Compounds
Chapter 3
Chapter 3 Educational Goals
1. Compare
and contrast the Bohr model and the quantum
mechanical model (Modern Atomic Theory) of the atom in terms
of how electrons are arranged about an atom’s nucleus.
2. Define the term valence electron and, given an element or ion,
draw the electron dot structures.
3. Define the term ion and explain how the electron dot structure of
a representative element atom (groups 1A-8A) can be used to
predict the charge of the monoatomic ion.
4. Given the symbol, the student will be able to name monoatomic
and cations and anions (and vice versa). I will give you a table
with the polyatomic anions and cations for the exams.
5. Explain the difference between an ionic bond and a covalent
bond.
6. Given the name, the student will be able to write the formulas of
simple ionic compounds and binary molecules (and vice versa).
7. Define the terms formula mass and molecular mass and use
these values in unit conversions involving moles and mass.
8. Given the formula, draw line bond structure of simple molecules
The Arrangement of Electrons
Before we learn about compounds, we must
build on our understanding of atoms and
their electrons.
1) Understand where electrons are located in
atoms.
2) Understand how the location of electrons
effect the energy of the atom.
Scientist used light to study how electrons
are arranged around the nucleus.
Light and Matter
•  Light (energy) can be absorbed by atoms and
molecules.
Consider a Hydrogen Atom
Energy is absorbed by moving an electron to a
new area.
+
p
Electrostatic
e-
Attractive
Force
Consider a Hydrogen Atom
Electrons can adsorb energy from heat or light
+
p
e-
Light and Matter
An example of light (energy) absorbed by atoms and
molecules is paint molecules.
Atoms and molecules can give off energy by
emitting light.
Energy is released when an excited electron is
moved back to a more stable area.
An example of light emitted by matter is fireworks.
Problems with Dalton and other’s
Model of the Atom
Did not explain the pattern of light given
off by excited atoms (atomic spectra).
After Dalton, scientists continued
to revise atomic theory
Atomic Theory = Atomic Model
The Modern Atomic Model
•  Quantum Mechanics is the laws of physics
developed in the early 1900’s that were able to
account for observations of particles that were
very small and very fast.
•  Quantum Mechanics predicts the 3-D areas
that the electrons occupy as well as the energy
of electrons observed.
The Quantum Mechanical Atomic Model:
The Modern Model of the Atom
•  The energy of atoms are quantized.
–  Electron can have only certain, discreet energies!
Sky-Scraper Model
Similar concept is
only being able to be
on certain floors
(levels) of a building.
-you can be on the 3rd
floor or 4th floor, but
not on the 3.5 floor!!!
The Quantum Mechanical Atomic Model:
The Modern Model of the Atom
•  Electrons can only occupy certain areas.
–  These 3-dimensional regions are called orbitals.
Sky-Scraper Model
Similar concept is
only being able to be
in a particular type of
room on any floor of a
building.
-example:
Office
Lobby
Kitchen
Living room
The Quantum Mechanical Atomic Model:
The Modern Model of the Atom
•  Each 3-D region (orbital) that an electron can
occupy can be described by two properties:
•  1) Average distance from the nucleus.
–  called quantum levels.
–  Labeled by numbers (n) ; 1, 2, 3, 4….. (like floors in the
building)
•  2) The shape of the area containing the electron
–  the various shapes are labeled: s, p, d, f… (like different
types of rooms in our sky-scraper)
The Quantum Mechanical Model of
Hydrogen
Electrons are assigned to various atomic orbitals labeled
by quantum level (n)
The n=1 level has one orbital.
Energy
It is called an s orbital (1s because
n=1).
This picture of an s-orbital represents
a region where one could find the
electron 90% of the time.
An electron is represented by an arrow
The lowest energy state is the n=1 level.
n=1
1s
The electron goes to the lowest energy
state unless it gains extra energy
Energy
Use this page (notes) to draw an
Energy Level Diagram
n=1
Electrons are assigned to various atomic orbitals
labeled by quantum level (n)
The next quantum level is the n=2 level.
The n=2 level has 4 orbitals
There is one s-orbital and three p-orbitals
Energy
-called 2s, 2px,2py,2pz
n=2
2s
n=1
1s
2px
2py
2pz
The shape of the 2s orbital is almost exactly the same as the
1s orbital, it is just larger
IMPORTANT: the larger n is, the further the electrons are
from the nucleus!!!
Energy
Here are pictures of where one can find the electrons 90% of
the time for the 2s and 2p orbitals
n=2
2s
n=1
1s
2px
2py
2pz
If energy is given to the electron
in hydrogen, it can be excited to
a higher energy orbital
Energy
It will stay there for a short
period of time, then loose the
energy (by giving off light or
colliding with another atom), and
return to it’s lowest energy
orbital.
n=2
2s
n=1
1s
2px
2py
2pz
3d orbitals
The next quantum level is the n=3 level.
