24 chemical reactions - irion county isd / overvie · 24 chemical reactions o hhhh! ahhhh! ......

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Chemical Reactions2424

Ohhhh! Ahhhh! The crowdgasps as fireworks explodeoverhead into patterns of

brightly colored sparks. Deafeningbooms split the air and childrensqueal with delight. All these effectsare produced by chemical reactionswith oxygen. The colors come fromsmall amounts of metal ions—bluefrom copper, red from lithium, goldfrom sodium, and green from barium.In this chapter, you will learn aboutchemical reactions—what they are,how to describe them, and how theyrelease or absorb energy.

What do you think?Science Journal Look at the picturebelow with a classmate. Discuss whatyou think this might be or what ishappening. Here’s a hint: These bubblesare lighter than air. Write down yourbest guess in your Science Journal.

Like exploding fireworks, rusting is a chemical reac-tion in which iron metal combines with oxygen.

Other metals combine with oxygen, too—somemore readily than others. In this activity, you willcompare how iron and aluminum react with oxygen.

Safety Precautions

Observe which metal reacts1. Place a clean, iron or steel nail in a

dish prepared by your teacher.

2. Place a clean, aluminum nail in asecond dish. These dishes contain agargel and an indicator that detects areaction with oxygen.

3. Observe both nails after one hour.Record any changes around the nailsin your Science Journal.

4. Carefully examine both of the dishes the next day. Record any differencesbetween the two dishes in your Science Journal.

Observe Did a reaction occur in either dish? How could you tell? What might havecaused any differences you observed between the two nails?

EXPLOREACTIVITY

737

Making a Cause and Effect Study Fold Make the following Foldable to help you understand the cause and effect relationshipof chemical reactions.

1. Place a sheet of paper in front of you so the short side is at thetop. Fold the paper in half from the left side to the right side.

2. Now fold the paper in half from top to bottom. Then fold it inhalf again top to bottom. Unfold the last two folds you did.

3. Through one thickness of paper, cut along each of the foldlines to form four tabs as shown.

4. Label the four sections “R” for reactants, “0” for produce,and “P” for products as shown. As you read the chapter, record examples of chemical reactions under the tabs.

FOLDABLESReading & StudySkills

FOLDABLESReading & Study Skills

RP

RP

RP

RP

Chemical Changes

738 CHAPTER 24 Chemical Reactions

S E C T I O N

Describing Chemical Reactions Dark mysterious mixtures react, gas bubbles up and

expands, and powerful aromas waft through the air. Where areyou? Are you in a chemical laboratory carrying out a crucialexperiment? No. You are in the kitchen baking a chocolate cake.Nowhere in the house do so many chemical reactions take placeas in the kitchen.

Actually, chemical reactions are taking place all around youand even within you. A chemical reaction is a change in whichone or more substances are converted into new substances. Thesubstances that react are called reactants. The new substancesproduced are called products. This relationship can be writtenas follows:

producereactants 0 products

Conservation of Mass By the 1770s, chemistry was changing from the art of

alchemy to a true science. Chemical reactions were studied care-fully. Through such study, the French chemist Antoine Lavoisierestablished that the mass of the products always equals the massof the reactants. This principle is demonstrated in Figure 1.

� Identify the reactants andproducts in a chemical reaction.

� Determine how a chemicalreaction satisfies the law of conservation of mass.

� Determine how chemists expresschemical changes using equations.

Vocabularychemical reaction reactantchemical equation productcoefficient

Chemical reactions cook our food,warm our homes, and provideenergy to our bodies.

Beforeburning

Afterburning

Figure 1The mass of the candles andoxygen before burning is exactlyequal to the mass of the remain-ing candle and gaseous products.

Lavoisier’s Contribution In one experiment, Lavoisier placed a carefully measured

mass of solid mercury(II) oxide, which he knew as mercury calx,into a sealed container. When he heated this container, he noteda dramatic change. The red powder had been transformed into asilvery liquid that he recognized as mercury metal, and a gas wasproduced. When he determined the mass of the liquid mercuryand gas, their combined masses were exactly the same as themass of the red powder he had started with.

mercury(II) oxide oxygen plus mercury10.0 g � 0.7 g � 9.3 g

Lavoisier also established that the gas produced by heatingmercury oxide, which we call oxygen, was a component of air. Hedid this by heating mercury metal with air and saw that a portionof the air combined to give red mercury oxide. He studied theeffect of this gas on living animals, including himself. Hundreds ofexperiments carried out in his laboratory, shown in Figure 2, con-firmed that in a chemical reaction, matter is not created ordestroyed, but is conserved. This principle became known as thelaw of conservation of mass. This means that the starting mass ofthe reactants equals the final mass of the products.

What does the law of conservation of mass state?

This principle is the basis for modern chemistry. It made itpossible for chemists to see clearly what happens in chemicalreactions. More importantly, the concept of balance betweenreactants and products led to a valuable tool of chemistry—chemical equations.

SECTION 1 Chemical Changes 739

Research Visit the Glencoe Science Web site atscience.glencoe.com formore information aboutAntoine Lavoisier and hiscontributions to chemistry.Communicate to your classwhat you learn.

Figure 2Antoine Lavoisier’s wife,Marie-Anne, drew thisview of Lavoisier in hislaboratory performingstudies on oxygen. Shedepicted herself at theright taking notes.

http://www.science.glencoe.com

Writing Equations If you wanted to describe the chemical reaction

shown in Figure 3, you might write something like this:

nickel(II) chloride, dissolved in water, plus sodium hydroxide, dissolved in water, produces solid nickel(II) hydroxide plus sodium chloride,

dissolved in water

This series of words is rather cumbersome, but all ofthe information is important. The same is true ofdescriptions of most chemical reactions. Many wordsare needed to state all the important information. As aresult, scientists have developed a shorthand method todescribe chemical reactions. A chemical equation is away to describe a chemical reaction using chemical for-mulas and other symbols. Some of the symbols used inchemical equations are listed in Table 1.

The chemical equation for the reaction describedabove in words and shown in Figure 3 looks like this:

NiCl2(aq) � 2NaOH(aq) 0 Ni(OH)2(s) � 2NaCl(aq)

It is much easier to tell what is happening by writ-ing the information in this form. Later, you will learnhow chemical equations make it easier to calculate thequantities of reactants that are needed and the quanti-ties of products that are formed.


