Nature of Chemical Reactions
Types of Chemical Reactions
Synthesis Reactions
Decomposition Reactions
Single-Replacement Reactions
Double-Replacement Reactions
Combustion Reactions
Identifying Oxidation-Reduction Reactions
Nature of Chemical Reactions Objectives
1. Define chemical Reactions.
2. Describe four indications that a chemical reaction
has occurred.
3. State the law of conservation of mass and describe
how it relates to a chemical reaction.
4. Identify and use the common symbols used in
writing chemical equations.
5. Translate chemical equations into word equations.
6. Translate word equations into chemical equations.
7. Balance chemical equations.
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Characteristics of Chemical Reactions
A chemical reaction is a process
in which the chemical and
physical properties of the
original substance changes as
new substances with different
physical and chemical
properties are formed.
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Characteristics of Chemical Reactions
In any chemical reaction, there are
always two kinds of substances; the
substances that are present before the
change, called the reactants, and the
substances that are formed by the
change, called the products.
Reactants → Products
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Identifying Reactants and Products
Identify the reactants and products
in the following equation:
2Al(s) + 3Cl2(g) → 2AlCl3(s)
reactant(s)
products(s)
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Al, Cl2
AlCl3
Indications of a Chemical Reaction
What are the four observations that generally
indicate that a chemical reaction has occurred?
1. Production of a gas.
2. Formation of a precipitate. (an insoluble solid
that forms in an aqueous reaction)
3. Change in energy.
Endothermic reaction – energy is absorbed
Exothermic Reaction – energy is released
4. Change in color or odor.
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Law of Conservation of Mass
Chemical equations are balanced to satisfy
the Law of Conservation of Mass
The Law of Conservation of Mass states that
mass cannot be created or destroyed by
ordinary physical or chemical means.
This means that the total mass of the
reactions must equal the total mass of the
products.
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Law of Conservation of Mass
Remember that atoms don’t
change in a chemical reaction; they
just rearrange. For a chemical
equation to accurately represent a
reaction, the same number of each
kind of atom must be on the left
side of the arrow as are on the right
side.
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Symbols Used in Chemical Equations
+ Used to separate two reactants or two
products
→ “yields”, separates reactants from products
⇄ Used in place of a → for reversible reactions
(s) Designates a reactant or product in the solid
state
(l) Designates a reactant or product in the liquid
state
(aq) Designates an aqueous solution, the
substance is dissolved in water.
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Symbols Used in Chemical Equations
(g) Designates a reactant or product in the
gaseous state
Indicates that heat is supplied to the
reaction.
A formula written above or below the yield
sign indicates its use as a catalyst (in this
example, platinum).
A catalyst is a substance that speeds up a reaction
without being used up itself or permanently
changed.
In biology, a catalyst is commonly called an enzyme.
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Writing Sentences to Describe Balanced Equations
Sentences can be written to describe
(interpret) what is happening in a balanced
equation.
Example 1.
2NaOH(aq) + CO2(g) → Na2CO3(s) + H2O(l)
Aqueous sodium hydroxide reacts with
carbon dioxide gas to produce solid sodium
carbonate and water.
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Writing Sentences to Describe Balanced Equations
Example 2.
2Al(s) + 3Cl2(g) → 2AlCl3(s)
Aluminum metal reacts with chlorine gas to
produce solid aluminum chloride.
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You Try It
Write sentences to describe each of the
following equations. (Do not worry about the
coefficients at this time.)
i. 2H2(g) + O2(g) → 2H2O(l)
Hydrogen gas reacts with oxygen gas to
produce water.
ii. Zn(s) + CuCl2(aq) → ZnCl2(aq) + Cu(s)
Zinc metal reacts with aqueous copper(II)
chloride to produce aqueous zinc
chloride and copper metal.
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Counting Atoms in Chemical Compounds
Before we can balance an equation, we must first make
sure that everyone can count the atoms present in a
compound. Here are some examples.