The n=3 level has 9 orbitals
Energy
There is one s-orbital, three p-orbitals, and 5 d-orbitals
n=3
3s
3px 3py 3pz
n=2
2s
2px 2py 2pz
n=1
1s
3d 3d 3d 3d 3d
3d orbitals
The shape of the 3s and 3p orbitals are almost exactly the
same as the 2s and 2p orbitals, they are just larger.
Energy
IMPORTANT: the larger n is, the further the electrons are
from the nucleus!!!
n=3
n=2
n=1
3s
2s
1s
3px 3py 3pz
2px 2py 2pz
3d 3d 3d 3d 3d
3d orbitals
Here are pictures of where one can find the electrons 90% of
the time for the 3d orbitals
Let’s take a look at what the orbitals look
like when we put them together in an atom!
+
4f orbitals
The next quantum level is the n=4 level.
The n=4 level has 16 orbitals
There is one s-orbital, three p-orbitals, five d-orbitals,
and seven-f orbitals.
Energy
n=4 4s
n=3
n=2
n=1
3s
2s
1s
4px 4py 4pz
3px 3py 3pz
2px 2py 2pz
4d 4d 4d 4d 4d
3d 3d 3d 3d 3d
4f
4f
4f
4f
4f
4f 4f
4f orbitals
The shape of the 4s,4p, and 4d orbitals are almost exactly
the same as the 3s, 3p, and 3d orbitals, they are just larger.
IMPORTANT: the larger n is, the further the electrons are
from the nucleus!!!
Energy
n=4 4s
n=3
n=2
n=1
3s
2s
1s
4px 4py 4pz
3px 3py 3pz
2px 2py 2pz
4d 4d 4d 4d 4d
3d 3d 3d 3d 3d
4f
4f
4f
4f
4f
4f 4f
4f orbitals
Here are pictures of where one can find the electrons 90% of
the time for the 4f orbitals
The next quantum level is the n=5 level.
The n=5 level has 25 orbitals
Energy
We do not worry about the g, h, i, etc orbitals because
they are never used.
n=5
5s
5px 5py 5pz
5d 5d 5d 4d 5d
5f
5f
5f
5f
5f
5f 5f
n=4
4s
4px 4py 4pz
4d 4d 4d 4d 4d
4f
4f
4f
4f
4f
4f 4f
n=3
3s
3px 3py 3pz
3d 3d 3d 3d 3d
n=2
2s
2px 2py 2pz
n=1
1s
Energy
The shape of the 5s,5p,5d and 5f orbitals are the same as the 4s, 4p,
4d,and 4f orbitals, they are just larger.
n=5
5s
5px 5py 5pz
5d 5d 5d 5d 5d
5f
5f
5f
5f
5f
5f 5f
n=4
4s
4px 4py 4pz
4d 4d 4d 4d 4d
4f
4f
4f
4f
4f
4f 4f
n=3 3s
3px 3py 3pz
3d 3d 3d 3d 3d
n=2
n=1
2s
1s
2px 2py 2pz
This is called an
“energy level diagram”
Sky-Scraper Model
n=3
3s
3p
3p
3p
Energy
2s
2p
2p
1s
3d
3d
3d
3d
n=2
2p
n=1
3d
•  These energy level diagrams are only for
hydrogen!
• This is how nature determined the energy and
location of a single electron atom.
• In all the other atoms, with more than one
electron, things change……
Energy Level Diagram for Multi-Electron
Atoms
6d 6d 6d 6d 6d
5f 5f 5f 5f 5f 5f 5f
7s
6px 6py 6pz
5d 5d 5d 5d 5d
4f 4f 4f 4f 4f 4f 4f
Energy
6s
5px 5py 5pz
5s
4s
4px 4py 4pz
3px 3py 3pz
3s
2s
1s
2px 2py 2pz
4d 4d 4d 4d 4d
3d 3d 3d 3d 3d
3d
3d
3d
3d
3d
Energy
4s
3p
3p
3p
3s
2p
2p
2s
1s
2p
Sky-Scraper
Model
How are the electrons configured
(arranged) into all these orbitals?
Electron Configuration
•  Nature wants everything to be at the lowest possible
energy.
•  Aufbau Principle ("aufbau" means "to build-up“)
–  An electron occupies the lowest energy orbital that can
receive it.
–  Electrons are configured in atoms in the order of the
lowest energy orbitals to highest energy orbitals
The Pauli Exclusion Principle (rule of
quantum mechanics):
An orbital can hold a maximum of two electrons
Filling Energy Level Diagrams
Let’s do Hydrogen first!
