Jacobson M.Z. Atmospheric Pollution

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people most likely to be exposed to asbestos in concentrations high enough to cause

health problems. In the United States, asbestos is deﬁned as a hazardous air pollutant

under CAAA90.

Sources and Sinks

Once insulation containing asbestos has been installed, the asbestos is not expected

to cause damage to humans unless the insulation is disturbed. At that time, ﬁbers can

be scattered into the air, where they can remain for minutes to days until they deposit

to the ground or are inhaled. Thus, the only source of asbestos in indoor air is turbulent

uplift, and the only sink is deposition.

Concentrations

Indoor concentrations of asbestos vary from building to building. Lee et al. (1992)

found an average of 0.02 structures per cubic centimeter in 315 public, commercial,

residential, school, and university buildings in the United States. The concentration of

ﬁbers longer than 5 m was only 0.00013 structures per cubic centimeter

Health Effects

The primary health effects of asbestos exposure are lung cancer, mesothelioma,

and asbestosis. Mesothelioma is a cancer of the mesothelial membrane lining the

lungs, and asbestosis is a slow, debilitating disease of the lungs.

The most dangerous

asbestos fibers are those longer than 5 to 10 m but with a diameter less than 1 m.

Exposure to 0.0004 structures per cubic centimeter is expected to cause a lifetime

cancer risk of 160 cases of mesothelioma per 1 million people (Turco, 1997). The

time between first exposure to asbestos and the appearance of tumors is estimated to

be 20 to 50 years (Jones, 1999). Cigarette smoking and e

xposure to asbestos are

believed to amplify the rates of lung cancer in comparison with the rates of cancer

associated with just smoking or just exposure to asbestos. Short-term acute exposure

to asbestos can lead to skin irritation (Spengler and Sexton, 1983).

9.1.11. Fungal Spores, Bacteria, and Viruses

Fungal spores, bacteria, and viruses are common indoor air contaminants. Sources

of fungal spores are soils, plants, foodstuff, internal surfaces, and outdoor air. Sources

of bacteria and viruses are humans, animals, plants, air conditioners,

and outdoor air.

Fungi grow well on damp surfaces. Common fungi in buildings include Penicillium,

Cladosporium, and Aspergillus. Common bacteria include Bacillus, Staphylococcus,

and Micrococcus (Jones, 1999). Some diseases associated with fungi, bacteria, and

viruses include rhinitis (a respiratory illness), asthma, humidiﬁer fever, extrinsic aller-

gic alveolitis, and atopic dermatitis. Most human illnesses due to viral and bacterial

infection are due to human-to-human transmission of microorganisms rather than to

building-to-human transmission (Ayars, 1997).

9.1.12. Environmental Tobacco Smoke

When a cigarette is smoked, some of the smoke is inhaled and swallowed, some is

inhaled and exhaled (mainstream smoke), and the rest is emitted from the cigarette

(sidestream smoke). The mainstream plus sidestream smoke is called environmental

tobacco smoke (ETS), a mixture of more than 4,000 aerosol particle components and

248 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

gases, at least 50 of which are known carcinogens. ETS, itself, has been classiﬁed by

the U.S. EPA as a carcinogen. Also called second-hand smoke, ETS, builds up in

enclosed spaces, increasing danger to others in the vicinity. Even in well-ventilated

indoor areas, particle and gas concentrations associated with ETS increase. Although

the cumulative effect of ETS on outdoor air pollution is relatively small compared

with the effects of other sources of pollution, such as automobiles, ETS concentrations

can build up outdoors in the vicinity of smokers.

Outdoor ETS emission rates are not regulated in the United States, although

regulations prohibiting smoking in many public and private indoor facilities cur-

rently exist at the local, state, and federal levels. Many chemical constituents of

ETS are classified as hazardous air pollutants under CAAA90,

but the Act controls

only sources emitting more than 10 tons of a hazardous substance per year, and

individual cigarettes emit less than 1 g of all pollutants combined. The product of

the number of cigarettes smoked per year multiplied by the emission rate per ciga-

rette is much larger than 10 tons per year, but this statistic is not recognized

by CAAA90.

