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 wellsuited 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\(]\)

One mystery in environmental science is the imbalance in the "carbon dioxide budget." Considering only human activities, scientists have estimated that 1.6 billion metric tons of \(\mathrm{CO}_{2}\) is added to the atmosphere every year because of deforestation (plants use \(\mathrm{CO}_{2},\) and fewer plants will leave more \(\mathrm{CO}_{2}\) in the atmosphere). Another 5.5 billion tons per year is put into the atmosphere because of burning fossil fuels. It is further estimated (again, considering only human activities) that the atmosphere actually takes up about 3.3 billion tons of this \(\mathrm{CO}_{2}\) per year, while the oceans take up 2 billion tons per year, leaving about 1.8 billion tons of \(\mathrm{CO}_{2}\) per year unaccounted for. Describe a mechanism by which \(\mathrm{CO}_{2}\) is removed from the atmosphere and ultimately ends up below the surface (Hint: What is the source of the fossil fuels?). [Sections \(18.1-18.3]\)

(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.

The Environmental Protection Agency (EPA) has established air quality standards. For ozone \(\left(\mathrm{O}_{3}\right),\) the 8 -hour average concentration permitted under the standards is 0.085 parts per million (ppm). (a) Calculate the partial pressure of ozone at 0.085 ppm if the atmospheric pressure is \(100 \mathrm{kPa}\). (b) How many ozone molecules are in \(1.0 \mathrm{~L}\) of air? Assume \(T=25^{\circ} \mathrm{C}\).

The average concentration of carbon monoxide in air in a city in 2007 was 3.0 ppm. Calculate the number of CO molecules in \(1.0 \mathrm{~L}\) of this air at a pressure of \(100 \mathrm{kPa}\) and a temperature of \(25^{\circ} \mathrm{C}\).

The dissociation energy of a carbon-iodine bond is typically about \(240 \mathrm{~kJ} / \mathrm{mol} .(\mathbf{a})\) What is the maximum wavelength of photons that can cause \(\mathrm{C}-\mathrm{I}\) bond dissociation? (b) Which kind of electromagnetic radiation-ultraviolet, visible, or infrared- does the wavelength you calculated in part (a) correspond to?

In \(\mathrm{CH}_{3} \mathrm{I}\) the \(\mathrm{C}\) - I bond-dissociation energy is \(241 \mathrm{~kJ} / \mathrm{mol}\). In \(\mathrm{C}_{6} \mathrm{H}_{5}\) I the \(\mathrm{C}-\) I bond-dissociation energy is \(280 \mathrm{~kJ} / \mathrm{mol}\). What is the range of wavelengths of photons that can cause \(\mathrm{C}-\mathrm{I}\) bond rupture in one molecule but not in the other?

(a) Distinguish between photodissociation and photoionization. (b) Use the energy requirements of these two processes to explain why photodissociation of oxygen is more important than photoionization of oxygen at altitudes below about \(90 \mathrm{~km}\).

Why is the photodissociation of \(\mathrm{N}_{2}\) in the atmosphere relatively unimportant compared with the photodissociation of \(\mathrm{O}_{2} ?\)

The dissociation energy of \(\mathrm{N}_{2}\) is very high, \(941 \mathrm{~kJ} / \mathrm{mol}\). (a) Calculate the wavelength of the photons that possess sufficient energy to dissociate \(\mathrm{N}_{2} .(\mathbf{b})\) In which region of the electromagnetic spectrum does this light fall? Does this light have enough energy to photoionize \(\mathrm{N}_{2}\) ?

The ultraviolet spectrum can be divided into three regions based on wavelength: UV-A (315-400 nm), UV-B (280-315 \(\mathrm{nm})\), and UV-C \((100-280 \mathrm{nm})\). (a) Photons from which region have the highest energy and therefore are the most harmful to living tissue? (b) In the absence of ozone, which of these three regions, if any, are absorbed by the atmosphere? (c) When appropriate concentrations of ozone are present in the stratosphere, is all of the UV light absorbed before reaching the Earth's surface? If not, which region or regions are not filtered out?

Do the reactions involved in ozone depletion involve changes in oxidation state of the O atoms? Explain.

