11.1 chemical quations - weebly · chapter 11. chemical reactions figure 11.4 dependent on the...

www.ck12.org Chapter 11. Chemical Reactions

CHAPTER 11 Chemical ReactionsChapter Outline

11.1 CHEMICAL EQUATIONS

11.2 TYPES OF CHEMICAL REACTIONS

11.3 REFERENCES

For millennia, humans have been fascinated with the composition of things and the workings of the chemical world.Over time we have come to understand that all matter is comprised of indivisible particles called atoms. Ourunderstanding of how matter works has been a long pursuit. It started with figuring out how to make things burn.Humans have been fascinated with chemical reactions that burn, explode, produce loud bangs, and have brilliantcolors. Early alchemists learned that throwing certain salts on a fire would produce different “magical” colors.Chinese alchemists created human kind’s first explosion with the invention of gunpowder. This chemical recipe waseventually shared across the medieval globe. Historically, humans have been fascinated with coaxing nature intodoing things, like burning. This fascination, coupled with our ability to observe, record, and share, has led us toour current understanding of matter. Our understanding of chemical reactions and the equations that describe themare based on many years of trial and error. The image above is an example of this. It is a star shell bursting overthe night sky. The technology of pyrotechnics, like composition of the propellant, the explosive charge, the colors,and the shapes of the burst, is a result of hundreds of years of intensive study of chemical reactions and chemicalequations. We are going to study chemical reactions and chemical equations in this chapter.Jon Sullivan. commons.wikimedia.org/wiki/File:Firework. jpg. Public Domain.

245

11.1. Chemical Equations www.ck12.org

11.1 Chemical Equations

Lesson Objectives

• Give examples of historically significant chemical recipes and equations.• Briefly describe the major milestones that took place in developing the chemical recipe of gunpowder.• Understand mass relations between reactants and products for a given chemical process.• Be able to use stoichiometric coefficients in chemical equations.• Be able to balance chemical equations.

Lesson Vocabulary

• stoichiometric coefficient: The letters a, b, c, and d where A and B are reactants, and C and D are products.The stoichiometric coefficients indicate the relative amounts of reactants and products.

• balanced chemical equation: An equation where the number of atoms of each element on the reactant sideis equal to the number of atoms on the product side.

Check Your Understanding

1. Which of the following are physical changes and which are chemical changes?

a. melting of iceb. a burning candlec. melting of candle waxd. sublimation of dry ice to CO2 gas.

Introduction

Ever since the 9th century, humans have been fascinated with the nature of explosions. Whether to scare away evilspirits, to light up the night sky in celebration, or to be used in warfare, our understanding of gunpowder is based onour understanding of chemical recipes. Our ability to modify, share, and replicate them has allowed us to developnew recipes and to refine existing ones. Chemical reactions can be described in terms of chemical equations. Theyare the foundation of our modern day chemical recipes.

246


FIGURE 11.1

Origins of Chemical Recipes: Gunpowder

The first people to discover gunpowder were 9th century Chinese alchemists. This discovery was made by accidentwhile they were creating various chemical mixtures in pursuit of an elixir that would make them immortal. Thefirst formulation of gunpowder was a thick toffee made from honey, saltpeter (a mixture composed primarily ofpotassium nitrate), and sulfur. They hoped that eating it would help them live forever. In reality, it burst into flamesand burnt down their homes.

Over time, Chinese alchemists refined the recipe and began to develop early pyrotechnic technology to help scareaway evil spirits. A more fully developed, and more explosive, formula called for 75 percent potassium nitrate, 15percent charcoal, and 10 percent sulfur.

FIGURE 11.2

247


Medieval Europe

This recipe made its way to Europe in medieval times. A Franciscan friar named Roger Bacon was particularlyfascinated by the properties of gunpowder. He discovered that a key factor in the energetics of the mixture was thepurity of the saltpeter. Bacon was responsible for developing early crystallization techniques to purify the mixture.He also discovered that the more tightly packed the powder, the larger the explosion.