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KClO3 ___K, ___Cl, ___O
Mg(OH)2 ___Mg, ___O, ___H
Al2(SO4)3 ___Al, ___S, ___O
2Al(NO3)3 ___Al, ___N, ___O
1 1 3
1 2 2
2 3 12
2 6 18
You Try It
Determine the number of each atom present
in each of the following.
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AlPO3 ___Al, ___P, ___O
Ba(NO3)2 ___Ba, ___N, ___O
HC2H3O2 ___H, ___C, ___O
3Fe3(PO4)2 ___Fe, ___P, ___O
1 1 3
1 2 6
4 2 2
9 6 24
Identifying Balanced Chemical Equations
An equation in which the number of atoms of
each element is the same on both sides of
the equation is called a balanced chemical
equation.
Pb(NO3)2 + 2KI → PbI2 + 2KNO3
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Reactant Side Element Product Side
Pb
N
O
K
I
1 1 2 2 6 6
2 2 2 2
You Try It
Classify each of the following equations as
balanced or unbalanced.
a. 2H2O2(aq) → 2H2O(l) + O2(g)
b. Mg(s) + HCl(aq) → MgCl2(aq) + H2(g)
c. Al4C3(s) + H2O(l) → CH4(g) + Al(OH)3(s)
d. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
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balanced
unbalanced
unbalanced
balanced
Balancing Chemical Equations
The easiest equations to balance are the
ones in which the skeleton equation is
already written.
Example 1. Cl2 + Na → NaCl
Add coefficients to balance the equation.
NEVER change subscripts.
What coefficient must be added in front of
NaCl in order to balance the chlorine?
What coefficient must be added in front of
Na in order to balance the sodium?
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2 2
Balancing Chemical Equations
Example 2.
Mg + HCl → MgCl2 + H2
What coefficient must be added in front of
HCl in order to balance the chlorine and the
hydrogen?
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2
Balancing Chemical Equations
Example 3.
Fe + O2 → Fe2O3
Although your first instinct might be to balance the
Fe first, you should actually balance the O first.
We have 2 O’s on the reactant side and 3 O’s on the
product side. What is the least common multiple of 2
and 3?
What number can we put in front of O2 to obtain 6?
What number can we put in front of Fe2O3 to obtain
6?
What coefficient must be added in front of Fe to
balance the Fe? Back to main menu
2 3 4
6
You Try It
Balance the following chemical equations.
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2 i. Cl2 + KI KCl + I2
ii. Li + O2 Li2O
iii. HC2H3O2 + CaCO3 → Ca(C2H3O2)2 + CO2 + H2O
iv. C2H6O(l) + O2(g) → CO2(g) + H2O(g)
v. Al(s) + HCl(aq) → AlCl3(aq) + H2(g)
vi. Pb(NO3)2 + K2CrO4 PbCrO4 + KNO3
vii C3H6 + O2 CO2 + H2O
viii. Al(OH)3 Al2O3 + H2O
ix. HCl + Fe2O3 FeCl3 + H2O
2
4 2
2
2 3 3
2 6 2 3
2
2 6 6 9
2 3
2 6 3
Balancing Chemical Equations
If the skeleton equation is not written for you,
you must write your own.
Example 1. When an electric current is
passed through water, the water molecules
break down to produce hydrogen and oxygen.
Bubbles of each gas are evidence of the
reaction.
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→ H2(g) H2O(aq) + O2(g) 2 2
Balancing Chemical Equations
Example 2. Solutions of aluminum sulfate
and barium chloride react to produce solid
barium sulfate and aqueous aluminum
chloride.
aluminum sulfate = Al3+, SO42- = Al2(SO4)3
barium chloride = Ba2+, Cl- = BaCl2
barium sulfate = Ba2+, SO42- = BaSO4
aluminum chloride = Al3+, Cl- = AlCl3
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main menu
→ BaSO4(s) Al2(SO4)3(aq) + BaCl2(aq)
2
3
+ AlCl3(aq)
3
You Try It
Write a balanced equation for each of the
following reactions.
a. When magnesium and oxygen react, the
product is solid magnesium oxide.
b. When nitrogen and hydrogen react, the
product is ammonia.
