Difference between Homogeneous and Heterogeneous Equilibria
In our daily lives, we witness several reactions such as iron rusting, paper burning, curd sourness, ozone generation, and so on. Many of these reactions require the presence of components in distinct phases, such as solid iron reacting with gaseous oxygen to generate solid iron oxide, which we call rust. Similarly, gaseous hydrogen and gaseous oxygen combine to make liquid water. Having to deal with such reactions is a time-consuming chore. When the components are in the same phase, their interaction is simple to understand; but, when the components are in different stages, the interaction becomes more complicated.
Homogeneous and Heterogeneous equilibrium
To simplify difficulties and grasp the notion, we divide such reactions into two categories: homogeneous reactions, in which the components involved are present in the same phase, and heterogeneous reactions, in which the components involved are present in separate phases. The methods for dealing with both reactions, as well as the determination of the equilibrium state, differ.
The reaction in which all of the products and reactants have the same phase. For instance, all of the products and reactants could be gases or all of the products and reactants could be liquids. An equilibrium reaction is one that can be reversed and forwarded while the concentrations of the reactants and products remain constant.
Equilibrium is a chemical reaction state in which the rate of forwarding and backward reaction is the same. Furthermore, there are two forms of equilibrium: homogeneous equilibrium and heterogeneous equilibrium. A homogeneous equilibrium in one phase is defined as a homogeneous mixture (reactants and products in a single solution). Remember that the reactants are on the left side of the equation and the products are on the right. As a result, the reaction between the solutes corresponds to a single homogeneous equilibrium. A heterogeneous equilibrium, on the other hand, is a reaction system in which the products and reactants exist in two or more phases.
A homogeneous equilibrium can be further classified into two types. The number of molecules in the product in the first category is the same as the number of molecules in the reactants in that particular equation. As an example:
N2 (g)+O2 (g) ⇌ 2NO (g)
From the above example, we can see that there are two molecules of reactants (one of each) and two molecules of product on the right side. In the second category of a homogeneous equilibrium equation, the opposite events occur. The product’s molecule count is not the same as or equal to the reactant’s molecule count. As an example:
2SO2 (g)+O2 (g) ⇌ 2SO3(g)
We can see from the preceding example that there are only three molecules of reactant and two molecules of product present in the reaction. The reactions in liquid solutions between solutes belong to one type of homogeneous equilibria in a homogeneous equilibrium, and the chemical species involved can be molecules, ions, or a mixture of both.
Difference between KC and KP
The two equilibrium constants are distinguished by the fact that they are applied to different concentrations. KP denotes the equilibrium constant at partial pressure during a reaction. These values can be calculated using the reactant and product values, the equations, and the specific values of those formulae. A relationship exists between the two equilibrium constants, which is shown below:
KP = KC(RT)Δn
Difference Between Homogeneous and Heterogeneous Equilibrium
- Equilibrium is defined as a situation in which the concentrations of reactants and products are constant. Homogeneous equilibrium and heterogeneous equilibrium are the two types of equilibria. The primary distinction between homogeneous and heterogeneous equilibrium is that in homogeneous equilibrium, the reactants and products are in the same phase of matter, but in heterogeneous equilibrium, they are in different phases.
- Furthermore, when determining the equilibrium constant for homogeneous equilibria, we must include the concentrations of all reactants and products; however, when determining the equilibrium constant for heterogeneous equilibria, we must exclude the concentrations of solids and pure liquids and use the concentrations of other reactants and products. As an illustration,
- 2SO2(g) + O2(g) ⇌ 2SO3(g) is a homogeneous equilibrium, and
- O2(g) + 2C(s) ⇌ 2CO(g) is an example for a heterogeneous equilibrium.
- To put it simply, equilibrium is a situation in which the concentrations of reactants and products remain constant. Homogeneous equilibrium and heterogeneous equilibrium are the two types of equilibria. The primary distinction between homogeneous and heterogeneous equilibrium is that in homogeneous equilibrium, the reactants and products are in the same phase of matter, but in heterogeneous equilibrium, they are in different phases. Furthermore, the equilibrium constant for homogeneous equilibria includes all reactant and product concentrations, whereas the equilibrium constant for heterogeneous equilibria must omit solid and pure liquid concentrations.
Question 1: What is meant by heterogeneous equilibria?
The heterogeneous equilibrium refers to a system in which the reactants and products exist in two or more phases. The system’s phases are any combination of liquids, gases, solids, and solutions. It is vital to remember that pure liquids and solids cannot appear as equilibrium constant expressions when dealing with heterogeneous equilibrium.
Question 2: What is meant by buffer solutions?
Buffer solutions are composed of either a weak base and its conjugate acid or a weak acid and its conjugate base. When extra ions are introduced to a buffer solution, the pH of the solution changes. According to Le Chatelier’s principle, when more ions are introduced, the equilibrium changes and the reactions shift to favour the solid or deionized form. In the case of an acidic buffer, the concentration of the hydrogen ion reduces, and the solution produced is less acidic than a solution containing pure weak acid.
Question 3: What is an equilibrium constant?
The word equilibrium constant can be defined as the expression that reflects the concentration of reactants and products after the chemical process has reached equilibrium. Temperature is critical in maintaining the equilibrium constant within the reactions. If the temperature remains constant, the equilibrium remains constant as well. This is evident throughout the equation, and it ultimately plays a critical role in maintaining a constant balance.
Question 4: What is the difference between a homogeneous mixture and a heterogeneous mixture?
When viewed at a macroscopic level, homogeneous mixes are frequently thought to be indistinguishable from the pure substance. The reaction that occurs between the solutes is part of a single homogeneous equilibrium. Sugar, salt, water, dye, air, and blood are examples of homogeneous mixes. A heterogeneous mixture has a distinct identifying quality in which the various components of the combination can be seen. It is a reaction system in which the reactants and products are found in two or more phases. Pizza, cookies, rocks, and other such items are examples of heterogeneous mixtures.
Question 5: What is meant by the common ion effect and its role?
The common ion effect depicts the changes that occur when ions are injected into a solution containing the same ion. When common ions are added to a solution, the solubility of a molecule decreases due to a shift in the equilibrium. The common ion effect is crucial in the modulation of buffers. A buffering solution contains either an acid or a base, either of which is accompanied by its conjugate counterpart. The pH of the solution changes as additional conjugate ions are added. This effect must be considered when evaluating solution equilibrium when common ions are introduced.
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