Chapter 18

Molecules in the upper atmosphere tend to contain double and triple bonds rather than single bonds. Suggest an explanation. [Section 18.1]

You are working with an artist who has been commissioned to make a sculpture for a big city in the eastern United States. The artist is wondering what material to use to make her sculpture because she has heard that acid rain in the eastern United States might destroy it over time. You take samples of granite, marble, bronze, and other materials, and place them outdoors for a long time in the big city. You periodically examine the appearance and measure the mass of the samples. (a) What observations would lead you to conclude that one or more of the materials are well-suited for the sculpture? (b) What chemical process (or processes) is (are) the most likely responsible for any observed changes in the materials? [Section 18.2]

Describe the properties that most clearly distinguish among salt water, freshwater, and groundwater. [Section 18,3]

(a) What is the primary basis for the division of the atmosphere into different regions? (b) Name the regions of the atmosphere, indicating the altitude interval for each one.

(a) How are the boundaries between the regions of the atmosphere determined? (b) Explain why the stratosphere, which is about \(35 \mathrm{~km}\) thick, has a smaller total mass than the troposphere, which is about \(12 \mathrm{~km}\) thick.

Air pollution in the Mexico City metropolitan area is among the worst in the world. The concentration of ozone in Mexico City has been measured at \(441 \mathrm{ppb}(0.441 \mathrm{ppm})\). Mexico City sits at an altitude of 7400 feet, which means its atmospheric pressure is only \(0.67\) atm. (a) Calculate the partial pressure of ozone at \(441 \mathrm{Ppb}\) if the atmospheric pressure is \(0.67 \mathrm{~atm}\). (b) How many ozone molecules are in \(1.0 \mathrm{~L}\) of air in Mexico City? Assume \(T=25^{\circ} \mathrm{C}\).

The average concentration of carbon monoxide in air in an Ohio city in 2006 was \(3.5 \mathrm{ppm}\). Calculate the number of CO molecules in \(1.0 \mathrm{~L}\). of this air at a pressure of 759 torr and a temperature of \(22^{\circ} \mathrm{C}\)

The dissociation energy of a carbon-bromine bond is typically about \(210 \mathrm{~kJ} / \mathrm{mol}\). (a) What is the maximum wavelength of photons that can cause C-Br bond dissociation? (b) Which kind of electromagnetic radiation-ultraviolet, visible, or infrared-does the wavelength you calculated in part (a) correspond to?

In \(\mathrm{CF}_{3} \mathrm{Cl}\) the \(\mathrm{C}-\mathrm{Cl}\) bond- dissociation energy is \(339 \mathrm{~kJ} / \mathrm{mol}\), In \(\mathrm{CCl}_{4}\) the \(\mathrm{C}-\mathrm{Cl}\) bond-dissociation energy is \(293 \mathrm{~kJ} / \mathrm{mol}\). What is the range of wavelengths of photons that can cause \(\mathrm{C}\) - Cl bond rupture in one molecule but not in the other?

Do the reactions involved in ozone depletion involve changes in oxidation state of the \(\mathrm{O}\) atoms? Explain.

Which of the following reactions in the stratosphere cause an increase in temperature there? (a) \(\mathrm{O}(\mathrm{g})+\mathrm{O}_{2}(g) \longrightarrow \mathrm{O}_{3}{\underline{\phantom{xx}}}^{*}(\mathrm{~g})\) (b) \(\mathrm{O}_{3}^{*}(g)+\mathrm{M}(g) \longrightarrow \mathrm{O}_{3}(g)+\mathrm{M}^{*}(g)\) (c) \(\mathrm{O}_{2}(\mathrm{~g})+h \mathrm{w} \longrightarrow 2 \mathrm{O}(\mathrm{g})\) (d) \(\mathrm{O}(\mathrm{g})+\mathrm{N}_{2}(\mathrm{~g}) \longrightarrow \mathrm{NO}(\mathrm{g})+\mathrm{N}(\mathrm{g})\) (e) All of the above

(a) What is the difference between chlorofluorocarbons and hydrofluorocarbons? (b) Why are hydrofluorocarbons potentially less harmful to the ozone layer than CFCs?

Draw the Lewis structure for the chlorofluorocarbon CFC-11, \(\mathrm{CFCl}_{2}\). What chemical characteristics of this substance allow it to effectively deplete stratospheric ozone?

(a) Why is the fluorine present in chlorofluorocarbons not a major contributor to depletion of the ozone layer? (b) What are the chemical forms in which chlorine exists in the stratesphere following cleavage of the carbon-chlorine bond?

