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AP Chapter 2
Chemical Equations and
Reactions
Empirical and Molecular Formulas
Substance
Water
Molecular
Formula
H2O
Empirical
Formula
H2O
Benzene
C6H6
CH
Acetylene
C2H2
CH
Glucose
C6H12O6
CH2O
Atomic mass unit
• The atomic mass unit (u), is a unit of
mass used to express atomic and
molecular masses.
• It is the approximate mass of a hydrogen
atom, a proton, or a neutron.
• By definition the atomic mass unit is equal
to one-twelfth of the mass of a carbon-12
atom.
The Atom
• The nucleus is very
small, dense, and
positively charged.
• Electrons surround
the nucleus.
• Most of the atom is
empty space
Subatomic Particles
PARTICLE
SYMBOL CHARGE
MASS
(u)
LOCATION
electron
e-
-1
0
orbit nucleus
proton
p+
+1
1
inside nucleus
neutron
n0
0
1
inside nucleus
Atomic Number (Z)
• The number of
protons in the nucleus
of an atom.
• The identifying
characteristic of an
element.
Mass Number
• The sum of the
protons and neutrons
in the nucleus of an
atom.
John Dalton
• In the first modern
atomic theory John
Dalton proposed
that atoms of the
same element
were identical.
Are atoms of a particular
element identical?
• What is different about
No
atoms of the same
element and how did we
discover this difference?
Atoms of the same element
can have different masses.
Isotopes
Isotopes
• This is the symbol for
carbon-12.
• Atomic number is 6.
• Mass number is 12.
Isotopes
• This is the symbol for
carbon-12.
• Atomic number is 6.
• Mass number is 12.
• Write the symbols for
carbon-13 and
carbon-14.
Isotopes
What is the average mass of a
carbon atom?
What is the average mass of a
carbon atom?
12.01
Atomic Mass
• The atomic mass of
carbon is 12.01u.
Atomic Mass
• The atomic mass of
carbon is 12.01u.
• Atomic mass is the
average mass of all
the isotopes of an
atom. It takes into
account the different
isotopes of an
element and their
relative abundance.
Data from mass spectrometry indicates that chlorine has
two isotopes, chlorine-35 and chlorine-37? The percent
abundance for chlorine-35 is 75.53%. The percent
abundance for chlorine-37 is 24.40%. What is the average
atomic mass for chlorine?
[(0.7553)(35.00 amu)] + [(0.2440)(37.00 amu)]
= [26.4355] + [9.028]
= 35.4635
the atomic mass of Cl = 35.47 amu
• How many electrons,
protons and neutrons
are in an atom of
actinium with a mass
number of 221?
• How many electrons,
protons and neutrons
are in an atom of
actinium with a mass
number of 221?
• 89p+
• 89e• 132n0
• How many electrons,
protons and neutrons
are in an atom of
rhodium-105?
• How many electrons,
protons and neutrons
are in an atom of
rhodium-105?
• 45p+
• 45e• 60n0
Families of the Periodic Table
The Noble Gases
•
•
•
•
Elements in group 18
All are gases.
VERY non-reactive.
Have a full outer
energy level.
The Octet Rule
• The octet rule states that an element's
outer energy level is full and most stable
when it contains eight electrons.
• This stability is the reason that the noble
gases are so non-reactive.
Exception to the Octet Rule
• The first energy level can only hold two
electrons and so elements such as
Hydrogen and Helium that only have one
energy level follow a “duet rule”.
Ion Vocabulary
• An ion is an atom or group of atoms that have a
charge.
• A monatomic ion is an atom with a charge.
• The charge on the atom is called an oxidation
number.
• A polyatomic ion is a group of atoms with a
charge.
• A cation is a positive ion.
• An anion is a negative ion.
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
2.5
Ionic
• When an element that easily loses
electrons (a metal) reacts with an element
that easily gains electrons (a nonmetal),
one or more electrons are transferred.
• This creates two ions which are held
together by an ionic bond.
