Chemistry
Unit 1 covered the atomic structure of elements and the periodic table of elements. There are
only 92 elements that occur in nature, many of these are extremely rare, and 26 elements have
been created in large, expensive experimental sites like the Stanford Linear Accelerator Center
and the Large Hadron Collider in Geneva. The millions of other pure substances both in
nature and created in labs are compounds. Compounds form when elements combine and
their atoms held together by chemical bonds. This unit is about how compounds form and the
properties of different types of compounds.
Ionization Energy and Electronegativity
Chemical bonds form when electrons are transferred or shared between atoms. The periodic
trends, of ionization energy and electronegativity, help explain why metals and nonmetals
react by transferring electrons and two nonmetals will share electrons when bonding.
Ionization energy is a measure of how hard it is to remove an electron. Metallic elements on
the left of the periodic table have low ionization energies and lose electrons easily.
Electronegativity measures the pull of the atomic nucleus for electrons towards itself. The
nonmetals in the upper right with a high electronegativity will form bonds by sharing
electrons between atoms of similar electronegativity as well as attract electrons from metallic
elements with a low electronegativity.
1
13.6
Ionization Energy (eV)
24.6
H
&
He
2.20
Electronegativity
––
5.4
2
9.3
8.3 11.3 14.5 13.6 17.4 21.6
Li Be
B
C
N
O
F
Ne
0.98 1.57
2.04 2.55 3.04 3.44 3.98 – –
5.1
7.6
6.0
8.2 10.5 10.4 13.0 15.8
3 Na Mg
Al
Si
4
K
6.1
S
Cl
Ar
1.61 1.90 2.19 2.58 3.16 – –
0.93 1.31
4.3
P
6.6 6.8 6.7 6.8 7.4 7.9 7.9 7.6 7.7 9.4
Ca Sc Ti
V
6.0
7.9
9.8
9.8 11.8 14.0
Cr Mn Fe Ni Co Cu Zn Ga Ge As Se
Br
Kr
0.82 1.00 1.36 1.54 1.63 1.66 1.55 1.83 1.88 1.91 1.90 1.65 1.81 2.01 2.16 2.55 2.96 – –
4.3
5.7
5 Rb Sr
6.2 6.6 6.8 7.1 7.3 7.4 7.5 8.3 7.6 9.0
Y
5.8
Zr Nb Mo Tc Ru Rh Pd Ag Cd In
7.3
8.6
9.0 10.5 12.1
Sn Sb Te
I
Xe
0.82 0.95 1.22 1.33 1.60 2.16 1.90 2.2 2.28 2.20 1.93 1.69 1.78 1.96 2.05 2.10 2.66 – –
Ionization Energy (top value) and Electronegativity of Selected Elements
Transferring Electrons
Ionic compounds are made by transferring electrons to become ions. A cation, which is a
positive ion, forms by losing an electron, and an anion, a negative ion, forms when the other
element gains an electron. The two ions will attract each other because of their opposite
charge by a bonding force called electrostatic attraction.
Page 1 of 28
Chemistry
Li
•
•
•
1s2
•
2s1
+
•
•
•
1s2
Ionization Energy
5.4 eV
Electronegativity
0.98 eV
F
Li1+
•
°
•
•
•
••
•

•
2s2 2p5
•
1s2
F1–
•
• • •
+• • • • •
• • •
1s2 2s2 2p6
Ionization Energy
17.4 eV
Electronegativity
3.98 eV
The Transfer of an Electron from a Lithium Atom to a Fluorine Atom
A lot of chemistry takes place in this diagram that’s common with most synthesis reactions
that combine elements to make an ionic compound. Let’s start with the reaction. The two
atoms at the left are called the reactants, which are the starting materials of the reaction, and
the two ions at the right are called the products, which are the results of a reaction. The arrow,
, means yields, forms, gives, reacts to make, etc.
The circles in the diagram represent the energy levels or shells of the electron configuration
and each dot, •, is an electron. This is called the Shell Model and can be a valuable tool to show
the movement of electrons and the formation of chemical bonds. Notice that the electron
configuration given below the shell model provides the same information about the number of
electrons on each shell.
The low ionization energy of lithium, Li, shows that it takes little energy to remove the
electron. The ionization energy of fluorine is higher and fluorine will not lose its electron. The
electronegativity of fluorine is high (the highest value actually) and its nucleus has a high
attraction for electrons. These two factors combine to cause the transfer of an electron to
fluorine.
The products of the reaction are a lithium cation, Li1+, which has lost a negative electron while
not changing the three positive protons in the nucleus so it has a positive charge, and a
fluorine anion, F1–, which has an extra electron so a negative charge. The superscript, 1+ and
1–, represent the charge of the ion and the number of electrons lost (for +) and electrons gained
(for –).
