1. Chapter 8:
Balancing Chemical
Reactions

2. Objectives
by the end of this chapter, you should be able to:
 Write equations describing chemical
reactions using appropriate symbols
 Write balanced chemical equations when
given the names or formulas of the
reactants and products in a chemical
reaction
 Define chemical
equation, catalyst, aqueous
solution, skeleton
equation, coefficients, and balanced
equation

3. Chemical Equations
 Chemical equations – using chemical formulas
to write equations
 Reactants (left side of arrow)
 Products (right side of arrow)
 Arrow means yields, gives, or reacts to produce
 Reactants  Products
 Catalyst (a substance that speeds up the rate of
the reaction but that is not used up in the
reaction) should be written above the arrow
C6H12O6 (s) + CO2 (g) O2 (g) + H2O (l) + energy

4.  Can indicate the physical state of a
substance in the equation by putting a
symbol after each formula
Solid – (s)
Liquid – (l)
Gas – (g)
Aqueous solution: a substance dissolved
in water – (aq)
 Refer to the Table on page 266 for
explanations of other symbols used in
chemical equations

5. Skeleton Equation
 A chemical equation that does not
indicate the relative amounts of the
reactants and products involved in the
reaction
 Examples:
a. Fe(s) + O2(g)  Fe2O3(s)
b. H2O2(aq)  H2O(l) + O2(g)
Manganese(IV) oxide is a catalyst, so
MnO2 should be written above the arrow.

6. Write a Skeleton Equation
 Solid sodium hydrogen carbonate reacts
with hydrochloric acid to produce aqueous
sodium chloride, water, and carbon
dioxide gas. Include appropriate symbols.
1. Write the correct formula for each
substance in the reaction.
2. Separate the reactants from the
products.
3. Indicate the physical state of each
substance.

7. Answer
 NaHCO3(s) + HCl(aq)  NaCl(aq) +
H2O(l) + CO2(g)

8. A Balanced Equation
 An equation that gives the correct quantity
of each reactant and product
 Coefficients (numbers placed in front of
the symbols) are used
 Must obey law of conservation of mass:
Each side of the equation has the same
number of atoms of each element
 Example: A standard bicycle is composed
of one frame, two wheels, one
handlebar, and two pedals
F + 2W + H + 2P  FW2HP2

9. Rules for Balancing Equations
 1. Determine the correct formulas for all of the
reactants and products. In some cases, also list
in parenthesis the physical state of matter.
 2. List reactants on the left side of the arrow
(Use plus sign (+) when there is more than one
reactant)
 3. List the products on the right side of the
arrow (Use plus sign (+) when there is more
than one product)
 4. Steps 1-3 provide a skeleton equation.
(Note: Sometimes this is also the balanced
equation. For example: C + O2  CO2)

10.  5. Count the number of atoms of each element
in the reactants and products. For simplicity, a
polyatomic ion appearing unchanged on both
sides of the arrow is counted as a single unit.
 6. Balance the elements one at a time by using
coefficients. DO NOT CHANGE THE
SUBSCRIPTS.
 7. Check each atom or polyatomic ion to be
sure that the equation is balanced.
 8. Make sure that all the coefficients are in the
lowest possible ratio that balances.

11. Problem: Hydrogen and oxygen react to
form water. Write a balanced equation.
Reactants: H2(g) + O2(g)
Products: H2O(l)
H2(g) + O2(g)  H2O(l)
Count the atoms
Left side Right side
H–2 H–2
O–2 O–1
 Use coefficient to get 2 oxygen on the right side:
H2(g) + O2(g)  2 H2O(l)
Left side Right side
H–2 H–4
O–2 O–2

12.  Need 4 hydrogen atoms, so place a
coefficient of 2 in front of H2
2H2(g) + O2(g)  2 H2O(l)
Left side Right side
H–4 H–4
O–2 O–2
 Check number of atoms
 Check that the coefficients are in the
lowest possible ratio
 The equation is balanced

13. Problems
 1. Balance the following equations:
a. SO2 + O2  SO3
b. Al + O2  Al2O3
 2. Rewrite the word equation as a
balanced chemical equation:
Aluminum sulfate and calcium
hydroxide react to form aluminum
hydroxide and calcium sulfate.

