Holt Chemistry

Explain the meaning of the terms reversible
reaction, completion reaction, and reaction at
equilibrium.

What is the distinction between a static
equilibrium and a dynamic equilibrium? In
which class do chemical equilibria fit?

Define complex ion and give two examples.

What are ligands?

What is an equilibrium constant?

Distinguish between solubility and solubility
product constant. Explain how one may be
calculated from the other.

In the context of Le Chatelier's principle,
what is a "stress"? List the stresses that
affect chemical equilibria.

Give one example of a stress on a reaction
in aqueous solution and at equilibrium.

Use Figure 3 to answer the following
questions:
a. What is happening to the rate of
formation of HI(g) before the system
reaches equilibrium?
b. When is the rate of the forward reaction
the greatest?

Give two examples of static equilibrium
and two examples of dynamic equilibrium
Your examples do not have to be chemical
examples.

Identify the ligands in the following
complex ions.

Why must a balanced chemical equation be
used when determining $K_{e q} ?$

Describe and explain how the concentrations
of $A, B, C,$ and $D$ change from the time when
A and $B$ first combine to the point at which
equilibrium is established for the reaction
$$A+B \rightleftarrows C+D$$

In general, which reaction (forward, reverse,
or neither) is favored if the value of $K$ at a
specified temperature is
a. equal to 1$?$
b. very small?
c. very small?

When nitrogen monoxide reacts with
oxygen to produce nitrogen dioxide, an
equilibrium is established.
a. Write the balanced equation.
b. Write the equilibrium constant
expression.

Write equilibrium constant expressions for
the following reactions:
$$\begin{array}{l}{\text { a. } 2 \mathrm{NO}_{2}(g) \rightleftarrows \mathrm{N}_{2} \mathrm{O}_{4}(g)} \\ {\text { b. } \mathrm{CO}(g)+\mathrm{Cl}_{2}(g) \rightleftarrows \operatorname{COCl}_{2}(g)} \\ {\text { c. } \mathrm{AgCl}(s) \rightleftarrows \mathrm{Ag}^{+}(a q)+\mathrm{Cl}^{-}(a q)}\end{array}$$
$$\begin{array}{c}{\mathrm{d} \cdot \mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftarrows} \\ {\mathrm \quad \quad \quad \quad \quad \quad \quad \quad\quad \quad \quad \quad \quad {H}_{3} \mathrm{O}^{+}(a q)+\mathrm{CH}_{3} \mathrm{COO}^{-}(a q)}\end{array}$$

Write the solubility product expressions for
the following slightly soluble salts: AgI,
$\mathrm{SrSO}_{4}, \mathrm{Ag}_{2} \mathrm{CO}_{3}, \mathrm{Ag}_{2} \mathrm{S}, \mathrm{PbI}_{2}, \mathrm{AgIO}_{3}$
$\mathrm{Mg}_{3}\left(\mathrm{PO}_{4}\right)_{2},$ and $\mathrm{Hg}_{2} \mathrm{Cl}_{2}$

How would you explain Le Chatelier's
principle in your own words to someone
who finds the concept difficult to
understand?

Predict the effect of each of the following
on the indicated equilibrium system in
terms of which reaction (forward, reverse,
or neither) will be favored.
$$\mathrm{H}_{2}(g)+\mathrm{Cl}_{2}(g) \rightleftarrows 2 \mathrm{HCl}(g)+184 \mathrm{kJ}$$
a. addition of $\mathrm{Cl}_{2}$
b. removal of $\mathrm{HCl}$
c. increased pressure
d. decreased temperature
e. removal of $\mathrm{H}_{2}$
g. decreased pressure
g. increased temperature
h. decreased system volume

What relative conditions (reactant concentrations, pressure, and temperature would
favor a high equilibrium concentration of
the substance in bold in each of the following equilibrium systems?
$$\begin{array}{c}{\text { a. } 2 \mathrm{CO}(g)+\mathrm{O}_{2}(g) \rightleftarrows 2 \mathrm{CO}_{2}(g)+167 \mathrm{kJ}} \\ {\text { b. } \mathrm{Cu}^{2+}(a q)+4 \mathrm{NH}_{3}(a q) \rightleftarrows} \\ {\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}(a q)+42 \mathrm{kJ}}\end{array}$$
$$\begin{array}{r}{\text { c. } 2 \mathrm{HI}(g)+12.6 \mathrm{kJ} \rightleftarrows \mathrm{H}_{2}(g)+\mathrm{I}_{2}(g)} \\ {\text { d. } 4 \mathrm{HCl}(g)+\mathrm{O}_{2}(g) \rightleftarrows} \\ {2 \mathrm{H}_{2} \mathrm{O}(g)+2 \mathrm{Cl}_{2}(g)+113 \mathrm{kJ}}\end{array}$$

