Ch. 18: Le Chatelier's Principle

Le Chatelier's Principle

Reversible Reactions

  • Reactions can bounce back and forth, forming products and reforming reactants; these are known as reversible reactions.

  • Reversible reactions occur simultaneously in both directions.

  • Example: A + B \rightleftharpoons C

Chemical Equilibrium

  • At chemical equilibrium, there is no net change in the actual amounts of the components of the system.

  • The rates of the forward and reverse reactions are equal at chemical equilibrium.

  • The concentrations of components on both sides of the chemical equation are not necessarily the same and can be dramatically different.

  • Analogy: Escalators where the number of people using the up escalator is the same as the number of people using the down escalator.

  • The number of people upstairs does not have to equal the number of people downstairs; only the transfer between floors must be consistent.

Which Direction is Favored?

  • The equilibrium position of a reaction is given by the concentrations of the system’s components at equilibrium.

  • The arrow direction indicates whether the components on the left or right side of a reversible reaction are at a higher concentration.

  • If A reacts to give B and the mixture at equilibrium contains more of B (e.g., 1% of A vs. 99% of B), the formation of B is favored.

  • If the mixture contains 99% of A and 1% of B at equilibrium, then the formation of A is favored.

  • Forward direction favored: A \rightleftharpoons B (99% A, 1% B)

  • Reverse direction favored: A \rightleftharpoons B (1% A, 99% B)

Reversibility vs. Reality

  • In principle, almost all reactions are reversible to some extent under the right conditions.

  • In practice, one set of components is often so favored at equilibrium that the other set cannot be detected.

  • If one set of components has established equilibrium by converting mostly into products, the reaction has gone to completion.

  • When no products can be detected, one can say there is no reaction.

  • Reversible reactions occupy a middle ground between the theoretical extremes of irreversibility and no reaction.

  • The addition of a catalyst will speed up forward and reverse reactions equally by reducing the energy needed to activate the reaction in both forward and reverse directions.

  • Catalysts do not affect the amount of reactants and products present at equilibrium; they simply decrease the time it takes to establish equilibrium.

Equilibrium Expression

  • Chemists can express the equilibrium position in terms of a numerical constant.

  • The equilibrium constant shows the relationship between the amount of product and reactant at equilibrium.

  • For a hypothetical reaction: aA + bB \rightleftharpoons cC + dD

Mass Action Expression

  • An expression to show the ratio of product concentrations to reactant concentrations.

  • \frac{[C]^c [D]^d}{[A]^a [B]^b}

  • The concentration of each substance is raised to a power equal to the number of moles of that substance in the balanced reaction equation.

  • Square brackets indicate concentration in Molarity (mol/L).

Equilibrium Constant

  • The constant is dependent on the temperature; if the temperature changes, so does the constant.

  • The resulting ratio of the equilibrium is called the equilibrium constant or K_{eq}.

  • K_{eq} = \frac{[C]^c [D]^d}{[A]^a [B]^b}

Examples of Mass Action Expressions

  • 2SO2(g) + O2(g) \rightleftharpoons 2SO_3(g)

  • Bi2S3(s) \rightleftharpoons 2Bi^{+3}(aq) + 3S^{-2}(aq)

Equilibrium Constant—Who Cares?

  • Equilibrium constants provide valuable chemical information.

  • They show whether the products or the reactants are favored in a reaction (spontaneous or non-spontaneous).

  • Always written as a ratio of products over reactants.

  • A value of K_{eq} > 1 means that products are favored.

  • K_{eq} < 1 means that reactants are favored.

  • K_{eq} > 1, products favored at equilibrium

  • K_{eq} \leq 1, reactants favored at equilibrium

Sample Problem 1

Dinitrogen tetroxide (N2O4), a colorless gas, and nitrogen dioxide (NO2), a brown gas, exist in equilibrium with each other according to the following equation: N2O4(g) \rightleftharpoons 2NO2(g)
A 1.0-liter gas mixture at 10°C at equilibrium contains 0.0045 mol N2O4 & 0.030 mol NO2. Write the mass action expression and calculate K{eq} for the reaction.

  • Known:

    • [N2O4] = 0.0045 \frac{mol}{1.0 L}

    • [NO_2] = 0.030 \frac{mol}{1.0 L}

  • Unknown:

    • Mass action expression = ?

    • K_{eq} = ?

  • At equilibrium, there is no net change in the amount of N2O4 or NO_2 at any given instant.

  • The only product of the reaction is NO_2, which has a coefficient of 2 in the balanced equation.