Would you expect the substance \(\mathrm{CFBr}_{3}\) to be effective in depleting the ozone layer, assuming that it is present in the stratosphere? Explain.

For each of the following gases, make a list of known or possible naturally occurring sources: (a) \(\mathrm{CH}_{4}\) (b) \(\mathrm{SO}_{2}\), (c) \(\mathrm{NO}\).

Why is rainwater naturally acidic, even in the absence of polluting gases such as \(\mathrm{SO}_{2}\) ?

(a) Write a chemical equation that describes the attack of acid rain on limestone, \(\mathrm{CaCO}_{3}\). (b) If a limestone sculpture were treated to form a surface layer of calcium sulfate, would this help to slow down the effects of acid rain? Explain.

The first stage in corrosion of iron upon exposure to air is oxidation to \(\mathrm{Fe}^{2+}\), (a) Write a balanced chemical equation to show the reaction of iron with oxygen and protons from acid rain. (b) Would you expect the same sort of reaction to occur with a silver surface? Fxplain.

Alcohol-based fuels for automobiles lead to the production of formaldehyde \(\left(\mathrm{CH}_{2} \mathrm{O}\right)\) in exhaust gases. Formaldehyde undergoes photodissociation, which contributes to photochemical smog: $$ \mathrm{CH}_{2} \mathrm{O}+h_{\mathrm{w}} \longrightarrow \mathrm{CHO}+\mathrm{H} $$ The maximum wavelength oflight that can cause this reaction is \(335 \mathrm{~nm}\). (a) In what part of the electromagnetic spectrum is light with this wavelength found? (b) What is the maximum strength of a bond, in \(\mathrm{kJ} / \mathrm{mol}\), that can be broken by absorption of a photon of 335 -nm light? (c) Compare your answer from part (b) to the appropriate value from Table 8.4. What do you conclude about \(\mathrm{C}-\mathrm{H}\) bond energy in formaldehyde? (d) Write out the formaldehyde photodissociation reaction. showing Lewis-dot structures.

An important reaction in the formation of photochemical smog is the photodissociation of \(\mathrm{NO}_{2}\) = $$ \mathrm{NO}_{2}+h w \longrightarrow \mathrm{NO}(g)+\mathrm{O}(g) $$ The maximum wavelength of light that can cause this reaction is \(420 \mathrm{~nm}\). (a) In what part of the electromagnetic spectrum is light with this wavelength found? (b) What is the maximum strength of a bond, in kJ/mol, that can be broken by absorption of a photon of 420 -nm light? (c) Write out the photodissociation reaction showing Lewis-dot structures.

Explain why an increasing concentration of \(\mathrm{CO}_{2}\) in the atmosphere affects the quantity of energy leaving Farth but does not affect the quantity of energy entering from the Sun.

(a) With respect to absorption of radiant energy, what distinguishes a greenhouse gas from a non-greenhouse gas? (b) \(\mathrm{CH}_{4}\) is a greenhouse gas, but \(\mathrm{N}_{2}\) is not. How might the molecular structure of \(\mathrm{CH}_{4}\) explain why it is a greenhouse gas?

Phosphorus is present in seawater to the extent of \(0.07 \mathrm{Ppm}\) by mass. Assuming that the phesphorus is present as dihydrogenphosphate, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\), calculate the corresponding molar concentration of phosphate in seawater.

The enthalpy of evaporation of water is \(40.67 \mathrm{~kJ} / \mathrm{mol}\). Sunlight striking Earthis surface supplies 168 W per square meter \((1 \mathrm{~W}=1 \mathrm{watt}=1 \mathrm{~J} / \mathrm{s})\). (a) Assuming that evaporation of water is due only to energy input from the Sun, calculate how many grams of water could be evaporated from a \(1.00\) square meter patch of ocean over a 12 -h day. (b) The specific heat capacity of liquid water is \(4.184 \mathrm{~J} / \mathrm{g}{\underline{\phantom{xx}}}^{\circ} \mathrm{C}\). If the initial surface temperature of a \(1.00\) square meter patch of ocean is \(26{ }^{\circ} \mathrm{C}\), what is its final temperature after being in sunlight for \(12 \mathrm{~h}\), assuming no phase changes and assuming that sunlight penetrates uniformly to depth of \(10.0 \mathrm{~cm}\) ?

Gold is found in seawater at very low levels, about \(0.05 \mathrm{Ppb}\) by mass. Assuming that gold is worth about \(\$ 1300\) per troy ounce, how many liters of seawater would you have to process to obtain \(\$ 1,000,000\) worth of gold? Assume the density of seawater is \(1.03 \mathrm{~g} / \mathrm{mL}\) and that your gold recovery process is \(50 \%\) efficient.