Which of the following reactions in the stratosphere cause an increase in temperature there? (a) \(\mathrm{O}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{O}_{3}^{*}(g)\) (b) \(\mathrm{O}_{3}^{*}(g)+\mathrm{M}(g) \longrightarrow \mathrm{O}_{3}(g)+\mathrm{M}^{*}(g)\) (c) \(\mathrm{O}_{2}(g)+h \nu \longrightarrow 2 \mathrm{O}(g)\) (d) \(\mathrm{O}(g)+\mathrm{N}_{2}(g) \longrightarrow \mathrm{NO}(g)+\mathrm{N}(g)\) (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, \(\mathrm{CFCl}_{3}\). What chemical characteristics of this substance allow it to effectively deplete stratospheric ozone?

The average bond enthalpies of the \(\mathrm{C}-\mathrm{C}\) and \(\mathrm{C}-\mathrm{H}\) bonds are \(348 \mathrm{~kJ} / \mathrm{mol}\) and \(413 \mathrm{~kJ} / \mathrm{mol}\), respectively. (a) What is the maximum wavelength that a photon can possess and still have sufficient energy to break the \(\mathrm{C}-\mathrm{H}\) and \(\mathrm{C}-\mathrm{C}\) bonds, respectively? (b) Given the fact that \(\mathrm{O}_{2}, \mathrm{~N}_{2},\) and \(\mathrm{O}\) in the upper atmosphere absorb most of the light with wavelengths shorter than \(240 \mathrm{nm}\), would you expect the photodissociation of \(\mathrm{C}-\mathrm{C}\) and \(\mathrm{C}-\mathrm{H}\) bonds to be significant in the lower atmosphere?

(a) When chlorine atoms react with atmospheric ozone, what are the products of the reaction? (b) Based on average bond enthalpies, would you expect a photon capable of dissociating a \(\mathrm{C}-\mathrm{Cl}\) bond to have sufficient energy to dissociate a \(\mathrm{C}-\mathrm{Br}\) bond? \((\mathbf{c})\) Would you expect the substance \(\mathrm{CFBr}_{3}\) to accelerate depletion of the ozone layer?

Nitrogen oxides like \(\mathrm{NO}_{2}\) and \(\mathrm{NO}\) are a significant source of acid rain. For each of these molecules write an equation that shows how an acid is formed from the reaction with water.

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

(a) It has been reported, that acid rain with a pH of 3.5 could corrode mild steel. Write a chemical equation that describes the attack of acid rain on an iron (Fe) material. (b) If the iron material were covered with a surface layer of copper, would this help to stop the effects of acid rain? Explain.

Copper exposed to air and water may be oxidized. The green oxidized product is referred to as "patina". (a) Write a balanced chemical equation to show the reaction of copper to copper (II) ions with oxygen and protons from acid rain. (b) Would you expect some kind of "patina" on a silver surface? Explain.

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: $$ \mathrm{CH}_{2} \mathrm{O}+h \nu \longrightarrow \mathrm{CHO}+\mathrm{H} $$ The maximum wavelength of light that can cause this reaction is \(335 \mathrm{nm} .(\mathbf{a})\) 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 8.3 . 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 \(\mathrm{NO}_{2}\) : $$ \mathrm{NO}_{2}+h \nu \longrightarrow \mathrm{NO}(g)+\mathrm{O}(g) $$ 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 \(\mathrm{kJ} / \mathrm{mol}\), that can be broken by absorption of a photon of 420 -nm light? \(?\) (c) Write out the photodissociation reaction showing Lewis-dot structures.

The atmosphere of Mars is \(96 \% \mathrm{CO}_{2}\), with a pressure of approximately \(0.6 \mathrm{kPa}\) at the surface. Based on measurements taken over a period of several years by the Rover Environmental Monitoring Station (REMS), the average daytime temperature at the REMS location on Mars is \(-5.7^{\circ} \mathrm{C},\) while the average nighttime temperature is \(-79^{\circ} \mathrm{C}\). This daily variation in temperature is much larger than what we experience on Earth. What factor plays the largest role in this wide temperature variation, the composition or the density of the atmosphere?

