The Chemical Equation
A chemical equation is a shorthand expression for a chemical change or reaction. A chemical equation uses
the chemical symbols and formulas of the reactants and products and other symbolic terms to represent a
chemical reaction. The equations are written according to this general format:
1.
Reactants are separated from products by an arrow → that indicates the direction of the reaction. The
reactants are placed to the left and the products to the right of the arrow. A plus sign (+) is placed
between reactants and between products when needed.
Al + Fe2O3reactants→Fe + Al2O3products
2.
Coefficients (whole numbers) are placed in front of substances to balance the equation and to indicate
the number of units (atoms, molecules, moles, ions) of each substance reacting or being produced.
When no number is shown, it is understood that one unit of the substance is indicated.
2 Al + Fe2O3→2 Fe + Al2O3
3.
Conditions required to carry out the reaction may, if desired, be placed above or below the arrow or
equality sign. For example, a delta sign placed over the arrow (→Δ) indicates that heat is supplied to
the reaction.
2 Al + Fe2O3→Δ2 Fe + Al2O3
4.
The physical state of a substance is indicated by the following symbols: (s) for solid state; (l) for liquid
state; (g) for gaseous state; and (aq) for substances in aqueous solution. States are not always given in
chemical equations.
2 Als + Fe2O3s→ΔFel + Al2O3s
Conservation of Mass
The law of conservation of mass states that no change is observed in the total mass of the substances involved
in a chemical change. This law, tested by extensive laboratory experimentation, is the basis for the quantitative
mass relationships among reactants and products.
The decomposition of water into hydrogen and oxygen illustrates this law. Thus, 100.0 g of water decompose
into 11.2 g of hydrogen and 88.8 g of oxygen:
water→hydrogen + oxygen100.0 g11.2 g88.8 g100.0 g→100.0 greactantproducts
In a chemical reaction mass of reactants = mass of products
PROBLEM-SOLVING STRATEGY
For Writing and Balancing a Chemical Equation
1.
Identify the reaction. Write a description or word equation for the reaction. For example, let’s consider
mercury(II) oxide decomposing into mercury and oxygen.
mercuryII oxide→Δmercury + oxygen
2.
Write the unbalanced (skeleton) equation. Make sure that the formula for each substance is correct and
that reactants are written to the left and products to the right of the arrow. For our example,
HgO→ΔHg + O2
The correct formulas must be known or determined from the periodic table, lists of ions, or experimental data.
3.
Balance the equation. Use the following process as necessary:
a.
Count and compare the number of atoms of each element on each side of the equation and determine
those that must be balanced:
b.
Hg is balanced
(1 on each side)
O needs to be balanced
(1 on reactant side, 2 on product side)
Balance each element, one at a time, by placing whole numbers (coefficients) in front of the formulas
containing the unbalanced element. It is usually best to balance metals first, then nonmetals, then
hydrogen and oxygen. Select the smallest coefficients that will give the same number of atoms of the
element on each side. A coefficient placed in front of a formula multiplies every atom in the formula
by that number (e.g., 2 H2SO4 means two units of sulfuric acid and also means four H atoms, two S
atoms, and eight O atoms). Place a 2 in front of HgO to balance O:
2 HgO→ΔHg + O2
c.
Check all other elements after each individual element is balanced to see whether, in balancing one
element, other elements have become unbalanced. Make adjustments as needed. Now Hg is not
balanced. To adjust this, we write a 2 in front of Hg:
2 HgO→Δ2 Hg + O2balanced
d.
Do a final check, making sure that each element and/or polyatomic ion is balanced and that the
smallest possible set of whole-number coefficients has been used:
2 HgO→Δ2 Hg + O2correct form4 HgO→Δ4 Hg + 2 O2incorrect form
When methane, CH4, undergoes complete combustion, it reacts with oxygen to produce carbon dioxide and
water. Write the balanced equation for this reaction.
SOLUTION
Use the Problem-Solving Strategy for Writing and Balancing a Chemical Equation.
1.
Word equation:
methane + oxygen→carbon dioxide + water
2.
Skeleton equation:
CH4 + O2→CO2 + H2Ounbalanced
3.
Balance:
a.
Carbon is balanced.
Hydrogen and oxygen are not balanced.
b.
Balance H atoms by placing a 2 in front of H2O:
CH4 + O2→CO2 + 2 H2Ounbalanced
Each side of the equation has four H atoms; oxygen is still not balanced. Place a 2 in front of
O2 to balance the oxygen atoms:
CH4 + 2 O2→CO2 + 2 H2Obalanced
c.
The other atoms remain balanced.
d.
Check: The equation is correctly balanced; it has one C, four O, and four H atoms on each
side.
Oxygen and potassium chloride are formed by heating potassium chlorate. Write a balanced equation for this
reaction.
SOLUTION
Use the Problem-Solving Strategy for Writing and Balancing a Chemical Equation.
1.
Word equation:
potassium chlorate→Δpotassium chloride + oxygen
2.
Skeleton equation:
KClO3→ΔKCl + O2unbalanced
3.
Balance:
a.
Potassium and chlorine are balanced.
Oxygen is unbalanced (three O atoms on the left and two on the right side).
b.
How many oxygen atoms are needed? The subscripts of oxygen (3 and 2) in KClO3 and
O2 have a least common multiple of 6. Therefore, coefficients for KClO3 and O2 are needed to
get six O atoms on each side. Place a 2 in front of KClO3 and a 3 in front of O2 to balance
oxygen:
2 KClO3→ΔKCl + 3 O2unbalanced
c.
Now K and Cl are not balanced. Place a 2 in front of KCl, which balances both K and Cl at the
same time:
2 KClO3→Δ2 KCl + 3 O2balanced
d.
Check: Each side now contains two K, two Cl, and six O atoms.
Silver nitrate reacts with hydrogen sulfide to produce silver sulfide and nitric acid. Write a balanced equation
for this reaction.
SOLUTION
Use the Problem-Solving Strategy for Writing and Balancing a Chemical Equation.
1.
Word equation:
silver nitrate + hydrogen sulfide→silver sulfide + nitric acid
2.
Skeleton equation:
AgNO3 + H2S→Ag2S + HNO3unbalanced
3.
Balance:
a.
Ag and H are unbalanced.
b.
Place a 2 in front of AgNO3 to balance Ag:
2 AgNO3 + H2S→Ag2S + HNO3unbalanced
c.
H and NO3 − are still unbalanced. Balance by placing a 2 in front of HNO3:
2 AgNO3 + H2S→Ag2S + 2 HNO3balanced
In this example, N and O atoms are balanced by balancing the NO3 − ion as a unit.
d.
The other atoms remain balanced.
e.
Check: Each side has two Ag, two H, and one S atom. Also, each side has two NO3 − ions.
When aluminum hydroxide is mixed with sulfuric acid, the products are aluminum sulfate and water. Write a
balanced equation for this reaction.
SOLUTION
Use the Problem-Solving Strategy for Writing and Balancing a Chemical Equation.
1.
Word equation:
aluminum hydroxide + sulfuric acid→aluminum sulfate + water
2.
Skeleton equation:
AlOH3 + H2SO4→Al2SO43 + H2Ounbalanced
3.
Balance:
a.
All elements are unbalanced.
b.
Balance Al by placing a 2 in front of Al(OH)3. Treat the unbalanced SO42 − ion as a unit and
balance by placing a 3 in front of H2SO4:
2 AlOH3 + 3 H2SO4→Al2SO43 + H2Ounbalanced
Balance the unbalanced H and O by placing a 6 in front of H2O:
2 AlOH3 + 3 H2SO4→Al2SO43 + 6 H2Obalanced
c.
The other atoms remain balanced.
d.
Check: Each side has 2 Al, 12 H, 3 S, and 18 O atoms.
When the fuel in a butane gas stove undergoes complete combustion, it reacts with oxygen to form carbon
dioxide and water. Write the balanced equation for this reaction.
SOLUTION
Use the Problem-Solving Strategy for Writing and Balancing a Chemical Equation.
1.
Word equation:
butane + oxygen→carbon dioxide + water
2.
Skeleton equation:
C4H10 + O2→CO2 + H2Ounbalanced
3.
Balance:
a.
All elements are unbalanced.
b.
Balance C by placing a 4 in front of CO2:
C4H10 + O2→4 CO2 + H2Ounbalanced
Balance H by placing a 5 in front of H2O:
C4H10 + O2→4 CO2 + 5 H2Ounbalanced
Oxygen remains unbalanced. The oxygen atoms on the right side are fixed because 4
CO2 and 5 H2O are derived from the single C4H10 molecule on the left. When we try to
balance the O atoms, we find that there is no whole number that can be placed in front of
O2 to bring about a balance, so we double the coefficients of each substance and then balance
the oxygen:
2 C4H10 + 13 O2→8 CO2 + 10 H2Obalanced
c.
The other atoms remain balanced.
d.
Check: Each side now has 8 C, 20 H, and 26 O atoms.
Combustion Reaction
In a combustion reaction a substance reacts with oxygen to form oxides. The most common combustion reactions are