(a) What is groundwater? (b) What is an aquifer?

The organic anion CCCCCC(C)c1ccc(S(=O)(=O)[O-])cc1 is found in most detergents. Assume that the anion undergoes aerobic decomposition in the following manner: $2 \mathrm{C}_{18} \mathrm{H}_{2} \mathrm{SO}_{3}^{-}(a q)+51 \mathrm{O}_{2}(a q) \longrightarrow$ $36 \mathrm{CO}_{2}(a q)+28 \mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{H}^{+}(a q)+2 \mathrm{SO}_{4}^{2-}(a q)$ What is the total mass of \(\mathrm{O}_{2}\) required to biodegrade $10.0 \mathrm{~g}$ of this substance?

The average daily mass of \(\mathrm{O}_{2}\) taken up by sewage discharged in the United States is \(59 \mathrm{~g}\) per person. How many liters of water at \(9 \mathrm{ppm} \mathrm{O}_{2}\) are \(50 \%\) depleted of oxygen in 1 day by a population of \(1,200,000\) people?

Magnesium ions are remeved in water treatment by the addition of slaked lime, \(\mathrm{Ca}(\mathrm{OH})_{2}\). Write a balanced chemical equation to describe what occurs in this process.

(a) Which of the following ionic species could be responsible for hardness in a water supply: \(\mathrm{Ca}^{2+}, \mathrm{K}^{+}, \mathrm{Mg}^{2+}, \mathrm{Fe}^{2+}, \mathrm{Na}^{+}\)? (b) What properties of an ion determine whether it will contribute to water hardness?

In the lime soda process at one time used in large scale municipal water softening, calcium hydroxide prepared from lime and sodium carbonate are added to precipitate \(\mathrm{Ca}^{2+}\) as \(\mathrm{CaCO}_{3}(s)\) and \(\mathrm{Mg}^{2+}\) as \(\mathrm{Mg}(\mathrm{OH})_{2}(\mathrm{~s})\) : $$ \begin{gathered} \mathrm{Ca}^{2+}(a q)+\mathrm{CO}_{3}{\underline{\phantom{xx}}}^{2-}(a q) \longrightarrow \mathrm{CaCO}_{3}(s) \\ \mathrm{Mg}^{2+}(a q)+2 \mathrm{OH}^{-}(a q) \longrightarrow \mathrm{MgOH}_{2}(a q) \end{gathered} $$ How many moles of \(\mathrm{Ca}(\mathrm{OH})_{2}\) and \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) should be added to soften \(1200 \mathrm{~L}\) of water in which $$ \begin{aligned} {\left[\mathrm{Ca}^{2+}\right] } &=5.0 \times 10^{-4} \mathrm{M} \text { and } \\\ {\left[\mathrm{Mg}^{2+}\right] } &=7.0 \times 10^{-4} \mathrm{M} \end{aligned} $$

The concentration of \(\mathrm{Ca}^{2+}\) in a particular water supply is \(5.7 \times 10^{-3} \mathrm{M}\). The concentration of bicarbonate ion, \(\mathrm{HCO}_{3}^{-}\). in the same water is \(1.7 \times 10^{-3} \mathrm{M}\). What masses of \(\mathrm{Ca}(\mathrm{OH})_{2}\) and \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) must be added to \(5.0 \times 10^{7} \mathrm{~L}\) of this water to reduce the level of \(\mathrm{Ca}^{2+}\) to \(20 \%\) of its original level?

Ferrous sulfate \(\left(\mathrm{FeSO}_{4}\right)\) is often tued as a coagulant in water purification. The iron(II) salt is dissolved in the water to be purified, then oxidized to the iron(III) state by dissolved exygen, at which time gelatinous \(\mathrm{Fe}(\mathrm{OH})_{3}\) forms, assuming the \(\mathrm{pH}\) is abeve approximately 6 . Write balanced chemical equations for the oxidation of \(\mathrm{Fe}^{2+}\) to \(\mathrm{Fe}^{3+}\) by dissolved oxygen and for the formation of \(\mathrm{Fe}(\mathrm{OH})_{3}(s)\) by reaction of \(\mathrm{Fe}^{3+}(a q)\) with \(\mathrm{HCO}_{3}^{-}(a q)\) -

What properties make a substance a good coagulant for water purification?

(a) What are trihalomethanes (THMs)? (b) Draw the Lewis structures of two example THMs.