We don’t need all these orbitals, hydrogen only has one
electron!
Draw the 1s Energy Level Diagram in your notes.
An electron is represented by an arrow.
Up arrow is for spin up (spin = +½)
Down arrow is for spin down (spin = -½)
Filling Energy Level Diagrams
Let’s do Hydrogen first!
Energy
We don’t need all the n=2 orbitals, hydrogen only has one
electron!
n=1
1s
Filling Energy Level Diagrams
Let’s do Helium (2 electrons) next!
Draw the 1s, 2s, and 2p Multi-Electron Energy Level Diagram in your
notes.
Draw the first electron in the n=1 level
Energy
The aufbau principle says that we must completely fill the lowest energy
orbitals before we put electrons in higher energy orbitals!
The Pauli exclusion principle says that within the same orbital,
electrons must have opposite spins!
2s
1s
2px 2py 2pz
Filling Energy Level Diagrams
Let’s do Carbon (6 electrons) next!
Draw the 1s, 2s, and 2p Multi-Electron Energy Level Diagram in your
notes.
Draw the first and second electron in the 1s orbital.
Draw the third and forth electron in the 2s orbital.
Energy
Draw the fifth electron any of the 2p orbitals.
They are identical and indistinguishable!
2s
1s
2px 2py 2pz
Filling Energy Level Diagrams
Energy
Orbitals of equal energy are each occupied by one electron
before any orbital is occupied by a second electron
2s
1s
2px 2py 2pz
2s
1s
2px 2py 2pz
Filling Energy Level Diagrams
You try one: Neon
Begin by drawing the 1s, 2s, and 2p Multi-Electron Energy
Level Diagram in your notes.
Energy
Then fill in the electrons
Filling Energy Level Diagrams
Now a hard one: Iodine
Begin by drawing the Multi-Electron Energy Level Diagram
(up to the 5p orbital) in your notes.
Energy
Then fill in the electrons
You now know our
modern theory of the
atom!
•  You know about the nucleus (protons and neutrons)
•  You know about the electrons (energies and how
they are arranged in orbitals)
Valence Electrons
•  Valence electrons are the electrons held in the
outermost shell (largest n).
– Valence electrons are furthest away from the
nucleus.
•  It is important to know how many valence electrons
are in an atom because:
– These are the electrons that are involved in
chemical bonding to other elements to form
molecules.
– These are the electrons that elements lose to
become ions
Valence Electrons
•  Understanding the arrangement of electrons
about an atom gives us some insight as to how
members of a given group or period in the
periodic table are related.
Example: Carbon’s Valence Electrons
How many valence electrons does carbon have?
Energy
First, fill the energy levels with electrons
2s
1s
2px 2py 2pz
4 Valence
Electrons
Carbon’s Electrons
Carbon has 2 electrons in the 1s orbital
These are not valence electrons and are not involved in bonding.
They are called core electrons
Carbon has 2 electrons in the 2s orbital and 2 electrons in the 2p orbitals.
These are the valence electrons that are furthest away from the nucleus.
They are the electrons involved in bonding to other atoms.
2s
1s
2px 2py 2pz
Determining the number of valence
electrons from the periodic table
– All elements in the same periodic
column (group) have the same number
of valence electrons as all others in
that column.
– We call this iso-electric.
• This is why elements in the same
group often react in similar ways.
Number of valence electrons in
groups 1-8A
1
2
8
3 4 5 6 7
Electron Dot Structures
•  Electron dot structures show the
number of valence electrons that an
atom carries.
•  In these structures, valence
electrons are represented by dots.
Electron Dot Structures
Consider the s-block elements.
We know that the first column elements are all isoelectric and have
one valence electron.
We draw one dot (1 dot = 1 electron)
1 2
3 4 5 6 7 8
Electron Dot Structures
We know that the second column elements are all isoelectric and
have two valence electrons.
We draw two dots (2 dot = 2 electrons)
1 2
3 4 5 6 7 8
Electron Dot Structures
Next, consider the p-block elements.
We know that the elements in each column are all isoelectric and
have the same number of valence electrons.
• Draw single dots until the fifth dot, then pair them up.
1 2
• That is because atoms in molecules like to have 8 electrons. (4 pairs of
electrons)
3 4 5 6 7 8
You try it!
Draw the electron dot structure of:
Br, Ga, Rb, Ba, Ge, Kr, As
Noble Gases
•  He, Ne, Ar, Kr, Xe and Rn belong to the noble
gas family, which gets its name from the fact
that these elements are resistant to change and ,
with few exceptions, do not lose or gain
electrons.