between hydrocarbons (compounds composed of hydrogen and carbon) and oxygen to form carbon dioxide and water.
These reactions are commonly used to generate energy from fossil fuels. Here are some examples of combustion
reactions of hydrocarbons that may be familiar to you:
Methane or CH4 is the fuel most often used in gas ovens and stoves. It is also the fuel generally used to power Bunsen
burners in the chemistry laboratory.
CH4g + 2 O2g→CO2g + 2 H2Og
(Notice that the water is generally vaporized.)
Octane or C8H18 is one of the major components of the gasoline used in automobiles.
2 C8H18l + 25 O2g→16 CO2g + 18 H2Og
Ethanol or C2H5OH is sometimes added to gasoline as an alternative to fossil fuels. It burns more cleanly than
hydrocarbons.
C2H5OHl + 3 O2g→2 CO2g + 3 H2Og
(Remember to account for the oxygen in the ethanol when balancing the equation.)
Notice that one of the products of the combustion of carbon-based fuels is carbon dioxide, which contributes to climate
change as described in Section 8.5.
Combustion reactions are not limited to hydrogen or carbon compounds. Two by-products of the combustion of fossil
fuels are oxides of nitrogen and sulfur. Often fossil fuels contain sulfur, which can also react with the oxygen when it is
exposed to the high temperatures in a gasoline engine. Atmospheric nitrogen will form an oxide as well under these
conditions. The reactions shown here could be classified as synthesis reactions or combustion reactions.
Ss + O2g→SO2gN2g + 2 O2g→2 NO2g
Decomposition Reaction
In a decomposition reaction, a single substance is decomposed, or broken down, to give two or more different
substances. This reaction may be considered the reverse of combination. The starting material must be a compound, and
the products may be elements or compounds. The general form of the equation is
AB→A + B
Predicting the products of a decomposition reaction can be difficult and requires an understanding of each individual
reaction. Heating oxygen-containing compounds often results in decomposition. Here are some reactions that fall into this
category:
2 HgOs→Δ2 Hgl + O2g2 PbO2s→Δ2 PbOs + O2gCaCO3s→ΔCaOs + CO2g2 NaHCO3s→ΔNa2CO3s + H2Og + CO2g
2 KClO3s→Δ2 KCls + 3 O2g2 NaNO3s→Δ2 NaNO2s + O2g2 H2O2l→Δ2 H2Ol + O2g
Note that a driving force in many of these reactions is the production of a gas.
Single-Displacement Reaction
In a single-displacement reaction, one element reacts with a compound to replace one of the elements of that
compound, yielding a different element and a different compound. The general forms of the equation follow.
Charles D. Winters/Science Source Images
When pieces of zinc metal are placed in hydrochloric acid, hydrogen bubbles form immediately.
If A is a metal, A will replace B to form AC, provided that A is a more reactive metal than B.
A + BC→B + AC
If A is a halogen, it will replace C to form BA, provided that A is a more reactive halogen than C.
A + BC→C + BA
A brief activity series of selected metals (and hydrogen) and halogens is shown in TABLE 8.2. This series is
listed in descending order of chemical activity, with the most active metals and halogens at the top. Many
chemical reactions can be predicted from an activity series because the atoms of any element in the series will
replace the atoms of those elements below it. For example, zinc metal will replace hydrogen from a
hydrochloric acid solution. But copper metal, which is below hydrogen on the list and thus less reactive than
hydrogen, will not replace hydrogen from a hydrochloric acid solution. Here are some reactions that fall into
this category:
CHECK YOUR UNDERSTANDING
8.5 Single-Displacement Reactions
WileyPLUS
1.
reactions with metals as A:
Zns + 2 HClaq→H2g + ZnCl2aq2 Als + 3 H2SO4aq→3 H2g + Al2SO43aq2 Nas + 2 H2Ol→H2g + 2 N
aOHaqCas + 2 H2Ol→H2g + CaOH2aq3 Fes + 4 H2Ogsteam→4 H2g + Fe3O4sFes + CuSO4aq→Cus
+ FeSO4aqCus + 2 AgNO3aq→2 Ags + CuNO32aq
2.
reactions with a halogen as A:
Cl2g + 2 NaBraq→Br2l + 2 NaClaqCl2g + 2 KIaq→I2s + 2 KClaq
A common chemical reaction is the displacement of hydrogen from water or acids. This reaction is a good
illustration of the relative reactivity of metals and the use of the activity series. For example,

