Equilibrium in chemical processes - class-XI

$N _2O _4\longrightarrow 2NO _2$

The intensity of color represents the concentration of either reactant or product. Thus when in a colored gaseous reversible reaction, the color has attained constant intensity, the concentrations of reactants and products have reached equilibrium values. In other words, an equilibrium is attained.

Given $H _2=4g$ & $I _2=127 g$

$CO _(g)^+\ H _2(g)\rightleftharpoons CO _{2(g)}+H _{2(g)}$

Equilibrium can be achieved in open vessel as well as in close vessel.

If CO is added 2$^{nd}$ equilibrium will proceed in the backward direction and concentration of Cl$ _2$ will decrease. This Cl$ _2$ will be further formed by the decomposition of PCl$ _5$.

Let reactions is started with a mole of $AB _2$ and b mole of $B _2C$
$\Rightarrow      AB _2 (g) + B _2C(g)  \rightarrow AB _3(g) + BC(g)$
                    a                 b               0               0
                  a - x           b - x - y        x              x - y
$BC(g) + B _2C(g)   \rightarrow B _2C _2 (g)$
    x - y            b - x - y        y                 As  x > y
Clearly $[AB _3] _{eq} > [B _3 C _2] _{eq} $ and  $[AB _3] _{eq}  >  [BC] _{eq}$

When the equilibrium is attained, the concentration of each of the reactants and products may or may not become equal.  It may happen that the concentrations of products are higher than the concentrations of the reactants. Or, it may happen that the concentrations of the reactants are higher than the concentrations of the products.

If you have a solution in equilibrium containing $Cl^-$ ions and you added $NaCl$ to the solution, the reaction will shift to the left to compensate for the increased concentration of $Cl^-$ ions. This is in accordance with the  Le Chatelier's Principle .
The Le Chatelier's Principle states that any change in a substance on one side of the equation, either in concentration, temperature, pressure or volume, results in an equilibrium shift to oppose the change until a new equilibrium is reached.

The following is true about the chemical equilibrium.
The chemical equilibrium is a state in which the rate of the forward reaction is equal to the rate of the reverse reaction. A chemical equilibrium is a state of balance in a reaction where the forward and reverse reaction speed is equal and the concentrations of the products and reactants remain unchanged.

At dynamic equilibrium, the reaction rate of the forward reaction is equal to the reaction rate of the backward reaction. A dynamic equilibrium exists once a reversible reaction ceases to change its ratio of reactants/products, but substances move between the chemicals at an equal rate, meaning there is no net change. It is a particular example of a system in a steady state. Equilibrium is the state of equal, opposite rates, not equal concentrations. The measurable properties like concentration, density, color, pressure etc remain constant at constant temperature.

Equilibrium is the state in which both reactants and products are present in concentrations which have no further tendency to change with time.Usually, this state results when the forward reaction proceeds at the same rate as the reverse reaction. The reaction rates of the forward and backward reactions are generally not zero, but equal. Thus, there are no net changes in the concentrations of the reactant(s) and product(s).

In reversible reactions, as equation the rates of forward and backward reactions are same so the concentrations of reactants and products are constant at equilibrium.

The rates of reaction for the forward and reverse reactions are equal in a reversible chemical reaction at equilibrium. When the concentrations of reactants and products have become constant, a reaction is said to have reached a point of equilibrium. The consistency of measurable properties such as concentration, color, pressure and density can show a state of equilibrium. The rates of the two opposing reactions have become equal. The amount of products and reactants produced are consistent, and there is no net change.

A reaction continues even after the attainment of equilibrium.
The reaction proceeds in forward as well as reverse direction. When equilibrium is attained, it is dynamic in nature. The reactants are converted into products through forward reaction and the products are converted into reactants through reverse reaction.