•  The resistance to change or stability is related to
having 8 electrons in their outermost shell.
Octet Rule
•  We observe atoms/ions that have 8 electrons in their
outermost shell, as in the case of the noble gas atoms, are
not “reactive”.
–  (not reactive = stable + low energy)
•  This is so common that chemist use the Octet Rule!!!
Octet Rule: Chemical compounds tend to form so that
each atom, by gaining, losing, or sharing electrons, has an
octet of electrons in its outermost shell (n)
Exception to the octet rule: Hydrogen and Helium
Hydrogen and Helium have filled outer shells with
just 2 electrons.
What is an ion?
An atom can gain or lose electron(s) to
become an ion.
Metals lose electrons to become positive ions
called - Cations
Non-metals gain electrons to become negative
ions
called -Anions
Octet rule in formation of ions
•
Octet Rule: Chemical compounds tend to
form so that each atom, by gaining, losing, or
sharing electrons, has an octet of electrons in its
outermost shell (n)
•  Very often, ions are formed such that the ion has
an octet in its outermost shell (n)
Let’s do a Cation
•  Example: Sodium (Na)
•
•
•
•
A sodium atom has ______
11 protons and _____
11 electrons.
Fill the energy level diagram with electrons.
How many valence electrons does the sodium atom have? _____
1
8
How many valence electrons does sodium want? _____
Na atom
3s
3p
2s
2p
1s
Let’s do a Cation
+
Na ion
3s
3p
2s
2p
1s
When sodium loses an electron,
it has an octet of electrons in its
outer shell.
Sodium will lose one electron to
become a sodium ion (Na+).
Draw electron dot structure for a Na atom.
•  Sodium has one valence electron
•  There are two ways to have an octet:
–  1) Add 7 electrons
–  2) Remove one electron
•  It is easier to remove one electron!
•  This is how we draw the electron dot structure for the ion:
Na
Na
+
Charge
Let’s do Another Cation
•  Example: Magnesium (Mg)
•
•
•
•
A magnesium atom has ______ protons and _____ electrons.
Fill the energy level diagram with electrons.
How many valence electrons does the magnesium atom have? ___
How many valence electrons does magnesium want? _____
Mg atom
3s
3p
2s
2p
1s
Let’s do Another Cation
Mg2+ ion
3s
3p
2s
2p
1s
2+
When magnesium loses two
electrons, it has an octet of
electrons in it’s outer shell.
Magnesium will lose two
electrons to become a
magnesium ion (Mg2+).
Draw the electron dot structure for a Mg atom.
•  Magnesium has two valence electrons
•  There are two ways to have a filled octet:
–  1) Add 6 electrons
–  2) Remove two electrons
•  It is easier to remove two electrons!
•  This is how we draw the electron dot structure for the ion:
Mg
2+
Charge
Octet rule in formation of ions
•  Example: Fluorine (F)
•
•
•
•
A fluorine atom has ______ protons and ______ electrons.
Fill the energy level diagram with electrons.
How many valence electrons does the fluorine atom have? _____
How many valence electrons does fluorine want? _____
F- ion
2s
1s
2p
When fluorine gains an
electron, it has an octet of
electrons in it’s outer shell.
Fluorine will gain one electron
to become a fluoride ion (F-).
Octet rule in formation of ions
•  The electron dot structure can give us the same conclusion!
•  Draw the electron dot structure for a F atom.
•  Fluorine has 7 valence electrons, if we add one electron, its outer
shell will have a full octet.
•  This is how we draw the electron dot structure for the ion:
F
-
Charge
Octet rule in formation of ions
•  Example: Oxygen (O)
•
•
•
•
A oxygen atom has ______ protons and ______ electrons.
Fill the energy level diagram with electrons.
How many valence electrons does the oxygen atom have? _____
How many valence electrons does oxygen want? _____
O atom
2s
1s
2p
Octet rule in formation of ions
O2- ion
2s
1s
2p
2-
When oxygen gains two
electrons, it has an octet of
electrons in it’s outer shell.
Oxygen will gain two electrons
to become an oxide ion (O2-).
Octet rule in formation of ions
•  The electron dot structure can give us the same conclusion!
•  Draw the electron dot structure for an O atom.
•  Oxygen has 6 valence electrons, if we add two electrons, its outer
shell will have a full octet.
•  This is how we draw the electron dot structure for the ion:
O
O
2-
Charge
The charge for s- and p- block elements can be
determined from their periodic column (group).
Hydrogen can lose it’s electron to form a 1+ cation
1+
H+
2+
Now you know why we put it in the column with
metals- it forms a cation
3+
3-
N3P3-
2-
1-
What about the transition metals?