What changes in conditions would favor the
reactants in the following equilibrium?
$$2 \mathrm{SO}_{2}(g)+\mathrm{O}_{2}(g) \rightleftarrows 2 \mathrm{SO}_{3}(g) \quad \Delta H=-198 \mathrm{kJ}$$

Write the equilibrium constant expression
for the reaction in problem $21 .$

What changes in conditions would favor the
products in the following equilibrium?
$$\mathrm{PCl}_{5}(g) \rightleftarrows \mathrm{Cl}_{2}(g)+\mathrm{PCl}_{3}(g) \quad \Delta H=88 \mathrm{kJ}$$

Write the equilibrium constant expression
for the reaction in problem $23 .$

Relate Le Chatelier's principle to the
common-ion effect.

Vinegar-a solution of acetic acid,
$\mathrm{CH}_{3} \mathrm{COOH},$ in water $-$ is used in varying
concentrations for different household tasks.
The following equilibrum exists in vinegar.
$$\begin{array}{c}{\mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftarrows} \\ {\mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{CH}_{3} \mathrm{COO}^{-}(a q)}\end{array}$$
If the concentration of the acetic acid
solution at equilibrium is 3.00 $\mathrm{M}$ and the
$\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right]=7.22 \times 10^{-3},$ what
is the $K_{e q}$ value for acetic acid?

Write the $K_{e q}$ expressions for the following
reactions:
$$\begin{array}{c}{\mathbf{a} \cdot 4 \mathrm{H}_{3} \mathrm{O}^{+}(a q)+2 \mathrm{Cl}^{-}(a q)+\mathrm{MnO}_{2}(s) \rightleftarrows} \\ {\mathrm{Mn}^{2+}(a q)+6 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{Cl}_{2}(g)} \\ {\text { b. } \mathrm{As}_{4} \mathrm{O}_{6}(s)+6 \mathrm{H}_{2} \mathrm{O}(l) \rightleftarrows 4 \mathrm{H}_{3} \mathrm{AsO}_{3}(a q)}\end{array}$$

The following equation shows an equilibrium reaction between hydrogen and carbon
dioxide.
$$\mathrm{H}_{2}(g)+\mathrm{CO}_{2}(g) \rightleftarrows \mathrm{H}_{2} \mathrm{O}(g)+\mathrm{CO}(g)$$
At $986^{\circ} \mathrm{C},$ the following data were obtained
at equilibrium in a 2.0 L reaction vessel.
Calculate the $K_{e q}$ for this reaction.

Determine the value of the equilibrium
constant for each reaction below assuming
that the equilibrium concentrations are as
specified.
$$\text {(a.)}A+B \rightleftarrows C ;[A]=2.0 ;[B]=3.0 ;[C]=4.0$$
$$\begin{array}{l}{\text { b. } \mathrm{D}+2 \mathrm{E} \rightleftarrows \mathrm{F}+3 \mathrm{G} ;[\mathrm{D}]=1.5 ;[\mathrm{E}]=2.0} \\ {[\mathrm{F}]=1.8 ;[\mathrm{G}]=1.2}\end{array}$$
$$\mathrm{c} \cdot \mathrm{N}_{2}(\mathrm{g})+3 \mathrm{H}_{2}(\mathrm{g}) \rightleftarrows 2 \mathrm{NH}_{3}(\mathrm{g}) ;\left[\mathrm{N}_{2}\right]=0.45$
$\left[\mathrm{H}_{2}\right]=0.14 ;\left[\mathrm{NH}_{3}\right]=0.62$$

An equilibrium mixture at a specific
temperature is found to consist of $1.2 \times 10^{-3}$
$\operatorname{mol} / \mathrm{L} \mathrm{HCl}, 3.8 \times 10^{-4} \mathrm{moll} \mathrm{O}_{2}, 5.8 \times 10^{-2}$
$\mathrm{mol} / \mathrm{L} \mathrm{H}_{2} \mathrm{O},$ and $5.8 \times 10^{-2} \mathrm{mol} / \mathrm{L} \mathrm{Cl}_{2}$
according to the following:
$$4 \mathrm{HCl}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g}) \rightleftarrows 2 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})+2 \mathrm{Cl}_{2}(\mathrm{g})$$
Determine the value of the equilibrium
constant for this system.

Analysis of an equilibrium mixture in
which the following equilibrium exists gave
$\left[\mathrm{OH}^{-}\right]=\left[\mathrm{HCO}_{3}^{-}\right]=3.2 \times 10^{-3} .$
$$\mathrm{HCO}_{3}^{-}(a q)+\mathrm{OH}^{-}(a q) \rightleftarrows \mathrm{CO}_{3}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$$
The equilibrium constant is $4.7 \times 10^{3} .$ What
is the concentration of the carbonate ion?