Bacon feared that bad things could happen if the mixture ended up in the wrong hands. He encoded the recipe in ananagram, which read (when translated from the original Latin) “And so thou wilt call up thunder and destruction ifthou know the art.” The secret recipe, however, did not stay secret for long.

FIGURE 11.3

Pyrotechnics, or fireworks, used in events recorded in 14th century Italy show that the recipe was no longer a secret.During the 15th and 16th centuries, the Italians continued refining the art of pyrotechnics. Then, in 1830, a majorleap forward in gunpowder technology occurred. It was discovered that replacing potassium nitrate with potassiumchlorate resulted in a more energetic mixture, and so the recipe was revised once again.

The modern day formulation of gun powder is called black powder. It is still commonly used today. Its formulationis still quite similar to what was used in 9th century China. Black powder is considered a low explosive. It is amixture that burns quickly, but the resulting shock wave travels at subsonic speeds. The speed at which it burns is

248


FIGURE 11.4

dependent on the accessibility of oxygen atoms to the carbon source. In contrast, high explosives like nitroglycerindetonate instead of burning, creating shock waves that are supersonic (faster than the speed of sound).

The Chemical Equation

Chemical equations describe the changes in composition that take place during a chemical reaction. Along withthe identities of the starting reactants and the final products, chemical equations show the ratios in which thesesubstances are consumed and produced. The reaction of iron with oxygen to form iron(III) oxide is shown in theFigure 11.5:

FIGURE 11.5The sparks from a steel grinder aremolten iron. The iron reacts with oxygento form iron(III) oxide.

We can describe this reaction with a chemical equation:

249


4Fe(s) +3O2(g) ! 2Fe2O3(s)

This equation is said to be balanced, because the amount of each element expressed on the reactants side is equal tothe amounts expressed on the products side. This is shown more explicitly in the following Table 11.1:

TABLE 11.1: 4Fe

Reactants ProductsFe 4 4O 6 6

Often times, the processes described by chemical equations do not represent a single reaction. For example, thefollowing equation shows the starting materials and the products for photosynthesis:

6CO2(g) +6H2O(l)light! C6H12O6(s) +6O2(g)

This process does not occur in a single step. A sequence of many individual reactions is required to make glucoseand oxygen gas out of carbon dioxide and water. Chemical equations can be used to represent individual reactionsor the net change that occurs after multiple sequential chemical processes.

The Balanced Chemical Equation

We can describe chemical reactions in terms of generic expressions like the following equation:

aA+bB ! cC+dD

where A and B are reactants, and C and D are products. The letters a, b, c, and d represent stoichiometriccoefficients, or the relative amount of each substance that is involved in the reaction. In this particular reaction,there are two reactants and two products, but others might have more or less. For example, the equation describingthe rusting of iron had two reactants (Fe and O2) and one product (Fe2O3).

In a balanced chemical equation, the number of atoms of each element on the reactant side is equal to the numberof atoms on the product side. This is necessary for all chemical equations, due to the law of conservation of mass.Atoms are neither created nor destroyed during a chemical reaction, only rearranged. Here are some examples ofgeneral expressions that will be applied to specific reactions in the next section.

Example 11.1

Substance A reacts with substance B to form substance AB. Write the balanced chemical equation for this process.

Answer:

Write the general expression.

A+B ! AB

Balance (it already is).

TABLE 11.2: A + B ! AB

Reactants ProductsA 1 1B 1 1

250


Example 11.2

Substance A reacts with substance B2 to form substance AB. Write the balanced chemical equation for this process.

Answer:


A+B2 ! AB

Balance.

2A+B2 ! 2AB

TABLE 11.3: Example 11.2


Example 11.3

Substance A2 reacts with substance B2 to form substance AB3. Write the balanced chemical equation for thisprocess.

Answer:


A2 +B2 ! AB3

Balance.

TABLE 11.4: A


Balancing Chemical Equations for Real Reactions

Now that we have studied the general process for describing and balancing chemical equations, we are going toapply this approach to examples that include actual chemicals. As we present the following reactions, we are goingto focus only on the changes in composition from reactants to products. In later chapters, we will look at otherreaction properties, such as states of matter, temperature, and the energy lost or gained by a given reaction. In thefollowing lesson, we will look at ways to classify different types of reactions. This knowledge will allow us to makereasonable predictions about the products that might be generated from a given set of reactants.