The equilibrium state can be attained from both sides of the chemical reaction.
Thus, if initially, the concentration of reactants is much higher than the equilibrium concentration and the concentration of products is much lower than the equilibrium concentration, the reaction will proceed in the forward direction till equilibrium is attained. On the other hand, if initially, the concentration of
reactants is much lower than the equilibrium concentration and the
concentration of products is much higher than the equilibrium
concentration, the reaction will proceed in the reverse direction till
equilibrium is attained.

We know that for a reaction if $\dfrac{K}{Q} < 1,$ the reaction proceeds in backward direction. Hence, the reaction mixture forms more reactant species.

Using the relation between Kp and Kc, namely Kp = Kc * (RT)^delta_n, we find delta_n = (4 + 2 + 1) - 4 = 3. Given P(HNO3) = 2 atm and total pressure 30 atm, we can find partial pressures at equilibrium, then compute Kp and subsequently Kc.

$\alpha $ from $=+150.{ 7 }^{ 0 }$, $\beta $ from $=+52.{ 8 }^{ 0 }$

                ${ N } _{ 2 }+{ 3H } _{ 2 }\rightleftharpoons 2{ NH } _{ 3 }$

Independent of density of gases, in equilibrium gases will mix homogenously, this comes from the fact that gases occupies entire volume of container. This also can be thought as gases tries to reduce energy. Hence they separate as far as possible.

$Q _c = \dfrac{[P _2(g)]^4}{[P _4 (g)]}$

According to the equation

To obtain the target equilibrium equation, multiply the first reaction by -1/2 (or reverse and take square root) and add the second reaction. This corresponds to combining the equilibrium constants as K3 = (1 / sqrt(K1)) * K2.

Chemical equilibria is dynamic in nature. It occurs in reversible reactions only. At equilibria all observable properties becomes constant. Chemical equilibrium is affected by a change in temperature, pressure, concentration, volume etc but is not altered by the addition of catalyst. Equilibria is of two types heterogeneous and homogeneous equilibria.

The equilibrium reaction containing $Fe(III)$ nitrate & $KSCN$ is given as :-

                                                  ${{S} _{8}}\rightleftharpoons 4{{S} _{2}}$

Initial pressure                          $1$atm     $0$

At equilibrium pressure    $(1-0.3)$     $4\times 0.3$

Therefore, equilibrium pressure of ${{S} _{8}}=(1-0.3)=0.7$ atm

And equilibrium pressure of ${{S} _{2}}=4\times 0.3=1.2$ atm

So, equilibrium constant $Kp=\dfrac{P _{S _2}^4}{P _{S _8}}=\dfrac{{{1.2}^{4}}}{0.7}=$ approximate $2.96$ atm$^3$. 

In chemistry, a dynamic equilibrium exists once a reversible reaction ceases to change its ratio of reactants/products, but substances move between the chemicals at an equal rate, meaning there is no net change. It is a particular example of a system in a steady state. In thermodynamics a closed system is in thermodynamic equilibrium when reactions occur at such rates that the composition of the mixture does not change with time. Reactions do in fact occur, sometimes vigorously, but to such an extent that changes in composition cannot be observed.

Hence, in a dynamic equilibrium, the concentrations of reactants and products remains constant.

The reaction is,

In PCl5 -> PCl3 + Cl2, one mole of gas produces two moles of gas. Thus, the total number of moles increases.

Kp = P_O2^3 / P_O3^2 = 4 * 10^14. Total pressure = P_O2 + P_O3 = 8. Since P_O2 >> P_O3, P_O2 approx 8. 8^3 / P_O3^2 = 4 * 10^14. 512 / P_O3^2 = 4 * 10^14. P_O3^2 = 128 * 10^-14. P_O3 = sqrt(128) * 10^-7 = 11.3 * 10^-7.

Given the reaction 3 C2H2 <=> C6H6 with Kc = 4 and 0.5 moles of benzene at equilibrium in a 1-liter vessel. Setting up an ICE table, if benzene is 0.5, acetylene reacted is 1.5, leaving 0 for acetylene at equilibrium, leading to a total of 0.5 + 0.5 = 1 mole if calculated correctly with initial values, or simply using stoichiometry and equilibrium concentrations.