How do we know the charge?
How do we name them?
We can not determine the charge from the position on the
periodic table.
It gets worse…..
Most of them can exist having various charges!!!
(They come in different flavors!!!!)
Consider Iron (Fe):
Iron (Fe) ions can come as Fe2+ or Fe3+
Iron(II)
Fe2+
Fe3+
Iron(III)
To name metal ions that can have multiple
charges, we write the charge in Roman numerals
after the name.
Consider Copper (Cu):
Copper (Cu) ions can come as Cu1+or Cu2+
Copper(I)
Cu1+
+
Cu2+
Copper(II)
Good News:
•  You do not need to memorize transitions metal
names and charges, I will give you a list for
exams or quizzes.
Ag+
silver
Cd2+
cadmium
Zn2+
zinc
Group Work
How many electrons will each element gain or lose in
forming an ion?
Ca
Al
Cl
S
K
N
Polyatomic Ions
•  Definition:
– A polyatomic ion is a charged group of covalently
bonded atoms.
– Can be positive (cation) or negative (anion)
• Polyatomic cations have a shortage of electrons
• Polyatomic anions have extra electrons
+
NH4
SO4
2-
Ammonium Ion
Sulfate Ion
Examples of Polyatomic Ions
I will give you this table on an exam, however, I suggest that
medical professionals or those continuing in science academically
or industrially memorize the following names and their symbols
and charges:
hydroxide, carbonate, bicarbonate, nitrate, sulfate, ammonium,
phosphate, acetate
Naming Ions
•  Cations are named the same as the metal
sodium
Na → Na+ + 1e- sodium ion
calcium
Ca → Ca+2 + 2e- calcium ion
• Anions are named by changing the ending of the name
to -ide
fluorine
F + 1e- → Ffluoride ion
oxygen
O + 2e- → O2oxide ion
Chemical
Compounds
Chemical Compounds
•  Compounds: matter that is constructed of two
or more chemically combined elements.
•  Each compound has the same proportion of the
same elements.
–  Water = 2 hydrogen atoms and 1 oxygen atom
(H2O)
–  Sodium chloride = 1 sodium atom for every 1
chlorine atom (NaCl)
Chemical Change
•  When matter is changed to a new substance, a
chemical change has taken place.
•  Chemical bonds are made, broken, or both.
•  The ability and rate of a substance to be
changed into a new substance is called its
chemical properties.
Chemical Bonding
The Molecule’s Structure
Determines it’s Properties!
•  A cardinal principle of chemistry is that the
macroscopic observed properties of a material
are related to its microscopic structure.
•  The microscopic structure entails
– the kinds of atoms
– the manner in which they are attached
– their relationship to other molecules (like and
dislike)
– the shape of the molecule
Chemical Bonds
•  Chemical bonds are the electrical attractive
forces that hold atoms together.
•  In this chapter, we will study two types of
chemical bonding:
–  1) Ionic Bonding
–  2) Covalent Bonding
Definitions
•  Ionic Bonding: Chemical bonding that
results from the electrical attraction between
large numbers of cations and anions.
Definitions
•  Covalent Bonding: Chemical bonding that
results from the sharing of electron pairs
between two atoms.
•  We will see two different types of compounds
based on how the atoms are bonded together:
Ionic Compounds
Molecular Compounds
(ionic bonding)
(covalent bonding)
Covalent Bonding: Molecular Compounds
•  A molecule is a neutral (no charge) group of
atoms that are held together by covalent bonds.
•  Examples:
Covalent bonding involves non-metal elements
only.
Covalent Bonding: Molecular Compounds
•  A chemical substance whose simplest units are
molecules is called a molecular compound.
•  The molecular formula shows the types and
numbers of atoms that make up a single
molecule.
Examples:
One nitrogen monoxide
molecule contains one
nitrogen atom and one atom
of oxygen atom.
Covalent Bonding: Molecular Compounds
One carbon dioxide
molecule contains 1 carbon
atom and 2 oxygen atoms
One glucose molecule
contains 6 carbons atoms,
12 hydrogen atoms, and 6
oxygen atoms
C6H12O6
Formation of a Covalent Bond
•  In chemistry, things happen because nature
wants everything to be at the lowest
possible energy!
•  Covalent bonding occurs because the bound
atoms are at a lower energy than the
unbound atoms.
Formation of a Covalent Bond
•  Let’s consider the simplest molecule: H2
•  Here is a graph of energy vs. distance between H atoms.
Energy
Bound
H atoms
Distance between H atoms
Separated
H atoms
Formation of a Covalent Bond
The two separated H atoms are higher in energy than the bound
H atoms.
