Chapter 18

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

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 seawater before it can be considered freshwater (dissolved salts \(<500 \mathrm{ppm}) ?[\) Section 18.3\(]\)

(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 \(3.0 \mathrm{ppm} .\) Calculate the number of \(\mathrm{CO}\) 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 \(240 \mathrm{~kJ} / \mathrm{mol}\). (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 (a) correspond to? part

In \(\mathrm{CH}_{3} \mathrm{I}\) the \(\mathrm{C}-\mathrm{I}\) bond-dissociation energy is \(241 \mathrm{~kJ} / \mathrm{mol}\). In \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{I}\) the \(\mathrm{C}-\mathrm{I}\) 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 \(\mathrm{N}_{2}\). (b) In which region of the electromagnetic spectrum does this light fall? Does this light have enough energy to photoionize \(\mathrm{N}_{2} ?\)

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

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?

Nitrogen oxides like \(\mathrm{NO}_{2}\) and 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 \(\mathrm{pH}\) 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.

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 \(420-\mathrm{nm}\) light? \((\mathbf{c})\) 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?

Sulfur is present in seawater to the extent of \(0.09 \%\) by mass. Assuming that the sulfur is present as sulfate, \(\mathrm{SO}_{4}^{2-}\), calculate the corresponding molar concentration of \(\mathrm{SO}_{4}^{2-}\) 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 \((1 \mathrm{~W}=1 \mathrm{watt}=1 \mathrm{~J} / \mathrm{s}) .\) (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 \(12-\mathrm{h}\) 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 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.

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: $$\begin{array}{l} \mathrm{C}_{18} \mathrm{H}_{35} \mathrm{O}_{2}^{-}(a q)+26 \mathrm{O}_{2}(a q) \longrightarrow \\ 17 \mathrm{CO}_{2}(a q)+17 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{HCO}_{3}^{-}(a q) \end{array}$$ 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 \(50.0 \mathrm{~g}\) 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.

(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) 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 in newspaper articles and would like an explanation: (a) acid rain, \((\mathbf{b})\) greenhouse gas, \((\mathbf{c})\) photochemical smog, (d) ozone depletion. Give a brief explanation of each term and 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?

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

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, \(\mathrm{C}_{3} \mathrm{H}_{8}(I)\) 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 \(0.50 \mathrm{~g} / \mathrm{mL}\)

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 \(\mathrm{SO}_{3}\), 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?

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?

In 2008 , the global average electricity consumption per head was \(3.0 \mathrm{MWh}\). The solar power striking Earth every day averages 168 watts per square meter. Considering that present technology for solar energy conversion is about \(10 \%\) efficient, from how many square meters of land must sunlight be collected in order to provide this power?

The CDC (Centers for Disease Control and Prevention) published a reference blood lead level (BLL), which is based on the BLL distribution among children. It is currently \(5 \mu \mathrm{g} / \mathrm{dL}\). (a) What is the molarity of an aqueous solution with this concentration? (b) Express this concentration in ppb.

As of the writing of this text, EPA standards limit atmospheric ozone levels in urban environments to 84 ppb. How many moles of ozone would there be in the air above Los Angeles County (area about 10,000 square kilometers; consider a height of \(100 \mathrm{~m}\) above the ground) if ozone was at this concentration?

The estimated average concentration of \(\mathrm{NO}_{2}\) in air in the United States in 2015 was 0.010 ppm. (a) Calculate the partial pressure of the \(\mathrm{NO}_{2}\) in a sample of this air when the atmospheric pressure is \(101 \mathrm{kPa}\). (b) How many molecules of \(\mathrm{NO}_{2}\) are present under these conditions at \(25^{\circ} \mathrm{C}\) in a room that measures \(10 \mathrm{~m} \times 8 \mathrm{~m} \times 2.50 \mathrm{~m} ?\)

A 500 megawatt electrical power plant typically burned (a) Assuming that 1,430,000 metric tons of coal in a year. the coal was \(80 \%\) carbon and \(3 \%\) sulfur and that combustion was complete, calculate the number of tons of carbon dioxide and sulfur dioxide produced by the plant during the year. \((\mathbf{b})\) If \(50 \%\) of the \(\mathrm{SO}_{2}\) could be removed by reaction with powdered \(\mathrm{CaO}\) to form \(\mathrm{CaSO}_{3},\) how many tons of \(\mathrm{CaSO}_{3}\) would be produced?

Common lab spectrometers can detect absorbance down to 0.0002 with good reliability. Consider a dissolved harmful organic substance with a molar mass of \(120.5 \mathrm{~g} / \mathrm{mol}\), which can be detected in this spectrometer. It shows an extinction coefficient of \(\varepsilon=1.43 \times 10^{3} \mathrm{M}^{-1} \mathrm{~cm}^{-1}\) at 320 \(\mathrm{nm}\), its absorption maximum (A Closer Look, p. 620). (a) Calculate the minimum concentration of the organic substance detectable by this spectrometer (path length \(1 \mathrm{~cm})\). (b) Convert the minimum observable molarity to ppb.

Bioremediation is the process by which bacteria repair their environment in response, for example, to an oil spill. The efficiency of bacteria for "eating" hydrocarbons depends on the amount of oxygen in the system, \(\mathrm{pH},\) temperature, and many other factors. In a certain oil spill, hydrocarbons from the oil disappeared with a first-order rate constant of \(2 \times 10^{-6} \mathrm{~s}^{-1}\). At that rate, how many days would it take for the hydrocarbons to decrease to \(10 \%\) of their initial value?

The following data were collected for the decomposition of \(\mathrm{O}_{3} \mathrm{by}\left(\mathrm{O}_{3}+\mathrm{H} \longrightarrow \mathrm{O}_{2}+\mathrm{OH}\right)\) at very low concentrations: $$\begin{array}{llll} \hline \text { Trial } & {\left[\mathrm{O}_{3}\right](M)} & {[\mathrm{H}](M)} & \text { Initial Rate }(\mathrm{M} / \mathrm{s}) \\ \hline 1 & 3.25 \times 10^{-33} & 2.25 \times 10^{-26} & 8.10 \times 10^{-15} \\\ 2 & 6.50 \times 10^{-33} & 4.50 \times 10^{-26} & 3.25 \times 10^{-14} \\ 3 & 6.48 \times 10^{-33} & 2.23 \times 10^{-26} & 1.62 \times 10^{-14} \\ \hline \end{array}$$ (a) Write the rate law for the reaction. (b) Calculate the rate constant.

The degradation of \(\mathrm{CF}_{3} \mathrm{CH}_{2} \mathrm{~F}\) (an \(\mathrm{HFC}\) ) by OH radicals in the troposphere is first order in each reactant and has a rate constant of \(k=2.1 \times 10^{8} \mathrm{M}^{-1} \mathrm{~s}^{-1}\) at \(10^{\circ} \mathrm{C}\). If the tropospheric concentrations of \(\mathrm{OH}\) and \(\mathrm{CF}_{3} \mathrm{CH}_{2} \mathrm{~F}\) are \(1.0 \times 10^{12}\) and \(7.5 \times 10^{14}\) molecules \(/ \mathrm{m}^{3}\), respectively, what is the rate of reaction at this temperature in \(M / \mathrm{s} ?\)