Tips for Balancing Equations

Before we get started with balancing chemical equations, here are some simple tips to consider:

1. If there are polyatomic ions that exist unchanged on both sides of the equation, it is often simpler to treat them

251


as single units than to break them down into their individual elements.2. It is often easier to leave elements that occur in their pure elemental form (on either side of the equation) for

last.3. If a reactant or product has a coefficient of 1, this number is not explicitly written.4. In a correctly balanced equation, all coefficients must be whole numbers. However, the use of fractions can

be helpful as a way of finding the correct coefficients. If all atoms in an equation are balanced but some havefractional coefficients, multiply all coefficients in the entire equation (including those not explicitly written!)by the lowest common denominator to get the final balanced equation.

Example 11.4

Liquid mercury is heated in the presence of oxygen to produce mercury(II) oxide. Write the balanced chemicalequation for this process.

Answer:

Start by writing the general expression.

Hg(l)+O2(g)! HgO(s)

Then, alter the coefficients to balance each element.

TABLE 11.5: 2Hg(l)+O

Reactants ProductsHg 2 2O 2 2

Notice that in this example, the formula for oxygen is the diatomic form O2. Many pure nonmetallic elementsare unstable as individual atoms and combine readily to make diatomic molecules. Hydrogen (H2), nitrogen (N2),oxygen (O2), and the halogens (F2, Cl2, Br2, and I2) exist as diatomic molecules when in their pure elemental forms.

Example 11.5

Hydrogen gas and fluorine gas react to form hydrogen fluoride gas. Write the balanced chemical equation for thisprocess.

Answer:

Start by writing the general expression.

H2(g)+F2(g)! HF(g)

Then balance each element.

TABLE 11.6: H

Reactants ProductsH 2 2F 2 2

Again, pure hydrogen and fluorine exist as diatomic gases.

252


Example 11.6

Ammonium nitrate decomposes to form nitrogen gas, water, and oxygen gas. Write the balanced chemical equationfor this process.

Answer:


NH4NO3(s)! N2(g)+H2O(l)+O2(g)

Balance.

Because this equation involves more than two elements, it is slightly less straightforward to balance. Since nitrogenand oxygen both occur in their pure elemental forms, we start by balancing hydrogen:

NH4NO3(s)! N2(g)+2H2O(l)+O2(g)

Hydrogen and nitrogen are now balanced, but oxygen is not. This can be fixed by changing the coefficient on itspure elemental form:

NH4NO3(s)! N2(g)+2H2O(l)+ 12 O2(g)

The atoms are now balanced, but to avoid having fractional coefficients, we must multiply all coefficients in theequation by 2:

2NH4NO3(s)! 2N2(g)+4H2O(l)+O2(g)

We can confirm that this equation is balanced by writing the following Table 11.7:

TABLE 11.7: 2NH

Reactants ProductsN 4 4H 8 8O 6 6

Example 11.7

Lead(II) nitrate reacts with sodium chloride to form lead(II) chloride and sodium nitrate. Write the balancedchemical equation for this process.

Answer:


Pb(NO3)2 +NaCl ! PbCl2 +NaNO3

Balance.

TABLE 11.8: Pb(NO

Reactants ProductsPb 1 1NO3 2 2Na 2 2Cl 2 2

253


By keeping the polyatomic nitrate ion intact as a single unit, balancing this equation becomes somewhat simpler.This was done because the ion exists unchanged on both sides of the equation. Note that this is in contrast to theprevious example, in which the nitrate ion decomposed to form other substances.

Lesson Summary

• The composition of gunpowder gradually changed as alchemists and scientists experimented with ways tomake it even more explosive.

• Chemical reactions are described using chemical equations.• Stoichiometric coefficients are used in chemical equations to indicate the amounts of reactants and products.• Because of the law of conservation of mass (matter can neither be created nor destroyed through chemical

reactions), chemical equations must have equal amounts of each specific atom on both sides of the equation.