Sources

In 1990, approximately 26 percent of the U.S. population, or 50 million people,

were estimated to smoke cigarettes regularly (Rando et al., 1997). In 1986, 70 percent

of all children in the United States li

ved in households in which at least one parent

smoked (Weiss, 1986). Table 9.2 shows mainstream and sidestream emissions for

some components emitted from cigarettes. The actual emission rate of a cigarette

depends on the type of tobacco, the density of its packing, the type of wrapping paper,

and the pufﬁng rate of the smoker (Hines et al., 1993).

Table 9.2 indicates that emission rates of sidestream smok

e are greater than are

those of mainstream smoke for many pollutants. Thus, in many cases, a person stand-

ing a short distance from a cigarette is exposed to more pollution than is the smoker

(Schlitt and Knöppel, 1989).

Table 9.3 compares emissions from cigarettes with those from vehicles. The table

indicates that driving a vehicle that meets current U.S. EPA emission standards 1 mile

results in particle emissions equivalent to emissions from about 1.4 cigarettes.

Emissions of CO(g) and NO

(g) from driving 1 mile are equivalent to those from

smoking 73 and 190 cigarettes, respectively. The total emissions per day of these sub-

stances from cigarettes is much less than that from mobile sources,

but ETS is often

emitted in enclosed spaces, where its concentrations build up.

Concentrations

ETS contributes to the buildup of gas and particle concentrations indoors. Spengler

et al. (1981) found that one pack of cigarettes per day contributes about 20 g m

3

particles over a 24-hour period. During the time a cigarette is actually smoked, particle

concentrations increase to 500 to 1,000 g m

3

. Leaderer et al. (1990) found that par-

ticle concentrations in homes with a cigarette smoker were up to three times those in

homes without a smoker.

Health Effects

Short-term exposure to ETS results in eye, nose, and throat irritation for most

individuals, and allergic skin reactions for some (Maroni et al., 1995). ETS also

elevates symptoms for people who have asthma, and may induce asthma in some

INDOOR AIR POLLUTION 249

children (Jones, 1999). A link between exposure to ETS and lower respiratory tract

illness has been found (Somerville et al., 1988). Children exposed to ETS are hos-

pitalized more often than are children not e

xposed (Harlap and Davies, 1974).

Several studies have linked long-term ETS exposure to lung cancer. Janerich et al.

(1990) found that ETS may cause up to 17 percent of lung cancers among non-

smokers. Rando et al. (1997) found a 30 percent increase in the cancer risk to

women whose husbands smoked. Ryan et al. (1992) linked brain cancer tumors

to ETS.

250 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

Table 9.2. Mainstream and Sidestream Emission Rates per Cigarette

Carbon dioxide 10,000–80,000 81,000–640,000

Carbon monoxide 500–26,000 1,200–65,000

Nitrogen oxides 16–600 80–3,500

Ammonia 10–130 400–9,500

Hydrogen cyanide 280–550 48–203

Formaldehyde 20–90 1,000–4,600

Acrolein 10–140 100–1,700

N-nitrosodimethylamine 0.004–0.18 0.04–149

Nicotine 60–2,300 160–7,600

Total par ticulate 100–40,000 130–76,000

Phenol 20–150 52–390

Catechol 40–280 28–196

Naphthalene 2.8 45

Benzo(a)pyrene 0.008–0.04 0.02–0.14

Aniline 0.10–1.20 3–36

2-Naphthylamine 0.004–0.027 0.02–1.1

4-Aminobiphenyl 0.002–0.005 0.06–0.16

N-nitrosonornicotine 0.2–3.7 0.02–18

Mainstream Smoke Sidestream Smoke

Substance (g per Cigarette) (g per Cigarette)

Source: Rando et al., 1997; Jones, 1999.

Table 9.3. Comparison of Emissions from Cigarettes and Mobile Sources

Carbon monoxide 0.0464 3.4 73.3 60 189,000

Nitrogen oxides 0.0021 0.4 190.5 2.7 32,000

Particulates 0.058 0.08 1.4 75 9300

Number of (d) Estimated

(a) (b) Cigarettes Estimated Mobile-Source

Average Average Resulting in Cigarette Direct

Cigarette Vehicle Same Emissions Emissions in the Emissions in the

Emissions Emissions as 1 Mile of United States United States

Substance (gCigarette)

(gmi)

Driving

(tonsday)

Sum of average mainstream and sidestream emissions in Table 9.2.

1994 federal light-duty vehicle emission standards.

Column (b) divided by column (a).

Assumes the current population of the United States is 250,000,000 and 26 percent of the population

smokes 20 cigarettes per day.