K, Ca, and Na displace hydrogen from cold water, steam (H2O), and acids.

Mg, Al, Zn, and Fe displace hydrogen from steam and acids.

Ni, Sn, and Pb displace hydrogen only from acids.

Cu, Ag, Hg, and Au do not displace hydrogen.
Double-Displacement Reaction
In a double-displacement reaction, two compounds exchange partners with each other to produce two
different compounds. The general form of the equation is
AB + CD→AD + CB
This reaction can be thought of as an exchange of positive and negative groups, in which A combines with D
and C combines with B. In writing formulas for the products, we must account for the charges of the
combining groups.
It’s also possible to write an equation in the form of a double-displacement reaction when a reaction has not
occurred. For example, when solutions of sodium chloride and potassium nitrate are mixed, the following
equation can be written:
NaClaq + KNO3aq→NaNO3aq + KClaq
When the procedure is carried out, no physical changes are observed, indicating that no chemical reaction has
taken place.
A double-displacement reaction is accompanied by evidence of reaction such as:
1.
The evolution of heat
2.
The formation of an insoluble precipitate
3.
The production of gas bubbles
Let’s look at some of these reactions more closely: The production of a molecule of water from an H+ and an
OH− ion is accompanied by a release of heat, which can be detected by touching the reaction container or by
using a thermometer.
HClaq + NaOHaq→NaClaq + H2Ol + heatH2SO4aq + BaOH2aq→BaSO4s + 2 H2Ol + heatCuOs + 2 HNO3aq
→CuNO32aq + H2Ol + heatCaOs + 2 HClaq→CaCl2aq + H2Ol + heat
CHECK YOUR UNDERSTANDING
8.6 Double-Displacement Reactions
WileyPLUS
Formation of a solid The solubilities of the products can be determined by consulting the solubility table
in Appendix 4 to see whether one or both of the products are insoluble in water. A solid is indicated by placing
(s) after its formula in the equation.
BaCl2aq + 2 AgNO3aq→2 AgCls + BaNO32aqPbNO32aq + 2 KIaq→PbI2s + 2 KNO3aq
Formation of a gas A gas such as HCl or H2S may be produced directly, as in these two examples:
H2SO4aq + NaCls→NaHSO4s + HClg2 HClaq + ZnSs→ZnCl2aq + H2Sg
Richard Megna/Fundamental Photographs
A double-displacement reaction results from pouring a clear, colorless solution of Pb(NO3)2 into a clear, colorless
solution of KI, forming a yellow solid of Pbl2.
Pb(NO3)2 + 2Kl → Pbl2 + 2KNO3.
A gas can also be produced indirectly. Some unstable compounds formed in a double-displacement reaction,
such as H2CO3, H2SO3, and NH4OH, will decompose to form water and a gas:
2 HClaq + Na2CO3aq→2 NaClaq + H2CO3(aq)→2 NaClaq + H2O(l)+CO2(g)2 HNO3aq + K2SO3aq→2 KNO3
aq + H2SO3(aq)→2 KNO3aq + H2O(l)+SO2(g)NH4Claq + NaOHaq→NaClaq + NH4OH(aq)→NaClaq + H2O(
l)+NH3(g)
Reactions are either exothermic or endothermic. Exothermic reactions liberate heat; endothermic
reactions absorb heat. In an exothermic reaction, heat is a product and may be written on the right side of the
equation for the reaction. In an endothermic reaction, heat can be regarded as a reactant and is written on the
left side of the equation. Here are two examples:
H2g + Cl2g→2 HClg + 185 kJexothermicN2g + O2g + 181 kJ→2 NOgendothermic
The quantity of heat produced by a reaction is known as the heat of reaction. The units used can be kilojoules
or kilocalories. Consider the reaction represented by this equation:
Cs + O2g→CO2g + 393 kJ
When the heat released is expressed as part of the equation, the substances are expressed in units of moles.
Thus, when 1 mol (12.01 g) of C combines with 1 mol (32.00 g) of O2, 1 mol (44.01 g) of CO2 is formed and 393
kJ of heat are released. In this reaction, as in many others, the heat energy is more useful than the chemical
products.
Aside from relatively small amounts of energy from nuclear processes, the sun is the major provider of energy
for life on Earth. The sun maintains the temperature necessary for life and also supplies light energy for the
endothermic photosynthetic reactions of green plants. In photosynthesis, carbon dioxide and water are
