Chemical Equilibrium

Overview of Chemical Equilibrium

When a chemical reaction can proceed in both directions, it does not usually go to complete consumption of all reactants. Instead, after some time it reaches a state in which the observable properties of the system no longer change with time, even though the forward and reverse reactions are still occurring. This state is called chemical equilibrium.

Equilibrium is central in chemistry because many important reactions—industrial syntheses, biological pathways, acid–base and redox reactions—operate under conditions where the forward and reverse processes compete.

This chapter introduces what is meant by chemical equilibrium at a general level. The detailed treatment of reversibility, the microscopic establishment of equilibrium, the law of mass action, and the connection to Gibbs free energy are covered in the subsequent chapters of this section.

Dynamic vs. Static View of Equilibrium

In everyday language, “equilibrium” often suggests rest or complete standstill. Chemical equilibrium is different: it is a dynamic equilibrium.

• Static equilibrium (for comparison, e.g. a book resting on a table):
• No motion.
• No processes going on that would change the state.
• Dynamic chemical equilibrium:
• The forward reaction (reactants → products) and the reverse reaction (products → reactants) both occur.
• At equilibrium, the rates of the forward and reverse reactions are equal.
• Because these rates are equal, the macroscopic quantities (concentrations, pressure, color, pH, etc.) remain constant over time.

So at equilibrium:

• The reaction has not stopped.
• The amount of reactants and products is no longer changing with time, although individual molecules keep reacting.

This is sometimes summarized as:

Chemical equilibrium is reached when the forward and reverse reaction rates are equal, leading to constant macroscopic composition.

Conditions for Chemical Equilibrium

Chemical equilibrium occurs only under certain conditions:

• The reaction must be reversible under the given conditions (so that products can form reactants again).
• The system must be closed with respect to the reacting species:
• No net exchange of reactants or products with the surroundings (no continuous loss of a component, no continuous feed).
• External conditions such as temperature and pressure must be kept constant (or change so slowly that the system can adjust).

If these conditions are met, a reversible reaction mixture evolves until it reaches a state where the macroscopic properties are constant in time: equilibrium.

Macroscopic Characteristics of Chemical Equilibrium

From a laboratory or industrial perspective, a reaction mixture at equilibrium can be recognized by several characteristic features:

• Constant observable properties
  After some time, measurable properties no longer change:
• concentrations (for solutions or gases at fixed volume),
• partial pressures (for gas mixtures),
• intensity of color,
• pH in acid–base systems,
• electrical conductivity, etc.
• Reproducibility from either direction
  Equilibrium is the same state regardless of whether we start from:
• pure reactants,
• pure products,
• or any mixture in between,
  as long as temperature, pressure, and total amounts of substances are the same.
• Dependence on external conditions
  Changing temperature, pressure, or certain other conditions shifts the equilibrium composition to a new state (this behavior is discussed in later chapters on shifting equilibria).

At this level, it is important to distinguish:

• Equilibrium state: No further macroscopic change at fixed conditions.
• Equilibrium composition: The particular distribution of reactants and products (e.g. “30% A, 70% B by amount”) that characterizes that equilibrium.

Different reactions, and the same reaction under different conditions, will have different equilibrium compositions.

Extent of Reaction and Position of Equilibrium

Not all equilibria have equal amounts of reactants and products. The position of equilibrium tells us qualitatively whether, at equilibrium, reactants or products are predominant.

For a general reversible reaction
$$
\text{a A} + \text{b B} \rightleftharpoons \text{c C} + \text{d D}
$$

one observes in practice:

• Product-favored equilibrium:
• The equilibrium mixture contains mostly products (C and D).
• Only small amounts of reactants (A and B) remain.
• One may loosely say the equilibrium “lies to the right”.
• Reactant-favored equilibrium:
• The equilibrium mixture contains mostly reactants (A and B).
• Only small amounts of products (C and D) are present.
• One may loosely say the equilibrium “lies to the left”.
• Intermediate equilibrium:
• Significant amounts of both reactants and products coexist.

The quantitative description of this position (using the equilibrium constant defined via the law of mass action) is addressed in a later chapter. It is sufficient here to recognize:

• Every reversible reaction has some characteristic equilibrium composition at given conditions.
• This position can be shifted by changing temperature or pressure, or by changing certain reaction conditions (treated later).

Types of Chemical Equilibria

Chemical equilibrium can appear in various forms, depending on the reaction system:

Homogeneous Equilibria

In homogeneous equilibrium, all reactants and products are in the same phase (all gas, or all in one liquid solution).

