UNIT 5
CHEMICAL EQUATIONS &
REACTIONS
Honors Chemistry
All chemical reactions
 have two parts:
 Reactants - the substances you start
with
 Products- the substances you end
up with
 The reactants turn into the products.
 Reactants  Products
In a chemical reaction
 The way atoms are joined is changed
 Atoms aren’t created of destroyed.
 Can be described several ways:
1. In a sentence
Copper reacts with chlorine to form copper (II) chloride.
2. In a word equation
Copper + chlorine  copper (II) chloride
Symbols in equations
 the arrow separates the reactants
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from the products
Read “reacts to form”
The plus sign = “and”
(s) after the formula = solid
(g) after the formula = gas
(l) after the formula = liquid
Symbols used in equations
 (aq) after the formula - dissolved in
water, an aqueous solution.
 used after a product indicates a gas
(same as (g))
 used after a product indicates a
solid (same as (s))
Symbols used in equations

indicates a reversible reaction

heat
   ,   shows that heat is
supplied to the reaction
Pt
   is used to indicate a catalyst
is supplied, in this case, platinum.
What is a catalyst?
 A substance that speeds up a
reaction, without being changed or
used up by the reaction.
 Enzymes are biological or protein
catalysts.
Skeleton Equation
 Uses formulas and symbols to
describe a reaction
 doesn’t indicate how many.
 All chemical equations are sentences
that describe reactions.
Convert these to equations
 Solid iron (III) sulfide reacts with gaseous
hydrogen chloride to form iron (III) chloride
and hydrogen sulfide gas.
 Nitric acid dissolved in water reacts with solid
sodium carbonate to form liquid water and
carbon dioxide gas and sodium nitrate
dissolved in water.
Now, read these:
 Fe(s) + O2(g)  Fe2O3(s)
 Cu(s) + AgNO3(aq) 
Pt
 NO2 (g)
Ag(s) + Cu(NO3)2(aq)
N2(g) + O2(g)
Balanced Equation
 Atoms can’t be created or destroyed
 All the atoms we start with we must
end up with
 A balanced equation has the same
number of each element on both
sides of the equation.
Rules for balancing:
 Assemble, write the correct formulas for all
the reactants and products
 Count the number of atoms of each type
appearing on both sides
 Balance the elements one at a time by adding
coefficients (the numbers in front) - save H
and O until LAST!
 Check to make sure it is balanced.
 Never change a subscript to balance an equation.
 If you change the formula you are describing a different
reaction.
 H2O is a different compound than H2O2
 Never put a coefficient in the middle of a formula
 2 NaCl is okay, Na2Cl is not.
Example
H2 + O2  H2O
R
P
2 H 2
2 O 1
Need twice as much O in the product
Example
H2 + O2 
R
P
2 H 2
2 O 1
Changes the O
2 H2O
Example
H2 + O2 
2 H2O
R
P
2 H 2
2 O 1 2
Also changes the H
Example
H2 + O2 
2 H2O
R
P
2 H 2 4
2 O 1 2
Need twice as much H in the reactant
Example
2 H2 + O2 
2 H2O
R
P
2 H 2 4
2 O 1 2
Recount
Example
2 H2 + O2 
2 H2O
R
P
4 2 H 2 4
2 O 1 2
The equation is balanced, has the same
number of each kind of atom on both sides
Example
2 H2 + O2 
2 H2O
R
P
4 2 H 2 4
2 O 1 2
This is the answer
Not this
Balancing Examples
 _AgNO3 + _Cu  _Cu(NO3)2 + _Ag
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_Mg + _N2  _Mg3N2
_P + _O2  _P4O10
_Na + _H2O  _H2 + _NaOH
_CH4 + _O2  _CO2 + _H2O
Balancing Examples
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2AgNO3 + Cu  Cu(NO3)2 + 2Ag
3Mg +N2  Mg3N2
4P + 5O2  P4O10
2Na + 2H2O  H2 + 2NaOH
CH4 + 2O2  CO2 + 2H2O
Types of Reactions
 There are millions of reactions.
 Can’t remember them all
 Fall into several categories.
 We will learn 5 major types.
 Will be able to predict the products.
 For some, we will be able to predict whether
they will happen at all.
 Will recognize them by the reactants
COMBINATION REACTION
A reaction in which two or more substances
combine to form a single product.
A +B + C  ABC
CaO(s) + SO2(g)  CaSO3(s)
Write and balance
Ca + Cl2 
Fe + O2 
Al + O2 
Remember that the first step is to
write the correct formulas
 Then balance by using coefficients
only
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DECOMPOSITION REACTION
A reaction in which a single compound reacts to
give two or more substances, usually requiring a
raise in temperature.