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N2(g) + H2(g) → NH3(g)
2 Mg(s) + O2(g) → MgO(s)
3 2
2
You Try It
Write a balanced equation for each of the
following reactions.
c. Aluminum reacts with oxygen to produce
aluminum oxide (rust).
d. Solutions of calcium chloride and sodium
sulfate react to produce aqueous sodium
chloride and solid calcium sulfate.
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CaCl2(aq) + Na2SO4(aq)→ NaCl(aq) + CaSO4(s)
2 Al(s) + O2(g) → Al2O3(s)
2
4 3
Synthesis Reactions
In a synthesis reaction, two
substances – either elements or
compounds – combine to form a
single compound.
The general equation for a
synthesis reaction is
A + B → C
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Synthesis Reactions – General Rules
a. Metal + Nonmetal → *Binary Ionic Compound
*You must remember to balance the charges
when writing the formula for the binary ionic
compound.
Na + O2 →
What product is going to be formed?
Sodium oxide
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Na2O 2 4
Na+, O2- Na2O
Synthesis Reactions – General Rules
b. Nonmetal + Nonmetal → Binary Covalent
Compound
S + O2 → SO2
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Synthesis Reactions – General Rules
c. Metal oxide + Water → *Base (metal hydroxide)
*You must remember to balance the charges
when writing the formula for the base formed.
CaO + H2O →
What product is going to be formed?
Calcium hydroxide
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Ca(OH)2
Ca2+, OH- Ca(OH)2
Synthesis Reactions – General Rules
d. Nonmetal oxide + Water → Acid
CO2 + H2O →
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H2CO3
Synthesis Reactions – General Rules
e. Metallic oxide + Nonmetallic oxide → Ternary Ionic
Compound (salt)
Na2O + CO2 →
Na2CO3
Synthesis Reactions – General Rules
f. Here are two synthesis reactions that must be
memorized.
N2(g) + 3H2(g) 2NH3(g)
NH3(g) + H2O(l) NH4OH(aq)
You Try It Balance the following equations. You may need to
complete the equation before balancing it.
a. Solid lithium metal reacts with oxygen gas.
b. Magnesium metal burns in oxygen gas.
Li + O2 Li2O 2 4
Mg + O2 MgO 2 2
You Try It c. Magnesium metal reacts with chlorine gas
d. Solid calcium oxide is heated in the presence
of sulfur trioxide.
e. Sulfur dioxide gas is bubbled into distilled water.
Mg + Cl2 MgCl2
CaO + SO3 CaSO4
SO2 + H2O H2SO3
You Try It f. Solid sodium oxide is added to water.
g. Calcium metal is heated strongly in nitrogen
gas.
h. Solid tetraphosphorus decoxide reacts with
water to produce phosphoric acid.
Na2O + H2O NaOH
Ca + N2 Ca3N2
P4O10 + H2O H3PO4
2
3
4 6
You Try It i. Iodine crystals react with chlorine gas to form
solid iodine trichloride.
j. Solid strontium oxide reacts with water.
I2 + Cl2 ICl3
SrO + H2O H2SrO2
2 3
Decomposition Reactions In a decomposition reaction, a compound breaks down
into two or more simpler substances.
The compound may break down into individual
elements, a compound and an element, or into simpler
compounds.
Many decomposition reaction take place only when
energy is added in the form of heat or light.
Electrolysis is the decomposition of a substance by an
electric current.
The general equation for a decomposition reaction is
C → A + B
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Decomposition Reactions – General Rules
a. Binary compounds break into their *elements.
