Chapter 16

How does a developing fetus get axygen in the womb?

What is dynamic equilibrium? Why is it called dynamic?

Glve the general expression for the equilibrium constant of the following generic reaction: $$ a A+b B=c C+d D $$

What is the significance of the equilibrium constant? What does a large equilibrium constant tell us about a reaction? A small one?

What happens to the value of the equilibrium constant for a reaction if the reaction equation is reversed? Multiplied by a constant?

If two reactions sum to an overall reaction, and the equilibrium constants for the two reactions are \(K_{1}\) and \(K_{2}\), what is the equlubrium constant for the overall reaction?

Explain the difference between \(K_{c}\) and \(K_{\mathrm{p}}\). For a given reaction, how are the two constants related?

Why are the concentrations of solids and liquids omitted from equilibrium expressions?

Does the value of the equilibrium constant depend on the inltial concentrations of the reactants and products? Do the equiIibrium concentrations of the reactants and products depend on their initial concentrations? Explain.

What is the definition of the reaction quotient \((Q)\) for a reaction? What does Q measure?

In what direction will a reaction proceed for each condition: a. \(QK ;\) and \(\mathrm{c} \quad Q=K ?\)

Many equllibelum calculations involve finding the equilibrium concentrations of reactants and products given their initial concentrations and the equllibrium constant. Outline the general procedure used in solving these kinds of problems.

What happens to a chemical system at equilibrium when that cqullibrium is disturbed?

What is the effect of a temperature change on a chemical reaction initially at equilabrium? How does the effect differ for an exothermic reaction compared to an endothermic one?

When this reaction comes to equilibrium, will the concentrations of the reactants or peoducts be greater? Does the answer to this question depend on the initial concentrations of the reac. tants and products? MISSED THIST Read Section 16.3; Watch KCV 16.3 $$ A(s)+B(s)=2 C(g) K_{c}-1.4 \times 10^{-5} $$

Consider the reaction: $$ 2 \mathrm{H}_{2} \mathrm{~S}(g) \rightleftharpoons 2 \mathrm{H}_{2}(\mathrm{~g})+\mathrm{S}_{2}(8) \quad K_{\mathrm{p}}=2.4 \times 10^{-4} \text {at } 1073 \mathrm{~K} $$ A reaction mixture contains 0.112 atm of \(\mathrm{H}_{2}, 0.055\) atm of \(\mathrm{S}_{2}\) and 0.445 atm of \(\mathrm{H}_{2} \mathrm{~S}\). Is the reaction mixture at equilibrium? If not, in what direction will the reaction proceed?

Consider the reaction and the associated equilibrium constant: $$ a A(g)+l B(g) \Longrightarrow c C(g) \quad K_{c}-5.0 $$ And the equilibrium concentrations of \(A, B,\) and \(C\) for each value of \(a, b,\) and \(c\). Assume that the initial concentrations of \(A\) and \(B\) are each \(1.0 \mathrm{M}\) and that no product is present at the begin. ning of the reaction. a. \(a-1 ; b-1 ; c-2\) b. \(a-1 ; b-1 ; c-1\) c. \(a-2 ; b-1 ; c=1\) (set up equation for \(x ;\) don't solve)

Consider the reaction: MISSED THIS? Read Section 16.8 : Watch KCV 16.8 , ME 16.9 $$ \begin{array}{r} \mathrm{NiO}(\mathrm{s})+\mathrm{CO}(g) \rightleftharpoons \mathrm{Ni}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{~g}) \\ K_{c}=4.0 \times 10^{3} \text { at } 1500 \mathrm{~K} \end{array} $$ If a mixture of solid nickel(II) axide and \(0.20 \mathrm{M}\) carbon monoxIde comes to equllibrium at \(1500 \mathrm{~K},\) what is the equilibrium concentration of \(\mathrm{CO}_{2} ?\)

