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]

The Earth's oceans have a salinity of \(35 .\) What is the concentration of dissolved salts in seawater when expressed in ppm? What percentage of salts must be removed from sea-water before it can be considered freshwater (dissolved salts \(<500\) ppm \() ?[\operatorname{Section} 18.3]\)

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.

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 ppb (0.441 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 ppb if the atmospheric pressure is 0.67 atm. (b) How many ozone molecules are in 1.0 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 ppm. Calculate the number of CO molecules in 1.0 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 276 \(\mathrm{kJ} / \mathrm{mol}\) . (a) What is the maximum wave-length of photons that can cause \(\mathrm{C}-\) 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 \(\mathrm{C}-\mathrm{Cl}\) 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 km.

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

The wavelength at which the \(\mathrm{O}_{2}\) molecule most strongly absorbs light is approximately 145 \(\mathrm{nm}\) . (a) In which region of the electromagnetic spectrum does this light fall? (b) Would a photon whose wavelength is 145 nm have enough energy to photodissociate \(\mathrm{O}_{2}\) whose bond energy is 495 \(\mathrm{kJ} / \mathrm{mol} ?\) Would it have enough energy to photoionize \(\mathrm{O}_{2} ?\)

The ultraviolet spectrum can be divided into three regions based on wavelength: UV-A \((315-400 \mathrm{nm}), \mathrm{UV}-\mathrm{B}(280-315\) \(\mathrm{nm} ),\) and \(\mathrm{UV}-\mathrm{C}(100-280 \mathrm{nm}) .\) (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? \begin{equation}\begin{array}{l}{\text { (a) } \mathrm{O}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{O}_{3}^{*}(g)} \\ {\text { (b) } \mathrm{O}_{3}^{\star}(g)+\mathrm{M}(g) \longrightarrow \mathrm{O}_{3}(g)+\mathrm{M}^{\star}(g)} \\ {\text { (c) } \mathrm{O}_{2}(g)+h \nu \longrightarrow 2 \mathrm{O}(g)}\\\\{\text { (d) } \mathrm{O}(g)+\mathrm{N}_{2}(g) \longrightarrow \mathrm{NO}(g)+\mathrm{N}(g)} \\\ {\text { (e) All of the above }}\end{array}\end{equation}

(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 \(\mathrm{CFC}-11, \mathrm{CFCl}_{3}\) . What chemical characteristics of this substance allow it to effectively deplete stratospheric ozone?

The average bond enthalpies of the \(\mathrm{C}-\mathrm{F}\) and \(\mathrm{C}-\) Cl bonds are 485 \(\mathrm{kJ} / \mathrm{mol}\) and 328 \(\mathrm{kJ} / \mathrm{mol}\) , respectively. (a) What is the maximum wavelength that a photon can possess and still have sufficient energy to break the \(\mathrm{C}-\mathrm{F}\) and \(\mathrm{C}-\mathrm{Cl}\) 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 \(\mathrm{C}-\mathrm{F}\) bonds to be significant in the lower atmosphere?

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) Write a chemical equation that describes the attack of acid rain on limestone, CaCO_. (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? Explain.

Alcohol-based fuels for automobiles lead to the production of formaldehyde (CH \(_{2} \mathrm{O} )\) in exhaust gases. Formaldehyde undergoes photodissociation, which contributes to photo-chemical smog: $$\mathrm{CH}_{2} \mathrm{O}+h v \longrightarrow \mathrm{CHO}+\mathrm{H}$$ The maximum wavelength of light 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 \(8.3 .\) 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 kJ/mol, that can be broken by absorption of a photon of \(420-\mathrm{nm}\) 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 \(6 \times 10^{-3}\) atm 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 \(-5.7^{\circ} \mathrm{C}\left(22^{\circ} \mathrm{F}\right),\) while the average nighttime temperature is \(-79^{\circ} \mathrm{C}\left(-109^{\circ} \mathrm{F}\right) .\) 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?

What is the molarity of \(\mathrm{Na}^{+}\) in a solution of \(\mathrm{NaCl}\) whose salinity is 5.6 if the solution has a density of 1.03 \(\mathrm{g} / \mathrm{mL}\) ?

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

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 \((1 \mathrm{W}=1\) 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 \(\mathrm{J} / \mathrm{g}^{\circ} \mathrm{C}\) . 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 \((1 \mathrm{W}=1\) 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 -h day. (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 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 lb of \(\mathrm{Mg}(\mathrm{OH})_{2} ?\)

Gold is found in seawater at very low levels, about 0.05 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 to prosity of seawater is 1.03 \(\mathrm{g} / \mathrm{mL}\) and that your gold recovery process is 50\(\%\) efficient.

(a) Explain why the concentration of dissolved oxygen in freshwater is an important indicator of the quality of the water. (b) Find graphical data in the text that show variations of gas solubility with temperature, and estimate to two significant figures the percent solubility of \(\mathrm{O}_{2}\) in water at \(30^{\circ} \mathrm{C}\) as compared with \(20^{\circ} \mathrm{C} .\) How do these data relate to the quality of natural waters?

The organic anion is found in most detergents. Assume that the anion undergoes aerobic decomposition in the following manner: $$\begin{aligned} 2 \mathrm{C}_{18} \mathrm{H}_{29} \mathrm{SO}_{3}^{-}(a q)+51 \mathrm{O}_{2}(a q) & \longrightarrow \\ & 36 \mathrm{CO}_{2}(a q)+28 \mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{H}^{+}(a q)+2 \mathrm{SO}_{4}^{2-}(a q) \end{aligned}$$ 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 ppm \(\mathrm{O}_{2}\) are 50\(\%\) depleted of oxygen in 1 day by a population of \(1,200,000\) people?

Magnesium ions are removed 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.

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 \(\mathrm{Mg}(\mathrm{OH})_{2}(s) :\) $$\begin{array}{c}{\mathrm{Ca}^{2+}(a q)+\mathrm{CO}_{3}^{2-}(a q) \longrightarrow \mathrm{CaCO}_{3}(s)} \\ {\mathrm{Mg}^{2+}(a q)+2 \mathrm{OH}^{-}(a q) \longrightarrow \mathrm{MgOH}_{2}(a q)}\end{array}$$ 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 \(\mathrm{Mg}^{2+} ) 1200 \mathrm{L}\) of water in which $$\begin{array}{l}{\left[\mathrm{Ca}^{2+}\right]=5.0 \times 10^{-4} M \text { and }} \\ {\left[\mathrm{Mg}^{2+}\right]=7.0 \times 10^{-4} \mathrm{M?}}\end{array}$$

(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?

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 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.

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 Chatelier's principle, why the equilibrium constant for the formation of 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, \(\mathrm{CH}_{4}(g)\) . (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 CO(g) as the only carbon-containing product. (c) At \(25^{\circ} \mathrm{C}\) and 1.0 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 \(S O_{2}\) into the atmosphere. Most of this \(S O_{2}\) underwent oxidation to \(S O_{3},\) 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?