ChemMatters, SEPTEMBER 2005 11

By Michael Tinnesand

It is a word you hear all the time. Get offa wild carnival ride and it takes a whileto restore your sense of equilibrium.The directions for setting up a new fishtank advise you to wait until the water

reaches equilibrium before you add fish. Whengrowing communities finally see populationincreases level off, the news might report thearea population is nearing equilibrium.

Have you heard the word in scienceclass? Physical science classes sometimestalk about levers and pivots at equilibrium.Astronomers debate whether the universe isexpanding, contracting, or reaching an equi-librium. And one of the most important andreported discussions in science is about theatmosphere and climate change: What is hap-pening to the Earth’s atmosphere that mightupset its equilibrium?

It’s easy to get confused about a wordused in so many ways. One of the most com-mon misconceptions about equilibrium is thatit means things have stopped changing. Butlet’s think about the community example—the one in which the population has reachedan equilibrium. When a population is grow-ing, the number of people newto the community is greaterthan people leaving the com-munity. Babies are born; newpeople move in; numbersgrow. But that’s not the wholestory. At the same time, peoplemove away and people die.

But, let’s do the math. When the population isgrowing, the rate of adding new people isgreater than the loss of people who leave ordie. At some point the population may reach astable number. Has change stopped? No.People still move into the community andthey still leave; it’s just that they are doing itat equal and opposite rates. Equal opposingactions. That’s equilibrium.

Take the example of thefish tank. Perhaps you fill thetank with water from your tapand find it is fairly cold andcontains a lot of dissolvedchlorine. As you let it sit for afew days, the temperature

gradually begins to rise and the smell of chlo-rine begins to decrease. Finally, the tankreaches a steady state where the temperatureis constant. You test for chlorine and find thatthe amount is staying constant as well. So,the tank has stopped changing, right?Wrong! Scientists call this situation adynamic equilibrium. And understanding howthis kind of balancing act works is importantto understanding a great deal about how theworld works.

Here’s what’s going on with the chlorineconcentrations in the fish tank. Two thingsare happening at once. Chlorine is leaving thetank to mix with and dissolve in the surround-ing room air. Initially, this rate is pretty large,because there’s a lot more chlorine in thewater than in the air. But it’s possible forsome atoms of chlorine in the air near thetank to find their way back into the aquarium.Possible. But, with a whole room full of air inwhich to escape, chlorine is very, veryunlikely to do that. Eventually, the amount ofchlorine in the tank stabilizes. Most of thechlorine has left the tank, so the rate of chlo-rine still leaving has become very small.Here’s the important point: Although the con-centration of chlorine in the tank is constant,there is still a very small amount leaving anda very small amount entering the water. Theleaving and entering rates are the same—that’s dynamic equilibrium.

Now think about the chemical reactionsyou’ve observed and described in class. It’s

PH

OT

OD

ISC

CRAI

G ER

ICKS

ON,W

WW

.CLO

QUET

MN.

NET

12 ChemMatters, SEPTEMBER 2005 http://chemistry.org/education/chemmatters.html

easy to get the impression that reactions onlygo in one direction. In fact, the way most text-books show and represent reactions in writtenequations just reinforces this one-way-streetidea. Take the reaction of hydrogen with oxy-gen. When hydrogen burns, it reacts rapidlywith oxygen, giving off a great deal of energy.The reaction is usually represented as

2H2 + O2 ➞ 2H2O + HeatThe way it’s written, it looks like the

hydrogen reacts with the oxygen and makeswater until the reactants run out. End of thestory. But, again, the system is not that sim-ple. It turns out that the reverse reaction canalso happen. Given enough energy input,water can break down into component hydro-gen and oxygen atoms. But since water is sostable, it doesn’t happen often. In this reac-tion, equilibrium is reached when virtually allthe reactants have been used up, lowering theformation of more water to almost nothing.This action is opposed by an equal vanishinglysmall number of water molecules breakingdown into separate atoms. It’s an equilib-rium—a dynamic equilibrium.

So why doesn’t this discussion come upwhen the textbook describes the reaction ofhydrogen with oxygen? It is because thisreaction, like many others you are learningabout, is so favorable to the formation ofproducts that it is barely worth mentioningthat the reverse reaction is also taking place.

Some chemical reactions, however, area different story! The reaction of hydrogenwith nitrogen to make ammonia is a goodexample. The formula for this reaction isalways written as

3H2 + N2 2NH3

Why the different reaction arrow? Undertypical standard conditions, equilibrium forthis reaction is reached well before all thereactants are consumed. With the doublearrow, the chemist shows that when a reac-tion “stops happening”, it has really just cometo an equilibrium of the forward and reversereactions. The reaction for making ammonia(NH3)—a component in fertilizers and otherproducts—is very important. Chemistsworked for many years to learn how to shiftthis equilibrium to favor the products. As aresult, industry can produce vital ammonia ina cost-effective manner.

The idea of a shifting equilibrium isanother important concept. Think back to ourexample of the population in a community. Thepopulation number may be in equilibrium for along time. But suppose something happens toshift one of the rates—like maybe a newindustry moving into the area. This mightcause the rate of people moving into the areato increase, raising the population. Eventually,this higher population will result in crowding,housing shortages, and even higher deathrates. As a result, the rate of people leaving will

finally balance the rate of people entering thearea, establishing a new dynamic equilibrium.