Review Questions

1. Early Chinese alchemists discovered an early form of gunpowder. What was the composition of this sub-stance?

2. What later developments were made to the gunpowder recipe that improved its pyrotechnic properties?3. Make an argument for why the burning of a candle is consistent with the law of conservation of matter/mass.4. Think of an experiment that you could conduct to demonstrate that mass is conserved for a given chemical

change.5. Balance the following chemical equations:

a. C+O2 ! COb. CO+O2 ! CO2c. H2 +Br2 ! HBrd. K+H2O ! KOH+H2e. O3 ! O2f. N2 +H2 ! NH3g. Zn+AgCl ! ZnCl2 +Agh. Cl2 +NaI ! NaCl+ I2i. P4O10 +H2O ! H3PO4j. Be2C+H2O ! Be(OH)2 +CH4k. S+HNO3 ! H2SO4 +NO2 +H2Ol. NH3 +CuO ! Cu+N2 +H2O

m. HCl+CaCO3 ! CaCl2 +H2O+CO2

Further Reading / Supplemental Links

• Youtube Video of Kaboom! The Sizzling Story of Explosions: http://www.youtube.com/watch?v=CShA52EKY80

• Gunpowder in Ancient China: http://www.historyforkids.org/learn/war/gunpowder.htm• Practice Balancing Chemical Equations:

– http://education.jlab.org/elementbalancing/index.html– http://www.files.chem.vt.edu/RVGS/ACT/notes/scripts/bal_eq1.html

254


– http://gregthatcher.org/Chemistry/BalanceEquation/S

• Chemical Equation Balances: http://www.personal.psu.edu/jzl157/balance.htm

Points to Consider

• What is the relationship between chemical equations and chemical reactions?• In this chapter, an argument was made that the human fascination with fire and explosions ultimately con-

tributed to our current understanding of chemical equations. Can you think of other aspects of nature forwhich further exploration has contributed to our current understanding of chemistry?

255

11.2. Types of Chemical Reactions www.ck12.org

11.2 Types of Chemical Reactions

Lesson Objectives

• Be able to classify a chemical reaction as a combination, decomposition, single replacement, double replace-ment, or combustion reaction.

• Be able to predict the products when given a set of reactants for a given chemical process.• Explain the concept of solubility and the process of precipitation.• Use solubility information to predict whether or not a given substance is soluble in water.• Use the general solubility rules to predict chemical behavior.• Be able to write molecular, ionic, and net ionic equations for a given chemical process.

Lesson Vocabulary

• combination reaction: A reaction where two or more chemical species combine to produce a single newcompound.

• decomposition reaction: A reaction where a single chemical species breaks down to produce two or morenew chemical species.

• single replacement reaction: Occurs when one chemical species (often a single element) replaces a portionof another compound to produce two new products.

• double replacement reaction: Occurs when the cations from the original two ionic compounds trade anionsto make two new ionic compounds.

• molecular equation: An equation that shows all ionic components as neutral compounds, but the ones thatare dissolved in water are denoted with "(aq)."

• ionic equation: A chemical equation in which the various reaction components are represented as theyactually exist in the reaction, for example, as individual ions.

• spectator ion: Ions that are present in solution but do not participate in the overall reaction.• net ionic equation: The simplified ionic equation in which all of the spectator ions are cancelled out.• combustion: Occurs when a hydrocarbon reacts in the presence of oxygen to produce water and carbon

dioxide.

Check Your Understanding

Study the Figure 11.6, which depicts the mass change that occurs when steel wool burns in air.

1. What happens to the mass of the steel wool as the reaction proceeds?2. Given that mass must be conserved in chemical reactions (it cannot come from nowhere), what might be your

explanation for the change in the mass of the steel wool?3. How might mass changes such as this help us identify and categorize a given chemical process?

256


FIGURE 11.6Mass changes for steel wool burning inair

Introduction

MEDIAClick image to the left for more content.

The video above at http://www.youtube.com/watch?v=_Y1alDuXm6A (1:12) shows the decomposition of mer-cury(II) oxide into liquid mercury and oxygen gas. This reaction was an important one in the history of chemistry,because it helped early chemists to understand the relationship between reactants and products.