• A compound that contains ions is called an
ionic compound.
a
Formula Unit
A formula unit is the empirical formula of an ionic compound.
It is the lowest whole number ratio of ions represented in an
ionic compound. Examples include ionic NaCl and K2O.
Ionic compounds do not exist as individual molecules; a
formula unit thus indicates the lowest reduced ratio of ions in
the compound.
Covalent
• When atoms share electrons the bond
created is said to be covalent. Covalent
bonds often form between nonmetal
atoms.
• These covalently bonded atoms act as
single units called molecules.
• A compound made up of molecules is a
molecular compound.
Polyatomic Ions
H2PO4C2H3O2HSO3HSO4HCO3NO2NO3CNOHMnO4ClOClO2ClO3ClO4-
Dihydrogen Phosphate
Acetate
Hydrogen Sulfite (Bisulfite)
Hydrogen Sulfate (Bisulfate)
Hydrogen Carbonate (Bicarbonate)
Nitrite
Nitrate
Cyanide
Hydroxide
Permanganate
Hypochlorite
Chlorite
Chlorate
Perchlorate
HPO42C2O42SO32SO42CO32CrO42Cr2O72SiO32-
Hydrogen Phosphate
Oxalate
Sulfite
Sulfate
Carbonate
Chromate
Dichromate
Silicate
PO33PO43-
Phosphite
Phosphate
NH4+
Hg22+
Ammonium
Mercury(I)
Note that these are charges
and not oxidation numbers.
Types of monatomic ions and the
rules for naming them
• The periodic table is useful in naming the monatomic
ions.
Monatomic cations with one
oxidation number
• The cations from the periodic table which have a
single oxidation number are as follows: Metals in
Group 1 (+1), Group 2 (+2), Ag+, Cd2+, Zn2+, and Al3+.
• These types of ions are named by using the name of
the element followed by the word ion.
• Na+ sodium ion
• Ba2+ barium ion
• Zn2+ zinc ion
• We can use the roman numeral from the periodic table
to identify the oxidation number for these ions.
Monatomic cations with multiple
oxidation numbers
• All other cations that are not listed in the previous
category are considered to have the possibility of
multiple oxidation numbers. (Remember that other
than fluoride all ions have the possibility of being
positive ions).
• These type of ions are named by using the name of
the element followed by a Roman numeral to indicate
the oxidation number.
• Cu2+ copper (II)
• Pb4+ lead (IV)
• N3+ nitrogen (III)
Monatomic anions
• All anions from the periodic table are
named by changing the ending of the
element’s name to –ide.
• F- fluoride ion
• O2- oxide ion
• N3- nitride ion
• Count back from the noble gases starting
at zero to determine the oxidation number.
Name these monatomic ions
•
•
•
•
•
•
Rb+
P3Fe3+
Br‾
Mn4+
Cd2+
Name these monatomic ions
•
•
•
•
•
•
Rb+
P3Fe3+
Br‾
Mn4+
Cd2+
rubidium ion
phosphide
iron (III)
bromide
manganese (IV)
cadmium ion
Write the formula for these
monatomic ions.
•
•
•
•
•
•
Sulfide
Lead (II)
Barium ion
Chromium (IV)
Aluminum ion
Carbide
Write the formula for these
monatomic ions.
•
•
•
•
•
•
Sulfide
Lead (II)
Barium ion
Chromium (IV)
Aluminum ion
Carbide
S2Pb2+
Ba2+
Cr4+
Al3+
C4-
Formulas of Ionic Compounds:
•
Formulas for ionic compounds can be
written by the following steps:
Formulas of Ionic Compounds:
•
Formulas for ionic compounds can be
written by the following steps:
• (1) Write the formula for the cation and anion
(Don’t forget to include the charge of each ion).
Formulas of Ionic Compounds:
•
Formulas for ionic compounds can be
written by the following steps:
• (1) Write the formula for the cation and anion
(Don’t forget to include the charge of each ion).
• (2) Decide how many cations and anions are
needed so that the sum of their charges
balances out to be zero.