Energy
Bound
Hatoms
Distance between H atoms
Separated
H atoms
Formation of a Covalent Bond
The difference in energy between the
separated atoms and the bound atoms is
called the bond energy.
Energy
Bond
Energy
Distance between H atoms
Remember…
•  It requires energy to break a bond
•  Energy is released when a bond is made.
Forces in Covalent Bonding
As the atoms get closer, the electrons feel an attractive
force from both nuclei (+) and vice versa.
This positive-negative-positive “sandwich” lowers the
energy and forms a covalent bond.
Forces in Covalent Bonding
Energy
Too Close!
Nuclei
Repel
Covalent
bond formed
Atoms get
closer
Atoms far
apart
Distance between H atoms
If the atoms continue to move closer together, the nucleus (+) of each
atom begin to repel each other.
This causes the energy to quickly rise if the nuclei get too close!
+
+
The forces in a bond are much like the forces in a
spring.
H
H
Quantum Mechanical Model of
Covalent Bonding
The stability of the H2 molecule can not be
entirely explained by the simple model of
the positive-negative-positive sandwich!
For example, it does not explain why He does not form
a He2 molecule.
Quantum Mechanical Model of
Covalent Bonding
•  Atoms tend to form stable molecules if
sharing electrons leads to having an octet of
electrons their outermost shell.
–  Example: H2 (recall that H and He are stable with
2 valence electrons)
H
1s
Non-bonded Hydrogen Atoms
Each atom has only 1 electron in the 1s
orbital.
H
1s
Quantum Mechanical Model of
Covalent Bonding
–  Example: H2
H
In the bonded H2 molecule, each atom
“feels” it has 2 electrons in its outer
shell.
H
1s
1s
Quantum Mechanical Model of
Covalent Bonding
•  The H2 covalent bond can also be illustrated
with valence electron dot structures.
HH
The two electrons between the atoms are shared in a covalent bond.
Chemist use a line to represent 2 electrons in a covalent bond.
These drawings are called line bond structures.
H H
Quantum Mechanical Model of
Covalent Bonding
•  Let’s do another example:
–  Hydrogen Chloride (HCl)
Cl
Non-bonded Hydrogen and Chlorine Atoms
3s
3p
Chlorine has 7 electrons in its outer shell (n=3).
Hydrogen has only 1 electron in its outermost
shell (n=1).
H
1s
Quantum Mechanical Model of
Covalent Bonding
•  Let’s do another example:
–  Hydrogen Chloride (HCl)
The chlorine feels 8 electrons in its outer shell.
The hydrogen feels 2 electrons in its outer shell.
Cl
3s
3p
H
1s
Quantum Mechanical Model of
Covalent Bonding
•  The HCl covalent bond can also be illustrated
with a line bond structure.
H
H Cl
H Cl
Cl
Quantum Mechanical Model of
Covalent Bonding
•  Let’s do another example:
–  Chlorine gas (Cl2)
Cl
3s
3p
Non-bonded Chlorine Atoms
Each chlorine has 7 electrons in its outer shell (n=3).
Cl
3s
3p
Quantum Mechanical Model of
Covalent Bonding
–  Chlorine gas (Cl2)
Cl
In the bonded Cl2 molecule, each atom
“feels” like it has 8 electrons in its outer
shell.
3p
3s
Cl
3s
3p
You try it:
•  Can you draw the line bond structure for Cl2?
–  Hint: Start with the Electron Dot Structure for two Cl atoms.
Cl
Cl
Cl Cl
Cl Cl
Quantum Mechanical Model of
Covalent Bonding
How about Oxygen Gas?
Quantum Mechanical Model of
Covalent Bonding
•  Let’s do oxygen gas (O2):
O
2s
2p
O
2s
2p
Non-bonded Oxygen Atoms
Each oxygen has 6 valence electrons in its outer
shell (n=2).
Quantum Mechanical Model of
Covalent Bonding
•  Let’s do oxygen gas (O2):
In the bonded O2 molecule, each atom
“feels” like it has 8 electrons in its outer
shell.
Here, 2 pairs of electrons are shared.
O
2s
2p
O
2s
2p
Line Bond Structure for O2
Let’s draw the line bond structure for oxygen
•  Oxygen atoms have 6 valence electrons.
•  We will rotate the electrons so they can form bonding pairs.
OO
We use lines to represent shared electron pairs.
When atoms are bonded with 2 pairs of electrons it is called a
double bond.
Double Bond
O O
Line Bond Structure for N2
Let’s draw the line bond structure for nitrogen.
•  Nitrogen atoms have 5 valence electrons.
•  We will align the electrons so they can form bonding pairs.
N N
We can use lines to represent shared electron pairs.