Figure 3A white precipitate of nickel(II) hydroxideforms when sodium hydroxide is added to a green solution of nickel(II) chloride.Sodium chloride, the other product formed, is in the solution.

Table 1 Symbols Used in Chemical Equations

Symbol Meaning

0 produces or forms

� plus

(s) solid

(l) liquid

(g) gas

(aq) aqueous, a substance is dissolved in water

heat the reactants are heated 0

light the reactants are 0 exposed to light

elec. an electric current is0 applied to the reactants

Unit ManagersWhat do the numbers to the left of the formulas for reac-

tants and products mean? Remember that according to the lawof conservation of mass, matter is neither made nor lost duringchemical reactions. Atoms are rearranged but never lost ordestroyed. These numbers, called coefficients, represent thenumber of units of each substance taking part in a reaction.Coefficients can be thought of as unit managers.

What is the function of coefficients in a chemicalequation?

Imagine that you are responsible for making sandwiches fora picnic. You have been told to make a certain number of threekinds of sandwiches, and that no substitutions can be made.You would have to figure out exactly how much food to buy sothat you had enough without any food left over. You might needtwo loaves of bread, four packages of turkey, four packages ofcheese, two heads of lettuce, and ten tomatoes. With these sup-plies you could make exactly the right number of each kind ofsandwich.

In a way, your sandwich-making effort is like a chemicalreaction. The reactants are your bread, turkey, cheese, lettuce,and tomatoes. The number of units of each ingredient are likethe coefficients of the reactants in an equation. The sandwichesare like the products, and the numbers of each kind of sandwichare like coefficients, also.

Knowing the number of units of reactants enables chemiststo add the correct amounts of reactants to a reaction. Also, theseunits, or coefficients, tell them exactly how much product willform. An example of this is the reaction of one unit of NiCl2with two units of NaOH to produce one unit of Ni(OH)2 andtwo units of NaCl. You can see these units in Figure 4.

SECTION 1 Chemical Changes 741

Designing a TeamEquationProcedure1. Obtain 15 index cards

and mark each as follows:five with guard, five withforward, and five with center.

2. Group the cards to form asmany complete basketballteams as possible. Eachteam needs two guards, twoforwards, and one center.

Analysis1. Write the formula for a

team. Write the formationof a team as an equation.Use coefficients in front ofeach type of player neededfor a team.

2. How is this equationlike a chemical equation?Why can’t you use theremaining cards?

Ni(OH)22NaOH� 0

� 0

NiCl2

Ni2�

OH�Na�

Na�

Cl�

Cl� OH�

OH�

OH�

Ni2�

2NaCl

Na�

Na�

Cl�

Cl�

�

�

Figure 4Each coefficient in the equationrepresents the number of unitsof each type in this reaction.


Metals and the AtmosphereWhen iron is exposed to air andmoisture, it corrodes or rusts, form-ing hydrated iron(III) oxide. Rust can

seriously damage iron structures because it crumbles and exposesmore iron to the air. This leads to more oxidation and eventuallycan destroy the structure. However, not all reactions of metalswith the atmosphere are damaging like rust. Some are helpful.

Aluminum also reacts with oxygen in the air to form alu-minum oxide. Unlike rust, aluminum oxide adheres to the alu-minum surface, forming an extremely thin layer that protectsthe aluminum from further attack. You can see this thin layer ofaluminum oxide on aluminum outdoor furniture. It makes theonce shiny aluminum look dull.

Copper is another metal that corrodes when itis exposed to air forming a blue-green coatingcalled a patina. You can see this type of corrosionon many public monuments and also on theStatue of Liberty, shown in Figure 5.

EnvironmentalScience

Figure 5The blue-green patina that coats the Statue ofLiberty contains copper(II) sulfate amongother copper corrosion products.

6. Recognizing Cause and Effect Lavoisierheated mercury(II) oxide in a sealed flask.Explain the effect on Lavoisier’s conclusionsabout the law of conservation of mass if he hadused an open container. For more help, refer tothe Science Skill Handbook.

7. Solving One-Step Equations When makingsoap, if 890 g of a specific fat react completelywith 120 g of sodium hydroxide, the productsformed are soap and 92 g of glycerin. Calculatethe mass of soap formed to satisfy the law of con-servation of mass. For more help, refer to theMath Skill Handbook.

Section Assessment

1. Identify the reactants and the products inthe following chemical equation.

Cd(NO3)2(aq) � H2S(g) 0 CdS(s) � 2HNO3(aq)

2. What is the state of matter of each sub-stance in the following reaction?

Zn(s) � 2HCI(aq) 0 H2(g) � ZnCl2(aq)

3. Why is the reaction of oxygen with iron aproblem, but the reaction of oxygen withaluminum is not?

4. What is the importance of the law of con-servation of mass?

5. Think Critically Why do you think thecopper patina was kept when the Statue ofLiberty was restored?

SECTION 2 Chemical Equations 743

Balanced Equations Lavoisier heated mercury(II) oxide—HgO—and obtained

mercury and oxygen. This reaction is shown in Figure 6. Whenwritten as a chemical equation, the reaction is:

heatHgO(s) 0 Hg(l) � O2 (g)

Notice that the number of mercury atoms is the same onboth sides of the equation but that the number of oxygen atomsis not the same. One oxygen atom appears on the reactant sideof the equation and two appear on the product side

Atoms HgO 0 Hg � O2Hg 1 1O 1 2

Now wait a minute, you say. The law of conservation of massstates that atoms are neither created nor destroyed in a chemicalreaction. One oxygen atom can’t just become two. Are the for-mulas wrong? Can you fix this problem simply by adding thesubscript 2 and writing HgO2 instead of HgO?

Think for a moment about water, H2O, and hydrogen perox-ide, H2O2. Although their formulas are similar, they are very dif-ferent compounds. Similarly, the formulas HgO2 and HgO donot represent the same compound. In fact, HgO2 does not exist.In Lavoisier’s experiment, he heated HgO, not HgO2. The formu-las in a chemical equation must accurately represent the com-pounds that react. They can’t be changed to something different.