converted to free oxygen and glucose:
6 CO2 + 6 H2O + 2519 kJ→C6H12O6glucose + 6 O2
Nearly all of the chemical energy used by living organisms is obtained from glucose or compounds derived
from glucose.
Glucose
C6H12O16
Paul Damien/National Geographic Stock
This cornfield is a good example of the endothermic reactions happening through photosynthesis in plants.
The major source of energy for modern technology is fossil fuel—coal, petroleum, and natural gas. The energy
is obtained from the combustion (burning) of these fuels, which are converted to carbon dioxide and water.
Fossil fuels are mixtures of hydrocarbons, compounds containing only hydrogen and carbon.
Natural gas is primarily methane, CH4. Petroleum is a mixture of hydrocarbons (compounds of carbon and
hydrogen). Liquefied petroleum gas (LPG) is a mixture of propane (C3H8) and butane (C4H10).
Here are some examples:
CH4g + 2 O2g→CO2g + 2 H2Og + 890 kJC3H8g + 5 O2g→3 CO2g + 4 H2Og + 2200 kJ
Richard Megna/Fundamental Photographs
Combustion of Mg inside a block of dry ice (CO2) makes a glowing lantern.
The combustion of these fuels releases a tremendous amount of energy, but reactions won’t occur to a
significant extent at ordinary temperatures. A spark or a flame must be present before methane will ignite. The
amount of energy that must be supplied to start a chemical reaction is called the activation energy. In an
exothermic reaction, once this activation energy is provided, enough energy is then generated to keep the
reaction going.
Be careful not to confuse an exothermic reaction that requires heat (activation energy) to get it started with an
endothermic process that requires energy to keep it going. The combustion of magnesium, for example, is
highly exothermic, yet magnesium must be heated to a fairly high temperature in air before combustion
begins. Once started, however, the combustion reaction goes very vigorously until either the magnesium or the
available supply of oxygen is exhausted. The electrolytic decomposition of water to hydrogen and oxygen is
highly endothermic. If the electric current is shut off when this process is going on, the reaction stops instantly.
The relative energy levels of reactants and products in exothermic and endothermic processes are presented
graphically in FIGURES 8.2 and 8.3.
@Science Photo Library/AGE
FIGURE 8.2 An endothermic process occurs when a cold pack containing an ampule of solid NH4CI is broken, releasing
the crystals into water in the pack. The dissolving of NH4CI in water is endothermic, producing a salt solution that is
cooler than the surroundings. This process is represented graphically to show the energy changes between the reactants
and the products.
Tom Pantages
FIGURE 8.3 Exothermic reaction between KCIO3 and sugar. A sample of KCIO3 and sugar is well mixed and placed on
a fireproof pad. Several drops of concentrated H2SO4 are used to ignite the mixture.
CHECK YOUR UNDERSTANDING
8.13 Activation Energy and Heat of Reaction
WileyPLUS
Examples of endothermic and exothermic processes can be easily demonstrated. In Figure 8.2, the products
are at a higher potential energy than the reactants. Energy has therefore been absorbed, and the reaction is
endothermic. An endothermic reaction takes place when you apply a cold pack to an injury. When a cold pack
is activated, ammonium chloride (NH4Cl) dissolves in water. For example, temperature changes from 24.5°C
to 18.1°C result when 10 g of NH4Cl are added to 100 mL of water. Energy, in the form of heat, is taken from
the immediate surroundings (water), causing the salt solution to become cooler.
EXAMPLE 8.11
Label the graph in Figure 8.2 showing the specific reactants and products as well as the heat of reaction for the reaction.
SOLUTION
For Figure 8.2 the chemical equation is NH4Cls + H2Ol→NH4 + aq + Cl − aq + H2Ol. The reactants
are NH4Cls + H2Ol and the products are NH4 + aq + Cl − aq + H2Ol. The graph is labeled and shown here:
© 2014 John Wiley & Sons, Inc. All right reserved.
8.1 The Chemical Equation