Finally, it is important to remember that lithium has a larger radius than fluorine since fluorine
has a greater effective nuclear charge (1 shell shielding the valence electrons and the nucleus
changes from +3 to +9). This is represented in the shell model. But when fluorine gains an
electron there are more electron-electron repulsions and the +9 charge is distributed among
ten electrons instead of nine so the radius of the anion increases. Likewise the lithium cation is
Page 2 of 28
Chemistry
much smaller than the lithium atom, since the electron in the outer shell is gone and the inner
shell experiences an unbalanced positive charge (nucleus is +3 and electrons are -2).
Noble Gas Electron Configuration is Stable
In the reaction between lithium and fluorine only one electron was transferred, but as many as
four electrons can be transferred in common reactions. The valence electrons, which are the
outermost electrons, are most easily removed because of shielding from the inner electrons. In
the figure below each atom from lithium to carbon shows the same result when all the valence
electrons are removed: two electrons in a 1s2 electron configuration, which is the same
electron configuration of helium.
Li
•
•
•
1s2
•
•
1s2
C
•
•
•
•
•
•
•
1s2
2
1s2
–
Be2+ lose 2 e
lose 1 e–
°
•
B
•2s
2s1
Li1+
Be
•
2
2p1
lose 3 e–
°
•
•
•
2
•
• •
2
1s2 2s2 2p2
C4+
lose 4 e–
°
• °
•
°
1s
°
1s
• •
•
•
2s
B3+
°
•
• •
•
° ••• °
°
1s2
Second Period Metals Losing Electrons to Form a Stable Noble Gas Electron Configuration
Metals with low ionization energies lose electrons until the valence electrons are removed. No
inner electrons are removed in the reaction because they are close to the nucleus which has an
extra positive charge. Removing these electrons would take too much energy compared to
reacting with other atoms that still have valence electrons. Since it is so common for the
valence electrons to be lost in a reaction and the inner electrons remain with s and p orbitals
filled, noble gas electron configurations are considered stable.
Page 3 of 28
Chemistry
Na
•
• • •
• ••• •
• • •
[Ne] 3s1
–
Na1+ lose 1 e
°
• • •
• ••• •
• •
•
Mg
•
• • •
• ••• •
• • •
•
[Ne] 3s
1
Mg2+
lose 2 e–
°
• ••
• ••• •
• • •
°
Al
•
• • •
• ••• • •
• • •
•
[Ne] 3s1
Al3+
lose 3 e–
°
• •• •
•• •• °
•••
°
1s 2s 2p
2
2
6
1s2 2s2 2p6
1s2 2s2 2p6
Third Period Metals Losing Electrons to Form a Stable Noble Gas Electron Configuration
When atoms gain electrons, they also stop adding electrons when they get to a noble gas
electron configuration, with a filled set of s and p orbitals, because gaining even more electrons
would require a new, more distant energy level or shell to hold the electrons. The atoms that
gain electrons have high values of electronegativity, which indicates the nucleus has a strong
pull toward electrons. Nonmetals on the upper right of the periodic table, other than noble
gases, have high electronegativities and form anions easily by gaining electrons.
Page 4 of 28
Chemistry
F
•
•
•
•
•
•
•
••
•
•
gain 1 e
F1–
gain 2 e
•
•
•
1s2
1s2 2s2 2p5
–
•
•
2s2
•
•
•
2s2
•
•
•
•
2p3
•
1s2 2s2 2p2
N3–
–
•
•
•
1s2
gain 3 e
•
•
•
2p4
•
•
•
•
•
•
O2–
–
C
N
O
gain 4 e
C4–
•
–
•
•
•
• •
• • • •
• • • • • • • • • •• • • • •• • • • •
•
• •
• • •
•
•
•
•
•
•
1s 2s 2p
1s 2s 2p
1s 2s 2p
1s 2s 2p
2
2
6
2
2
6
2
2
6
2
2
6
Second Period Nonmetals Gaining Electrons to form Stable Noble Gas Electron Configurations
Because the electron configurations for families of elements are similar, their chemical
reactions are similar. For example, the alkali metals the electron configurations is ns1 , where n
is the highest energy level or shell. Thus, all alkali metals easily lose one electron in a reaction
where there is a transfer of electrons. Likewise, alkaline earth metals with electron
configurations of ns2 will lose two electrons. On the other side of the periodic table the
halogens will gain one electron since their electron configuration is ns2 np5, while the oxygen
family will gain two electrons because these atoms have electron configurations of ns2 np4.
Similar Electron
Family
For example
Configuration
alkali metals
ns1
K = [Ne] 3s1 (lose 1e–) [Ne] = 1s2 2s2 2p6 = K1+
alkaline earth
metals
ns2
Ba = [Xe] 6s2 (lose 2e–) [Xe] = Ba2+
oxygen family
ns2 np4
S = [Ne] 3s2 3p4 (gain 2e–) [Ne] 3s2 3p6 = S2–
halogen
ns2 np5
Br = [Ar] 4s2 4p5 (gain 1e–) [Ar] 4s2 4p6 = Br1–
Most elements gain or lose electrons until a noble gas electron configuration is reached
because any further changes are energetically unfavorable. Gaining more electrons requires
adding electrons to a more distant energy level and losing more electrons would remove
electrons from inner energy levels. Transition metals, however, with their d-orbitals, usually
Page 5 of 28
Chemistry
have multiple types of ions (tin has
and
This is because the d–orbital electrons are
2
x
in an inner shell, ns (n – 1)d , and have energies very close to the outer s-orbital electrons. So
while the outer s-orbital electrons are usually removed, some d-orbital electrons can also be
removed. For example, manganese, [Ar] 4s2 3d5, can have cations of Mn2+, Mn3+, Mn4+, Mn6+,
and Mn7+.