ABC  A + B + C
2KClO3(s)  2KCl(s) + 3O2(g)
#2 - Decomposition Reactions
 decompose = fall apart
 one reactant falls apart into two or
more elements or compounds.
electricity
 NaCl   Na + Cl2
 CaCO3  CaO + CO2
 Note that energy is usually required
to decompose
#2 - Decomposition Reactions
 Can predict the products if it is a
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binary compound
Made up of only two elements
Falls apart into its elements
electricity
H2O 


HgO 
2H2O(l) →
•Can predict the products if it is a binary compound
•Made up of only two elements
•Falls apart into its elements
Fig. 8-13, p. 215
#2 - Decomposition Reactions
 If the compound has more than two
elements you must be familiar with
page 10 of your packet
 NiCO3 
 KClO3(aq) 

 
Single Displacement (or Replacement)
Reactions
pp. 218, 220
Fig. 8-15, p. 218
Single Displacement (or
Replacement) Reactions
PREDICT THE PRODUCT
1. Ca + HCl
2. ZnBr2 + I2
3. Cu + AgNO3
Answers are on the next slide.
#3 Single Replacement
 Metals replace other metals (and
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hydrogen)
K + AlN 
Zn + HCl 
Think of water as HOH
Metals replace one of the H, combine
with hydroxide.
Na + HOH 
 We can tell whether a reaction
will happen
 Some chemicals are more
“active” than others
 More active replaces less active
 There is a list on page 11 - called
the Activity Series of Metals
 Higher on the list replaces lower.
#3 Single Replacement
 Note the
* concerning Hydrogen
 H can be replaced in acids by everything higher
 Li, K, Ba, Ca, & Na replace H from acids and
water
 Fe + CuSO4 
 Pb + KCl 
 Al + HCl 
#3 - Single Replacement
 What does it mean that Hg and Ag are on
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the bottom of the list?
Nonmetals can replace other nonmetals
Limited to F2 , Cl2 , Br2 , I2 (halogens)
Higher replaces lower.
F2 + HCl 
Br2 + KCl 
Double Displacement (or Replacement)
Reactions
pp. 220, 223
Double Displacement (or
Replacement) Reactions
 PREDICT THE PRODUCT & BALANCE
1. MgSO4 + LiOH
___________
2. Pb(NO3)2 + Na2CO3
3. HNO3 + Ba(OH)2
____________
___________
Answers are on the next slide.
Double Displacement (or
Replacement) Reactions
 ANSWERS
1. MgSO4 + 2 LiOH
Mg(OH)2 + Li2SO4
Exchange cations
2. Pb(NO3)2 + Na2CO3
3. 2 HNO3 + Ba(OH)2
PbCO3 + 2 NaNO3
Ba(NO3)2 + 2 H2O
Double Replacement
 Two things replace each other.
 Reactants must be two ionic compounds or
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acids.
Usually in aqueous solution
NaOH + FeCl3 
The positive ions change place.
NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1
NaOH + FeCl3 Fe(OH)3 + NaCl
Double Replacement
 Has certain “driving forces”
 Will only happen if one of the products:
 doesn’t dissolve in water and forms a solid (a
“precipitate”), or
 is a gas that bubbles out, or
 is a covalent compound (usually water).
Oxidation and Reduction
Reactions
 Early chemists saw oxidation as the
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combination of a material with oxygen
Today, many of these reactions may not even
involve oxygen
Redox currently says that electrons are
transferred between reactants
Always occur simultaneously.
Called redox reactions.
Oxidation and Reduction
Reactions
 Oxidation – refers to a loss of electrons, or
gain of oxygen
 Reduction – refers to a gain of electrons, or
loss of oxygen
 Oxidizing agent – the substance being
reduced
 Reducing agent – the substance being
oxidized
Oxidation and Reduction
Reactions
 Loss of Electrons is Oxidation
 Gain of Electrons is Reduction
LEO SAYS GER
Rules For Assigning Oxidation States
Oxidation numbers are used to keep track of how many electrons are lost or
gained by each atom
Balancing Redox Reactions
 If the oxidation number changes, that element
has undergone oxidation or reduction
 Reaction as a whole must be redox
 Half-reaction – equation showing just the
oxidation or just the reduction
Balancing Redox Reactions
(in an acid)
+4 -2
+7 -2
+6 -2
+2
SO2 + MnO4- →SO42- + Mn2+
Divide the equation into half-reactions.