A chemical equation is shorthand for expressing a chemical change or reaction.

In a chemical reaction atoms are neither created nor destroyed.

All atoms in the reactants must be present in the products.
KEY TERMS
reactants
products
chemical equation
law of conservation of mass
8.2 Writing and Balancing Chemical Equations

To balance a chemical equation:

Identify the reaction.

Write the unbalanced (skeleton) equation.

Balance the equation:

Count the number of atoms of each element on each side and determine which need
to be balanced.

Balance each element (one at a time) by placing whole numbers (coefficients) in
front of the formulas containing the unbalanced element:


Begin with metals, then nonmetals, and then H and O.
Check the other elements to see if they have become unbalanced in the process of
balancing the chosen element. If so, rebalance as needed.


Do a final check to make sure all elements are balanced.
The following information can be found in a chemical equation:

Identity of reactants and products

Formulas for reactants and products

Number of formula units for reactants and products

Number of atoms of each element in the reaction

Number of moles of each substance
KEY TERMS
balanced equation
8.3 Why Do Chemical Reactions Occur?

Formation of a solid

Formation of a gas

Formation of water

Transfer of electrons
8.4 Types of Chemical Equations

Synthesis reactions
A + B→AB

Decomposition reactions
AB→A + B

Single-displacement reactions:

In which A is a metal

In which A is a halogen
A + BC→C + BA
AB + CD→AD + CB

Double-displacement reactions

Evidence for a chemical reaction:

Evolution of heat

Formation of an insoluble precipitate

Production of a gas
KEY TERMS
synthesis reaction
combustion reaction
decomposition reaction
single-displacement reaction
double-displacement reaction
8.5 Heat in Chemical Reactions

A + BC→B + AC
Exothermic reactions release heat.

Endothermic reactions absorb heat.

The amount of heat released or absorbed in a chemical reaction is called the heat of reaction:

It can be written as a reactant or product in the chemical equation.

Units are joules (J) or kilojoules (kJ).

The major source of energy for modern technology is hydrocarbon combustion.

To initiate a chemical reaction, activation energy is required:

In exothermic reactions, this energy is returned and more is released, which allows the
reaction to continue on its own.

In endothermic reactions, energy must be added to start and continue to be added to sustain
the reaction.
KEY TERMS
exothermic reactions
endothermic reactions
heat of reaction
hydrocarbons
activation energy