Formulas of Ionic Compounds:
•
Formulas for ionic compounds can be written by
the following steps:
• (1) Write the formula for the cation and anion (Don’t
forget to include the charge of each ion).
• (2) Decide how many cations and anions are needed
so that the sum of their charges balances out to be
zero.
• (3) Write the formula of the compound by writing the
number of cations followed by the number of anions
which you used in step #2. Remember not to include
the charges of the ions since now they balance out to
be neutral. (*Note when using more than one
polyatomic ion the polyatomic ion must be written in
parentheses).
Write the formula for barium chloride
Write the formula for iron (II) oxide
Write the formula for calcium phosphate
Write the formula for ammonium carbonate
Write the formulas for the following
compounds
(a) cobalt (II) chloride
(b) lithium sulfate
(c) ammonium dichromate
(d) aluminum oxide
(e) boron (III) phosphide
(f) Chromium (V) nitrate
Naming Ionic Compounds
• When naming ionic compounds the
following steps are followed:
Naming Ionic Compounds
• When naming ionic compounds the
following steps are followed:
(1) Separate the compound into its positive
and negative parts (Note that the positive
part of a compound will be only the first
element with the exception of ammonium
which is NH4+)
Naming Ionic Compounds
• When naming ionic compounds the
following steps are followed:
(1) Separate the compound into its positive
and negative parts (Note that the positive
part of a compound will be only the first
element with the exception of ammonium
which is NH4+)
(2) Write the name of the cation followed by
the name of the anion.
Oxidation Number Rules
• The oxidation number of all Group 1 metals (+1), Group 2
metals (+2), Ag+, Zn2+, Cd2+, and Al3+ within compounds
is a set value.
• Hydrogen (H) has two possible oxidation numbers:
– +1 when bonded to a nonmetal
– -1 when bonded to a metal
• The oxidation number of fluorine (F) is always -1.
• In ionic compounds the nonmetal closest to fluorine is
negative.
• The sum of the oxidation numbers of all atoms (ions) in a
neutral compound = 0.
• The sum of the oxidation numbers of all atoms (ions) in a
polyatomic ion = charge on the polyatomic ion.
Write the name of ZnO and
determine the oxidation numbers of
the elements within this compound.
Write the name of CuO and
determine the oxidation numbers of
the elements within this compound.
Write the name of MnCO3 and
determine the oxidation numbers of
the elements within this compound.
Write the name of Fe2(SO4)3 and
determine the oxidation numbers of
the elements within this compound.
Name the following compounds
and determine the oxidation
numbers of each element.
(a) SrCl2
(b) Cr(OH)2
(c) KClO4
(d) NH4MnO4
(e) CuP
Binary Molecular Compounds
• Binary molecular compounds are
composed of two different nonmetals
– examples: CO, SO2, N2H4, P4Cl10
• These compounds are named by using a
prefix to indicate the number of atoms of
each element present.
• The prefix mono- is often
omitted especially when
the first element would
have the prefix monoCO
• (example: CO is named
carbon monoxide, not
monocarbon monoxide).
Name the following compounds:
NF3
N2O4 P4S10
• NF3 is nitrogen trifluoride
• N2O4 is dinitrogen tetraoxide
• P4S10 is tetraphosphorous
decasulfide
Write formulas for the following compounds:
• dichlorine heptaoxide
• carbon hexasulfide
• octaphosphorous pentaoxide
• dichlorine heptaoxide is
Cl2O7
• carbon hexasulfide is
CS6
• octaphosphorous
pentaoxide is P8O5
Acids
• Acids are compounds that give off hydrogen
ions, (H+) when dissolved in water. When a
compound has hydrogen as its cation the
substance is generally an acid
– Examples: HCl, H2SO4, H3PO3
• The rules for naming acids are based on the
anion portion of the acid formula.
Rules for Naming Acids
• The names of acids are based on the
ending of the anion name.