Estimated 1997 emissions from U.S. EPA (1997).

9.2. SICK BUILDING SYNDROME

In some workplaces, employees experience an unusually high rate of headaches, nau-

sea, nasal congestion, chest congestion, eye problems, throat problems, fatigue, fever,

muscle pain, dizziness, and dry skin. These symptoms, present during working hours,

often improve after a person leaves work. The situation described is called sick build-

ing syndrome (SBS), the cause of which is unknown, although it may be due to

certain VOCs, building ventilation systems, or the exposure to many pollutants in low

doses simultaneously. SBS may also be caused by enhanced stress levels and heavy

workloads or a combination of psychological and chemical factors (Jones, 1999).

9.3. REGULATION OF INDOOR AIR POLLUTION

In the United States, National Ambient Air Quality Standards (NAAQS), initiated under

the Clean Air Act Amendments, control outdoor air pollutants only. No regulations control

air pollution concentrations in indoor residences. Standards for pollutant concentrations in

indoor workplaces are set by a government agency, the Occupational Safety and Health

Administration (OSHA), which obtains recommendations for standards from another go

ernment agency, the National Institute for Occupational Safety and Health (NIOSH), and

an independent professional society, the American Conference of Governmental Industrial

Hygienists, Inc. (ACGIH). OSHA and NIOSH were created by the 1970 U.S.

Occupational Safety and Health Act. OSHA is in the U.S. Department of Labor and is

responsible for setting and enforcing workplace standards, whereas NIOSH is in the U.S.

Department of Health and Human Services and is responsible for researching workplace

health issues. ACGIH’s primary mission is to promulgate workplace safety standards.

NIOSH recommends permissible exposure limits (PELs), short-time exposure

limits (STELs), and ceiling concentrations. A PEL is the maximum allowable con-

centration of a pollutant in an indoor workplace over an 8-hour period during a day.

An STEL is the maximum allowable concentration of a pollutant over a 15-minute

period. A ceiling concentration is a concentration that may never be exceeded. ACGIH

sets 8-hour time-weighted average threshold limit values (TWA-TLVs), which are

similar to PELs. For most pollutants, PELs and TWA-TLVs are the same. When differ-

ences occur, they are small.

Indoor standards exist for more than 150 compounds. Table 9.4 compares indoor

with outdoor standards for a few compounds. For ozone and carbon monoxide, the

INDOOR AIR POLLUTION 251

Table 9.4. Comparison of Indoor Workplace Standards with Outdoor Federal and California

State Standards for Selected Gases

Carbon monoxide 35 — 200 9.5 (8 hour) 9 (8 hour)

Nitrogen dioxide — 1 — 0.053 (annual) 0.25 (1 hour)

Ozone 0.1 0.3 — 0.08 (8 hour) 0.09 (1 hour)

Sulfur dioxide 2 5 — 0.14 (24 hour) 0.05 (24 hour)

Indoor 8-hour Indoor Outdoor

PEL and TWA-TLV 15-min STEL Indoor Ceiling Outdoor NAAQS California

Gas (ppmv)

(ppmv)

(ppmv) Standard (ppmv)

National Institute for Occupational Safety and Health (2000).

8-hour PEL/TWA-TLV standards are less stringent than are the outdoor standards

because outdoor standards are designed to protect the entire population, particularly

infants and people afﬂicted with disease or illness. Indoor standards are designed to

protect workers, who are assumed to be healthier than is the average person. The table

also indicates that the 15-minute STEL for nitrogen dioxide is four times higher (less

stringent) than is the 1-hour outdoor California standard. Stringent outdoor standards

for nitrogen dioxide are set because it is a precursor to photochemical smog. Because

UV sunlight does not penetrate indoors, nitrogen dioxide does not produce ozone

indoors, and indoor regulations of nitrogen dioxide as a smog precursor are not neces-

sary. Indoor standards for nitrogen dioxide are based solely on health concerns.