In the last lesson, we began investigating how a chemical equation can represent a given chemical reaction. In thislesson, we are going to study the ways in which chemical reactions are classified. There are literally thousandsof chemical reactions that take place every day in our lives. Some reactions take place in the atmosphere, suchas the combustion of fossil fuels. Others occur in solution, like the reactions responsible for photosynthesis or thereactions that break down our food to give us energy. Chemical reactions can take place in a variety of environments.Reactions happen on the sea floor, in our cells, and in the upper atmosphere. As we look at chemical reactions, wenotice some commonalities and trends. When we studied the elements, we saw characteristics that allowed us tocategorize them by family. There are also various ways to categorize chemical reactions. Some reactions produceheat, while others consume it. Some reactions are spontaneous, while others are not. Some reactions happen innanoseconds, while others happen over longer spans of time. Some produce electricity, some emit light, and somerelease gaseous products. The products of chemical reactions tell us a lot about the chemistry of the process. In theabove figure, we see mercury(II) oxide decomposing into elemental mercury and oxygen gas. Decomposition wasone of the first reaction types to be identified by chemists. Decomposition is one type of reaction you’ll learn aboutin this lesson.

Combination Reactions

The first type of reaction that we will investigate is the combination reaction, which is sometimes also referred toas a synthesis reaction. In combination reactions, two or more chemical species combine to produce a single newcompound. A generic combination reaction might have the following form:

A+B !C

257


Substances in all states of matter can participate in combination reactions. For example, oxygen in the air can reactwith iron to produce rust. Rusting is a common occurrence, especially in regions of the world where precipitation isrelatively high. Although rust tends to be a mixture of compounds, its primary component is iron(III) oxide (Fe2O3).Rusting is generally a very slow process, but when the iron has a very high surface area, as in the case of steel wool,it can happen at a much faster rate, as shown in the following video:

http://www.youtube.com/watch?v=5MDH92VxPEQ

MEDIAClick image to the left for more content.

The balanced chemical equation for this process is shown below:

4Fe(s)+3O2(g)! 2Fe2O3(s)

Decomposition Reactions

A decomposition reaction is the exact opposite of a combination reaction. In decomposition reactions, a singlechemical species breaks down to produce two or more new chemical species. A generic decomposition reactionmight take the following form:

C ! A+B

Again, substances in all states of matter commonly participate in decomposition reactions. For example, hydrogenperoxide will decompose over time to produce water and oxygen gas according to the following equation:

2H2O2(l)! 2H2O(l)+O2(g)

Another common type of decomposition reaction involves the process of electrolysis, in which an electrical currentis passed through a substance to break apart a compound. One example of a decomposition reaction requiring theuse of electrolysis is the decomposition of molten sodium chloride, as shown by the following equation:

2NaCl(s)! 2Na(s)+Cl2(g)

Single Replacement Reactions

A single replacement reaction (sometimes called a single displacement reaction) occurs when one chemical species(often a single element) replaces a portion of another compound to produce two new products. The general form ofa single replacement reaction is shown below:

AB+C ! AC+B

Two common types of single replacement reactions involve pure metals reaction with aqueous solutions of either anacid or an ionic compound. When a reactive metal is placed in an acid solution, the following reaction is likely tooccur:

Metal + acid ! ionic solution + hydrogen gas

258


An example of this would be the reaction between zinc and hydrochloric acid, which produces zinc chloride andhydrogen gas. Here is an image of this reaction:

FIGURE 11.7Zinc metal reacting with a solution of hy-drochloric acid

The balanced chemical equation for this single replacement reaction is shown below:

Zn(s)+2HCl(aq)! ZnCl2(aq)+H2(g)

Another type of single replacement reaction involves a solid metal replacing the metal cation in an ionic compoundthat has been dissolved in water. If the solid metal is more reactive than the dissolved metal cations, the followingtype of reaction can occur:

Metal + ionic solution ! different metal + different ionic solution

A common example of this reaction is when iron is replaced by the more reactive zinc metal. The balanced chemicalequation for this process is shown below.