– Examples: HCl, H2SO4, H3PO3
• Cl‾ = chloride
• SO42‾ = sulfate
• PO33‾ = phophite
Rules for Naming Acids
Name the acids
• HNO2
• HCN
• H3PO4
Write formulas for the following acids
• chromic acid
• hydroiodic acid
• chlorous acid
How do we know when a
chemical change has occurred?
• Noticeable energy change (gets hot or
cold, produces light).
• Formation of a gas (may notice a new
odor).
• Formation of a precipitate.
• Color change.
Balanced Chemical Equation
• A chemical equation is a written
representation of a chemical reaction.
2Na + 2H2O → H2 + 2NaOH
• Reactants
• Products
• Coefficients
– You should be able to balance equations
using coefficients.
Balance the following equations
KClO3 
C3H8 +
KCl +
O2 
O2
CO2 +
H 2O
Mg(NO3)2 + K3PO4  Mg3(PO4)2 +
KNO3
Symbols Used in Equations
2Na(s) + 2H2O(l) → H2(g) + 2NaOH(aq)
•
•
•
•
Solid
Liquid
Gas
Aqueous solution
Other Symbols Used in Equations
• Solid (cr) or (s)
• Precipitate (↓)
• “Heated” ∆
• Escaping gas ()
H2SO4
• Catalyst
• A word may be written above an arrow
to indicate something is necessary for
the reaction to occur. electricity
When a chemical
change has occurred we
often indicate the
change by writing a
chemical reaction.
Net Ionic Equations
• Solutions of sodium chloride and silver
nitrate are mixed.
• Step 1: Change the word equation into a
chemical equation by writing the formulas
for the reactants.
NaCl + AgNO3
NaCl + AgNO3
• Step 2: Classify each reactant as a particular type of substance.
•
•
•
•
•
•
Classification of Substances
Acids: compounds with formulas that begin with H. (Examples:
HCl, H2SO4).
Bases: compounds that end with OH. (Examples: NaOH, Ba(OH)2)
Metal Oxides: binary compounds of a metal and oxygen.
(Examples: CaO, Na2O).
Nonmetal Oxides: binary compounds of a nonmetal and oxygen.
(Examples: SO2, P4O10).
Salts: Ionic compounds other than bases and metal oxides.
(Examples: NaCl, Mg3(PO4)2, NH4NO3).
Other Compounds: All compounds not classified as one of the five
types above. (Examples: CH4, NH3).
• NaCl and AgNO3 are salts
Salt
NaCl
Salts
• A salt is an ionic compound other than
a base or oxide.
K2Cr2O7
CuSO4
NaCl + AgNO3
(salt + salt)
• Step 3: Based on your classification of the
substances determine the type of reaction.
Types of Net Ionic Equations
1. Double Replacement Reactions: These reactions start with two compounds and produce two different compounds. Such reactions can be
expected when the two reactants are some combination of acid, base, or salt. The products can be predicted by exchanging the positive parts of the
two reactants.
If carbonic acid, H2CO3 is produced as a product it should be written as H2O and CO2. If ammonium hydroxide, NH4OH is
produced as a product it should be written as NH3 and H2O.
2. Single Replacement Reactions: The reactants are an element and a compound and the products are a different element and compound. A metallic
element will replace the positive part of a compound or a nonmetallic element will replace the negative part of a compound.
3. Synthesis (Combination) Reactions: (a) Two elements combine to form a binary compound. (b) A metal oxide and water combine to form a base.
(c) A nonmetal oxide and water combine to form an acid. (d) a metal oxide and a nonmetal oxide combine to form a salt. In these reactions it is
necessary to know the charges of certain ions in order to predict the formulas of your products. You should determine these ion charges by using
their charges within the reacting substances. If this is impossible use your prior experience or the periodic table to make a prediction.
4. Decomposition Reactions: These reactions begin with a single compound and decompose into more that one product. In general they are simply
the reverse of the synthesis reactions listed in #3 above. There are however a few other common decomposition reactions that you should learn: (a)
Hydrogen peroxide, H2O2 will decompose into water, H2O and oxygen, O2. (b) Potassium chlorate, KClO3 will decompose into potassium chloride,
KCl and oxygen O2.
