What Is Homogeneous Equilibrium Definition Equilibrium Constant and Examples
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 in Chemical Reactions
1. What is homogeneous equilibrium in chemistry?
Homogeneous equilibrium is a chemical equilibrium in which all reactants and products are in the same physical state, usually all gases or all aqueous solutions. In this type of equilibrium:
- All species exist in a single phase (g) or (aq).
- The forward and reverse reaction rates are equal at equilibrium.
- The composition of the system remains constant over time.
2. What is the difference between homogeneous and heterogeneous equilibrium?
The main difference is that homogeneous equilibrium involves one phase, while heterogeneous equilibrium involves two or more phases.
- Homogeneous equilibrium: All reactants and products are in the same phase (e.g., all gases).
- Heterogeneous equilibrium: Reactants and products are in different phases (solid, liquid, gas).
3. What is the equilibrium constant expression for homogeneous equilibrium?
The equilibrium constant (K) expression for a homogeneous equilibrium is written as the ratio of product concentrations to reactant concentrations, each raised to their stoichiometric coefficients. For a general reaction: aA + bB ⇌ cC + dD
- Kc = [C]c[D]d / [A]a[B]b
4. How do you write the Kc expression for a homogeneous gaseous equilibrium?
To write the Kc expression for a homogeneous gaseous equilibrium, place product concentrations over reactant concentrations with exponents equal to coefficients. Example: For N2(g) + 3H2(g) ⇌ 2NH3(g)
- Kc = [NH3]2 / ([N2][H2]3)
5. What is Kp and how is it related to homogeneous equilibrium?
Kp is the equilibrium constant expressed in terms of partial pressures for homogeneous gaseous equilibria. For a gaseous reaction:
- Kp = (Pproducts) / (Preactants), each raised to their coefficients.
- Relationship: Kp = Kc(RT)Δn
6. Can you give an example of homogeneous equilibrium in aqueous solution?
An example of homogeneous equilibrium in aqueous solution is the ionization of acetic acid in water: CH3COOH(aq) ⇌ CH3COO-(aq) + H+(aq). In this system:
- All species are dissolved in water (single aqueous phase).
- The equilibrium constant is the acid dissociation constant Ka.
- Ka = [CH3COO-][H+] / [CH3COOH]
7. Why does homogeneous equilibrium not change at equilibrium?
Homogeneous equilibrium does not appear to change because the forward and reverse reaction rates are equal at equilibrium.
- The system is in dynamic equilibrium, not static equilibrium.
- Molecules continue reacting in both directions.
- Concentrations remain constant because rates are equal.
8. How does Le Chatelier’s principle apply to homogeneous equilibrium?
Le Chatelier’s principle states that a homogeneous equilibrium shifts to oppose any change in concentration, pressure, or temperature. For gaseous equilibria:
- Increasing pressure shifts equilibrium toward fewer moles of gas.
- Increasing concentration shifts equilibrium to consume the added substance.
- Increasing temperature favors the endothermic direction.
9. What are the conditions required to establish homogeneous equilibrium?
Homogeneous equilibrium is established when a reversible reaction occurs in a closed system under constant temperature. Key conditions:
- The reaction must be reversible.
- The system must be closed (no loss of reactants or products).
- Temperature must remain constant.
- Sufficient time must be allowed for equilibrium to be reached.
10. What is Δn in homogeneous gaseous equilibrium and why is it important?
Δn in homogeneous gaseous equilibrium is the difference between moles of gaseous products and moles of gaseous reactants.
- Δn = (moles of gaseous products) − (moles of gaseous reactants)
- It is used in the relation Kp = Kc(RT)Δn.