The enthalpy of evaporation of water is \(40.67 \mathrm{~kJ} / \mathrm{mol}\). Sunlight striking Earth's surface supplies \(168 \mathrm{~W}\) per square meter \((1 \mathrm{~W}=1 \mathrm{watt}=1 \mathrm{~J} / \mathrm{s}) .(\mathbf{a})\) 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 \(4.184 \mathrm{~J} / \mathrm{g}^{\circ} \mathrm{C}\). 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 \(10.0 \mathrm{~cm} ?\)

The enthalpy of fusion of water is \(6.01 \mathrm{~kJ} / \mathrm{mol}\). Sunlight striking Earth's surface supplies \(168 \mathrm{~W}\) per square meter \((1 \mathrm{~W}=1 \mathrm{watt}=1 \mathrm{~J} / \mathrm{s}) .(\) a) Assuming that melting of ice is due only to energy input from the Sun, calculate how many grams of ice could be melted from a 1.00 square meter patch of ice over a \(12-\mathrm{h}\) day. \((\mathbf{b})\) The specific heat capacity of ice is \(2.032 \mathrm{~J} / \mathrm{g}^{\circ} \mathrm{C}\). If the initial temperature of a 1.00 square meter patch of ice is \(-5.0^{\circ} \mathrm{C},\) what is its final temperature after being in sunlight for \(12 \mathrm{~h}\), assuming no phase changes and assuming that sunlight penetrates uniformly to a depth of \(1.00 \mathrm{~cm} ?\)

A first-stage recovery of magnesium from seawater is precipitation of \(\mathrm{Mg}(\mathrm{OH})_{2}\) with \(\mathrm{CaO};\) $$ \mathrm{Mg}^{2+}(a q)+\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Mg}(\mathrm{OH})_{2}(s)+\mathrm{Ca}^{2+}(a q) $$ What mass of \(\mathrm{CaO}\), in grams, is needed to precipitate \(1000 \mathrm{~kg}\) of \(\mathrm{Mg}(\mathrm{OH})_{2} ?\)

Platinum is found in seawater at very low levels, about 0.23 ppt (parts per trillion) by mass. How much platinum can be found in the entire ocean \(\left(1.3 \times 10^{21} \mathrm{~L}\right)\) ? Assume the density of seawater is \(1.03 \mathrm{~g} / \mathrm{mL}\). Estimate the price of the following amount of platinum: \(\$ 1,600\) per troy ounce.

Although there are many ions in seawater, the overall charges of the dissolved cations and anions must maintain charge neutrality. Consider only the six most abundant ions in seawater, as listed in Table \(18.5\left(\mathrm{Cl}^{-}, \mathrm{Na}^{+},\right.\) \(\mathrm{SO}_{4}^{2-}, \mathrm{Mg}^{2+}, \mathrm{Ca}^{2+},\) and \(\left.\mathrm{K}^{+}\right),\) calculate the total charge in Coulombs of the cations in \(1.0 \mathrm{~L}\) of seawater. Calculate the total charge in Coulombs of the anions in \(1.0 \mathrm{~L}\) of seawater. To how many significant figures are the two numbers equal?

List the common products formed when an organic material containing the elements carbon, hydrogen, oxygen, sulfur, and nitrogen decomposes (a) under aerobic conditions, (b) under anaerobic conditions.

Sodium stearate \(\left(\mathrm{C}_{18} \mathrm{H}_{35} \mathrm{O}_{2} \mathrm{Na}\right)\) is the most common soap. Assume that the stearate anion undergoes aerobic decomposition in the following manner: $$ \begin{aligned} \mathrm{C}_{18} \mathrm{H}_{35} \mathrm{O}_{2}^{-}(a q)+26 \mathrm{O}_{2}(a q) & \longrightarrow \\ & 17 \mathrm{CO}_{2}(a q)+17 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{HCO}_{3}^{-}(a q) \end{aligned} $$ What is the total mass of \(\mathrm{O}_{2}\) required to biodegrade \(3.0 \mathrm{~g}\) of this substance?

Sewage causes removal of oxygen from the fresh water into which the sewage is discharged. For a town with a population of 100,000 people, this effluent causes a daily oxygen depletion of 50.0 g per person. How many liters of water at \(8 \mathrm{ppm} \mathrm{O}_{2}\) are \(50 \%\) depleted of oxygen in a day by the population of this town?