9.4. SUMMARY

People spend most of their time indoors; thus, most of their exposure to pollution occurs

indoors. Indoor air contains many of the same pollutants as does outdoor air, but pollu-

tion concentrations in indoor and outdoor air usually differ. Indoor mixing ratios of

ozone and sulfur dioxide are usually less than are those outdoors, but indoor mixing

ratios of formaldehyde are usually greater than are those outdoors. Indoor mixing ratios

of carbon monoxide and nitrogen dioxide are generally the same as or less than are those

outdoors, unless appliances or other indoor combustion sources are turned on. Indoor

concentrations of radon and environmental tobacco smoke, when present, are usually

greater than are those outdoors, giving rise to potentially serious health problems to peo-

ple exposed to these pollutants indoors. Indoor air also contains v

olatile organic

compounds, allergens, fungi, bacteria, and viruses. In the United States, indoor air is reg-

ulated only in the workplace and in public buildings; residential air is not regulated.

9.5. PROBLEMS

9.1. Why are ozone mixing ratios almost always lower indoors than outdoors?

9.2. Why are workplace standards for pollutant concentrations generally less strin-

gent than standards for outdoor air?

9.3. What would be the volume mixing ratio (ppmv) of carbon monoxide and the

mass concentration (g m

3

) of particulates if ten cigarettes were smoked in a

5 m  10 m  3 m room. How do these values compare with the California

state 1-hour standard for CO(g) and the 24-hour average standard for PM

Based on the results, which pollutant do you think is more of a cause for concern

with respect to indoor air quality? Assume the dry air partial pressure is 1013

mb, the temperature is 298 K, and use cigarette emission rates from Table 9.3.

9.4. Why is radium less of a concern than radon?

9.5. Why is removing asbestos from buildings often more dangerous than leaving it

alone?

9.6. Why does the gas-phase chemical decay of organic compounds generally take

longer indoors than outdoors?

252 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

ACID DEPOSITION

254

cid deposition occurs when sulfuric acid, nitric acid, or hydrochloric acid,

emitted into or produced in the air as a gas or liquid, deposits to soils,

lakes, farmland, forests, or buildings. Deposition of acid gases is dry acid

deposition, and deposition of acid liquids is wet acid deposition. Wet acid deposition

can be through rain (acid rain), fog (acid fog), or aerosol particles (acid haze). On the

Earth’s surface, acids have a variety of environmental impacts, including damage to

microorganisms, ﬁsh, forests, agriculture, and structures. In the air, acids in high con-

centrations are harmful to humans. Acid deposition problems have existed since coal

was ﬁrst combusted, but were exacerbated during the Industrial Revolution in the eigh-

teenth century. The problem became more severe with the growth of the alkali industry

in 19th-century France and the United Kingdom. In this chapter, the history, science,

and regulation of acid deposition problems are discussed.

10.1. HISTORICAL ASPECTS OF ACID DEPOSITION

Acid deposition is caused by the emission or atmospheric formation of gas- or

aqueous-phase sulfuric acid (H

), nitric acid (HNO

), or hydrochloric acid

(HCl). Historically, coal was the first and largest source of anthropogenically

produced atmospheric acids. Coal combustion produces sulfur dioxide gas [SO

(g)]

and hydrochloric acid gas [HCl(g)]. SO

(g) produces gas-and aqueous-phase

sulfuric acid. Humans have combusted coal for thousands of years. In the 1200s,

sea coal was brought to London and used in lime

kilns and forges (Section 4.1). It was later burned

in furnaces to produce glass and bricks, in brew-

eries to produce beer and ale, and in homes to

provide heating. During the Industrial

Revolution, which started in the eighteenth centu-

ry, coal was used to provide energy for the steam

engine.

In the late eighteenth century, a second major

source of atmospheric acids emerged. Around 1780,

the demand for sodium carbonate [Na

(s)]

(also known as soda ash, washing soda, and salt

cake), used in the production of soaps, detergents,

cleansers, glass, paper, bleaches, and dyes, increa-

sed. Although this chemical could be extracted from

barilla plants, which grew near the Mediterranean

Sea, and from sea kelp, a more efﬁcient method of

producing it was desired. In 1781, the French

Academy of Sciences offered a prize to the person

who could develop the most efﬁcient and economic

method of producing soda ash. The prize was never

awarded because of the French Revolution, but

Nicolas Leblanc (1742–1806; Fig. 10.1), encour-

aged by the competition, began experimenting in 1784 to ﬁnd a new process. In 1789,

he came up with a two-step set of reactions to produce soda ash. In the ﬁrst step, he

Figure 10.1. Nicolas Leblanc (1742–1806).

dissolved common salt into a sulfuric acid solution at a high temperature in an iron

pan to produce dissolved sodium sulfate (Glauber’s salt) and hydrochloric acid

gas by

High

temperature

2NaCl(s)  H

(aq) Na

(aq)  2HCl(g)