When a sample of $\mathrm{NO}_{2}$ gas equilibrated in a
closed container at $25^{\circ} \mathrm{C},$ a concentration of
0.0187 $\mathrm{mol} / \mathrm{L}$ of $\mathrm{N}_{2} \mathrm{O}_{4}$ was found to be
present. Use the data in Table 2 to calculate
the $\mathrm{NO}_{2}$ concentration.

Methanol, $\mathrm{CH}_{3} \mathrm{OH},$ can be prepared in the
presence of a catalyst by the reaction of $\mathrm{H}_{2}$
and $\mathrm{CO}$ at high temperatures according to
the following equation:
$$\mathrm{CO}(g)+2 \mathrm{H}_{2}(g) \rightleftarrows \mathrm{CH}_{3} \mathrm{OH}(g)$$
What is the concentration of $\mathrm{CH}_{3} \mathrm{OH}(\mathrm{g})$
in moles per liter if the concentration of
$\mathrm{H}_{2}=0.080 \mathrm{mol} / \mathrm{L}$ , the concentration of
$\mathrm{CO}=0.025 \mathrm{mol} / \mathrm{L},$ and $K_{e q}=290$ at 700 $\mathrm{K} ?$

At $450^{\circ} \mathrm{C}$ the value of the equilibrium constant for the following system is $6.59 \times 10^{-3}$ .
If $\left[\mathrm{NH}_{3}\right]=1.23 \times 10^{-4}$ and $\left[\mathrm{H}_{2}\right]=2.75 \times 10^{-3}$
at equilibrium, determine the concentration
of $\mathrm{N}_{2}$ at that point.
$$\mathrm{N}_{2}(\mathrm{g})+3 \mathrm{H}_{2}(\mathrm{g}) \rightleftarrows 2 \mathrm{NH}_{3}(\mathrm{g})$$

The solubility of cobalt $(\mathrm{II})$ sulfide, coS, is
$1.7 \times 10^{-13} \mathrm{M}$ . Calculate the solubility
product constant for $\mathrm{CoS} .$

What is the solubility product for copper $(\mathrm{I})$
sulfide, $\mathrm{Cu}_{2} \mathrm{S},$ given that the solubility of
$\mathrm{Cu}_{2} \mathrm{S}$ is $8.5 \times 10^{-17} \mathrm{M} ?$

The ionic substance $\mathrm{T}_{3} \mathrm{U}_{2}$ ionizes to form
$\mathrm{T}^{2+}$ and $\mathrm{U}^{3-}$ ions. The solubility of $\mathrm{T}_{3} \mathrm{U}_{2}$ is
$3.77 \times 10^{-20} \mathrm{mol} / \mathrm{L}$ . What is the value of the
solubility-product constant?

The ionic substance EJ dissociates to form
$\mathrm{E}^{2+}$ and $\mathrm{J}^{2-}$ ions. The solubility of EJ is
$8.45 \times 10^{-6} \mathrm{mol} / \mathrm{L}$ . What is the value of the
solubility-product constant?

Four lead salts and their solubility products
are as follows: $\mathrm{PbS}, 3.0 \times 10^{-28} ; \mathrm{PbCrO}_{4}, 1.8$
$\times 10^{-14} ; \mathrm{PbCO}_{3}, 7.4 \times 10^{-14} ;$ and $\mathrm{PbSO}_{4}, 2.5$
$\times 10^{-8} .$ Calculate the $\mathrm{Pb}^{2+}$ concentration in
a saturated solution of each of these salts.

Review problem $39,$ and note that all of the
lead salts listed have anions whose charges
are $2-.$ Why is the calculation more difficult
for salts such as $\mathrm{PbCl}_{2}$ or $\mathrm{PbI}_{2} ?$

What is the concentration of $\mathrm{F}^{-}$ ions in a
saturated solution that is 0.10 $\mathrm{M}$ in $\mathrm{Ca}^{2+}$ ?
The $K_{s p}$ of $\mathrm{CaF}_{2}$ is $1.6 \times 10^{-10} .$

Silver bromide, AgBr, is used to make
photographic black-and-white film. Calculate
the concentration of $\mathrm{Ag}^{+}$ and $\mathrm{Br}^{-}$ ions in a
saturated solution at $25^{\circ} \mathrm{C}$ using table 3 .

The figure below shows the results of
adding three different chemicals to
distilled water and stirring well.
a. Which substance(s) are completely
soluble in water?
b. Is it correct to say that AgCl is
completely insoluble? Explain your
answer. Is Ba(OH)_ completely insoluble?