Each reaction is allowed to come to equilibrium, and then the volume is changed as indicated. Predict the effect (shift right, shift left, or no effect) of the indicated volume change. MissED THis? Read Section 16.9 ; Watch KCV 16.9 a. \(\mathrm{I}_{2}(g) \rightleftharpoons 2 \mathrm{l}(\mathrm{s})\) (volume is increased) b. \(2 \mathrm{H}_{2} \mathrm{~S}(\mathrm{~g})=2 \mathrm{H}_{2}(\mathrm{~s})+\mathrm{S}_{2}(\mathrm{~g})\) (volume is decreased) c. \(\mathrm{I}_{2}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{ICl}(\mathrm{g})\) (volume is decreased)

Each reaction is allowed to come to equilibrium, and then the volume is changed as indicated. Predict the effect (shift right, shift left, or no effect) of the indicated volume change. a. \(\mathrm{CO}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightleftharpoons \mathrm{CO}_{2}(\mathrm{~g})+\mathrm{H}_{2}(g)\) (volume is decreased) c. \(\mathrm{CaCO}_{3}(\mathrm{~s}) \rightleftharpoons \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{~g})\) (volume is increased)

A system at equilatium contains \(I_{2}(\xi)\) at a pressure of 0.21 atm and \(1(g)\) at a pressure of 0.23 atm. The system is then comptessed to half its volume. Find the pressure of each gas when the system returns to equilabrium.

Consider the exothermic reaction: $$ \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{Q}_{2}(g) \rightleftharpoons \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}(\mathrm{~g}) $$ If you were trying to maximize the amount of \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) produced, which tactic might you try? Assume that the reaction mixture reaches equllubrium. a. increasing the reaction volume b. removing \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\) from the reaction mixture as it forms c. lowering the reaction temperature d. adding \(\mathrm{Cl}_{2}\)

Consider the endothermic reaction: $$ \mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{~g})+\mathrm{I}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{l}_{2}(\mathrm{~g}) $$ If you were trying to maximize the amount of \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{l}_{2}\) produced, which tactic might you try? Assume that the reaction mixture reaches equilibrium. a. decreasing the reaction volume b. removing \(1_{2}\) from the reaction mixture c. raising the reaction temperature d. adding \(\mathrm{C}_{2} \mathrm{H}_{4}\) to the reaction mixture

A sample of \(\mathrm{SO}_{3}\) is introduced into an evacuated sealed container and heated to \(600 \mathrm{~K}\). The following equilibrium is established: $$ 2 \mathrm{SO}_{3}(\xi) \rightleftharpoons 2 \mathrm{SO}_{2}(g)+\mathrm{O}_{2}(\mathrm{~g}) $$ The total pressure in the system is \(3.0 \mathrm{~atm}\), and the mole fraction of \(\mathrm{O}_{2}\) is 0.12 . Find \(\mathrm{K}_{\mathrm{P}}\)

100\. A reaction \(A(g) \rightleftharpoons B(s)\) has an equilibrium constant of \(1.0 \times 10^{-4}\). For which of the inital reaction mixtures is the \(x\) is small approximation most likely to apply? a. \([A]-0.0010 M ;[B]-0.00 M\) b. \([A]-0.00 \mathrm{M} ;[\mathrm{B}]-0.10 \mathrm{M}\) \begin{tabular}{l} c. \([A]-0.10 M ;[B]-0.10 M\) \\ \hline \end{tabular} d. \([\mathrm{A}]-0.10 \mathrm{M} ;[\mathrm{B}]-0.00 \mathrm{M}\)

Consider the simple one-step reaction: $$ \mathcal{A}(g) \rightleftharpoons B(b) $$ Since the reaction occurs in a single step, the forward reaction What happens to the rate of the forward reaction when we increase the concentration of \(\mathrm{A}\) ? How does this explain the reason bchind Le Chatelier's principle?

Consider the reaction: \(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)\). a. Write the equilibrium constant expression for this reaction. If some hydrogen is added, before the reaction shifts, b. How will the numerator and denominator of the expression In part a compare to the value at equilibrium? c. Wul Q be larger or smaller than \(K\). Why? d. Will the reaction have to shift forwand or backward to retain equilibrium? Explain. e. Are your answers above consistent with le Chateller's principle? Explain.

Have cach group member explain to the group what happens if a system at equilibrium is subject to one of the following changes and why: a. the concentration of a reactant is increased b. a solid product is added c. the volume is decreased d. the temperature is ralsed