7.1.1. Reversible Reactions and Establishing Equilibrium

● Reversible reactions are chemical processes that can proceed in both the forward and the reverse directions. In a simple example, consider a reaction in which A reacts with B to form C and D: A + B ⇌ C + D At any moment, the forward reaction rate (rate at which A and B form C and D) and the reverse reaction rate (rate at which C and D revert to A and B) may be different.

Reversible reactions are unique in that they can go both ways—reactants can turn into products and products can revert back into reactants. For example, if we have a reaction where substances A and B create substances C and D, the reaction can move from left to right (A & B turning into C & D) or from right to left (C & D turning back into A & B). At any given moment, the speed of formation of products doesn’t always equal the speed at which they revert to reactants.

Think of a reversible reaction like cooking pasta. You boil water (A) and add pasta (B), which eventually cooks into delicious spaghetti (C) and sauce (D). If you don’t like the outcome, you can also return to the initial state by draining the pasta and adding different ingredients. Similar to the way pasta can be cooked or uncooked, in chemical reactions, substances can switch back and forth.

● Dynamic equilibrium occurs in a closed system when the rate of the forward reaction equals the rate of the reverse reaction. At that point: ○ The concentrations (or partial pressures) of all species remain constant in time, even though reactant molecules continue to form products and vice versa. ○ There is no net change in composition, but molecular exchange continues.

Dynamic equilibrium is a state that occurs in a closed system where two opposing processes happen at the same rate. For instance, as A and B transform into C and D, the newly created products C and D can also revert back to A and B. The key here is that, although the individual molecules are continuously transforming between states, the overall amounts of each substance remain constant. This state of balance is crucial in chemical processes as it shows that while reactions are still taking place, there’s no overall change in concentration.

Imagine a busy airport. Planes are constantly landing and taking off—just like molecules exchanging between states in a chemical reaction. When the number of planes arriving equals the number of planes leaving, the overall number of planes in the airport remains constant, despite the constant movement. This is akin to how dynamic equilibrium works in chemical systems.

● Example: In a sealed container at a fixed temperature, let nitrogen monoxide decompose into nitrogen and oxygen: 2 NO(g) ⇌ N₂(g) + O₂(g) Initially, only NO might be present, so the forward reaction (2 NO → N₂ + O₂) proceeds and produces N₂ and O₂. As N₂ and O₂ accumulate, however, they react to form NO. Eventually, the rates of “2 NO → N₂ + O₂” and “N₂ + O₂ → 2 NO” become equal. At that point, concentrations of NO, N₂, and O₂ no longer change, and the system is at equilibrium.

In this example, we start with only nitrogen monoxide (NO) which breaks down into nitrogen (N₂) and oxygen (O₂) in a closed system. Initially, the forward reaction dominates as NO molecules transform into N₂ and O₂. As these products form, they can revert back to NO. Over time, the rates of formation of N₂ and O₂ start matching the rates of NO reproduction from N₂ and O₂, leading to an equilibrium state where amounts of each substance remain stable, illustrating the concept of reversible reactions in dynamic equilibrium.

Consider a simple balance scale with weights. Initially, you might have more weight on one side, tipping it. As you slowly distribute weights more evenly, the scale would eventually balance out, much like the equilibrium where the rates of forward and reverse reactions equalize over time.

● Homogeneous versus heterogeneous equilibria ○ Homogeneous equilibrium: All reactants and products are in the same phase (all gases or all dissolved in the same solvent). ○ Heterogeneous equilibrium: The reacting species exist in two or more phases (for instance, a solid metal reacting with a gas, or a solid in equilibrium with its dissolved ions in water). In heterogeneous equilibria, only the concentrations of species in the fluid phase (gas or aqueous) appear in the equilibrium expression; solids and pure liquids are omitted (treated as having constant activity = 1).

Equilibrium can either be homogeneous or heterogeneous based on the states of the reacting substances. In a homogeneous equilibrium, all substances are in the same phase, like gases or solutions; for example, mixing gases in a container. In contrast, a heterogeneous equilibrium involves different phases; for example, when a solid and gas react or when a solid dissolves in water. In the case of heterogeneous equilibria, we focus on concentrations of phases that interact to establish equilibrium, focusing on gases or ions in solution since solids and pure liquids do not change concentration significantly, thus assumed constant.

Think of homogeneous equilibrium like a smoothie where all the ingredients (fruits, yogurt, and milk) are mixed seamlessly into one phase. Heterogeneous equilibrium is like a salad, where you have distinct layers of greens, croutons, and dressing, all physically separate yet interacting in a bowl. In chemistry, we focus on the active parts—liquids and gases—when calculating equilibria in mixtures.

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