Sodium Sulfuric Sodium Hydrochloric (10.1)

chloride (salt) acid sulfate acid

In the second step, he heated sodium sulfate together with charcoal and chalk in a kiln

to form sodium carbonate by

High

temperature

(aq)  2C  CaCO

(s) Na

(aq)  CaS(s)  2CO

(g)

Sodium Carbon Calcium Sodium Calcium Carbon (10.2)

sulfate from carbonate carbonate sulﬁde dioxide

charcoal from chalk

By-products of the second reaction, when complete, included calcium sulﬁde

[CaS(s)], an odorous, yellow to light-gray powder, and carbon dioxide gas. In practice,

the reaction was incomplete and produced gas-phase sulfuric acid and soot as well.

The sodium carbonate residue was separated from calcium sulﬁde by adding water,

which preferentially dissolved the sodium carbonate. The resulting solution was then

dried, producing sodium carbonate crystals. Sodium carbonate w

as then combined

with animal fat to produce soap.

A necessary ingredient for the production of sodium carbonate was aqueous sulfu-

ric acid (Reaction 10.1). This was obtained by burning elemental sulfur (S) powder

with saltpeter [KNO

(s)], then dissolving the resulting H

(g) in water, as was done

by Libavius in 1585. This process was inefﬁcient, releasing volumes of gas-phase sul-

furic acid and nitric oxide [NO(g)] (which converts to nitrogen dioxide, then to nitric

acid in air).

In sum, although the production of sodium carbonate from common salt and sul-

furic acid allowed the alkali industry to become self-sufﬁcient and escape reliance on

the import of natural sodium carbonate, it resulted in the release of HCl(g), H

(g),

HNO

(g), and soot, causing widespread acid deposition and air pollution in France

and the United Kingdom. The Leblanc process also produced large amounts of

impure solid “alkali waste” containing calcium sulﬁde. The waste was piled near each

factory. When rainwater fell on the waste, some of the calcium sulﬁde dissolved, pro-

ducing dissolved hydrogen sulﬁde H

S(aq), which evaporated as a harmful, odorous,

colorless gas. Much of the rest of the calcium sulﬁde oxidized over time to form cal-

cium sulfate [CaSO

-2H

O(s), gypsum]. Piles of gypsum from soda ash factories can

still be seen today.

Although Leblanc patented the soda ash technique and started a soda ash factory in

1791, he lost his patent and factory to the state, which nationalized patents and facto-

ries in 1793. Leblanc ultimately committed suicide in 1806 because of his inability to

recover from his business losses.

In 1863, an estimated 1.76 million tons of raw material were burned to form soda ash,

producing only 0.28 million tons of useful products. Most of the remaining 1.48 million

ACID DEPOSITION 255

tons was emitted as HCl(g) or other gases and produced as solid waste (Brock, 1992).

Environmental damage due to the alkali industry was severe. HCl(g) and other pollutants

rained down onto agricultural property and cities in France and the U.K., causing proper-

ty and health damage. In the United Kingdom, St. Helens, Newcastle, and Glasgow

countrysides were decimated (Brimblecombe, 1987).

Early complaints against alkali factories were in the form of civil litigation. For

example, in 1838, a Liverpool landowner ﬁled a complaint against an alkali factory,

charging that it destroyed his crops and interfered with his hunting. Ultimately, HCl(g)

from alkali manufacturers was regulated in France and the United Kingdom. In

France, regulation took the form of planning laws that controlled the location of alkali

factories. In the United Kingdom, the 1863 Alkali Act required alkali manufacturers to

reduce 95 percent of their HCl(g) emissions. The act also called for the appointment of

an alkali inspector to watch over the industry.

Impetus for the 1863 Alkali Act came in the early 1860s, when William Gossage

invented a technique to wash HCl(g) from waste gases before it was released from

chimneys. Gossage built his own soda ash factory in Wo

rcestershire in 1830. Spurred

by his neighbors’ complaints about the hydrochloric acid emitted from his factory, he

worked to mitigate the problem. His solution was to convert a windmill into a tower,

ﬁll the tower with brushwood, and spray water down the top of the tower as smoke

rose from the bottom. The water dissolved most of the hydrochloric acid (just as rain

does) and drained it into a nearby waterway (which he was not too concerned about).

