What is Homogeneous Equilibrium?
The reaction in which the phase of all the products and the reactants is the same. For example- all the products and reactants are gases or all the products and reactants are liquids. An equilibrium reaction is the one in which the reaction can be reversed and forwarded and the concentrations of the reactants and products remain the same.
Equilibrium is a state of chemical reaction where the rate of forwarding 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.
Homogeneous Reaction- Equilibrium Constant
Suppose a homogeneous system is given by the reaction given below-
\[N_{2}(g) + 3H_{2}(g) \leftrightarrow 2NH_{3} (g)\]
It is observed that gaseous nitrogen when reacted with gaseous hydrogen gives gaseous ammonia.
This reaction can be written in terms of molar concentration in the following way-
\[K_{c} = \frac{[NH_{3}]^{2}}{[N_{2}][H_{2}]^{3}}\]
The equilibrium concentration, for the reactions involving gases, is expressed in terms of parietal pressure.
By using the ideal gas equation,
PV = nRT
P and V here are the pressure and volume of the system respectively
The number of moles of the component present is expressed as n
R is for the universal gas constant
T is for the temperature
\[C = \frac{n}{v}RT = cRT\]
C is denoted for the concentration of the system which can also be written as
\[C = \frac{P}{RT}\]
Therefore the equilibrium constant for a homogeneous equilibrium can be written as
\[K_{p} = K_{c} (RT)^{\Delta n}\]
where \[K_{p}\] is the equilibrium constant being calculated from the partial pressure of a reaction equation
and
\[{\Delta n}\]represents the number of moles of the gaseous products
Homogeneous Equilibrium- Example
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:
\[N_{2} (g) + O_{2} (g) = 2NO (g)\]
We can observe from the example above that there are two molecules of reactants (one of each) and the product also has 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:
\[2SO_{2} (g) + O_{2} (g) \rightleftharpoons 2SO_{3} (g)\]
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.
In a homogeneous equilibrium, the reactions in liquid solutions between solutes belong to one type of homogeneous equilibria and the chemical species which are involved can be either molecules or ions or a mixture of both.
For the following homogeneous reaction
\[C_{2}H_{2}(aq) + 2Br_{2}(aq) C_{2}H_{2}B_{4}(aq)\]
\[K= \frac{[C_{2}H_{2}B_{4}]}{[C_{2}H_{2}][Br_{2}]^{2}}\]
where K is the equilibrium constant
Homogeneous Chemical Equilibrium
It is important to note that the two different equilibriums are dealt with in a different way and so are the calculations related to them.
In order to understand this further, a chemical equation is provided below about the homogeneous equilibrium.
\[C_{2}H_{2} (aq) + 2Br_{2} (aq) \rightleftharpoons C_{2}H_{2}Br_{2} (aq)\]
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:
\[H_{2}O (s) \rightleftharpoons H_{2}O (l)\]
Calculate Equilibrium Constant
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, the 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:
\[wW + xX \rightleftharpoons yY +zZ\]
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:
\[K_{c}= \frac{Y^{Y}Z^{Z}}{W^{W}X^{X}}\]
Using this formula, one can calculate the equilibrium constant of any equation. Another example of a chemical equation that can be used to explain is given below:
Equation:
\[2SO_{2} (g) + O_{2} (g) \rightleftharpoons 2SO_{3} (g)\]
Equilibrium Constant:
\[K_{C} = \frac{ [SO_{3}]^{2} }{[SO_{2}]^{2}[O_{2}]}\]
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:
\[K_{P} = K_{C}(RT)^{\Delta n}\]
\[{\Delta n}\]- is the number of moles of gaseous products,
\[K_{P}\]- is defined as the equilibrium constant that is calculated from the partial pressure of a reaction equation.
Difference between \[K_{C} and K_{P}\]
The main difference between the two equilibrium constants is that they are used for the different concentrations. \[K_{P}\] specifically represents the 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 constants has also been derived, which is given below:
\[K_{P} = K_{C}(RT)^{\Delta n}\]
FAQs on Homogeneous Equilibrium
1. What is meant by heterogeneous equilibria?
The system in which there are two or more phases of the reactants and the products is known as the heterogeneous equilibrium. The phases in the system mean any combination of liquids, gases, solids, and solutions. While dealing with heterogeneous equilibrium it is important to note that pure liquids and solids can not appear as equilibrium constant expressions. An example of a heterogeneous reaction would be \[Br_{2}(liq) \leftrightarrow Br_{2}(gas)\]. The equilibrium constant K for this equation would be \[Br_{2}\]. The concentration of the pure liquid \[Br_{2}\] will be excluded since they cannot appear as equilibrium constant expressions.
2. What is meant by the common ion effect and its role?
The changes occurring when ions are introduced into a solution containing the same ion are displayed in the common ion effect. When common ions are added into a solution, there is a decrease in the solubility of a compound as there is a shift in the equilibrium. The common ion effect plays an important role in the regulation of buffers. There is either an acid or a base in a buffering solution which is accompanied by its conjugate counterpart. When more conjugate ions are added there is a shift in the pH of the solution. While determining solution equilibrium when common ions are added, this effect must be taken into consideration.
In order to learn more about pH change by common ion effect, students can visit Vedantu.
3. What is meant by buffer solutions?
Solutions that are either composed of a weak base and its conjugate acid or a weak acid and its conjugate base are known as buffer solutions. If excess ions are added into a buffer solution, the pH of the solution will be altered. Le Chatelier’s principle states that when more ions are added there is an alteration in the equilibrium and the reactions shift to favor the solid or deionized form. The concentration of the hydrogen ion decreases in the case of an acidic buffer and the solution that is produced will be less acidic as compared to a solution that contains pure weak acid.
4. What is meant by solubility?
The amount of the material which, in a particular solvent, is able to be dissolved is known as the solubility of that material. Let us take an example of table salt (NaCl). When table salt is placed in water it dissolves in water in seconds. If we keep on adding table salt to the solution till when a point comes where no more salt can be dissolved in the water, we say that the solution is saturated. This means that the table salt has reached its solubility limit. Students can visit Vedantu for an in-depth understanding of this concept.
5. What is an equilibrium constant?
The symbolic representation of the 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.
6. What are some of the differences between a homogeneous and a heterogeneous equilibrium?
The primary difference between homogeneous equilibrium and heterogeneous equilibrium is that in heterogeneous equilibrium the phases of the products and reactants are different while the phase of product and reactants in the case of homogeneous equilibrium reactions is the same. Another important difference is that in the case of heterogeneous equilibrium the concentrations of solids and pure liquids have to be excluded while in the case of homogeneous equilibrium concentrations of reactants and products have to be included.
7. 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 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.