When atoms are bonded with 3 pairs of electrons it is called a
triple bond.
Triple Bond
N N
Naming Binary Covalent Compounds
•  Binary covalent compounds always involve:
One Nonmetal
Another Nonmetal
Examples:
O
A
O=C=O
Carbon dioxide
H
H
water
Naming Binary Covalent Compounds
•  Binary covalent compounds contain only two
types of nonmetal elements.
– There may be more than one of each
element.
• For example CO2 contains just two types
of elements: carbon and oxygen.
•  We will discuss naming covalent compounds
that contain more than two types of elements,
like glucose C6H12O6, in later chapters.
Naming Binary Covalent Compounds
•
•
•
•
Rules:
1) List the name of the first element in the formula.
2) List the second element and add the –ide suffix.
3) Use Greek prefixes to indicate the number of each
atom in the formula.
–  Exception: do not use mono- for the first element in the
name.
•  Example: CO2
•  monocarbon dioxide
carbon dioxide
–  The o or a at the end of the Greek pre-fix is usually
dropped when the element name begins with a vowel.
•  Example : CO
•  carbon monooxide
carbon monoxide
Naming Binary Covalent Compounds
PREFIX
monoditritetrapentahexaheptaoctanonadeca-
NUMBER
1
2
3
4
5
6
7
8
9
10
Naming Binary Covalent Compounds
•  Example: Name the following compound CCl4
•  1) List the name of the first element in the formula.
•  2) List the second element and add the –ide suffix.
•  3) Use Greek prefixes to indicate the number of each atom in the
formula.
–  Exception: do not use mono- for the first element in the
name.
mono carbon tetrachloride
Examples
Write the name for the following molecular formulas:
•  CF4
•  N2O
•  SF6
More Examples
Write the molecular formula for the covalent
compounds:
•  arsenic trichloride
•  dinitrogen pentoxide
•  tetraphosphorus decoxide
Naming Covalent Compounds
•  For covalent compounds with more than
two types of atoms, we use common names
or IUPAC system names.
–  Examples of common names:
•  Glucose (C6H12O6)
•  Acetone (C3H6O)
Ionic Compounds
•  Ionic Bonding: Chemical bonding that results from the electrical
attraction between large numbers of cations and anions.
•  Results from:
–  1) combining metal ions with non-metal ions
–  2) a polyatomic ion with any other ion
Example: NaNO3 (sodium nitrate)
Ions are arranged in a pattern called a crystal lattice.
•  The arrangement maximizes attractions between + and - ions
For an example of an ionic compound, let’s
take a magnified look at Sodium Chloride
(table salt)!
Example of a Crystal Lattice:
Table salt is an ionic compound made from
Na and Cl ions.
Ionic Compounds
•  The cations and anions will combine in a
ratio such that the total of the + and –
charges equals ZERO!
•  Example: NaCl
Sodium ions have a charge of 1+
Chloride ions have a charge of 1They combine in a 1-to-1 ratio in
the crystal
For every sodium ion, there is
exactly one chloride ion!
The charges add up to ZERO!
Ionic Compounds
We write the formula unit for the compound
in the ratio that the ions combine:
Na1Cl1
= NaCl
For ionic compounds,
we write the formula
unit, not to be confused
with the molecular
formula of covalent
compounds!
What if the charges of the ions are
not the same?
Ionic Compounds
•  The cations and anions will combine in a
ratio such that the total of the + and –
charges equals ZERO!
•  Example: CaCl2
Calcium ions have a charge of 2+
Chloride ions have a charge of 1Ca2+
Cl-
They combine in a 1-to-2 ratio in
the crystal
For every calcium ion, there are
exactly two chloride ions!
The charges add up to ZERO!
Balancing the charge to get the
correct formula unit
1+
Na
1-
Cl
Hint:
Total Charge = ZERO
Balancing the charge to get the
correct formula unit
2+
1-
Mg Cl
Hint:
Total Charge = ZERO
Balancing the charge to get the
correct formula unit
1+
K
O
2-
Hint:
Total Charge = ZERO
Balancing the charge to get the
correct formula unit
2+
Mg N
3-
Hint:
Total Charge = ZERO
Dr. Zoval’s Caveman Style, Works
Every Time Method:
2+
Mg N
3-
The Criss-Cross Method
Group Work:
Write the formula for the ionic compound formed
between each of the following pairs of ions:
Cu+ and O2Fe3+ and S2Cu2+ and ClMg2+ and O2Sn4+ and S2V3+ and Cl-
Properties of Ionic Compounds
•  All ionic compounds are solids at room temperature
–  Melting points greater than 300°C
•  Liquid state conducts electricity, solid state does not
–  Liquid = molten
•  Brittle and Hard
•  Often dissolve in water, and when dissolved, the
solution becomes an electrical conductor
–  When ionic compounds containing polyatomic ions dissolve,
the covalent bonds holding the polyatomic ion do not break,
the ion stays together even though it separates from the other
ion
Formula Unit vs. Molecular Formula
Formula Unit: ex: NaCl
Lowest RATIO of atoms
Na = 1 Cl = 1
Molecular Formula:
ex: H2O
Actual # of atoms
O=1 H=2
O
A
H
H
Naming Ionic Compounds
Just follow these two simple rules:
1) Write the cation name first, then name
the anion.