*You must remember to watch out for the diatomic
elements (H2, Br2, O2, N2, Cl2, I2, F2).
K2O →
What products are going to be formed?
Potassium
Oxygen
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K 2 4 + O2
Decomposition Reactions – General Rules
b. Metal carbonates generally break down to
produce a metal oxide and carbon dioxide.
K2CO3 →
What products are going to be formed?
Potassium oxide
Carbon dioxide
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K2O + CO2
K+, O2- K2O
Decomposition Reactions – General Rules
c. Metal hydroxides generally break down to
produce a metal oxide and water.
Sr(OH)2 →
What products are going to be formed?
Strontium oxide
Water
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SrO + H2O
Sr2+, O2- SrO
Decomposition Reactions – General Rules
d. Metal chlorates generally break down to produce
a metal chloride and oxygen.
KClO3 →
What products are going to be formed?
Potassium chloride
Oxygen
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KCl + O2
K+, Cl- KCl
2 3 2
Decomposition Reactions – General Rules
e. Many acids decompose to produce water and a
nonmetal oxide.
H2CO3 →
What products are going to be formed?
Water
Nonmetal oxide
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+ CO2 H2O
Decomposition Reactions – General Rules
f. Ternary Ionic compounds (salts) decompose to
produce a metal oxide and a nonmetal oxide.
CaSO3 →
What products are going to be formed?
Metal oxide
Nonmetal oxide
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+ SO2 CaO
Decomposition Reactions – General Rules
g. Here are four decomposition reactions that must
be memorized.
2H2O2 2H2O + O2
NH4OH NH3 + H2O
(NH4)2CO3 2NH3 + H2O + CO2
NH4NO3 N2O + 2H2O
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Balance the following decomposition reactions.
Use the examples above to help you predict the
products when necessary. You do not have to
indicate the states of the reactants and products.
a. Solid barium carbonate is heated.
BaCO3 →
b. Solid silver oxide is heated.
Ag2O
You Try It!
BaO + CO2
Ag + O2 2 4
c. Solid sodium chlorate is heated.
NaClO3 →
d. An aqueous solution of sulfuric acid is
heated.
H2SO4 →
e. Aluminum can be obtained from
aluminum oxide with the addition of a
large amount of electrical energy.
Al2O3 →
You Try It!
NaCl + O2
H2O(l) + SO3(g)
2
4 Al + O2
2 3
2 3
f. Heating tin(IV) hydroxide gives tin(IV)
oxide and steam.
Sn(OH)4 →
g. Solid sodium carbonate is heated.
Na2CO3 →
h. Solid magnesium sulfite is heated.
MgSO3 →
You Try It!
SnO2 + 2H2O
Na2O + CO2
MgO + SO2
i. Powdered magnesium carbonate is
heated strongly.
MgCO3 →
j. Solid calcium hydroxide decomposes.
Ca(OH)2(s) →
You Try It!
MgO(s) + CO2(g)
CaO(s) + H2O(g)
Single-Replacement Reactions
In a single-replacement reaction, an uncombined
element replaces an element in a compound.
There are three general equations for a single
replacement reaction.
A + BC → AC + B (If A is a metal)
A + BC → BA + C (If A is a halogen)
A + BC → Base + H2
(If A is an active metal and BC is H2O)
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Single-Replacement Reactions
Examples
Zn + 2HCl → ZnCl2 + H2
Cl2 + 2LiBr → 2LiCl + Br2
2K + 2H2O → 2KOH + H2
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Single-Replacement Reactions
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Whether one metal will
replace another from a
compound can be
determined by using the
activity series of metals.
The activity series of
metals lists metals in order
of chemical reactivity.
A reactive metal will
replace any metal found
below it in the activity
series.
Single-Replacement Reactions
The halogens can also take place
in single-replacement reactions.
The order of reactivity for the
halogens from increasing to
decreasing reactivity is fluorine,
chlorine, bromine, iodine.