Why can’t you change the subscripts in a chemical equation?

If you can’t change the subscripts in an equation, how can youfix Lavoisier’s equation so that the number of oxygen atoms is thesame on both sides? Fixing this equation requires a process calledbalancing. Balancing an equation doesn’t change what happens ina reaction—it simply changes the way the reaction is represented.The balancing process involves changing coefficients in a reactionto achieve a balanced chemical equation, which has the samenumber of atoms of each element on both sides of the equation.

� Identify what is meant by abalanced chemical equation.

� Determine how to write balancedchemical equations.

Vocabularybalanced chemical equation

Chemical equations are the languageused to describe chemical change.

Chemical Equations S E C T I O N

Figure 6Mercury metal forms when mercury oxide is heated. Becausemercury is poisonous, thisreaction is never performed in aclassroom laboratory.

Mercury(II) oxide

Mercury metal

Choosing Coefficients Finding out which coefficients touse to balance an equation is often a trial-and-error process, butwith practice the process becomes easier. In the equation forLavoisier’s experiment, the number of mercury atoms is bal-anced, but one oxygen atom is on the left and two are on theright. Therefore, this equation isn’t balanced. If you put a coeffi-cient of 2 before the mercury(II) oxide on the left, the oxygenatoms will be balanced. However, you have two mercury atomson the left and only one on the right. To correct this, put a 2 infront of mercury on the right. The equation is now balanced.

Atoms 2HgO 0 2Hg � O2Hg 2 2O 2 2

Try Your Balancing Act Magnesium burns with such abrilliant, white light that it is often used in emergency flares asshown in Figure 7. Burning leaves a white powder called magne-sium oxide. To write a balanced chemical equation for this andmost other reactions, follow these four steps.

Step 1 Write a chemical equation for the reaction usingformulas and symbols. Review how to write formulas for com-pounds. Recall that oxygen gas is a diatomic molecule.

Mg(s) � O2(g) 0 MgO(s)

Step 2 Check the equation for atom balance.

Atoms Mg � O2 0 MgOMg 1 1O 2 1

The magnesium atoms are balanced, but the oxygen atomsare not. Therefore, this equation isn’t balanced.

Step 3 Choose coefficients that balance the equation.Remember, never change subscripts of a correct formula to bal-ance an equation. Try putting a coefficient of 2 before MgO.

Mg(s) � O2(g) 0 2MgO(s)

Step 4 Recheck the numbers of each atom on each side ofthe equation and adjust coefficients again if necessary.

Now two Mg atoms are on the right side and only one is onthe left side. So a coefficient of 2 is needed for Mg also.

2Mg(s) � O2(g) 0 2MgO(s)

The equation is now balanced.


Figure 7 Magnesium combines withoxygen giving an intensewhite light.

Research Visit theGlencoe Science Web site at science.glencoe.com formore information on balanc-ing chemical equations. Com-municate to your class whatyou learned.


Polish Your Skill When lithium metal is treated with water,hydrogen gas and lithium hydroxide are produced, as shown inFigure 8.

Step 1 Write the chemical equation.

Li(s) � H2O 0 H2(g) � LiOH(aq)

Step 2 Check for balance by counting the atoms.

Atoms Li � H2O 0 LiOH � H2Li 1 1H 2 1 2O 1 1

This equation is not balanced. There are three hydrogenatoms on the right and only two on the left. Complete steps 3and 4 to balance the equation. After each step, count the atomsof each element. When equal numbers of atoms of each elementare on both sides, the equation is balanced.

Just as in the first example, this balancing does not changethe reaction that took place between lithium and water. It juststates what happened more accurately. This accurate statementtells chemists how much lithium metal to use to produce a certain amount of hydrogen gas.

Section Assessment

1. Give two reasons for balancing equationsfor chemical reactions.

2. Write a balanced chemical equation foreach of the following reactions.a. Iron metal plus oxygen produces

iron(II) oxide.b. Sodium metal plus water produces

sodium hydroxide plus hydrogen gas.

3. Explain why oxygen gas must always bewritten as O2 in a chemical equation.

4. What is understood if no coefficient is writ-ten before a formula in a chemical equation?

5. Think Critically Explain why the sum ofthe coefficients on the reactant side of abalanced equation does not have to equalthe sum of the coefficients on the productside of the equation.

6. Predicting Silver tarnish, Ag2S, forms from sulfur compounds in the air and in some foods, such as eggs. Polishing removes this tarnish. Why might people not wish to polish their silverware very frequently? Howmight they prevent tarnish from forming and thereby reduce how often they need to polish? For more help, refer to the Science Skill Handbook.

7. Using an Electronic Spreadsheet Use aspreadsheet and design a table for the reac-tion of CaCO3 with HCl to form CO2, CaCl2, andwater. Remember that equal numbers ofeach element must be on both sides of a bal-anced equation. For more help, refer to theTechnology Skill Handbook.

Figure 8When lithium metal is added towater, it reacts, producing asolution of lithium hydroxide andbubbles of hydrogen gas.

SECTION 2 Chemical Equations 745

Classifying Chemical Reactions

S E C T I O N

Types of ReactionsYou might have begun to notice that there are all sorts of

chemical reactions. You could just memorize them, but thatwould be difficult. Fortunately there is a better way.

Imagine if the public library shelved its books without anyorder. Every time you wanted a particular book you would haveto search through all the books or remember where you hadseen it last. Fortunately libraries classify books into groups, suchas biography, history, and fiction. This way you know exactlywhere to look.

You can do the same thing with chemical reactions. Also,once a chemical reaction is placed into a certain type, you canlearn a great deal by comparing it to others of that type. Onesystem is based upon the way the compounds interact. Mostreactions can be divided into four main types—synthesis,decomposition, single displacement, and double displacement.

Synthesis Reactions One of the easiest reaction types torecognize is a synthesis reaction. In a synthesis reaction, two ormore substances combine to form another substance. The gen-eralized formula for this reaction type is as follows.

A � B 0 AB

The reaction in which hydro-gen burns in oxygen to formwater is an example of a synthesisreaction.