Hydrogen phosphate \(\left(\mathrm{HPO}_{4}^{2-}\right)\) can be removed in water treatment by the addition of slaked lime, \(\mathrm{Ca}(\mathrm{OH})_{2} .\) Write a balanced chemical equation for the reaction (using ions as reactant), in which \(\mathrm{Ca}_{5}(\mathrm{OH})\left(\mathrm{PO}_{4}\right)_{3}\) forms as a precipitate.

In the lime soda process once 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 \(\mathrm{Mg}(\mathrm{OH})_{2}(s);\) $$ \begin{aligned} \mathrm{Ca}^{2+}(a q)+\mathrm{CO}_{3}^{2-}(a q) & \longrightarrow \mathrm{CaCO}_{3}(s) \\ \mathrm{Mg}^{2+}(a q)+2 \mathrm{OH}^{-}(a q) & \longrightarrow \mathrm{Mg}(\mathrm{OH})_{2}(s) \end{aligned} $$ How many moles of \(\mathrm{Ca}(\mathrm{OH})_{2}\) and \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) should be added to soften (remove the \(\mathrm{Ca}^{2+}\) and \(\mathrm{Mg}^{2+}\) ) 1000 L of water in which $$ \begin{array}{l} {\left[\mathrm{Ca}^{2+}\right]=3.5 \times 10^{-4} \mathrm{M}} \\ {\left[\mathrm{Mg}^{2+}\right]=7.5 \times 10^{-4} \mathrm{M}} \end{array} $$

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

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.

A reaction for converting ketones to lactones, called the Baeyer-Villiger reaction, is used in the manufacture of plastics and pharmaceuticals. 3-Chloroperbenzoic acid is shock-sensitive, however, and prone to explode. Also, 3 -chlorobenzoic acid is a waste product. An alternative process being developed uses hydrogen peroxide and a catalyst consisting of tin deposited within a solid support. The catalyst is readily recovered from the reaction mixture. (a) What would you expect to be the other product of oxidation of the ketone to lactone by hydrogen peroxide? (b) What principles of green chemistry are addressed by use of the proposed process?

The hydrogenation reaction shown here was performed with an iridium catalyst, both in supercritical \(\mathrm{CO}_{2}\left(\mathrm{scCO}_{2}\right)\) and in the chlorinated solvent \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\). The kinetic data for the reaction in both solvents are plotted in the graph. In what respects is the use of \(\operatorname{scC} \mathrm{O}_{2}\) a good example of a green chemical reaction?

In the following three instances, which choice is greener in each situation? Explain. (a) Petroleum as a raw material or vegetable oil as a raw material. (b) Toluene as a solvent or water as a solvent. (c) Catalyzed reaction at \(600 \mathrm{~K}\) or uncatalyzed reaction at \(800 \mathrm{~K}\).

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 i newspaper articles and would like an explanation: \((\mathbf{a}\) acid rain, \((\mathbf{b})\) greenhouse gas, \((\mathbf{c})\) photochemical smog (d) ozone depletion. Give a brief explanation of each term an identify one or two of the chemicals associated with each.

Suppose that on another planet the atmosphere consists of \(10 \% \mathrm{Kr}, 40 \% \mathrm{CH}_{4},\) and \(50 \% \mathrm{O}_{2} .\) What is the average molar mass at the surface? What is the average molar mass at an altitude at which all the \(\mathrm{O}_{2}\) is photodissociated?

If an average \(\mathrm{O}_{3}\), 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.

Liquefied petroleum gas (LPG) consists primarily of propane, \(\mathrm{C}_{3} \mathrm{H}_{8}(l)\) or butane \(\mathrm{C}_{4} \mathrm{H}_{10}(l)\) (a) Write a balanced chemical equation for the complete combustion of propane to produce \(\mathrm{CO}_{2}(g)\) as the only carbon-containing product. (b) Write a balanced chemical equation for the incomplete combustion of propane to produce \(\mathrm{CO}(g)\) as the only carbon-containing product. (c) At \(25^{\circ} \mathrm{C}\) and \(101.3 \mathrm{kPa}\) pressure, what is the minimum quantity of dry air needed to combust \(10.0 \mathrm{~mL}\) of \(\mathrm{C}_{3} \mathrm{H}_{8}(l)\) completely to \(\mathrm{CO}_{2}(g)\) ? The density of the LPG is \(0.50 \mathrm{~g} / \mathrm{mL}\).