5. Combustion Reactions: Generally involve a hydrocarbon and oxygen, and if so will produce CO2 and H2O. Some form of the word “burn”
usually identifies combustion reactions. The most common hydrocarbons are alkanes and alcohols. Some examples are listed below:
Alkane
Formula
Alcohol
Formula
Methane
CH4
Methanol
CH3OH
Ethane
C2H6
Ethanol
C2H5OH
Propane
C3H8
Propanol
C3H7OH
Butane
C4H10
Butanol
C4H9OH
Pentane
C5H12
Pentanol
C5H11OH
Hexane
C6H14
Hexanol
C6H13OH
NaCl + AgNO3
(salt + salt)
• Step 3: Based on your classification of the
substances determine the type of reaction.
• This is a double replacement reaction.
NaCl + AgNO3 →
• Step 4: Predict the products of the
reaction based on the reaction type.
NaCl + AgNO3 → NaNO3 + AgCl
NaCl + AgNO3 → NaNO3 + AgCl
• Step 5: Use solubility rules if necessary.
• Solutions of sodium chloride and silver
nitrate are mixed.
Updated Solubility Rules
• Strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4
• Strong bases: Group I and Group II hydroxides
• Soluble salts: All salts containing the cations
sodium, potassium or ammonium. All salts
containing the anions nitrate or acetate. All chlorides
except silver, lead and mercury(I).
NaCl + AgNO3 → NaNO3 + AgCl
• Step 5: Use solubility rules if necessary.
Na+ + Cl- + Ag+ + NO3- → Na+ + NO3- + AgCl
Na+ + Cl- + Ag+ + NO3- → Na+ + NO3- + AgCl
• Step 6: Eliminate all spectator ions.
– A spectator ion appears as both a reactant
and a product in a chemical equation.
Na+ + Cl- + Ag+ + NO3- → Na+ + NO3- + AgCl
• Step 7: Write the final net ionic equation
Ag+ + Cl- → AgCl
+
Ag
+
Cl
→ AgCl
The AgCl is a precipitate
• A precipitate is a insoluble solid
formed when solutions are mixed.
• A precipitation reaction normally
occurs by reacting two soluble salts
to form an insoluble salt.
Precipitate
Solutions of lead nitrate and potassium iodide are mixed.
What is the yellow precipitate?
Types of Net Ionic Equations
1. Double Replacement Reactions: These
reactions start with two compounds and produce
two different compounds. Such reactions can be
expected when the two reactants are some
combination of acid, base, or salt. The products
can be predicted by exchanging the positive
parts of the two reactants.
• If carbonic acid, H2CO3 is produced as a product
it should be written as H2O and CO2.
Solid calcium phosphate is added
to excess hydrochloric acid.
Ca3(PO4)2 + H+ → H3PO4 + Ca2+
Equal volumes of 0.1M sulfuric acid
and 0.1M sodium hydroxide are mixed.
H+ + OH- → H2O
• This is an acid-base reaction.
• Acid + Base → Salt + Water
• The salt formed is soluble and
therefore does not appear in
the net ionic equation.
Solid barium carbonate is added to
an excess of dilute nitric acid.
BaCO3 + H+ → Ba2+ + H2CO3
Solid barium carbonate is added to
an excess of dilute nitric acid.
BaCO3 + H+ → Ba2+ + H2CO3
BaCO3 + H+ → Ba2+ + H2O + CO2
Demonstration
Single Replacement Reactions:
2. The reactants are an element and a
compound and the products are a
different element and compound. A
metallic element will replace the positive
part of a compound or a nonmetallic
element will replace the negative part of
a compound.
• These are really redox reactions.
Oxidation-Reduction
(Redox) Reaction
Sodium burns in air.
Na + O2 → NaO
In a redox reaction the
oxidation numbers of
at least some of the
substances change.