Zn(s)+FeSO4(aq)! Fe(s)+ZnSO4(aq)

Double Replacement Reactions

Double replacement reactions typically include two water-soluble salts that react with one another in solution. Thegeneral form of a double replacement reaction would look something like the following:

AB+CD ! AD+CB

In double replacement reactions, the cations from the original two ionic compounds trade anions to make two newionic compounds. In general, at least one of the new compounds must precipitate (form an insoluble solid) for us toconclude that a reaction has occurred. An example of such a process is shown below with the double replacementreaction between solutions of potassium iodide and lead(II) nitrate.

At the molecular level, our model for the way in which a precipitate forms can be described in an animation:

259


FIGURE 11.8A double replacement reaction is used toform lead(II) iodide. The reactants shownhere are colorless solutions of potassiumiodide and potassium nitrate. When com-bined, these produce a yellow precipitateof lead(II) iodide.

http://www.crescent.edu.sg/crezlab/webpages/PptReaction_PbI2.htm

Representing Ionic Reactions as Chemical Equations

For reactions that involve ions dissolved in water, there are several different ways to express the overall process as achemical equation. For example, the overall molecular equation shows all ionic components as neutral compounds,but the ones that are dissolved in water are denoted with "(aq)." Note that the ionic substances do not exist asmolecules, but we write them out as though they were. In the following example, two water-soluble compoundstrade partners to produce one dissolved ionic compound and one solid precipitate:

AB(aq)+CD(aq)! AD(aq)+CB(s)

In reality, the aqueous substances do not exist as molecules or ionic crystal lattices. Instead, the individual ions aredissolved and distributed throughout the solution. If the reaction above were written as an ionic equation, it wouldlook something like the following:

A+(aq)+B�(aq)+C+(aq)+D�(aq)! A+(aq)+D�(aq)+CB(s)

In this example, the various reaction components are presented in a form that is closer to the way they actually existduring the reaction. The aqueous components are separated into ions, and the precipitate is found as a combinedsolid. We are assuming in this example that A and C form cations with a charge of 1+, while B and D form anionswith a charge of 1-. In real examples, we would look at which group each element is found in on the periodic tableto determine its likely charge.

Notice that in the ionic equation, A+ and D� were unchanged over the course of the reaction; they exist as aqueousions on both the reactant and product sides. In other words, these species did not experience any net change. Ions

260


that are present in solution but do not participate in the overall reaction are known as spectator ions. The ionicequation can be simplified to the net ionic equation by canceling out all the spectator ions.

⇠⇠⇠⇠A+(aq)+B�(aq)+C+(aq)+⇠⇠⇠⇠D�(aq)!⇠⇠⇠⇠A+(aq)+⇠⇠⇠⇠D�(aq)+CB(s)B�(aq)+C+(aq)!CB(s)

Let’s look at these three types of equations again using a real example. If we were to mix aqueous solutions ofpotassium iodide and lead(II) nitrate, lead(II) iodide would precipitate as a solid, and potassium nitrate would remaindissolved. This can be represented by any of the three following equations:

Molecular Equation

2KI(aq)+Pb(NO3)2(aq)! 2KNO3(aq)+PbI2(s)

Ionic Equation

2K+(aq)+2I�(aq)+Pb2+(aq)+2NO�3 (aq)! 2K+(aq)+2NO�

3 (aq)+PbI2(s)

Net Ionic Equation

Pb2+(aq)+2I�(aq)! PbI2(s)

Predicting Solubility of Ionic Compounds

How do we determine which ions are likely to form an insoluble precipitate and which will remain dissolved inwater? By combining various ionic solutions, chemists have come up with some general guidelines for whether agiven cation-anion pairing is likely to be soluble or insoluble in water. It should be noted that such an approach isan oversimplification. Each compound has its own solubility value, so two "soluble" compounds might have verydifferent abilities to dissolve in water. Additionally, even "insoluble" salts can dissolve in water to a very limitedextent. We will take a more quantitative approach to solubility in the chapter on solutions. However, qualitativerules like the ones in the Table 11.9 are useful for predicting whether a precipitate is likely to form when combiningmoderate amounts of specific cations and anions.