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Single-Replacement Reactions
Examples
a. Li + KCl →
b. Mg + K2SO4 →
c. Cl2 + KBr →
d. Li + H2O →
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+ K LiCl
no reaction
KCl + Br2
LiOH + H2
2 2
2 2 2
You Try It
Complete and balance the following
equations.
a. Zn + Pb(NO3)2 →
b. K + Ba(C2H3O2)2 →
c. Al + NiSO4 →
d. Na + H2O →
e. Zn + BaCl2 →
f. F2 + AlCl3 →
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+ Pb Zn(NO3)2
KC2H3O2 + Ba
Al2(SO4)3
no reaction
+ Ni
2 2
2 3
2
3
NaOH + H2 2 2
AlF3 + Cl2 3 2 2 3
Double-Replacement Reactions
In a double-replacement reaction, the
negative ions of two compounds
exchange places in an aqueous
solution to form two new compounds.
The general equation for a double
replacement reaction is
AB + CD → AD + CB
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Double-Replacement Reactions For a double-replacement reaction to occur, one of
the following statements is usually true concerning
at least one of the products of the reaction.
a. It is a gas that bubbles out of the mixture.
FeS(s) + HCl(aq) →
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2 FeCl2(aq) + H2S(g)
Other gases commonly formed include:
H2S
CO2 (formed from the decomposition of H2CO3)
SO2 (formed from the decomposition of H2SO3
NH3 (formed from the decomposition of NH4OH)
Double-Replacement Reactions
b. It is a molecular compound such as water.
This is common in acid-base neutralization
reactions.
HCl(aq) + NaOH(aq) →
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H2O(l) + NaCl(aq)
Double-Replacement Reactions c. It is only slightly soluble and precipitates from
solution.
KI(aq) + Pb(NO3)2(aq) →
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2 KNO3(aq) + PbI2(s) 2
Identifying Precipitates A solubility chart can be used to identify insoluble
solids (precipitates).
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Identifying Precipitates
Example. Label the precipitate formed in the
following reaction.
Na2CO3(aq) + Ca(NO3)2(aq) → 2NaNO3( ) + CaCO3( )
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aq s
You Try It
Label the precipitate formed in each of the following
reactions.
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aq s i. NH4Cl(aq) + AgNO3(aq) → NH4NO3( ) + AgCl( )
ii. BaCl2(aq) + Na2SO4(aq) → BaSO4( ) + 2NaCl( ) aq s
a. AgNO3 + NaCl →
b. Mg(NO3)2 + KOH →
c. LiOH + Fe(NO3)3 →
d. Pb(NO3)2 + NaOH →
e. NH4Cl(aq) + NaOH(aq) →
f. K2SO3(aq) + HBr(aq) →
You Try It
Complete and balance the following
equations.
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+ NaNO3 AgCl
Mg(OH)2 + KNO3
LiNO3 + Fe(OH)3
2 2
3 3
2 2 Pb(OH)2 + NaNO3
NH3 + H2O + NaCl KBr + SO2 + H2O 2 2
NH4OH
+ H2SO3
Combustion Reactions
In a combustion reaction, a substance combines
with oxygen, releasing a large amount of energy in
the form of heat and light.
For our purpose, we will talk about hydrocarbons
burning in the presence of oxygen.
The general equation for a complete combustion
reaction is
CxHy + O2 → CO2 + H2O
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Types of Combustion Reactions
Complete Combustion – the products
are CO2 and H2O.
CH4 + 2O2 → CO2 + 2H2O
Incomplete Combustion – the products
are C and/or CO and H2O.
2C3H8 + 7O2 → 6CO + 8H2O
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Combustion Reactions
Example Problem. Write a
balanced equation for the complete
combustion of ethane, C2H6.
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C2H6 + O2 → CO2 + H2O 2 4 6 7
You Try It
Write balanced equations for the
complete combustion of the
following substances.