Teacher Example:
Calcium metal is added to dilute nitric acid.
Aluminum metal is added to a solution
of copper (II) chloride.
Al + Cu2+ → Al3+ + Cu
Liquid bromine is added to a solution of
potassium iodide.
Br2 + I- → Br- + I2
Solid calcium is added to warm water.
Ca + HOH → H2 + Ca2+ + OH-
Synthesis (Combination) Reactions:
3.
•
(a) Two elements combine to form a binary compound.
(b) A metal oxide and water combine to form a base.
(c) A nonmetal oxide and water combine to form an
acid.
(d) a metal oxide and a nonmetal oxide combine to
form a salt.
In these reactions it is necessary to know the charges
of certain ions in order to predict the formulas of your
products. You should determine these ion charges by
using their charges within the reacting substances. If
this is impossible use your prior experience or the
periodic table to make a prediction.
Teacher Example:
Magnesium metal is heated strongly
in nitrogen gas.
Teacher Example:
Calcium oxide is added to water.
Teacher Example:
Dinitrogen trioxide gas is bubbled
through water.
Calcium metal is heated strongly in
nitrogen gas.
Ca + N2 → Ca3N2
Teacher Example:
Excess chlorine gas is passed over
hot iron filings.
Cl2 + Fe → FeCl3
A piece of lithium metal is dropped
into a container of nitrogen gas.
Li + N2 → Li3N
Solid barium oxide is added to
distilled water.
BaO + HOH → Ba2+ + OH-
Lithium oxide powder is added to
excess water.
Li2O + H2O → Li+ + OH-
Solid dinitrogen pentoxide is added
to water.
N2O5 + H2O → H+ + NO3-
Phosphorus (V) oxide powder is
sprinkled over distilled water.
P2O5 + HOH → H3PO4
Metal oxide + Nonmetal oxide
Solid calcium oxide is exposed to a
stream of carbon dioxide gas.
CaO + CO2 → CaCO3
Solid calcium oxide is heated in the
presence of sulfur trioxide gas.
CaO + SO3 → CaSO4
Decomposition Reactions:
4. Decomposition Reactions: These reactions
begin with a single compound and decompose
into more than one product. In general they
are simply the reverse of the synthesis
reactions listed in #3 above. You should learn
the decomposition reaction for hydrogen
peroxide. Hydrogen peroxide, H2O2 will
decompose into water, H2O and oxygen, O2.
A solution of hydrogen peroxide is
exposed to an iron catalyst.
H2O2 → H2O + O2
• A catalyst speeds up a reaction but
remains unchanged at the end of the
process.
Solid calcium sulfite is heated in a
vacuum.
CaSO3 → CaO + SO2
Combustion Reactions:
5. Generally involve a hydrocarbon and oxygen, and if so will produce CO2 and H2O. Some form of
the word “burn” usually identifies combustion reactions. The most common hydrocarbons are
alkanes and alcohols. Some examples are listed below:
Alkane
Formula
Alcohol
Formula
Methane
CH4
Methanol
CH3OH
Ethane
C2H6
Ethanol
C2H5OH
Propane
C3H8
Propanol
C3H7OH
Butane
C4H10
Butanol
C4H9OH
Pentane
C5H12
Pentanol
C5H11OH
Hexane
C6H14
Hexanol
C6H13OH
Combustion
•
Hexane is burned in excess oxygen.
•
Propanol is burned completely in air.
Oxidation and Reduction
(Redox)
• Many reactions including single
replacement, combustion, and many
synthesis and decomposition reactions are
really redox reactions.
• You should be able to recognize if a
reaction is a redox reaction and identify
the substance that is reduced and the
substance which is oxidized.
Redox
• If a substance is reduced its oxidation
number decreases.
• If a substance is oxidized its oxidation
number increases.
Redox Reactions?
• Cl2 + Fe → FeCl3
• H2O2 → H2O + O2
• P2O5 + HOH → H3PO4
Hydrates
CoCl2 ∙ 6H2O
CoCl2