Examples of typical situations:

• Gas phase reactions:
• All components gaseous; equilibrium described in terms of partial pressures.
• Solution reactions:
• All substances are dissolved (e.g. in water or an organic solvent); equilibrium described in terms of concentrations.

Homogeneous equilibria are particularly convenient for mathematical treatment and are the typical setting in which the law of mass action is introduced.

Heterogeneous Equilibria

In heterogeneous equilibrium, reactants and/or products are in different phases, e.g.:

• solid + gas,
• solid + liquid,
• multiple solids and a solution, etc.

In such systems:

• The composition of the fluid phase (gas or solution) changes until equilibrium is reached.
• The amounts of pure solids or pure liquids present may change, but their activities (a formal thermodynamic measure) are often treated as constant under typical conditions; this has consequences for the form of the equilibrium expression (discussed later).

Heterogeneous equilibria commonly occur in:

• dissolution and precipitation processes,
• gas–solid reactions,
• phase transitions (e.g. vapor–liquid equilibrium).

Chemical vs. Physical Equilibria

The concept of equilibrium is not limited to reactions forming new substances. A helpful comparison is with physical equilibria, in which only the phase or state of a substance changes, not its chemical identity.

Examples include:

• Vapor–liquid equilibrium:
  A liquid in a closed container partially evaporates; vapor also condenses back to liquid. At equilibrium, the rates of evaporation and condensation are equal and the vapor pressure is constant.
• Solid–liquid equilibrium:
  Ice in water at constant temperature and pressure can reach a state where melting and freezing occur at equal rates; the amount of ice and water remains constant in time.

From the standpoint of equilibrium:

• Both physical and chemical equilibria are dynamic.
• Both involve opposing processes happening at equal rates.
• Both result in constant macroscopic properties at fixed conditions.

Chemical equilibrium is specifically concerned with changes in chemical composition (bond-breaking and bond-forming processes), whereas physical equilibrium involves only changes in aggregation state or arrangement of molecules.

Equilibrium as a Consequence of Competing Processes

Many systems in chemistry can be thought of in terms of competing processes:

• formation vs. breakdown of a species,
• association vs. dissociation,
• protonation vs. deprotonation,
• oxidation vs. reduction, etc.

Where both directions can occur repeatedly and reversibly, and the system is closed under constant conditions, it tends to a state where the net effect of all competing processes is zero. That state is equilibrium.

For chemical reactions:

• The direct process is the forward reaction.
• The reverse process is the backward reaction.
• Equilibrium is the balance point between these two.

Understanding equilibrium in these general terms helps unify different topics later in the course—acid–base equilibria, solubility equilibria, redox equilibria, complex formation equilibria, and more.

Multiple Equilibria and Coupled Reactions

In realistic systems, more than one reversible process often occurs at the same time. For instance:

• A substance may undergo both acid–base reactions and complex formation.
• A gas may dissolve in water, react with the solvent, and partition between gas and aqueous phases.

In such cases:

• Each individual reaction has its own equilibrium that it would reach if isolated.
• When multiple reactions share components, they are coupled, and the final composition must satisfy all relevant equilibrium conditions simultaneously.

This leads to equilibrium networks, where:

• The presence of one equilibrium can influence the apparent position of another.
• Adding or removing a species can shift several equilibria at once.

The systematic treatment (using equilibrium constants and mass balances) is handled in later chapters. Here, it is enough to note that:

• Real systems often involve overlapping equilibria.
• The overall observed behavior results from the combination of several underlying equilibria.

Conceptual Summary

Key points for a foundational understanding of chemical equilibrium:

• Chemical equilibrium is a dynamic state:
• Forward and reverse reactions occur continuously.
• Their rates are equal, so there is no net change in composition.
• At equilibrium:
• Macroscopic properties are constant in time at fixed external conditions.
• The system is in a balance determined by reaction and conditions.
• Equilibrium is:
• Condition-dependent (changes in temperature, pressure, and certain other conditions alter the equilibrium composition).
• Reproducible from any starting mixture under the same conditions.
• The position of equilibrium indicates whether reactants or products predominate at equilibrium; this position is later quantified by the equilibrium constant.
• Chemical equilibrium is conceptually similar to physical equilibrium (e.g. phase equilibria), but involves changes in chemical identity rather than just phase.

Subsequent chapters in this section will build on this conceptual foundation by examining:

• how reversibility appears in reaction schemes,
• how equilibrium is established in time,
• how the law of mass action is formulated and used,
• and how equilibrium is connected to thermodynamics through Gibbs free energy.