2) If the cation is one of the transition
metals with various charges, write the
charge using parenthesis and Roman
numerals after the metal name.
Example:
Name the following compound:
MgCl2
Name the metal ion first: magnesium
Name the anion next: chloride
magnesium chlor ide
Name the following ionic compounds
l  Examples:
NaCl
ZnI2
Al2O3
Let’s do some examples with
polyatomic ions
l  Examples:
Na2SO4
Zn(NO3)2
IMPORTANT: When there is more than one of a
polyatomic ion in the formula unit we use parenthesis!
Al2(CO3)3
NH4Cl
Example:
Name the following compound:
2+
1-
CuBr2
Name the anion next
Name the metal ion first:
Note: The Caveman, works
every time method can
always be used!
What is the charge of the bromide ion?
What must the charge of the copper ion be?
copper(II )brom ide
Complete the names of the following ionic
compounds with variable metal ions:
FeBr2
iron (____) bromide
CuCl
copper (__) chloride
SnO2
_______(__
) _____________
tin
Fe2O3
________________________
Naming Compounds Summary
Determine if the compound is ionic or binary covalent
Does the compound contain only two types of non-metal elements?
Yes
Binary Covalent Compound
1) List the name of the first element in the
formula.
2) List the second element and add the –ide
suffix.
3) Use Greek prefixes to indicate the number
of each atom in the formula.
• Exception: do not use mono- for the
first element in the name.
• The o or a at the end of the Greek prefix is dropped when the element name
begins with a vowel
No
Ionic Compound
1)  Write the cation name first, then name
the anion.
•  monoatomic anions use the “ide” suffix
2) If the cation is one of the transition metals
with various charges, write the charge using
parenthesis and Roman numerals after the
metal name.
What about the reverse!
Writing the formula if we are given the
name of the ionic compound.
We have already done this:
Criss-Cross Method!
•  The cations and anions will combine in a
ratio such that the total of the + and –
charges equals ZERO!
Given the Name, Writing the Formula:
Determine if the compound is ionic or binary covalent
Does the compounds contain only two types of non-metal elements?
Yes
No
Binary Covalent Compound
Ionic Compound
1)  Write the symbol/formula of the first
element in the compound’s name, then
the symbol/formula of the second ion in
the compound’s name.
2) Indicate how many of each element the
molecule contains using subscripts after
the atomic symbol.
• The numbers of atoms are given in the
molecule’s name in Greek prefixes
• NOTE: If there is no Greek prefix in
front of the first element in the name,
that means the number is 1.
1)  Write the symbol/formula of the first ion
in the compound’s name, then the
symbol/formula of the second ion in the
compound’s name.
2) Indicate the ratio of each ion in the
compound using subscripts after each
ion.
• The ratio of the ions are deduced by
balancing the charges of the ions.
IMPORTANT: When there is more than one
of a polyatomic ion in the formula
unit we use parenthesis. Example Mg(NO3)2
Remember to use roman numerals or
alternative names for the metals that can
form more than one type of ion:
Molecular Mass
And
Formula Mass
Do you remember this?
# Atoms
Or
# Molecules
Mass (grams)
Use
Avogadro’s
Number
1 mole =
6.022 x 1023
Use
Molar Mass
# Moles
grams
mole
Well, that works for atoms, but how
about molecules?
Molar Mass of Compounds
•  The molar mass is the mass, in grams, of one
mole of a compound
•  In the case of molecules and compounds, it is
often called molecular mass or formula mass.
•  The molar masses of compounds can be
calculated from atomic molar masses:
Water = H2O = 2 x (1.01 g/mole) + 16.00
g/mole
= 18.02 g/mole 1 mole of H2O will weigh
18.02 g
You try it!
Example: Calculate the Molar Mass of methane (CH4)
1 Carbon
1 X 12.01 g/mole = 12.01 g/mole
4 Hydrogen
4 X 1.01 g/mole = 4.04 g/mole
Total= MM of CH4
= 16.05 g/mole
Group Work
•  Example: Calculate the Molar Mass of Glucose.