Year 12 Chemistry ATAR: Chemical Equilibrium Systems and Reversible Reactions

Comprehensive practice flashcards covering open, closed, and isolated systems, activation energy barriers, reversibility, dynamic equilibrium characteristics, collision theory, and graphical interpretations from Year 12 Chemistry ATAR.

System

All the chemical species (reactants and products) involved in a reaction. Everything outside the system is defined as the surroundings.

Open System

A chemical system that allows both matter and energy to be exchanged freely with the surroundings.

Closed System

A chemical system that allows energy to be exchanged with the surroundings, but does not allow matter to enter or leave.

Isolated System

A chemical system in which neither matter nor energy can move into or out of the system (such as an ideal thermos flask).

Reversible Reaction

A reaction that can proceed simultaneously in both directions—reactants forming products and products reacting to reform reactants—indicated by a double arrow $\rightleftharpoons$.

Forward Reaction

The reaction that proceeds from left to right as written in a chemical equation ($\text{reactants} \rightarrow \text{products}$).

Reverse Reaction

The reaction that proceeds from right to left as written in a chemical equation ($\text{products} \rightarrow \text{reactants}$).

Dynamic Equilibrium

A state established in a closed system when the rate of the forward reaction equals the rate of the reverse reaction, keeping reactant and product concentrations constant while both reactions continue occurring.

Macroscopic Property

A visible or measurable property of bulk matter—such as colour, pressure, temperature, density, and mass—which remains constant once a system reaches dynamic equilibrium.

Position of Equilibrium

The relative amounts or concentrations of reactants and products present in a system at dynamic equilibrium, indicating how far a reaction proceeds.

Steady State

A condition in which concentrations and macroscopic properties remain constant over time; it includes dynamic equilibrium in closed systems as well as open systems where substances enter and leave at identical rates.

Readily Reversible Reaction

A reaction that occurs in both directions at significant rates because the activation energies for both the forward and reverse reactions are low and the enthalpy change ($\Delta H$) is relatively small.

Practically Irreversible Reaction

A reaction that effectively proceeds in one direction only because the activation energy in the reverse direction is extremely high (usually coupled with a large $\Delta H$), or because a product continuously escapes an open system.

Homogeneous Equilibrium

An equilibrium system in which all participating reactants and products are in the same physical phase (for example, all gases or all aqueous solutions).

Heterogeneous Equilibrium

An equilibrium system in which the participating species exist in two or more different physical phases (such as a solid in equilibrium with a gas or solution).

Dynamic Nature (Equilibrium)

The characteristic of chemical equilibrium describing how forward and reverse reactions never cease (rates are non-zero), demonstrated experimentally when isotopic tracers exchange between reactants and products.

Concentration-Time Relationship at Equilibrium

On a concentration-time graph, reactant curves fall and product curves rise proportionally according to stoichiometric coefficients until time $E_1$, after which all curves become completely flat (constant, though not necessarily equal).

Catalyst Effect on Equilibrium

A substance that lowers the activation energies of both the forward and reverse reactions by the exact same amount without altering $\Delta H$, $K_c$, or the final position of equilibrium, merely reducing the time required to reach equilibrium.

How are a chemical "system" and its "surroundings" defined?

The system includes all the chemicals involved in the reaction (both reactants and products). The surroundings include everything else outside the system.

What is the difference between an open system, a closed system, and an isolated system?

An open system allows both matter and energy to be exchanged with the surroundings. A closed system allows energy to be exchanged, but not matter. An isolated system allows neither matter nor energy to move in or out.

Why can dynamic chemical equilibrium only be established in a closed system?

In an open system, products or matter can escape. If a product escapes, the reverse reaction cannot keep pace with or catch up to the forward reaction, causing concentrations to continually change rather than remain constant.

What distinguishes a physical change from a chemical change regarding reversibility?

In a physical change, no new chemical substances form (e.g., changes of state or dissolving) and these processes are reversible. In a chemical change, existing bonds break and new bonds form to produce new substances, and only some chemical reactions are reversible.

What is a reversible reaction, and how is it symbolized in a chemical equation?

A reversible reaction is one where reactants convert into products (forward reaction: left to right) while products simultaneously react to reform reactants (reverse reaction: right to left). It is denoted with a double arrow ($\rightleftharpoons$).

How is the activation energy of a reverse reaction ($E_a\text{(reverse)}$) determined from the forward activation energy ($E_a\text{(forward)}$) and enthalpy change ($\Delta H\text{(forward)}$)?

$E_a\text{(reverse)} = E_a\text{(forward)} - \Delta H\text{(forward)}$. For an exothermic forward reaction, this means $E_a\text{(reverse)}$ equals $E_a\text{(forward)}$ plus the absolute value of $\Delta H$.