TABLE 11.9: Solubility Properties to Predict Products of Chemical Reactions

Type of Particle Soluble InsolubleCommon Cations Alkali metal cation (Li+, Na+, K+,

Rb+, or Cs+) or the NH4+ cation

Common Anions ClO4� and NO3

� compoundsHalides Most Cl�, Br�, and I� compounds Compounds that include the Ag+,

Pb2+, or Hg22+ cations

Sulfates Most SO42� compounds PbSO4, Ag2SO4, Hg2SO4, CaSO4,

SrSO4, and BaSO4Sulfides Compounds with NH4

+ or a metalfrom group IA or IIA as cation

Most S2� compounds

Hydroxides Compounds with NH4+, Ba2+, or a

metal from group IA as cationMost OH� compounds

Carbonates, phosphates, and sul-fites

Compounds with NH4+ or a metal

from group IA as cationMost CO3

2� and PO43�, and

SO32� compounds

261


Combustion

Combustion occurs when a hydrocarbon reacts in the presence of oxygen to produce water and carbon dioxide.These reactions are very exothermic, which means that they produce a large amount of heat. Combustion reactionsare quite common in our everyday lives, such as the burning of gasoline to fuel a car. The chemical equation for acombustion reaction has the following generic form:

CxHy +O2 ! H2O+CO2

FIGURE 11.9Combustion reaction of a marshmallow(sucrose) and wood (cellulose).

The process of cellular respiration can be thought of as a highly controlled version of a combustion reaction. We donot literally burn hydrocarbons in our body, but the overall reactants and products are the same. Hydrocarbons, suchas sucrose (C12H22O11), are combined with oxygen in a series of enzymatic steps to product water, carbon dioxide,and energy, which is stored in the form of reactive molecules. The unbalanced chemical equation for this overallprocess is shown below:

C12H22O11 +O2 ! CO2 +H2O

Lesson Summary

• Combination reactions occur when two or more reactants combine to produce a single compound.• Decomposition reactions involve one compound decomposing into two or more products.• Single replacement reactions occur when one reactant replaces part of another compound to form new sub-

stances.• A common type of double replacement reaction occurs when two ionic reactants exchange anions, making

two new ionic compounds. The precipitation of a solid is a common result for this type of reaction.• Combustion reactions involve the reaction of a hydrocarbon with oxygen gas to produce water and carbon

dioxide.

262


Review Questions

1. Categorize the following chemical reactions as single replacement, double replacement, combustion, combi-nation, or decomposition.

a. Equimolar (having the same number of moles) solutions of silver nitrate and potassium chloride aremixed to produce solid silver chloride and aqueous potassium nitrate.

b. Magnesium metal is added to hydrochloric acid to produce hydrogen gas and aqueous magnesiumchloride.

c. Ethanol is burned in air to produce water and carbon dioxide gas.d. Water is electrolyzed to produce hydrogen and oxygen gas.e. Hydrogen gas and oxygen gas are ignited to produce water.

2. Write the balanced chemical equation for the following combination and decomposition reactions.

a. Magnesium carbonate is heated strongly to produce magnesium oxide and carbon dioxide gas.b. Hydrogen peroxide decomposes to produce water and oxygen gas.c. Solid potassium chlorate is heated in the presence of manganese dioxide as a catalyst to produce potas-

sium chloride and oxygen gas. (Catalysts speed up reactions but are not expressed in the overall balancedequation)

d. Molten aluminum oxide is electrolyzed using inert (non-reactive) electrodes to produce aluminum metaland oxygen gas.

3. Write the balanced chemical equations for the following replacement reactions:

a. Zinc metal is added to a solution of iron(II) sulfate.b. Equimolar solutions of lead(II) nitrate and sodium chloride are mixed to produce solid lead(II) chloride

and aqueous sodium nitrate.c. Solutions of potassium phosphate and zinc nitrate are mixed.