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+ O2 + H2O 3 4 5
2
a. C3H8
b. C5H10
c. C6H14
+ O2
+ O2
→ CO2
→ CO2
→ CO2
+ H2O
+ H2O
10 10 15
2 12 14 19
Oxidation-Reduction Reactions (Redox)
Oxidation-Reduction reactions are
the chemical changes that occur
when electrons are transferred
between reactions.
Examples include the burning of
gasoline and the rusting of a nail.
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Oxidation Oxidation originally meant the
combination of an element with
oxygen to give oxides.
However, today it is also defined as
the loss of electrons. (Oxygen
does not have to be present for
oxidation to occur.)
Example: 4Fe + 3O2 → 2Fe2O3
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Reduction Reduction originally meant the loss
of oxygen from a compound.
Today it is also defined as the gain
of electrons.
Example: 2Fe2O3 + 3O2 → 4Fe + 3CO2
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Oxidizing Agent The oxidizing agent in a redox
reaction gains electrons.
It is the substance being reduced.
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Reducing Agent The oxidizing agent loses
electrons.
It is the substance being oxidized.
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Helpful Mnemonics Leo the Lion says Ger
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L – Loss
e – of electrons
o – is oxidation
G – Gain
e – of electrons
r – is reduction
Oil Rig
O – Oxidation
i – is
l – loss of electrons
R – Reduction
i – is
g – gain of electrons
Rules for Assigning Oxidation Numbers
1. The oxidation number of a
monatomic ion is equal to its
charge.
Examples:
Br- equals
Fe3+ equals
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-1
+3
Rules for Assigning Oxidation Numbers
2. For a polyatomic ion, the sum
of the oxidation numbers must
equal the ionic charge of the
ion.
Examples:
SO42- equals
NO3- equals
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-2
-1
Rules for Assigning Oxidation Numbers
3. The oxidation number of a
metal cation is the same as its
ionic charge.
Examples:
sodium is
calcium is
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+1
+2
Rules for Assigning Oxidation Numbers
4. The oxidation number of
hydrogen in a compound is +1
except in metal hydrides, for
example, NaH, where it is -1.
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Rules for Assigning Oxidation Numbers
5. The oxidation number of
oxygen in a compound is -2
except in peroxides, for
example, H2O2, where it is -1.
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Rules for Assigning Oxidation Numbers
6. The oxidation number of an
uncombined element is 0.
For example, the oxidation
number of the potassium atoms
in potassium metal, K, and of
the nitrogen atoms in nitrogen
gas, N2, is zero.
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Rules for Assigning Oxidation Numbers
7. For any neutral compound, the
sum of the oxidation numbers
of the atoms in the compound
must equal 0.
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Rules for Assigning Oxidation Numbers
8. Assign the oxidation numbers in
the following order.
a. Metals (determined by looking at their group number)
b. Hydrogen (usually +1)
c. Oxygen (usually -2)
d. Transition Metals and Everything Else
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What is the oxidation number of each element in the following?
SO2
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+4 -2
-4 +4 Start with oxygen. According to rule 5, the oxidation
number of oxygen in a compound is usually -2.
What number does sulfur need to be in order for the
overall compound to have a sum of zero?
=0
S = +4
O= -2
What is the oxidation number of each element in the following?
KClO3
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+1 -2
-6 +1 Start with potassium. According to rule 3, the oxidation
number of potassium is the same as its ionic charge,
which is +1.
Next do oxygen. According to rule 5, the oxidation
number of oxygen in a compound is usually -2.
What number does chlorine need to be in order for the
overall compound to have a sum of zero?
=0
K = +1
O= -2
Cl = +5 +5
+5
What is the oxidation number of each element in the following?
KClO2
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+1 -2
-4 +1 Start with potassium. According to rule 3, the oxidation
number of potassium is the same as its ionic charge,
which is +1.