2H2(g) � O2(g) 0 2H2O(g)

This reaction is used to powersome types of rockets. Anothersynthesis reaction is the combina-tion of oxygen with iron in thepresence of water to form hy-drated iron(II) oxide or rust. Thisreaction is shown in Figure 9.

� Identify the four general types ofchemical reactions.

� Predict which metals will replaceother metals in compounds.

Vocabularysynthesis reactiondecomposition reactionsingle-displacement reactiondouble-displacement reactionprecipitate

Classifying reactions helps youunderstand what is happening andpredict the outcome of reactions.

Figure 9 Rust has accumulated on theTitanic since it sank in 1912.

746 CHAPTER 24

Decomposition Reactions A decomposition reaction isjust the opposite of a synthesis. Instead of two substances com-ing together to form a third, a decomposition reaction occurswhen one substance breaks down, or decomposes, into two ormore substances. The general formula for this type of reactioncan be expressed as follows:

AB 0 A � B

Most decomposition reactions require the use of heat, light,or electricity. For example, an electric current passed throughwater produces hydrogen and oxygen as shown in Figure 10.

elec.2H2O(l) -02H2(g) � O2(g)

Single Displacement When one element replaces anotherelement in a compound, it is called a single-displacementreaction. There are two generalized types of this reaction.

In the first case, A replaces B as follows:

A � BC 0 AC � B

In the second case, D replaces C:

D � BC 0 BD � C

The first case is illustrated in Figure 11, where a copper wire isput into a solution of silver nitrate. Because copper is a moreactive metal than silver, it replaces the silver, forming a blue cop-per(II) nitrate solution. The silver, which is not soluble, formson the wire.

Cu(s) � 2AgNO3(aq) 0 Cu(NO3)2 (aq) � 2Ag(s)

Describe a single-displacement reaction.

Sometimes single-displace-ment reactions can causeproblems. For example, if iron-containing vegetables such asspinach are cooked in alu-minum pans, aluminum candisplace iron from the veg-etable. This causes a blackdeposit of iron to form on thesides of the pan. For this rea-son, it is better to use stainlesssteel or enamel cookware whencooking spinach.

SECTION 3 Classifying Chemical Reactions 747

Figure 10Water decomposes into hydrogenand oxygen when an electric cur-rent is passed through it. A smallamount of sulfuric acid is addedto increase conductivity. Noticethe proportions of the gasescollected. How is this related tothe coefficients of the products inthe equation?

Figure 11Copper in a wire replaces silver in silvernitrate, forming a blue solution ofcopper(II) nitrate.

O2 H2

The Activity Series We can predict which metalwill replace another using the diagram shown in Figure 12, which lists metals according to how activethey are. A metal will replace any less active metal.Notice that copper, silver, and gold are the least activemetals on the list. That is why these elements oftenoccur as deposits of the relatively pure element. Forexample, gold is sometimes found as veins in quartzrock, and copper is found in pure lumps known asnative copper. Other metals occur as compounds.

LithiumPotassiumCalciumSodiumAluminumZincIronTinLead(Hydrogen)CopperSilverGold

MOST ACTIVE

LEAST ACTIVE

Figure 12This figure shows the activity series of metals. Ametal will replace any other metal that is less active.

Math Skills Activity

Example ProblemA sample of barium sulfate is placed on a piece of paper, which is then ignited.

Barium sulfate reacts with the carbon from the burned paper producing barium sulfideand carbon monoxide. Write a balanced chemical equation for this reaction.

Solution

Write a chemical equation using formulas and symbols:BaSO4(s) � C(s) 0 BaS(s) � CO(g)

Check the equation for atom balance by setting up a table:Kind of Number of Atoms Number of Atoms Atom Before Reaction After Reaction

Ba 1 1S 1 1O 4 1C 1 1

The oxygen atoms are not equal—this equation is not balanced.

Choose coefficients that balance the equation. Try putting a 4 in front of CO. Now you have 4 oxygens, but 4 carbons also, so add another 4 in front of the C in the reactants to balance the remaining elements.

BaSO4(s) � 4C(s) 0 BaS(s) � 4CO(g)

Using Coefficients to Balance Chemical Equations


Practice Problem

1. HCl is slowly added to aqueous Na2CO3 forming NaCl, H2O, and CO2. Follow thesteps above to write a balanced equation for this reaction.

For more help with solving equations, refer to the Math Skill Handbook.

Double Displacement In a double-displacementreaction, the positive ion of one compound replaces thepositive ion of the other to form two new compounds. Adouble-displacement reaction takes place if a precipitate,water, or a gas forms when two ionic compounds insolution are combined. A precipitate is an insolublecompound that comes out of solution during this type of reaction. The generalized formula for this type ofreaction is as follows.

AB � CD → AD � CB

The reaction of barium nitrate with potassium sulfateis an example of this type of reaction. A precipitate—barium sulfate—forms, as shown in Figure 13. The chemicalequation is as follows:

Ba(NO3)2(aq) � K2SO4(aq) 0 BaSO4(s) � 2KNO3(aq)

These are a few examples of chemical reactions classifiedinto types. Many more reactions of each type occur around you.

What type of reaction produces a precipitate?

SECTION 3 Classifying Chemical Reactions 749

Section Assessment

1. Classify each of the following reactions:a. CaO(s) � H2O → Ca(OH)2 (aq)b. Fe(s) � CuSO4(aq) → FeSO4(aq) � Cu(s)c. NH4NO3(s) → N2O(g) � 2H2O(g)

2. Sulfur trioxide, (SO3), a pollutant releasedby coal-burning plants, can react withwater in the atmosphere to produce sulfu-ric acid, H2SO4. Write a balanced equationfor this reaction.

3. Explain the difference between synthesisand decomposition reactions.

4. Why can’t you change the subscripts ofreactants and products when balancing achemical equation?

5. Think Critically Balance the followingequation and identify the reaction type.

Au(CN)2- (aq) � Zn(s) → Au(s) � Zn(CN)4

2-(aq)

6. Forming Hypotheses Group 1 metals replacehydrogen in water in reactions that are oftenviolent. This sample equation shows what hap-pens when potassium reacts with water.

2K(s) � 2HOH(l) → 2KOH(aq) � H2(g)

Use Figure 12 to help you hypothesize whythese metals replace hydrogen in water. Formore help, refer to the Science Skill Handbook.