Because the technique was so simple and inexpensive and because the consequences

of not implementing the technique were so severe, the United Kingdom easily passed

the 1863 Alkali Act shortly after Gossage’s invention. Subsequent to the Act, Walter

Weldon (in 1866) and Hugh Deacon (in 1868) developed processes for converting

HCl(g) to chlorine that could be used for bleaching powder. These inventions allowed

chlorine, which otherwise would have been wasted, to be recycled.

Despite the success of the Alkali Act at reducing HCl(g), the alkali industry contin-

ued to emit sulfuric acid, nitric acid, soot, and other pollutants in abundance. Because

the act did not control emissions from factories other than those producing soda ash,

pollution problems in the U.K. worsened. Although the Alkali Act was modiﬁed in

1881 to regulate other sections of the chemical industry, the number of factories and

the volume of emissions from them had grown so much that the new law had little

effect. In 1899, one writer described St. Helens as

a sordid ugly town. The sky is a low-hanging roof of smeary smoke. The

atmosphere is a blend of railway tunnel, hospital ward, gas works and open

sewer. The features of the place are chimneys, furnaces, steam jets, smoke

clouds and coal mines. (Blatchford, 1899)

The ﬁrst Alkali Act inspector in the United Kingdom was Scottish chemist Robert

Angus Smith (1817–1884; Fig. 10.2). He was charged with ensuring industry reduced

HCl(g) emissions by 95 percent. Smith was also a ﬁeld experimentalist. In 1872, he

published Air and Rain: The Beginnings of a Chemical Climatology, in which he dis-

cussed results of the ﬁrst monitoring network for air pollution in Great Britain. As part

of the analysis, he recorded the gas-phase mixing ratios of molecular oxygen and car-

bon dioxide and measured the composition of chloride, sulfate, nitrate, and ammonium

in rainwater in the British Isles. In his book, Smith introduced the term acid rain to

describe the high sulfate concentrations in rain near coal-burning facilities.

256 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

Leblanc’s soda ash process dominated the alkali industry until the 1880s. In 1861,

Belgian chemist Ernst Solvay (1838–1922) developed a more efﬁcient process of pro-

ducing soda ash by raining a salt produced from

sodium chloride and ammonia gas down a tower

(Solvay tower) over an upcurrent of carbon dioxide

gas. The technique required no sulfuric acid or

potassium nitrate (saltpeter). The disadvantages were

the cost of the tower and the fact that the ﬁnal chlo-

ride was tied up in a form that could not easily be

reused. Only by the 1880s did the Solvay process

overtake the LeBlanc process in economic efﬁciency

and popularity.

Throughout the twentieth century, acid deposi-

tion problems continued to plague cities in the

United Kingdom and municipalities in other coun-

tries where coal and chemical b

urning occurred. The

fatal pollution episodes in London, discussed in

Chapter 4, included contributions from acidic com-

pounds in smoke.

Acid deposition became an issue of international

interest in the 1950s and 1960s, when a relationship

was found between sulfur emissions in continental

Europe and acidiﬁcation of Scandinavian lakes.

These studies led to a 1972 United Nations

Conference on Human Environment in Stockholm

that called for an international ef

fort to reduce acidi-

ﬁcation. Since then, numerous studies on acid

deposition have been carried out. As a result of these studies, national governments

have intervened to control acid deposition problems. Such efforts are discussed at the

end of this chapter. First, the basic science of acid deposition is discussed.

10.2. CAUSES OF ACIDITY

In Chapter 5, pH was deﬁned as

pH log



] (10.3)

where [H



] is the molarity (moles per liter) of H



in a solution containing a solvent and

one or more solutes. The pH scale, shown in Fig. 10.3 for a limited range, varies from less

than 0 (lots of H



and very acidic) to greater than 14 (very little H



and very basic or

alkaline). Neutral pH, the pH of distilled water, is 7.0. At this pH, the molarity of H



7

mol L

1

. A pH of 4 means that the molarity of H



is 10

4

mol L

1

. Thus, water at a

pH of 4 is 1,000 times more acidic (contains 1,000 times more H



ions) than is water at a

pH of 7.

The molarity of H



is related to that of OH



by the equilibrium relationship,

O(aq)



 OH



Liquid Hydrogen Hydroxide (10.4)

water ion ion

ACID DEPOSITION 257

Figure 10.2. Robert Angus Smith (1817–1884).