Next do oxygen. According to rule 5, the oxidation
number of oxygen in a compound is usually -2.
What number does chlorine need to be in order for the
overall compound to have a sum of zero?
=0
K = +1
O= -2
Cl = +3 +3
+3
What is the oxidation number of each element in the following?
CO32-
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+4 -2
-6 +4 Start with oxygen. According to rule 5, the oxidation
number of oxygen in a compound is usually -2.
What number does carbon need to be in order for the
polyatomic ion to have an overall charge of -2?
= -2
C = +4
O= -2
You Try It. Determine the oxidation number of each element in each of the following.
1. Na2Cr2O7 Na = O = Cr =
2. BaH2 Ba = H =
(Hint: BaH2 is a metal hydride.)
3. Li2O2 Li = O =
(Hint: Li2O2 is a peroxide.)
4. ClO3- Cl = O =
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+2
-2 +6 +1
-1
+1 -1
+5 -2
Oxidation Number Changes
The changes in oxidation number can be
used to determine which elements are
oxidized and which elements are reduced.
Remember – an increase in the oxidation
number of an atom signifies oxidation and a
decrease in the oxidation number signifies
reduction.
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Oxidation Number Changes
Use the change in oxidation number to
identify which elements are oxidized and
reduced in each of these reactions. Also
identify the oxidizing (OA) and reducing (RA)
agents.
a. F2(g) + 2HBr(aq) → 2HF(aq) + Br2(l)
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0 +1
-1
+1
-1
0
Fluorine: 0 to -1; reduced, OA
Bromine: -1 to 0; oxidized, RA
+1 =0
-1
+1 =0
-1
Oxidation Number Changes
Use the change in oxidation number to
identify which elements are oxidized and
reduced in each of these reactions. Also
identify the oxidizing (OA) and reducing (RA)
agents.
b. 2KClO3(s) → 2KCl(s) + 3O2(g)
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+1 -2 +5 +1 -1 0
Chlorine: +5 to -1; reduced, OA
Oxygen: -2 to 0; oxidized, RA
-6 +1 +5 =0 +1 =0 -1
Oxidation Number Changes
Use the change in oxidation number to
identify which elements are oxidized and
reduced in each of these reactions. Also
identify the oxidizing (OA) and reducing (RA)
agents.
c. SO3(g) + H2O(l) → H2SO4(aq)
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+6 -2 +1 -2 +1
Since there is no change in oxidation
number, this is not a redox reaction.
-6 +6 +2
-2
-8 +6
+6
+2 -2 =0 =0 =0
You Try It
Use the change in oxidation number to
identify which elements are oxidized and
reduced in each of these reactions. Also
identify the oxidizing (OA) and reducing (RA)
agents.
a. 2H2(g) + O2(g) → 2H2O(g)
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0 0 +1 -2
+2
hydrogen: 0 to +1; oxidized, RA
oxygen: 0 to -2; reduced, OA
-2 =0
You Try It
Use the change in oxidation number to
identify which elements are oxidized and
reduced in each of these reactions. Also
identify the oxidizing (OA) and reducing (RA)
agents.
b. 2Al(s) + 6HCl(aq) → 2AlCl3(s) + 3H2(g)
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0 0 +1 +3
+3
aluminum: 0 to +3; oxidized, RA
hydrogen: +1 to 0; reduced, OA
-3
-1 -1
=0 =0 +1 -1
You Try It
Use the change in oxidation number to
identify which elements are oxidized and
reduced in each of these reactions. Also
identify the oxidizing (OA) and reducing (RA)
agents.
c. 2NaCl(aq) + Li2S(aq) → Na2S(s) + LiCl(aq)
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+1 +1 +1 -2
+2
Since there is no change in oxidation
number, this is not a redox reaction.
-2
-1
+2
-2
-2
+1 -1
=0 =0 =0 =0 +1 -1 +1 -1