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Equilibrium is a state of chemical reaction where the rate of forward and backward reaction is the same. Moreover, an equilibrium can be of two types: homogeneous equilibrium and heterogeneous equilibrium. A homogeneous equilibrium is defined as a homogeneous mixture (reactants and products in a single solution) in one phase. Remember, the reactants are on the left side of the equation and the products are on the right side of the equation. Therefore, the reaction which takes place between the solutes belongs to a single homogeneous equilibrium. A heterogeneous equilibrium, on the other hand, can be defined as a reaction system where the products and the reactants are found in two or more phases.

It is important to note that the two different equilibriums a dealt with in a different and so are the calculations related to them.

In order to understand this further, a chemical equation is provided below about the homogeneous equilibrium.

C2H2 (aq) + 2Br2 (aq) ⇌ C2H2Br2 (aq)

Moreover, a heterogeneous equilibrium example is also provided in order to learn about the difference between homogeneous and heterogeneous equilibrium. The example is given below:

H2O (s) ⇌ H2O (l)

A homogeneous equilibrium can further be divided into two categories. In the first category, the number of molecules of the product is the same as the number of molecules of the reactants of that particular equation.

For example:

N2 (g) + O2 (g) = 2NO (g)

We can observe from the example above that there are two molecules of reactants (one of each) and the product also had two molecules on the right side.

In the second category of a homogeneous equilibrium equation, opposite circumstances are observed. The number of molecules of the product is not the same or equal to the number of molecules of the reactant.

For example:

2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

From this above example, we can notice that there are only three molecules of reactant and only two molecules of the product present in the reaction.

You can solve or calculate the equilibrium constant for a given reaction. To explain this, let us take a hypothetical example of W, X, Y, and Z as reactants and products. Their coefficient, number in front of the compound or molecule, is represented by w, x, y, and z. The equation with these molecules and their coefficient is given below:

wW + xX ⇌ yY +zZ

In order to find the equilibrium constant of this equation and similar equations, the products of the equation go in the numerator, with the coefficient as their exponent. The reactants, on the other hand, go in the denominator, with their coefficients as their exponents. The written expression will look similar to the one below:

KC = [Y]Y[Z]Z / [W]W[X]X

Using this formula, one can calculate for the equilibrium constant of any equation. Another example of a chemical equation which can be used to explain is given below:

Equation

2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

Equilibrium Constant

KC = [SO3]2 / [SO2]2[O2].

Moreover, in terms of finding equilibrium constant for gases, partial pressure is taken into account. The first step involved here is to use the ideal gas equation. The formula for the ideal gas equation is given below:

PV = nRT; where

P- is the pressure,

V- is the volume

n- is the number of moles of components

R- is the universal gas constant

T- is the temperature

The above relationship can also be written as:

P = (n/v)RT; where

n/v = c (concentration of the system)

P = cRT

Therefore, the equilibrium constant can also be written using the concentration of the system in the formula. This relationship can be derived through the following steps

KP = KC(RT)Δn; where

Δn - is the number of moles of gaseous products,

KP - is defined as the equilibrium constant that is calculated from the partial pressure of a reaction equation.

The main difference between the two equilibrium constants is that they are used for the different concentrations. KP specifically represents equilibrium constant at partial pressure during a reaction. In terms of calculation, these values can be found from the reactant and products, using the formulae and using the specific values of those formulae. A relationship between the two equilibrium constant has also been derived, which is given below:

KP = KC(RT)Δn

FAQ (Frequently Asked Questions)

1.What is the difference between a homogeneous mixture and a heterogeneous mixture?

Homogeneous mixtures are often considered to be indistinguishable from the pure substance when we introspect at a macroscopic level. The reaction which takes place between the solutes belongs to a single homogeneous equilibrium.

Some examples of the homogeneous mixtures are known as sugar, salt, water, dye, air, and blood.

A heterogeneous mixture has a clear identifying property where one can see various different components of the mixture. It is a reaction system where the products and the reactants are found in two or more phases.

Some examples of a heterogeneous mixture are pizza, cookies, rocks, etc.

2. What is an equilibrium constant?

The symbolic representation of equilibrium constant is either denoted as K or KC. The term equilibrium constant can be defined as the expression which denotes the concentration of the reactants and the products which is achieved after the chemical reaction has reached the state of equilibrium. Temperature plays a very important role in maintaining the equilibrium constant within the reactions. If the temperature remains constant then the equilibrium also remains constant. This can be seen throughout the equation which eventually plays a very vital role in maintaining a constant equilibrium.