14. Answers
1a) 2SO2 + O2  2SO3
1b) 4Al + 3O2  2Al2O3
2) Word equation to balanced chemical
equation:
Al2(SO4)3 + 3Ca(OH)2  2Al(OH)3 + 3CaSO4

15. Practice Problems
 1. __NaCl + __BeF2 __NaF + __BeCl2
 2. __FeCl3 + __Be3(PO4) 2 __BeCl2 + __FePO4
 3. __AgNO3 + __LiOH __AgOH + __LiNO3
 4. __CH4 + __O2 __CO2 + __H2O
 5. __Mg + __Mn2O3 __MgO + __Mn

16. Types of Chemical Reactions
Objectives:
 1. Identify a reaction as
combination, decomposition, single-
replacement, double-replacement, or
combustion
 2. Predict the products of
combination, decomposition, single-
replacement, double-replacement, and
combustion reactions

17. Classifying Reactions
 For combination (synthesis: combination of parts
into a whole) and decomposition, compare the
number of reactants and products
 For combustion, check for oxygen (O2)
 For single- and double-replacement, look for a
cation swap or the formation of a precipitate
 Not all chemical reactions fit uniquely into only
one of these classes

18. Combination Reactions
(AKA- Synthesis Reaction)
 Two or more substances combine to form
a single substance
A + B AB
 Reactants are usually either two elements
or two compounds
 The product is always a compound (Can
be an ionic compound or a molecular
compound)

19. Decomposition Reactions
 A single compound is broken down into two or
more products
AB A + B
 The products can be any combination of
elements and compounds
 Most decomposition reactions require energy in
the form of heat, light, or electricity
 Extremely rapid decomposition reactions that
produce gaseous products and heat are often
the cause of explosions

20. Single-replacement Reactions
(Also called single-displacement reactions)
 One element replaces a second element in a
compound
A + BC AC + B
 Remember: either the anions or cations will
switch with each other. They cannot be paired
together since their charges repel each other.
How do we tell which is which? Use the periodic
table to predict their oxidation numbers.

21.  Whether one metal will displace another
metal from a compound can be
determined by the relative reactivities of
the two metals. (Memorize the symbols
and activity series of metals on page 286.)
 A reactive metal will replace any metal
listed below it in the activity series
 Examples:
Iron will displace copper from a copper
compound in solution.
Magnesium does not replace lithium
from aqueous solutions of their
compounds.

22. The Activity Series of Metals

23. Refer to Table 8.2 on page 217
 Will magnesium displace zinc from a zinc
compound in solution?
 Will magnesium displace silver from a silver
compound in solution?
 Important Note:
1. Metals from lithium to lead will replace
hydrogen from acids.
2. Metals from lithium to sodium will also
replace hydrogen from water.

24. Single-Replacement (cont’d)
 A nonmetal can also replace another nonmetal
from a compound
 This replacement is usually limited to the
halogens (Group 7A):
F2 (most activity)
Cl2 .
Br2 .
I2 (least activity)
 The activity of the halogens decreases as you go
down group 7A on the periodic table

25. Double-replacement Reactions
 Involves an exchange of positive ions between
two reacting compounds
AB + CD AD + CB
 Often characterized by the production of a
precipitate (ppt.-insoluble substance that “falls
out” of a solution)
 Product may be a gas that “bubbles” out of the
mixture
 Product may be a molecular compound, such as
water

26. Combustion Reactions
 An hydrocarbon reacts with oxygen (often
producing energy) with water and carbon
dioxide as products
CxHy + O2 CO2 + H2O
 Commonly involve hydrocarbons (compounds of
hydrogen and carbon)
 The complete combustion of a hydrocarbon
produces carbon dioxide and water
 If the supply of oxygen during a reaction is
insufficient, combustion will be incomplete

27.  During incomplete combustion, elemental
carbon and toxic carbon monoxide may be
additional products
 Reaction between some elements and
oxygen
Example: Both magnesium and sulfur will
burn by reaction with oxygen
 Refer to worksheet handout.
 Identify the combustion reactions.