SO2 → SO42Oxidation
MnO4- → Mn2+
Reduction
Balance all atoms except O and H
Balancing Redox Equations
Balance O by adding H2O.
2H2O + SO2 →SO42-
MnO4- →Mn2+ + 4H2O
Balance H by adding H+.
2H2O + SO2 → SO42- + 4H+
8H+ + MnO4- → Mn2+ + 4H2O
Balance net charge by adding e-.
2H2O + SO2 →SO42- + 4H+ + 2e5e- + 8H+ + MnO4- →Mn2+ + 4H2O Step 6.
Make e- gain equal e- loss; then add halfreactions.
5(2H2O + SO2 →SO42- + 4H+ + 2e-)
2(5e- + 8H+ + MnO4- →Mn2+ + 4H2O)
______________________________________
10H2O + 5SO2 + 10e- + 16H+ + 2MnO4→2Mn2+ + 8H2O + 5SO42- + 20H+ + 10e-
Cancel anything that's the same on
both sides.
2H2O + 5SO2 +2MnO4- →2Mn2+ + 5SO42- + 4H+
Balancing Redox Reactions
(in a base)
Follow steps 1 – 7 then:
8. Identify the number of proton (H+) in the acidic
answer. Add the same number of OH+ ions to
BOTH sides of the equation.
9. If H+ and OH+ appear on the same side of the
equation, they will react in a 1:1 ratio to form H2O.
10. Cancel out water molecules that appear on both
sides of the equation.
11. Check that all is balanced
PRECIPITATION REACTION
A reaction where an insoluble solid is formed during a
reaction between two aqueous solutions.
(aq) + (aq)  (aq) + (s)
2KI(aq) + Pb(NO3)2(aq)  2KNO3(aq) + PbI2(s)
Exchange cations
Ionic Equations
 Many reactions occur in water- that is, in aqueous
solution
 Many ionic compounds “dissociate”, or separate,
into cations and anions when dissolved in water
 AgNO3 + NaCl  AgCl + NaNO3
1. this is the full equation
2. Write it as an ionic equation
3. Eliminating ions not directly involved
(spectator ions) = net ionic equation
Molecular, Ionic, Net Ionic
Equations
How Do We Predict Solubility?
NEUTRALIZATION REACTION
A reaction between an acid and a base which
results in the production of a salt and
water.
HA + BOH  (metal/nonmetal) + H2O
HNO3(aq) + KOH(aq)  KNO3(aq) + H2O(l)
COMBUSTION REACTION
 Means “add oxygen”
 A compound composed of only C, H, and maybe O is
reacted with oxygen
 If the combustion is complete, the products will be
CO2 and H2O.
 If the combustion is incomplete, the products will be
CO (possibly just C) and H2O.
 A reaction of a substance with oxygen, usually the
rapid release of heat produces a flame.
CH + O2  CO2 + H2O
2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)
Combustion
 Many times in a combustion reaction,
heat energy is given off. In chemical terms
this is called an exothermic reaction.
 Thermochemistry is field of chemistry
which studies the transfer of heat in a
reaction.
 The thermodynamic equation
representing this exothermic reaction is:
2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g) +
heat (in Joules)
Examples
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C4H10 + O2  (assume complete)
C4H10 + O2  (incomplete)
C6H12O6 + O2  (complete)
C8H8 +O2  (incomplete)
GAS FORMATION REACTIONS
 A reaction that produces a gas from
reactants not in the gaseous state.
2 HCl (aq) + ZnS (s)  ZnCl2 (aq) + H2S (g)
Zn (s) + 2 HCl (aq)  ZnCl2 (aq) + H2 (g)
 Many gas formation reactions involve two steps,
first the double displacement reaction then the
decomposition reaction of an unstable substance.
Na2CO3 + 2HCl  2 NaCl + H2CO3
H2CO3  CO2 + H2O
 Besides carbonic acid (H2CO3), sulfurous acid
(H2SO3) also decomposes into SO2 and water.
COMMON GAS FORMATION REACTIONS YOU
SHOULD REMEMBER
 NH4OH → NH3 (g) + H2O (l)
 H2CO3 → CO2 (g) + H2O (l)
 H2SO3 → SO2 (g) + H2O (l)
Remember an equation...
 Describes a reaction
 Must be balanced in order to follow the Law
of Conservation of Mass
 Can only be balanced by changing the
coefficients.
 Has special symbols to indicate physical
state, and if a catalyst or energy is required.
Table 8-3, p. 223
Examples
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H2 + O 2 
H2 O 
Zn + H2SO4 
HgO 
KBr +Cl2 
AgNO3 + NaCl 
Mg(OH)2 + H2SO3 