7. Using Proportions The following equation,showing the formation of iron oxide, is balanced, but the coefficients used are largerthan necessary. Rewrite this equation using the smallest coefficients that give a balancedequation. For more help, refer to the MathSkill Handbook.

9Fe(s) � 12H2O(g) → 3Fe3O4(s) � 12H2(g)

Figure 13A precipitate of barium sulfatesettles to the bottom of a testtube containing potassiumnitrate. Because it is opaque toX rays, barium sulfate is used totake X-ray photographs of theintestinal tract.


Chemical Reactions and Energy

S E C T I O N

Chemical Reactions—Energy ExchangesOften a crowd gathers to watch a building being demolished

using dynamite. In a few breathtaking seconds, tremendous struc-tures of steel and cement that took a year or more to build arereduced to rubble and a large cloud of dust. A dynamite explosion,as shown in Figure 14, is an example of a rapid chemical reaction.

Most chemical reactions proceed more slowly, but all chemi-cal reactions release or absorb energy. This energy can take manyforms, such as heat, light, sound and electricity. The heat pro-duced by a wood fire and the light emitted by a candle are twoexamples of reactions that release energy.

Chemical bonds are the source of this energy. When mostchemical reactions take place, some chemical bonds in the reac-tants must be broken, and breaking these bonds takes energy. Inorder for products to be produced, new bonds must form. Bondformation releases energy. Reactions such as dynamite combus-tion require much less energy to break chemical bonds than theenergy released when new bonds are formed. The result is arelease of energy and sometimes a loud explosion. Anotherrelease of energy is used to power rockets, as shown in Figure 15.

� Identify the source of energychanges in chemical reactions.

� Compare and contrast exergonicand endergonic reactions.

� Examine the effects of catalystsand inhibitors on the speed ofchemical reactions.

Vocabularyexergonic reactionexothermic reactionendergonic reactionendothermic reactioncatalystinhibitor

Chemical reactions provide energy tocook your food, keep you warm, andtransform the food you eat into sub-stances you need to live and grow.

Figure 14When its usefulness is over, abuilding is sometimes demol-ished using dynamite. Dynamitecharges must be placed carefullyso that the building collapsesinward, where it cannot harmpeople or property.

751

Figure 15

Rockets burn fuel to provide the thrust necessary to propel them upward. In 1926,engineer Robert Goddard used gasoline and

liquid oxygen to propel the first ever liquid-fueledrocket. Although many people at the time ridiculedGoddard’s space travel theories, his rockets eventu-ally served as models for those that have gone to theMoon and beyond. A selection of rockets—includingGoddard’s—is shown here. The number below eachcraft indicates the amount of thrust—expressed innewtons (N)—produced during launch.

SPACE SHUTTLE Themain engines produceenormous amounts ofenergy by combiningliquid hydrogen andoxygen. Coupled withsolid rocket boosters,they produce over 32.5million newtons (N) of thrust to lift the sys-tem’s 2 million kg off the ground.

LUNAR MODULE Smaller rocketengines, like those used by theLunar Module to leave the Moon,use hydrazine-peroxide fuels. Thenumber shown below indicates the fixed thrust from one of themodule’s two engines; the otherengine’s thrust was adjustable.

GODDARD’S MODELROCKET Although his first rocket rose only 12.6 m, Goddard success-fully launched 35 rocketsin his lifetime. The high-est reached an altitude of 2.7 km.

JUPITER C This rocketlaunched the first UnitedStates satellite in 1958. Itused a fuel called hydyneplus liquid oxygen.

369,350 N 15,920 N32,500,000 N400 N

VISUALIZING CHEMICAL ENERGY

�

�

�

�

More Energy Out You have probably seen many reactions that release energy.

Chemical reactions that release energy are called exergonic (ek sur GAH nihk) reactions. In these reactions less energy isrequired to break the original bonds than is released when newbonds form. As a result, some form of energy, such as light orheat, is given off by the reaction. The familiar glow from thereaction inside a glow stick, shown in Figure 16, is an exampleof an exergonic reaction, which produces visible light. In otherreactions however, the energy given off can produce heat. This isthe case with some heat packs that are used to treat muscleaches and other problems that require heat.

When the energy given off in a reaction is primarily in theform of heat, the reaction is called an exothermic reaction. Theburning of wood and the explosion of dynamite are exothermicreactions. Iron rusting is also exothermic, but the reaction pro-ceeds so slowly that it’s difficult to detect any temperature change.

Why is a log fire considered to be an exothermic reaction?

Exothermic reactions provide most of the power used inhomes and industries. Fossil fuels that contain carbon, such ascoal, petroleum, and natural gas, combine with oxygen to yieldcarbon dioxide gas and energy. Unfortunately impurities inthese fuels, such as sulfur, burn as well, producing pollutantssuch as sulfur dioxide. Sulfur dioxide combines with water inthe atmosphere, producing acid rain.


Figure 16Glow sticks contain three differ-ent chemicals—an ester and adye in the outer section andhydrogen peroxide in a centerglass tube. Bending the stickbreaks the tube and mixes thethree components.

Colorful ChemicalReactionProcedure1. Pour 5 mL of water into a

test tube.2. Sprinkle a few crystals of

copper(II) bromide intothe test tube and observethe color change of the crystals.

3. Slowly add more water andobserve what happens.

Analysis1. What color were the cop-

per(II) bromide crystalsafter you added them to thetest tube of water?

2. What color were they whenyou added more water?

3. What caused this colorchange?

Hydrogenperoxide

Solution ofdye and ester

More Energy InSometimes a chemical reaction requires more energy to break

bonds than is released when new ones are formed. These reac-tions are called endergonic reactions. The energy absorbed canbe in the form of light, heat, or electricity.

Electricity is often used to supply energy to endergonic reac-tions. For example, electroplating deposits a coating of metalonto a surface, as shown in Figure 17. Also, aluminum metal isobtained from its ore using the following endergonic reaction.

elec.2Al2O3(l) -04Al(l) � 3O2(g)

In this case, electrical energy provides the energy needed to keepthe reaction going.