28. Make a Chemistry Foldable
 1. Fold a sheet of notebook paper to the red
margin line.
 2. Using scissors, cut the folded section into five
equal parts.
 3. Label each section with the name of one of
the five types of reactions.
 4. Open each flap and put in three
characteristics and one example (include
balanced equation).
 5. Write the title : Types of Chemical Reactions
on the top of the sheet. Add your name and
class.
 6. Use the chemistry foldable as a study guide.

29. Predicting Products of a Chemical
Reaction
 Recognize the possible type of reaction that
the reactants can undergo
 Some reactions do not fit any one of the five
general types (Example: redox reactions)
 Oxidation-reduction (redox) reactions will be
discussed during the second semester
OIL RIG – oxidation is the loss of electrons
and reduction is the gain of electrons
LEO the lion says GER – loss of electrons
is oxidation and gain of electrons is reduction

30. Reactions in Aqueous Solution
 Objectives:
1. Write and balance net ionic equations
2. Use solubility rules to predict the
precipitate formed in double replacement
reactions

31. Net Ionic Equations
 Most ionic compounds dissociate, or separate,
into cations and anions when they dissolve in
water.
 Refer to question #21 on the worksheet
handout. Use this equation to answer #22 on
the worksheet handout.
 Write a complete ionic equation that shows
dissolved ionic compounds as their free ions.
 Eliminate ions that do not participate in the
reaction by canceling ions that appear on both
sides of the equation. These are called spectator
ions.

32.  Ions that are not directly involved in a reaction
are called spectator ions.
 Rewrite the equation, leaving out the canceled
spectator ions.
 Balance the atoms and the charges of the ions.
(In this case, the number of atoms and the net
ionic charge on each side of the equation is zero
and it is therefore balanced.)
 A net ionic equation indicates only those
particles that actually take part in the reaction.
 Record your answer to #23 on the worksheet
handout.

33. Practice Problem
 Write a balanced net ionic equation for the following
reaction:
Pb(s) + AgNO3 (aq)  Ag (s) + Pb(NO3)2 (aq)
Answer:
1. The nitrate ion is the spectator ion.
2. The number of atoms balance, but the charges on
the ions do not balance.
3. Place a coefficient 2 in front of Ag+ (aq) to balance
the charges.
4. A coefficient of 2 in front of Ag (s) rebalances the
atoms.
5. Pb(s) + 2Ag+ (aq)  2Ag (s) + Pb2+ (aq) is the
balanced net ionic equation

34. Predicting the Formation of a
Precipitate
 Use the general rules for solubility of ionic
compounds (Table 8.3 on page 227)
 Examples:
1. Sodium nitrite will not form a
precipitate because alkali metal salts and
nitrate salts are soluble (Rules 1 and 2)
2. Rule 3 (Exceptions) indicates that
barium sulfate is insoluble and therefore
will precipitate.

35. Solubility Rules for Ionic Compounds
Compounds Solubility Exceptions
1. Salts of alkali metals Soluble Some lithium
and ammonia compounds
2. Nitrate salts and Soluble Few Exceptions
chlorate salts
3. Sulfate salts Soluble Compounds of Pb, Ag,
Hg, Ba, Sr, and Ca
4. Chloride salts Soluble Compounds of Ag and
some compounds of Hg
and Pb
5. Carbonates, Most are insoluble Compounds of the
phosphates, chromates, alkali metals and of
sulfides, and ammonia
hydroxides

36. Practice Problem
 Identify the precipitate formed and write
the net ionic equation for the reaction of
aqueous potassium carbonate with
aqueous strontium chloride.
1. Write the reactants showing each as
dissociated free ions. Balance the charges.
2. Using solubility rules, look at possible
new pairings of cation and anion that give
an insoluble substance.
3. Eliminate the spectator ions and write
the net ionic equation.

37. Answer
 1. Reactants as dissociated free ions
2K+ (aq) + CO32- (aq) + Sr2+ (aq) + 2Cl- (aq)
Charges must be balanced to equal 0.
 2. Of the two possible combinations, KCl is soluble
(Rules 1 and 4) and SrCO3 is insoluble (Rule 5)
 3. The net ionic equation must be balanced for the
number of atoms of each element and the charges
on the ions.
Sr2+ (aq) + CO32- (aq)  SrCO3 (s)
 Note: Ignore Sample Problem 8-11 on page 228.
There is a textbook error.