Question

Define homogeneous equilibrium and heterogeneous equilibrium. Give two examples of each.

Short answer

Homogeneous equilibria involve same-phase components; heterogeneous equilibria involve different-phase components. Examples include ammonia synthesis (homogeneous) and calcium carbonate decomposition (heterogeneous).

Step-by-step solution

Define Homogeneous Equilibrium

A homogeneous equilibrium occurs in a chemical reaction where all the reactants and products are in the same phase. This means that the substances involved are all either gaseous or all in solution. For example, the equilibrium established in the reaction of nitrogen and hydrogen gases to form ammonia is homogeneous:\[N_2 (g) + 3H_2 (g) ightleftharpoons 2NH_3 (g)\]

Examples of Homogeneous Equilibrium

An example of homogeneous equilibrium is the Haber process:\[N_2 (g) + 3H_2 (g) ightleftharpoons 2NH_3 (g)\]Another example is the reaction between acetic acid and ethanol to form ethyl acetate and water, which all occur in a liquid phase:\[CH_3COOH (l) + C_2H_5OH (l) ightleftharpoons CH_3COOC_2H_5 (l) + H_2O (l)\]

Define Heterogeneous Equilibrium

A heterogeneous equilibrium occurs in a chemical reaction where the reactants and products are in different phases. Commonly, this involves a combination of solids, liquids, and gases. An example is the equilibrium between calcium carbonate, calcium oxide, and carbon dioxide:\[CaCO_3 (s) ightleftharpoons CaO (s) + CO_2 (g)\]

Examples of Heterogeneous Equilibrium

An example of heterogeneous equilibrium is the decomposition of calcium carbonate:\[CaCO_3 (s) ightleftharpoons CaO (s) + CO_2 (g)\]Another example is the formation of sulfur trioxide from sulfur dioxide and oxygen in the presence of a solid catalyst, which typically involves solids and gases:\[2SO_2 (g) + O_2 (g) ightleftharpoons 2SO_3 (g)\]

Key concepts

Homogeneous Equilibrium

In chemistry, equilibrium is a vital concept where the rates of the forward and reverse chemical reactions are equal, resulting in stable concentrations of reactants and products. **Homogeneous equilibrium** specifically occurs when all the reactants and products are in the same phase of matter, such as all being gases or all being in solution.

This type of equilibrium is commonly observed in reactions where substances can freely mix with each other.
Here are key points to understand about homogeneous equilibrium:

• All reacting species exist in the same phase, typically either gaseous or liquid.
• The concentration of reactants and products remains constant over time.
• It is common in reactions involved in industrial applications, such as in the production of ammonia.

For example, the **Haber process** is a homogeneous equilibrium as all species involved exist in the gaseous state. The reaction can be represented as:\[N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)\]Another common example is the reaction of **acetic acid with ethanol** to form ethyl acetate and water, where all components are in a liquid phase:\[CH_3COOH(l) + C_2H_5OH(l) \rightleftharpoons CH_3COOC_2H_5(l) + H_2O(l)\]Understanding homogeneous equilibrium helps chemists predict how changes in conditions like pressure and temperature can affect the balance of a reaction.

Heterogeneous Equilibrium

While homogeneous equilibrium involves all components being in the same state, **heterogeneous equilibrium** involves two or more phases.

This could involve solids and gases, or liquids and solids, among others. Mixed phases are a key characteristic of heterogeneous equilibrium scenarios.

• It often involves solids that are insoluble in the liquid or gaseous reactants/products.
• The amount of the solid phase usually does not affect the equilibrium position.
• These reactions are important in areas like metallurgy and geology.

One classic example is the **decomposition of calcium carbonate (chalk):**\[CaCO_3(s) \rightleftharpoons CaO(s) + CO_2(g)\]Here, both solid and gaseous states are involved. In another example, consider the formation of sulfur trioxide, which requires a solid catalyst and involves gases:\[2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)\]Heterogeneous equilibria play a crucial role in processes such as the evolution of carbon dioxide in open carbonate systems like lakes and soils, impacting environmental and engineering systems.

Reaction Examples

Understanding real-world applications and examples of chemical reactions helps to contextualize theoretical knowledge.
Here are common examples for both homogeneous and heterogeneous equilibria.
Homogeneous equilibrium examples include:

• **Haber Process:** A well-known method for synthesizing ammonia from nitrogen and hydrogen gases. This is a critical reaction for agricultural fertilizers.
• **Esterification Reaction:** Mixing acetic acid with ethanol to produce ethyl acetate and water, all in a liquid phase.

Heterogeneous equilibrium examples include:

• **Calcium Carbonate Decomposition:** This reaction involves multiple phases where calcium carbonate decomposes into calcium oxide (solid) and carbon dioxide (gas).
• **Sulfur Trioxide Formation:** Occurring in the presence of a solid catalyst but with gaseous reactants and products, it demonstrates the interaction between different phases.

These examples are not just academic exercises—they are crucial in industrial applications, environmental science, and more.
Recognizing the type of equilibrium helps determine how one might control or predict the reaction's behavior under varying conditions.

Phase of Matter

To fully grasp the idea of chemical equilibria, it's essential to understand the **phase of matter** each component involved occupies.

The phase—whether solid, liquid, or gas—affects not just the equilibrium process but also the calculations and implications of the reactions.
For homogeneous equilibria:

• Reactions occur fully within a single phase, typically gas or liquid.
• Changes in pressure or concentration can significantly influence the equilibrium.

For heterogeneous equilibria:

• Multiple phases interact, often involving solids.
• The concentration of substances solely in solid form is considered constant.
• These systems often involve catalysts that remain in the solid phase.

Understanding the phase characteristics is crucial in predicting reaction behaviors and designing processes for chemical manufacturing and other applications.
It shows why some reactions reach equilibrium quickly, while others may take longer, depending on the complexity of the phase interactions involved.