4. Write the balanced chemical equations for the following combustion reactions.

a. Propane (C3H8) is ignited in air to produce water and carbon dioxide gas.b. Methanol(CH4O) is ignited in air to produce water and carbon dioxide gas.c. Ethanol (C2H5OH) is burned in air.

5. Write the molecular equation, ionic equation, and net ionic equation for each of the following double replace-ment reactions.

a. Silver nitrate reacts with potassium iodide to produce potassium nitrate and silver iodide.b. Silver nitrate reacts with iron(III) chloride to produce iron(III) nitrate and silver chloride.c. Lead(II) nitrate reacts with potassium iodide to produce potassium nitrate and lead(II) iodide.d. Iron(III) chloride reacts with lead(II) nitrate to produce lead(II) chloride and iron(III) nitrate.e. Calcium chloride reacts with sodium hydroxide to produce calcium hydroxide and sodium chloride.

6. Would it be possible to have a double precipitate formed for a double replacement process? Can you write anequation where a double precipitate forms?

7. What is meant when we describe a compound as (aq) or (s)? Explain the similarities and differences betweenthese terms.

8. Write the balanced chemical equation for the combination reaction in which hydrogen and oxygen gases reactexplosively to produce water. (Remember that hydrogen and oxygen exist as diatomic gases in their mostcommon elemental form.)

9. Write the balanced chemical equation for the reaction that occurs when a piece of aluminum metal is placedin a solution of silver nitrate.

10. Using the solubility rules given above, predict whether or not the following compounds are soluble or insolublein water.

263


a. Potassium nitrateb. Lead(II) chloridec. Barium sulfated. Aluminum sulfidee. Calcium carbonate

Further Reading / Supplemental Links

• Chemical reaction library: http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/STILLS/VOLTAGE/VOLTAGE3/64JPG48/3.JPG

• Solubility concepts video: http://www.khanacademy.org/science/chemistry/states-of-matter/v/solubility• How to use a solubility chart: http://www.sophia.org/solubility-table/solubility-table-tutorial

Points to Consider

1. In an earlier section, we discussed the origins of the chemical recipe for gunpowder, one of the earliestchemical formulas to be described. The recipe for gun powder is 75 percent potassium nitrate, 15 percentcharcoal, and 10 percent sulfur. How might one measure out these amounts in a predictable and reliable way?

2. So far, we have discussed the characteristics of a variety of reactions. However, we have spent little timediscussing how we might measure and calculate amounts of reactants and products. The steel wool reaction isas follows: 4Fe(s) +3O2(g) ! 2Fe2O3(s). How might you measure the amounts of each reactant used and theproduct that forms?

3. In the chemical reactions that we have already studied, we have assumed that all reactants are transformedinto products (the reaction "goes to completion"). Are there reactions that do not go to completion? How doyou know whether you will have reactants left over?

4. What are some factors that control whether or not a chemical reaction takes place?

264


11.3 References

1. Courtesy of NASA. http://commons.wikimedia.org/wiki/File:Chinese_rocket.gif . Public Domain2. User:PericlesofAthens/Wikimedia Commons. http://commons.wikimedia.org/wiki/File:Chinese_Gunpowder

_Formula.JPG . Public Domain3. Joseph Wright. http://commons.wikimedia.org/wiki/File:JosephWright-Alchemist.jpg . Public Domain4. Oliver H.. http://commons.wikimedia.org/wiki/File:Spk-RZ.jpg . Public Domain5. Jared Tarbel. http://www.flickr.com/photos/generated/5554654375/ . CC BY 2.06. Jodi So. CK-12 Foundation . CC BY-NC 3.07. User:Chemicalinterest/Wikimedia Commons. http://commons.wikimedia.org/wiki/File:Zn_reaction_with_-

HCl.JPG . Public Domain8. Der Kreole. http://commons.wikimedia.org/wiki/File:Lluvia_de_oro.JPG . CC BY 3.09. Flickr:webhamster. http://www.flickr.com/photos/26316553@N07/2897369014/ . CC BY 2.0

265

11.1 chemical quations - weebly · chapter 11. chemical reactions figure 11.4 dependent on the...

Documents