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Chemical Equilibrium
In a chemical reaction, chemical equilibrium is the state in which both reactants and products are present in concentrations which have no further tendency to change with time.
The chemical equilibrium is achieved when the rate of forward reaction is same as the reverse reaction. Since the rates are equal, there are no net changes in the concentrations of the reactant(s) and product(s). This state is known as dynamic equilibrium.
The entire process can be graphically represented by this diagram.
Law of Chemical Equilibrium & Equilibrium Constant Kc
aA(g) + bB(g) ⇌ cC(g) + dD(g)
Using the Law of mass action,
Forward reaction rate = k+ [A]a[B]b
Backward reaction rate = k– [C]c[D]d
where, [A], [B], [C] and [D] being the active masses and k+ and k− are rate constants.
aA(g) + bB(g) ⇌ cC(g) + dD(g)
Forward reaction rate = Backward reaction rate
k+ [A]a[B]b = k– [C]c[D]d
We have,
Kc = k+ / k–
Kc = [ C ]c·[ D ]d / [ A ]a·[ B ]b
Types of Chemical Equilibria
There are two types of chemical equilibria:
Homogeneous Equilibria
Heterogeneous Equilibria
So,
Applications of Equilibrium Constant
For a general reaction,
The concentrations are not necessarily equilibrium values.
Relation between K, Q, and G (Gibbs Energy)
Factors Affecting Chemical Equilibrium
According to this principle, if a system is in equilibrium and it is subjected to any change in any of the factors that determine the equilibrium conditions of the system, then it will shift the equilibrium in such a way to reduce or to counteract the effect of the change.
To help you understand the effects of various changes better, here is the summary:
Chemical equilibrium is a dynamic process that consists of a forward reaction, in which reactants are converted to products, and a reverse reaction, in which products are converted to reactants. At equilibrium, the forward and reverse reactions proceed at equal rates. Consider, for example, a simple system that contains only one reactant and one product, the reversible dissociation of dinitrogen tetroxide (N2O4) to nitrogen dioxide (NO2). You may recall from Chapter 14 “Chemical Kinetics” that NO2 is responsible for the brown color we associate with smog. When a sealed tube containing solid N2O4 (mp = −9.3°C; bp = 21.2°C) is heated from −78.4°C to 25°C, the red-brown color of NO2appears (Figure 15.1 “The “). The reaction can be followed visually because the product (NO2) is colored, whereas the reactant (N2O4) is colorless:
Equation 15.
The double arrow indicates that both the forward and reverse reactions are occurring simultaneously; it is read “is in equilibrium with.”
System at Different Temperatures
(left) At dry ice temperature (−78.4°C), the system contains essentially pure solid N2O4, which is colorless. (center) As the system is warmed above the melting point of N2O4 (−9.3°C), the N2O4 melts and then evaporates, and some of the vapor dissociates to red-brown NO2. (right) Eventually the sample reaches room temperature, and a mixture of gaseous N2O4 and NO2 is present. The composition of the mixture and hence the color do not change further with time: the system has reached equilibrium at the new temperature.
Figure 15.2 “The Composition of N” shows how the composition of this system would vary as a function of time at a constant temperature. If the initial concentration of NO2 were zero, then it increases as the concentration of N2O4 decreases. Eventually the composition of the system stops changing with time, and chemical equilibrium is achieved. Conversely, if we start with a sample that contains no N2O4 but an initial NO2 concentration twice the initial concentration of N2O4 in part (a) in Figure 15.2 “The Composition of N”, in accordance with the stoichiometry of the reaction, we reach exactly the same equilibrium composition, as shown in part (b) in Figure 15.2 “The Composition of N”. Thus equilibrium can be approached from either direction in a chemical reaction.
(a) Initially, this idealized system contains 0.0500 M gaseous N2O4 and no gaseous NO2. The concentration of N2O4 decreases with time as the concentration of NO2 increases. (b) Initially, this system contains 0.1000 M NO2 and no N2O4. The concentration of NO2 decreases with time as the concentration of N2O4 increases. In both cases, the final concentrations of the substances are the same: [N2O4] = 0.0422 M and [NO2] = 0.0156 M at equilibrium.
Figure 15.3 “The Forward and Reverse Reaction Rates as a Function of Time for the ” shows the forward and reverse reaction rates for a sample that initially contains pure NO2. Because the initial concentration of N2O4 is zero, the forward reaction rate (dissociation of N2O4) is initially zero as well. In contrast, the reverse reaction rate (dimerization of NO2) is initially very high (2.0 × 106 M/s), but it decreases rapidly as the concentration of NO2 decreases. (Recall from Chapter 14 “Chemical Kinetics” that the reaction rate of the dimerization reaction is expected to decrease rapidly because the reaction is second order in NO2: rate = kr[NO2]2, where kr is the rate constant for the reverse reaction shown in Equation 15.1.) As the concentration of N2O4 increases, the rate of dissociation of N2O4 increases—but more slowly than the dimerization of NO2—because the reaction is only first order in N2O4 (rate = kf[N2O4], where kf is the rate constant for the forward reaction in Equation 15.1). Eventually, the forward and reverse reaction rates become identical, kF = kr, and the system has reached chemical equilibrium. If the forward and reverse reactions occur at different rates, then the system is not at equilibrium.
The rate of dimerization of NO2 (reverse reaction) decreases rapidly with time, as expected for a second-order reaction. Because the initial concentration of N2O4 is zero, the rate of the dissociation reaction (forward reaction) at t = 0 is also zero. As the dimerization reaction proceeds, the N2O4 concentration increases, and its rate of dissociation also increases. Eventually the rates of the two reactions are equal: chemical equilibrium has been reached, and the concentrations of N2O4 and NO2 no longer change.
ASSIGNMENT : THE CONCEPT OF CHEMICAL EQUILIBRIUM ASSIGNMENT MARKS : 10 DURATION : 1 week, 3 days