerisation reactions such as the setting of epoxy resin  Thermite reaction

Examples of endothermic reactions

 Photosynthesis  Melting ice