Why might an industrial process be operated at high temperature even though the reaction is more favorable at lower temperatures?

How does temperature affect equilibrium
constants?

The reaction below has an equilibrium
constant of $4.9 \times 10^{11}$ .
$$\mathrm{Fe}(\mathrm{OH})_{2}(s)+2 \mathrm{H}_{3} \mathrm{O}^{+}(a q) \rightleftarrows \mathrm{Fe}^{2+}(a q)+4 \mathrm{H}_{2} \mathrm{O}(l)$$
Write the equilibrium constant expression,
and determine the concentration of $\mathrm{Fe}^{2+}$
ions in equilibrium when the hydronium ion
concentration is $1.0 \times 10^{-7} \mathrm{mol} / \mathrm{L}$

When does a pressure change affect a
chemical equilibrium?

What information may be conveyed by the
knowledge that the equilibrium constant of
a reaction is very small?

The $K_{s p}$ of potassium periodate, KIO $_{4, \text { is } 3.7}$
$\times 10^{-4} .$ Determine whether a precipitate will
form when 2.00 g of $\mathrm{KCl}$ and 2.00 $\mathrm{g}$ of
$\mathrm{NaIO}_{4}$ are dissolved in 1.00 $\mathrm{L}$ of water.

Calculate the solubility of a substance MN
that ionizes to form $\mathrm{M}^{2+}$ and $\mathrm{N}^{2-}$ ions given
that $K_{s p}=8.1 \times 10^{-6} .$

A student wrote, "The larger the equilibrium
constant, the greater the rate at which reactants convert to products." How was he
wrong?

Imagine the following hypothetical reaction,
taking place in a sealed, rigid container, to
be neither exothermic nor endothermic.
$$\mathrm{A}(g)+\mathrm{B}(g) \rightleftarrows \mathrm{C}(g) \quad \Delta H=0$$
Would an increase in temperature favor the
forward reaction or the reverse reaction?
(Hint: Recall the gas laws.)

A very dilute solution of silver nitrate is
added dropwise to a solution that contains
equal concentrations of sodium chloride
and potassium bromide. What salt will
precipitate first?

Changes in the concentrations of the reactants and products at equilibrium have no
effect on the value of the equilibrium
constant. Explain this statement.

Research the practical uses of Le Chatelier's
principle. Present your results to your class.

Develop a model that shows the concept of
equilibrium. Be sure that your model
includes the impact of Le Chatelier's
principle on equilibrium.

Use the following terms to create a concept
map: chemical equilibrium, equilibrium constant, solubility product constant, reversible
reactions, and Le Chatelier's principle.

What is the change in the rate of the
forward reaction after the reaction is at
equilibrium?

Does the reverse reaction rate ever equal
zero? Why or why not?

Does the forward reaction rate ever equal
zero? Why or why not?

Calculating the Equilibrium Constant, $K$
for a System
The graphing calculator can run a program
that calculates $K$ for a system, given the
concentrations of the products and the
concentrations of the reactants.
Given the balanced chemical equation
$$\mathrm{H}_{2}(g)+\mathrm{I}_{2}(g) \rightarrow 2 \mathrm{HI}(g)$$
and the equilibrium mixture at $425^{\circ} \mathrm{C},$ you
can calculate the equilibrium constant for
the system. Then you can use the program
to make calculations.
Go to Appendix $\mathbf{C}$ If you are using a TI- $-83$
Plus, you can download the program
CONSTANT and data and run the application
as directed. If you are using another calculator, your teacher will provide you with key-
strokes and data sets to use. After you have
graphed the data, answer the questions
below.
a. What is the equilibrium constant given the
following equilibrium concentrations:
0.012840 $\mathrm{mol} / \mathrm{L}$ of $\mathrm{H}_{2}, 0.006437 \mathrm{mol} / \mathrm{L}$ of
$\mathrm{I}_{2},$ and 0.066807 $\mathrm{mol} / \mathrm{L}$ of $\mathrm{HI}$ ?
b. What is the equilibrium constant given the
following equilibrium concentrations:
0.000105 $\mathrm{mol} / \mathrm{L}$ of $\mathrm{H}_{2}, 0.000107 \mathrm{mol} / \mathrm{L}$ of
$\mathrm{I}_{2},$ and 0.000779 $\mathrm{mol} / \mathrm{L}$ of $\mathrm{HI}$ ?
c. What is the equilibrium constant given the
following equilibrium concentrations:
0.000527 $\mathrm{mol} / \mathrm{L}$ of $\mathrm{H}_{2}, 0.000496 \mathrm{mol} / \mathrm{L}$ of
$\mathrm{I}_{2},$ and 0.003757 $\mathrm{mol} / \mathrm{L}$ of HI?