In a similar way, a chemical equilibriumcan shift. It may be influenced by factors likeincreasing the concentration of reactants,removing products, and changing the pressureor temperature. This principle was summa-rized in the late 19th century by Frenchchemist and engineer Henri-Louis Le Chatelier.

The chemistry of the Earth’s atmospherecontains many examples of various equilib-rium systems, each complex on its own. Butthey become even more complex when therelationships among all of these chemical sys-tems are taken into account. We would havefewer worries if the atmosphere retained itslong-term equilibrium. But such is not thecase. Natural and human-generated changesare constantly influencing our atmosphere,making it essential for us to study and under-stand the complex web of chemical reactionsthat define our atmosphere.

Examples of various systems and factorsthat influence their equilibrium abound.

Consider the ozone (O3) layer in thestratosphere. The ozone layer is importantbecause it absorbs and screens out a portionof the ultraviolet light coming from the sun.Without the ozone layer, the amount of UVreaching the earth’s surface would reach dan-gerous levels. This thin layer of O3 reachespeak concentration in the upper atmospherebetween altitudes of 19 and 23 km.

Maps of chemical concentrations from Aura's Microwave Limb Sounder instrument help us better understand the complex equilibria of the atmosphere. Shownabove are the concentrations of hydrogen chloride (top), chlorine monoxide (center), and ozone (bottom) for selected days during the 2004-2005 Arctic winter.

➞➞

NA

SA

/JP

L

Ozone forms when ultraviolet light fromthe sun breaks oxygen molecules apart. Thisatomic oxygen can join existing O2 mole-cules to form O3. By a reverse reaction,chemicals in the upper atmosphere causethe ozone to break apart. During times ofincreased sunspot activity the amount ofultraviolet light from the sun increases. Thisincrease in UV light increases the rate atwhich ozone is formed.

By these naturally opposing reactions,the net effect is this: More ozone eventuallyleads to more ozone breaking down. Theequilibrium is reestablished at a high con-centration of ozone. But the opposite shiftcan also happen. An increase in the concen-tration of substances that react with ozoneto break it down can increase the rate ofozone destruction. If the rate of ozone for-mation stays constant, then the total amountof ozone will fall until the rate of destructionfalls and a new equilibrium is established ata lower ozone concentration.

This is exactly what was observed inthe 1970s. Chemists noticed that the con-centrations of ozone were decreasing—wellbeyond what could be explained by naturalprocesses. It was soon discovered that theconcentration of ozone-depleting substanceswas increasing, mostly due to human use ofgases for refrigeration and spray can prod-ucts. Fortunately, worldwide efforts toreduce production and use of these damag-ing substances seem to be helping, andrecent measurements indicate that the ozoneconcentrations may be stabilizing.

At the same time, scientists haveobserved a cooling in the stratosphere. Thismay be, in part, due to an increase in theinsulating layer of greenhouse gases (H2O,CO2, CH4, O3, N2O, and others). Theirincrease in concentration means less solarenergy is radiated back into the upper atmos-phere as heat. When temperatures dropbelow –88 °C, thin clouds form. The presenceof stratospheric clouds in the polar regions

appears to increase the rate of ozone destruc-tion. The surface of ice crystals in the strato-spheric polar clouds can accelerate reactionsbetween O3 and substances that destroyozone. The overall effect is a shift in the equi-librium and a decrease in the ozone levels atthe Earth’s poles.

On a global scale, equilibrium gets verycomplicated. System after system, changesoccur in intricate and sometimes unexpectedways, both in response to natural conditionsand also because of human action. Aura andthe other scientific satellites in orbit are gather-ing data that will play a crucial role in develop-ing an understanding of our atmosphere. Ashard as it is, understanding these equilibria andhow one thing leads to another is a vital step indeciding strategies for preserving the lives andwell-being of all who call this planet home.

Michael Tinnesand is associate director forAcademic Programs at ACS and formerly taughtchemistry in Hillsboro, OR.

ChemMatters, SEPTEMBER 2005 13

Stra

tosp

here

Trop

osph

ere

Polar ice clouds

O2

O2

O3

O3

Catalysts: -hydroxyl radical -bromine -chlorine -nitrogen oxides

Combustion

Oxidationof

organicsubstances

NOx

Hydrocarbons

Sunlight

CFCs

Volcanos

SolarUltraviolet

Light

SolarUltraviolet

Light

SolarUltraviolet

Light

UltravioletLight

Lightning

10%of ozoneontroposphere

90%of ozoneonstratosphere

10%of ozoneontroposphere

90%of ozoneonstratosphere

C=CC=C

The concentration of various substances in the atmosphere depends on complex and interrelated reactions. In the stratosphere, ultraviolet light creates ozone.Catalysts such as bromine and chlorine break ozone down to form oxygen. If the rate of ozone formation is the same as the rate of ozone destruction, the amount ofozone in the stratosphere will remain constant. In the troposphere we have a slightly different story. Most of the ozone is caused by combustion, either from humanactivities or natural sources. The reaction products react in the sunlight to form ozone. Ground level ozone is destroyed when it reacts with substances such as organicmaterial, or when it is exposed to UV light.

CE

SA

R C

AM

INE

RO