  • The only reactant N2O4 has a coefficient of 1 in the balanced equation.

  • The mass action expression is:
    \frac{[NO2]^2}{[N2O_4]^1}

  • K_{eq} = \frac{[0.030M]^2}{[0.0045M]^1} = 0.20

  • K_{eq} < 1, therefore the reaction doesn’t favor products.

Reaction Quotient

  • We can also determine if a reaction has reached equilibrium by calculating a reaction quotient (Q).

  • It’s like taking a snapshot of a reaction at a given time and interpreting how far along the reaction is.

  • Once the reaction quotient is solved, it is compared to the equilibrium constant.

    • If Q < K, movement toward equilibrium is towards the products.

    • If Q > K, movement toward equilibrium is towards the reactants.

    • If Q = K, reactants and products are at equilibrium.

Sample Problem

At a certain temperature, K{eq} = 55, and a reaction vessel contains a mixture with the following concentrations: [SO3] = 0.85 M, [NO] = 1.2 M, [SO2] = 1.5 M, [NO2] = 2.0 M.
SO3(g) + NO(g) \rightleftharpoons SO2(g) + NO_2(g)
Is the reaction at equilibrium, and if not, which direction will the reaction proceed?

  • Solve just as if you were solving for the equilibrium constant.

  • Then analyze the resulting quotient with the given K_{eq}.

  • Q = \frac{[SO2][NO2]}{[SO_3][NO]} = \frac{[1.5][2.0]}{[0.85][1.2]} = 2.94

  • Q = 2.94, which is < K_{eq} (55).

  • If Q < K{eq}, then the numerator of our quotient must increase; therefore, the reaction continues in order to increase [products] until it reaches K{eq}.

Le Chatelier’s Principle

  • There is a principle that can be studied to govern changes in equilibrium: Le Chatelier’s Principle.

  • Le Chatelier’s Principle states: “If a stress is applied to a system in dynamic equilibrium, the system changes to relieve the stress.”

  • Stresses are changes in temperature, pressure, concentration of reactants, or concentration of products.

Concentration & Equilibrium

  • Adjusting the concentrations of either reactants or products can have a dramatic impact on the equilibrium.

  • If we add more of reactant A to a system at equilibrium, the system will strive to reestablish equilibrium at a new equilibrium position.

  • The reaction will push to use up the extra A and generate more C.

  • [A] \uparrow, reaction will shift toward products.

  • Adjusting the concentrations of either reactants or products can have a dramatic impact on the equilibrium.

  • If we add more of product C to a system at equilibrium, the system will strive to reestablish equilibrium at a new equilibrium position.

  • The reaction will push to use up the extra C and generate more A and B.

  • [C] \uparrow, reaction will shift toward reactants.

Temperature Effects on Equilibrium

  • The impact of temperature changes on an equilibrium is dependent on if the process is endothermic or exothermic.

  • Endothermic processes use energy as a reactant, while exothermic processes produce energy.

  • K_{eq} is temperature-dependent.

  • Exothermic: If T↑, the equilibrium shifts left (250 kJ is a product).

  • Endothermic: If T↑, the equilibrium shifts right (energy is a reactant).In summary, understanding these shifts in equilibrium can help predict how changes in temperature will affect the direction of a reaction and its eventual products. This knowledge is crucial for optimizing reaction conditions in industrial applications and for understanding natural processes. This principle not only aids in maximizing yield in chemical manufacturing but also provides insight into temperature regulation in biological systems. Furthermore, applying Le Chatelier's Principle allows chemists to manipulate reaction conditions effectively, adjusting concentrations, pressure, and temperature to achieve desired outcomes. By recognizing how these variables interact, researchers can design experiments and processes that are more efficient and tailored to specific needs. This adaptability is particularly important in areas such as pharmaceuticals, where precise control over reaction conditions can lead to higher purity and efficacy of drug products.

Pressure & Equilibrium

  • If A, B, and C are all gases, then the equilibrium they establish is pressure-dependent.

  • When the pressure is increased, the system relieves the pressure by favoring the direction that produces fewer gas molecules.

  • Pressure is # of particles dependent; the more particles, the higher the pressure.

  • Fewer gas molecules will exert less pressure, so more product is formed, which overall reduces the pressure; this is a shift right.

  • Conversely, a decrease in pressure will favor the reaction that produces the most molecules, so we have a shift to the left.

  • P↑, this equilibrium shifts right

  • If P↓, this equilibrium shifts left