Heat Absorption When the energy needed is in the form ofheat, the reaction is called an endothermic reaction. The termendothermic is not just related to chemical reactions. It also candescribe physical changes. The process of dissolving a salt inwater is a physical change. If you ever had to soak a swollenankle in an Epsom salt solution, you probably noticed that whenyou mixed the Epsom salt in water, the solution became cold.The dissolving of Epsom salt absorbs heat. Thus it is a physicalchange that is endothermic.

Cold packs also use endothermic processes, asshown in Figure 18. They contain ammoniumnitrate crystals and a water container. When thewater container is ruptured, water mixes with theammonium nitrate. Almost immediately the packcools dramatically.

SECTION 4 Chemical Reactions and Energy 753

Figure 18Cold packs are often used insports. Because they cool soquickly, it’s like carrying aninstant bag of ice. They are handyfor treating athletic injuries onthe field. They start to workonly when you press the pack,breaking a container of waterthat mixes with ammoniumnitrate crystals in a process thatabsorbs heat.

Figure 17Electroplating of a metal is an endergonic reaction thatrequires electricity. A coating of copper was plated onto thiscoin.

Cold pack

How to use the Instant Cold Pack

170 ml

1. SQUEEZE. Squeeze pack inthe palm of your hand until

you hear a pop andpack begins to feel cold.

2. SHAKE. Hold pack at eitherend to mix the contents.

4. IMPORTANT. Do not attempt to refreeze. Dispose of

3. APPLY. Apply chilled packimmediately to the affected

area for fifteen minute intervals.

after use.

Catalysts and Inhibitors Some reactions proceed tooslowly to be useful. To speed them up, a catalyst can be added. Acatalyst is a substance that speeds up a chemical reaction with-out being permanently changed itself. When you add a catalyst toa reaction, the mass of the product that is formed remains thesame, but it will form more rapidly. The catalyst remainsunchanged and often is recovered and reused. Catalysts are usedto speed many reactions in industry, such as polymerization tomake plastics and fibers.

Why would a catalyst be needed for a chemicalreaction?

At times, it is worthwhile to prevent certain reactions fromoccurring. Substances called inhibitors are used to combinewith one of the reactants. This ties up the reactant and preventsit from undergoing the original reaction. The food preservativesBHT and BHA are inhibitors that prevent spoilage of certainfoods, such as cereals and crackers.

Sometimes you might want to inhibit a catalyst. For exam-ple, the cut surfaces of fruits darken rapidly due to an enzyme-catalyzed reaction with air. Brushing or spraying the surfaceswith lemon juice prevents the catalyst from acting.

One thing to remember when thinking about catalysts andinhibitors is that they do not change the amount of product pro-duced. They only change the rate of production. Catalystsincrease the rate and inhibitors decrease the rate. Therefore, foodscontaining preservatives will spoil eventually, but preservatives areinhibitors that often add months or even years to the shelf time.


Section Assessment

1. Discuss the difference between exergonicand endergonic reactions.

2. What happens to a catalyst in a reaction?

3. Crackers containing BHT stay fresh longerthan those without it. Explain why.

4. Classify the reaction that makes a fireflyglow in terms of energy input or output.

5. Think Critically To develop a productthat warms people’s hands, would youchoose an exothermic or endothermic reaction to use? Why?

6. Concept Mapping Construct a concept map to show the relationship between energy and bond formation and bond breakage.For more help, refer to the Science SkillHandbook.

7. Communicating In your Science Journal,compare the profit or loss in an investment or abusiness deal with the gain or loss of energy ina chemical reaction. For more help, refer tothe Science Skill Handbook.

Metals, such as platinumand palladium, are used ascatalysts in the exhaustsystems of automobiles.What reactions do youthink they catalyze?

ACTIVITY 755

Catalyzed Reaction

4. Place all three tubes in a beaker of hot water.Heat on a hot plate until all of the remainingH2O2 is driven away and no liquid remains.

Conclude and Apply1. What changes did you observe when the

solids were added to the tubes?

2. In which tube was oxygen produced rapidly,and how do you know?

3. Which substance, sand or MnO2, caused therapid production of gas from the H2O2?

4. What remained in each tube after the H2O2was driven away?

Abalanced chemical equation tells nothingabout the rate of a reaction. One way to

affect the rate is to use a catalyst.

What You’ll InvestigateHow does the presence of a catalyst affect therate of a chemical reaction?

Materialstest tubes (3) sand (1⁄4 teaspoon)test-tube stand hot plate3% hydrogen peroxide, wooden splint

H2O2 (15 mL) beaker of hot water10-mL graduated cylinder manganese dioxide,small plastic teaspoon MnO2 (1⁄4 teaspoon)

Goals� Observe a catalyst on the rate of reaction.� Conclude based on your observations

whether the catalyst remained unchanged.

Safety Precautions

WARNING: Hydrogen peroxide can irritate skinand eyes. Wipe up spills promptly. Point test tubesaway from other students.

Procedure1. Set three test tubes in a test-tube stand. Pour

5 mL of hydrogen peroxide into each tube.

2. Place about 1⁄4 teaspoon of sand in tube 2 andthe same amount of MnO2 in tube 3.

3. In the presence of a catalyst, H2O2 decom-poses rapidly producing oxygen gas, O2. Testthe gas produced as follows: Light a woodensplint, blow out the flame, and insert theglowing splint into the tube. The splint willrelight if oxygen is present.

Compare your results with those of your class-mates and discuss any differences observed.For more help refer to the Science SkillHandbook.

You’ve probably heard a lot about global warming and the greenhouse effect.According to one theory, certain gases in the atmosphere might be causing

Earth’s average global temperature to rise. The gases carbon dioxide, nitrous oxide,and methane, known as greenhouse gases, result from chemical reactions with oxy-gen when fossil fuels, such as coal, oil, and gas, are burned. What are some everydayactivities that you do that might produce greenhouse gases?

Recognize the ProblemWhat do you do to produce greenhouse gases?

Form a HypothesisThink of the different ways you use fossil fuels every day. Why are fossil fuels important?Form a hypothesis about how certain activities add greenhouse gases to our atmosphere.

