Chemical Equilibrium Practice Flashcards

A set of 70 vocabulary-style flashcards covering Chemical Equilibrium, Le Châtelier’s Principle, and the relationship between Free Energy and Equilibrium based on Chapter 14.

Reversible reaction

A reaction that can easily travel in either direction, often denoted by the double arrow symbol.

Chemical equilibrium

Achieved when the rates of the forward and reverse reactions are equal and the concentrations of the reactants and products remain constant.

Equilibrium state

A state in which there are no observable changes as time goes by.

Law of Mass Action

For the reaction $aA + bB \rightleftharpoons cC + dD$, the equilibrium constant is expressed as $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$.

$$K \gg 1$$

Indicates that at equilibrium, the products are much more favored than the reactants.

$$K \ll 1$$

Indicates that at equilibrium, the reactants are much more favored than the products.

Molar concentration ($C$) of a gas

Represented as $C = \frac{n}{V}$, or the number of moles of gas per unit volume.

Relationship between Pressure and Concentration

Expressed by the equation $P = CRT$, derived from the ideal gas equation.

$$K_c$$

The equilibrium constant calculated using molar concentrations of the reacting species.

$$K_p$$

The equilibrium constant calculated using the partial pressures of the gases involved.

Homogeneous equilibrium

Applies to chemical reactions in which all reacting species are in the same phase.

Heterogeneous equilibrium

Applies to reactions in which reactants and products are in different phases.

Relationship between $K_p$ and $K_c$

Expressed as $K_p = K_c(RT)^{\Delta n}$, where in most cases they are not equal.

Equilibrium constant units

It is general practice not to include units; the equilibrium constant is a dimensionless quantity.

Non-included species in $K$

The concentrations of pure solids and pure liquids are not included in the equilibrium constant expression.

Reciprocal rule

When the equation for a reversible reaction is written in the opposite direction, the new equilibrium constant is the reciprocal of the original ($1/K$).

Multiple equilibrium rule

If a reaction is the sum of two or more reactions, the overall equilibrium constant is the product of the individual constants ($K_c = K'_c \times K''_c$).

Reaction quotient ($Q_c$)

Calculated by substituting the initial concentrations of reactants and products into the equilibrium constant expression.

$$Q_c > K_c$$

The system proceeds from right to left (towards reactants) to reach equilibrium.

$$Q_c = K_c$$

The system is already at equilibrium.

$$Q_c < K_c$$

The system proceeds from left to right (towards products) to reach equilibrium.

Le Châtelier’s Principle

If an external stress is applied to a system at equilibrium, the system adjusts to partially offset the stress.

Increase concentration of product

Causes the equilibrium to shift to the left.

Decrease concentration of product

Causes the equilibrium to shift to the right.

Increase concentration of reactant

Causes the equilibrium to shift to the right.

Decrease concentration of reactant

Causes the equilibrium to shift to the left.

Increase pressure (gas system)

Shifts the equilibrium toward the side with the fewest moles of gas.

Decrease pressure (gas system)

Shifts the equilibrium toward the side with the most moles of gas.

Decrease volume (gas system)

Shifts the equilibrium toward the side with the fewest moles of gas.

Increase volume (gas system)

Shifts the equilibrium toward the side with the most moles of gas.

Increase temperature (Exothermic)

Causes the equilibrium constant $K$ to decrease.

Decrease temperature (Exothermic)

Causes the equilibrium constant $K$ to increase.

Increase temperature (Endothermic)

Causes the equilibrium constant $K$ to increase.

Decrease temperature (Endothermic)

Causes the equilibrium constant $K$ to decrease.

Catalyst effect on $K$

Adding a catalyst does not change the equilibrium constant.

Catalyst effect on position

Adding a catalyst does not shift the position of an equilibrium system.

Catalyst effect on rate

Lowers the activation energy ($E_a$) for both forward and reverse reactions, allowing the system to reach equilibrium sooner.

$$\Delta G < 0$$

The reaction is spontaneous in the forward direction.

$$\Delta G = 0$$

The system is at equilibrium.

$$\Delta G > 0$$

The reaction is nonspontaneous in the forward direction but spontaneous in reverse.

Standard Free Energy relationship ($\Delta G^\circ$ and $K$)

Expressed as $\Delta G^\circ = -RT \ln(K)$.

$$K > 1, \Delta G^\circ < 0$$

Products are more abundant at equilibrium.

$$K < 1, \Delta G^\circ > 0$$

Reactants are more abundant at equilibrium.

$$K = 1, \Delta G^\circ = 0$$

Reactants and products are comparably abundant at equilibrium.

Scenario: $\Delta H > 0$ and $\Delta S > 0$

Process is spontaneous at high temperatures and nonspontaneous at low temperatures.

Scenario: $\Delta H < 0$ and $\Delta S < 0$

Process is spontaneous at low temperatures and nonspontaneous at high temperatures.

Scenario: $\Delta H > 0$ and $\Delta S < 0$

Process is nonspontaneous at all temperatures.

Scenario: $\Delta H < 0$ and $\Delta S > 0$

Process is spontaneous at all temperatures.

Dimensionless quantity

The equilibrium constant has no units.

Initial concentrations

The concentrations given before the system reaches equilibrium, used to calculate $Q_c$.

Change in concentration ($x$)

A single unknown used in calculations to represent the amount of a species consumed or produced to reach equilibrium.

Gas phase concentrations

Can be expressed in molarity ($M$) or atmospheres ($atm$).

Ideal Gas Constant ($R$)

In free energy calculations, it is equal to $8.314\,J/mol\,K$.

Kinetics vs. Equilibrium

Equilibrium does not have anything to do with the speed of the reaction (kinetics).

Static vs. Dynamic

Equilibrium is not static or unchanging; it involves active forward and reverse processes at equal rates.

Solvent in equilibrium expressions

Concentrations of solvents do not appear in the equilibrium constant expression.

Balanced equation requirement

In quoting a value for the equilibrium constant, one must specify the balanced equation and the temperature.

$$n/V$$

The formula used to define the molar concentration of a gas.

Temperature Dependence ($K$)

Temperature is the only factor listed that changes the value of the equilibrium constant itself.

Pressure Dependence ($K$)

Changes in pressure may shift equilibrium but do not change the equilibrium constant.

Volume Dependence ($K$)

Changes in volume may shift equilibrium but do not change the equilibrium constant.

Ammonia Synthesis ($K_c$ expression)

For $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, $K_c = \frac{[NH_3]^2}{[N_2][H_2]^3}$.

$$(RT)^{\Delta n}$$

The conversion factor used between $K_c$ and $K_p$ where $\Delta n$ is moles of gaseous products minus moles of gaseous reactants.

Standard conditions

Defined as $1\,atm$ for gases and $1\,M$ for solutions in thermodynamic context.

$$-RT \ln(K)$$

The formula for $\Delta G^\circ$ which links thermodynamics to the equilibrium constant.

Le Châtelier Summary: Concentration Change

Shifts equilibrium: yes; Changes equilibrium constant: no.

Le Châtelier Summary: Temperature Change

Shifts equilibrium: yes; Changes equilibrium constant: yes.

Le Châtelier Summary: Catalyst Change

Shifts equilibrium: no; Changes equilibrium constant: no.

Le Châtelier Summary: Volume Change

Shifts equilibrium: yes (if gaseous moles differ); Changes equilibrium constant: no.

$$K = e^{-\Delta G^\circ / RT}$$

The rearranged formula used to solve for the equilibrium constant $K$ from the standard free energy change.