(a) Suppose that tests of a municipal water system reveal the presence of bromate ion, \(\mathrm{BrO}_{3}^{-}\). What are the likely origins of this ion? (b) Is bromate ion an oxidizing or reducing agent? Write a chemical equation for the reaction of bromate ion with hyponitrite ion. Green Chemistry (Section 18.5)

One of the principles of green chemistry is that it is better to use as few steps as possible in making new chemicals. In what ways does following this rule advance the goals of green chemistry? How does this principle relate to energy efficiency?

Discuss how catalysts can make processes more energy efficient.

In the following three instances which choice is greener in each situation? Explain. (a) Benzene as a solvent or water as a solvent. (b) The reaction temperature is \(500 \mathrm{~K}\), or \(1000 \mathrm{~K}\). (c) Sodium chloride as a by-product or chloroform \(\left(\mathrm{CHCl}_{3}\right)\) as a by-product.

In the following three instances which choice is greener in a chemical process? Explain. (a) A reaction that can be run at \(350 \mathrm{~K}\) for \(12 \mathrm{~h}\) without a catalyst or one that can be run at \(300 \mathrm{~K}\) for \(1 \mathrm{~h}\) with a reusable catalyst. (b) A reagent for the reaction that can be obtained from corn husks or one that is obtained from petroleum. (c) A process that produces no by-products or one in which the by-products are recycled for another process.

A friend of yours has seen each of the following items in newspaper articles and would like an explanation: (a) acid rain, (b) greenhouse gas, (c) photochemical smog, (d) ozone depletion. Give a brief explanation of each term and identify one or two of the chemicals associated with each. s in this respect?

If an average \(O_{3}\) molecule "lives" only 100-200 seconds in the stratosphere before undergoing dissociation, how can \(\mathrm{O}_{3}\) offer any protection from ultraviolet radiation?

What properties of CFCs make them ideal for various commercial applications but also make them a long-term problem in the stratosphere?

(a) What is the difference between a CFC and an HFC? (b) It is estimated that the lifetime for HFCs in the stratosphere is 2-7 years. Why is this number significant? (c)Why have HFCs been used to replace CFCs? (d) What is the major disadvantage of HFCs as replacements for CFCs?

Explain, using Le Châtelier's principle, why the equilibrium constant for the formation of \(\mathrm{NO}\) from \(\mathrm{N}_{2}\) and \(\mathrm{O}_{2}\) increases with increasing temperature, whereas the equilibrium constant for the formation of \(\mathrm{NO}_{2}\) from \(\mathrm{NO}\) and \(\mathrm{O}_{2}\) decreases with increasing temperature.

Natural gas consists primarily of methane, \(\mathrm{CH}_{4}(\mathrm{~g})\). (a) Write a balanced chemical equation for the complete combustion of methane to produce \(\mathrm{CO}_{2}(g)\) as the only carbon-containing product. (b) Write a balanced chemical equation for the incomplete combustion of methane to produce \(\mathrm{CO}(\mathrm{g})\) as the only carbon-containing product. (c) At \(25^{\circ} \mathrm{C}\) and \(1.0 \mathrm{~atm}\) pressure, what is the minimum quantity of dry air needed to combust \(1.0 \mathrm{~L}\) of \(\mathrm{CH}_{4}(\mathrm{~g})\) completely to \(\mathrm{CO}_{2}(\mathrm{~g})\) ?

It was estimated that the eruption of the Mount Pinatubo volcano resulted in the injection of 20 million metric tons of \(\mathrm{SO}_{2}\) into the atmosphere. Most of this \(\mathrm{SO}_{2}\) underwent oxidation to \(\mathrm{SO}_{2}\), which reacts with atmospheric water to form an aerosol. (a) Write chemical equations for the processes leading to formation of the aerosol. (b) The aerosols caused a \(0.5-0.6^{\circ} \mathrm{C}\) drop in surface temperature in the northern hemisphere. What is the mechanism by which this occurs? (c) The sulfate aerosols, as they are called, also cause loss of ozone from the stratosphere. How might this occur?

One of the possible consequences of climate change is an increase in the temperature of ocean water. The oceans serve as a "sink" for \(\mathrm{CO}_{2}\) by dissolving large amounts of it. (a) The figure below shows the solubility of \(\mathrm{CO}_{2}\) in water as a function of temperature. Does \(\mathrm{CO}_{2}\) behave more or less similarly to other gase

The rate of solar energy striking Earth averages 168 watts per square meter. The rate of energy radiated from Earth's surface averages 390 watts per square meter. Comparing these numbers, one might expect that the planet would cool quickly, yet it does not. Why not?