Data SourceGo to the Glencoe

Science Web site at science.glencoe.com for more information about fossilfuels, the chemical reactions that pro-duce greenhouse gases, uses of fossilfuels, their effects on the environment,and data from other students.

Fossil Fuels and Greenhouse Gases


Goals� Observe how you use fossil fuels in

your daily life.

� Gather data on the processof burning fossil fuels and how greenhouse gases are released.

� Research the chemical reactions that produce green-house gases.

� Identify theimportance of fossil fuels and their effect on theenvironment.

� Communicateyour findings to other students.


Test Your Hypothesis

Analyze Your Data

Draw Conclusions

2. Research the different greenhousegases and the chemical reactionsthat produce them.

3. Compare the different reactionsand their products.

4. Record your data in your ScienceJournal.

Plan1. Observe the activities of your daily

life. How do you use energy fromfossil fuels each day?

2. Develop a way to categorize thedifferent chemical reactions and thegreenhouse gases they produce.

3. Search reference sources to learnwhat chemical reactions producegreenhouse gases.

4. What are some of the most com-mon uses of fossil fuels? Is it possi-ble to never use fossil fuels?

Do1. Make sure your teacher approves

your plan before you start.

3. Compare your results with otherstudents. What greenhouse gas ismost often released?

4. Make a table of your data.

1. Record in your Science Journalwhat activities contribute to thegreatest amount of greenhousegases in our atmosphere.

2. Analyze the types of chemical reac-tions that produce greenhouse gases.What types of reactions are they?

ACTIVITY 757

1. How do you think your data wouldbe affected if you had performedthis experiment 100 years ago?

2. What processes in nature might alsocontribute to the release of green-house gases? Compare their impactto that made by fossil fuels.

Find this Use the Internetactivity on the Glencoe

Science Web site at science.glencoe.com.Post your data in the table provided.Compare your data to that of other students.Combine your data with that of otherstudents and write an entry in your ScienceJournal that explains how the production ofgreenhouse gases could be reduced.


Accidentsin SCIENCE

SOMETIMES GREAT DISCOVERIES HAPPEN BY ACCIDENT!

Most people might not think that a

person who had the two qualities

described above would make a

great scientific discovery. But they never met

a chemist named Hilaire de Chardonnet

(hee LAYR • duh • shar doh NAY). In 1878,

Chardonnet accidentally knocked over some

nitrate chemicals. He put off cleaning up the

mess—and ended up inventing artificial silk.

Silk is produced naturally by silkworms. In

the mid-1800s, though, silkworms were dying

from disease and the silk industry was suffer-

ing. Businesses were going under and people

were put out of work. Many scientists were

working to develop a solution to this problem.

AClumsyMove

PaysOffHilaire de

Chardonnet Clumsy (KLUM zee)—lackingphysical coordinationProcrastinate (proh KRAS tuhnayt)—to put off doing something until later

758

Silkworms atwork—spinningtheir cocoons

To help prevent counterfeit-ing, dollars are printed onpaper which contains red andblue rayon fibers. If you canscratch off the red or blue,that means it’s ink—and yourbill is counterfeit. If you canpick out the red or blue fiberwith a needle, it’s a real bill.

Asanteweavers userayon as one of the threads in the Kentecloth.

PolymersRayon is a polymer. A

polymer is a big moleculemade up of smaller molecules that

are linked together to form a straightchain, a twisting ladder, or a branching

tree, depending on how the smaller molecules link together. The cellulose inplant cells is a natural polymer, as is the

DNA in your body. To make artificialpolymers, such as rubber or plastics,scientists studied natural polymers

and then tried to mimic theirstructure in the lab.

For more information, visitscience.glencoe.com

CONNECTIONS Create Make a data table. Work with a partnerto examine the labels on the inside collars of your clothes. These labelslist the different materials found in the clothes. Research the materials,then make a data table that identifies their characteristics.

Chardonnet had been searching

for a silk substitute for years—he just didn’t

plan to find it by knocking it over!

A Messy DiscoveryChardonnet was in his darkroom develop-

ing photographs when the accidental spill

took place. He decided to clean up the spill

later—and finish what he was working on. By

the time he returned to wipe up the spill, the

chemical solution had turned into a thick,

gooey mess. When he pulled the cleaning

cloth away, the goop formed long, thin

strands of fiber that stuck to the cloth. The

chemicals had reacted with the cellulose in

the wooden table and liquefied it. The strands

of fiber looked just like the raw silk made by

silkworms.

Within six years, Chardonnet had devel-

oped a way to make the fibers into an artifi-

cial silk. Other scientists extended his work,

developing a fiber called rayon—“ray” for its

shiny appearance, and “on” to remind people

of the word “cotton.” The rayon we wear on

our backs today, called viscose rayon, is made

with sodium hydroxide. The sodium hydrox-

ide is mixed with wood fibers. The wood turns

into a liquid that can be squeezed through tiny

holes to make strands or fibers. The fibers are

then woven together to make cloth.


2. Bleach—sodium hypochlorite (NaClO)—decomposes in two ways. In one way, it forms oxygen and another product.What common household substance is the other product?

3. In single-displacement reac-tions, one element replaces another in acompound.

4. In double-displacement reactions, ions intwo compounds switch places, often form-ing a gas or insoluble compound.

Section 4 Chemical Reactions and Energy

1. Energy in the form oflight, heat, sound orelectricity is releasedfrom some chemicalreactions known asexergonic reactions.What types of energyare produced in thereaction shown here?

2. Reactions that need energy to proceed arecalled endergonic reactions.

3. Reactions may be sped up by adding cata-lysts and slowed down by adding inhibitors.

4. When energy is released in the form of heat,the reaction is exothermic.

760 CHAPTER STUDY GUIDE

Section 1 Chemical Changes1. In a chemical reaction, one or more sub-

stances are changed to new substances.

2. The substances that react are called reactants,and the new substances formed are calledproducts. Charcoal, the reactant shownbelow, is almost pure carbon. What is a majorproduct formed?

3. The law of conservation of mass states thatin chemical reactions, matter is neither cre-ated nor destroyed, just rearranged.

4. Chemical equations efficiently describewhat happens in chemical reactions.

Section 2 Chemical Equations1. A balanced chemical equation has the same

number of atoms of each element on bothsides of the equation.