Why is the neutralisation reaction $HCl(aq) + NaOH(aq) \rightarrow H_2O(l) + NaCl(aq)$ considered practically irreversible at Standard Laboratory Conditions (SLC)?

Neutralisation is highly exothermic ($\Delta H \approx -57\,kJ\,mol^{-1}$). The forward reaction has a very low activation energy, while the reverse activation energy ($E_a\text{(reverse)} = E_a\text{(forward)} + 57\,kJ\,mol^{-1}$) is extremely high, meaning almost no collisions between product molecules possess sufficient kinetic energy to react in reverse at SLC.

How does the addition of a catalyst affect the forward activation energy, reverse activation energy, $\Delta H$, and $K_c$ of a reversible reaction?

A catalyst lowers both $E_a\text{(forward)}$ and $E_a\text{(reverse)}$ by the exact same amount by providing an alternative pathway with a lower peak. It causes no change to $\Delta H$ or the equilibrium constant $K_c$.

Using collision theory, describe the changes occurring in the forward and reverse reaction rates from Time 0 until dynamic equilibrium is established at Time E1.

At Time 0, reactant concentration is at its highest, so the forward collision frequency and forward rate are at their peak, while the reverse rate is zero. As reactants react, their concentration drops, lowering forward collision frequency and slowing the forward rate. Concurrently, product concentrations increase, raising product collisions and increasing the reverse rate. At E1, the forward and reverse rates become equal.

What are the four primary characteristics of a chemical system at dynamic equilibrium?

1. It is a closed system with a reversible change. 2. The forward reaction rate equals the reverse reaction rate (neither is zero). 3. The concentrations of all reactants and products remain constant (not necessarily equal). 4. Macroscopic properties (colour, pressure, temperature, density, mass) remain constant.

How does a "steady state" differ from "dynamic equilibrium"?

In both states, macroscopic properties and concentrations remain constant. However, dynamic equilibrium requires a closed system where forward and reverse rates are equal. A steady state can also occur in an open system if matter enters and leaves at the same rate (e.g., an industrial flow reactor or an open lake).

Explain how the concentration-time curves in the grid below reflect the stoichiometry of the reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$.

Following the $1 : 3 : 2$ mole ratio, the concentration of $H_2$ decreases by three times the amount that $N_2$ decreases, and the concentration of $NH_3$ increases by exactly double the decrease in $N_2$, before all species flatten into constant concentrations at equilibrium.

In the concentration-time graph shown below for $H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$, why does $[HI]$ increase by double the magnitude that $[H_2]$ decreases?

The changes in concentration follow the stoichiometric mole ratio of the chemical equation ($1 : 1 : 2$). For every $1\text{ mol}$ of $H_2$ (and $I_2$) consumed, $2\text{ mol}$ of $HI$ are produced.

What is the difference between homogeneous and heterogeneous equilibria, and how are pure solids and liquids represented?

In a homogeneous equilibrium, all reactants and products exist in the same phase. In a heterogeneous equilibrium, species exist in more than one phase. Because pure solids and pure liquids have constant concentrations, they are not plotted on concentration-time graphs and are omitted from equilibrium constant ($K$) expressions.

In the Haber process ($N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$), why does injecting nitrogen enriched with isotope $^{15}N$ result in $^{15}N$ being detected in both $N_2$ and $NH_3$?

This proves the equilibrium is dynamic: the forward and reverse reactions continue to occur simultaneously. $^{15}N$ from the injected $N_2$ reacts forward to form $^{15}N$-labelled $NH_3$, which subsequently decomposes via the reverse reaction to regenerate $N_2$, eventually establishing constant isotopic proportions.

For the closed system reaction $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) + \text{heat}$, state how pressure and temperature change from Time 0 to Time E1.

Pressure decreases because $3\text{ mol}$ of gaseous reactants convert to $2\text{ mol}$ of gaseous product, reducing the total moles of gas. Temperature increases because the forward reaction is exothermic, releasing thermal energy into the closed system.

For the physical equilibrium $I_2(s) \rightleftharpoons I_2(g)$ in a sealed jar, explain why the purple vapour darkens initially and then stays constant.

Initially, the forward rate of sublimation exceeds the reverse rate of deposition, causing $[I_2(g)]$ to rise and darkening the purple colour. As $[I_2(g)]$ rises, the deposition rate increases until sublimation and deposition rates become equal at equilibrium, maintaining a constant $[I_2(g)]$ and colour intensity.

What does the "position of equilibrium" signify, and what do large versus small values of $K$ convey about it?

The position of equilibrium describes the relative concentrations of reactants and products present at equilibrium. A large $K$ (e.g., $10^4$) indicates the equilibrium lies to the right (mostly products), while a small $K$ (e.g., $10^{-8}$) indicates it lies to the left (mostly reactants). It reflects how far a reaction proceeds, not how fast.