2. A balanced equation tells the ratio of prod-ucts and reactants.

3. When balancing equations, change onlythe coefficients of the formulas, never thesubscripts.

Section 3 Classifying Chemical Reactions1. In synthesis reactions, two or more sub-

stances combine to form another substance.

Under the tabs of yourFoldable write the chemi-cal reactions you recorded

using a sentence or chemical equation.

After You ReadFOLDABLESReading & StudySkills

FOLDABLESReading & Study Skills

Study GuideChapter 2424

Study GuideChapter 2424

Complete the spider map about chemical reactions.

ChemicalReactions

When energyis released from the

reaction

Endergonicreaction

Single-displacement

reaction

AB 0A � B

Two substancescombine to form

one substance

Decompositionreaction

AB � CD 0AD � CB

One element replacesanother element in

a compound

A � BC 0AC � B

CHAPTER STUDY GUIDE 761

a. balanced chemicalequation

b. chemical equationc. catalystd. chemical reactione. coefficientf. decomposition

reaction

g. double-displacementreaction

h. endergonic reactioni. endothermic reactionj. exergonic reactionk. exothermic reactionl. inhibitorm. precipitate

n. product p. single-displacement reactiono. reactant q. synthesis reaction

Using VocabularyFor each set of vocabulary words below,

explain the relationship that exists.

1. coefficient, balanced chemical equation

2. synthesis reaction, decomposition reaction

3. reactant, product

4. catalyst, inhibitor

5. exothermic reaction, endothermic reaction

6. chemical reaction, product

7. endergonic reaction, exergonic reaction

8. single-displacement reaction, double-displacement reaction

9. chemical reaction, synthesis reaction

Vocabulary Words

Use acronyms to help you remember importantfacts. For example, the acronym SDSD will helpyou remember the types of chemical reactions—Synthesis, Decomposition, Single Displacement,and Double Displacement.

Study Tip

Choose the word or phrase that best answersthe question.

1. Oxygen gas is always written as O2 in chemical equations. What term is used todescribe the two in this formula?A) product C) catalyst B) coefficient D) subscript

2. What law is based on the experiments ofLavoisier?A) chemical reaction C) coefficientsB) conservation D) gravity

of mass

3. What must an element be in order toreplace another element in a compound?A) more active C) more inhibitingB) a catalyst D) more soluble

4. How do you indicate that a substance in anequation is a solid?A) (l) C) (s)B) (g) D) (aq)

5. What term is used to describe the 4 in theexpression 4 Ca(NO3)2?A) coefficient C) subscriptB) formula D) symbol

6. What type of compound is the food addi-tive BHA?A) catalyst C) inhibitorB) formula D) CFC

7. How do you show that a substance is dis-solved in water when writing an equation?A) (aq) C) (g)B) (s) D) (l)

8. What word would you use to describe HgOin the reaction that Lavoisier used to showconservation of mass?A) catalyst C) productB) inhibitor D) reactant

9. When hydrogen burns, what isoxygen’s role?A) catalyst C) productB) inhibitor D) reactant

10. Give an example of a chemical reaction.A) bending C) meltingB) evaporation D) photosynthesis

11. Chromium is produced by reacting its oxidewith aluminum. If 76 g of Cr2O3 and 27 gof Al completely react to form 51 g ofAl2O3, how many grams of Cr are formed?

12. Propane, C3H8(g), burns in oxygen to formcarbon dioxide and water vapor. Write abalanced equation for burning propane.

13. Use the balanced chemical equation fromquestion 12 to explain the law of conserva-tion of mass.

14. Zn is placed in a solution of Cu(NO3)2 andCu is placed in a Zn(NO3)2 solution. Inwhich of these will a reaction occur?

15. If lye, NaOH(s), is put in water, the solutiongets hot. What kind of energy process is this?

16. Recognizing Cause and Effect Sucrose, ortable sugar, is a disaccharide. This means thatsucrose is composed of two simple sugarschemically bonded together. Sucrose can beseparated into its components by heating itin an aqueous sulfuric acid solution. Researchwhat products are formed by breaking upsucrose. What role does the acid play?

17. Classifying Make an outline with the gen-eral heading “Chemical Reactions.” Includethe four types of reactions, with a descrip-tion and example of each.

762 CHAPTER ASSESSMENT

AssessmentChapter 2424

18. Concept Mapping The arrow in a chemicalequation tells the reaction direction. Somereactions are reversible. Sometimes, thebond formed is weak, and a product breaksapart as it’s formed. Double arrows are usedin the equation to indicate this reaction. Fillin the concept map, using the words prod-uct(s) and reactant(s). In the blank in thecenter, fill in the formulas for the substancesappearing in the reversible reaction.

H2(g) � I2(g) 7 2HI(g)

19. Interpreting Data When 46 g of sodiumwere exposed to dry air, 62 g of sodiumoxide formed. How many grams of oxygenfrom the air were used?

20. Portfolio List five chemical reactions youhave observed. Illustrate them using dia-grams or photos. Include a brief descriptionof the compounds involved.

Test Practice

CHAPTER ASSESSMENT 763

Go to the Glencoe Science Web site at science.glencoe.com or use theGlencoe Science CD-ROM for additionalchapter assessment.

TECHNOLOGY

Students were studying how to balancechemical equations. The table below showssome equations they wrote for the reactionof sulfuric acid with potassium hydroxide.

Study the table above and answer thefollowing questions:

1. Which of the equations in the table iscorrectly balanced?A) oneB) twoC) threeD) four

2. All of the chemical reactions in thetable show _____.F) reactants combining with products

to produce new substancesG) reactants not being changed by the

chemical reactionH) reactants combining to produce

new productsJ) products combining to produce

new reactants

AssessmentChapter 2424

1

2

3

4

H2SO4 � KOH 0 K2SO4 � H2O

H2SO4 � 2KOH 0 K2SO4 � H2O

H2SO4 � 2KOH 0 K2SO4 � 2H2O

2H2SO4 � 4KOH 0 K2SO4 � H2O

are

are

forms

is

is

become becomes

ProductReactants

2HI


24 chemical reactions - irion county isd / overvie · 24 chemical reactions o hhhh! ahhhh! ......

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