Sn2+
Sn4+).
Charge
Up to now the charge has been added to the cation or anion without explanation. Still the
convention for showing charge is straightforward. When aluminum, Al, loses 3 electrons its
cation is show as Al3+. The 3 represents the number of electrons lost and the + indicates there
is an excess of positive charge. Similarly, when oxygen, O, gains 2 electrons the anion is O2–.
The 2 represents the 2 gained electrons and the – is for the negative charge the oxygen gained.
Valence Electrons and Octets
The number of valence electrons, those electrons in the outermost energy level, determines the
number of electrons an element will gain or lose. Since the families have the same pattern of
outer electrons, the number of valence electrons is the same in each family.
1 2
1
⇐
Number of
⇒
Valence Electrons
3 4 5 6 7 8
H
He
2 Li B
e
3 Na Mg
*d electrons generally don’t count as
valence electrons.
4
V Cr Mn Fe Ni Co Cu Zn Ga Ge As Se Br Kr
K Ca Sc Ti
B
Al
C
Si
N
P
O
S
F N
e
Cl Ar
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
1+ 2+
⇐Usual Charge⇒
I
3+ 4± 3- 2- 1-
Xe
NR
NR = no reaction
The table above is a result of atoms losing and gaining electrons to form stable noble gas
electron configurations. When an element loses electrons it commonly does so until all valence
electrons are gone; thus, magnesium with 2 valence electrons will react to become Mg2+,
magnesium cation by losing the 2 valence electrons. Likewise, when an element gains
electrons, it will gain only enough electrons to fill the outer shell s and p-orbitals; such as,
nitrogen with 5 valence electrons will gain 3 electrons to become N3–, nitrogen anion, so that
there are 8 electrons in the outer shell. The full outer shell of 8 electrons is called an octet, and
ions are said to be stable when they have an octet. Here are two more examples:
K = [Ne] 3s1(lose 1e–)[Ne] = 1s2 2s2 2p6 = K1+ or S = [Ne] 3s2 3p4 (gain 2e–) [Ne] 3s2 3p6 =
S2–.
Noble gases do not react because they already have full shells, ns2 np6. The energy change
when electrons are added or removed is too large because the effective nuclear charge is too
high for an electron to be removed easily and because adding electrons would require using a
new shell which is not favorable.
Page 6 of 28
Chemistry
For nonmetals in the second period that lose electrons, the number of electrons lost will result
in the noble gas electron configuration of helium 1s2, so only two valence electrons are in the
outer shell not an octet. Hydrogen usually loses one electron to become a single proton in the
nucleus without any electrons in the electron cloud, H+; however, it very occasionally gains
one electron to become H– with an electron configuration similar to helium.
Lewis Structures
To show the movement of electrons in a reaction, Gilbert Lewis introduced method of writing
atoms and their valence electrons using chemical symbols and dots for electrons. Lewis dot
diagrams, also called Lewis structures, electron dot structures, electron dot diagrams, etc.,
simply show the valence electrons about the chemical symbol of an element placed on the top
and bottom and on the left and right.
1
•
H
2
•
Li
•
3 Na
4
•
K
Lewis Dot Structures
•
•
•
Be
B• •C•
•
•
•
•
•
•
*d electrons generally don’t
Mg
Al•
•Si•
count as valence electrons.
•
•
•
•
•
•
Ca Sc Ti V Cr Mn Fe Ni Co Cu Zn Ga• •Ge•
•
•
•
••
•N•
•
••
•P•
•
••
•As•
•
••
••
•O• •F••
••
••
••
••
•S• •Cl••
••
••
••
••
•Se• •Br••
••
••
•
He
•
••
••Ne••
••
••
•• Ar••
••
••
••Kr••
••
Forming Ionic Compounds
To form an ionic compound, elements must transfer electrons from one another. So if a metal
and nonmetal react, then the metal will lose electrons and the nonmetal will gain electrons to
make a cation and an anion.
Model Used to Show the
Reaction
Reactants
Chemical Symbols
Be
+
S
Products

Be2+
+
S2–
Shell Model
Electron Configuration
Electron Dot Structure
1s2 2s2 + [Ne] 3s2 3p4
•
Be
•
+
••
S
••
Page 7 of 28

1s2
+ [Ne] 3s2 3p6
2+

Be +
•• 2–
•• S••
••
Chemistry
By convention the cation, usually a metal, is written before the anion. This is particularly
important for writing the chemical formula of ionic compounds.
Balancing the Electrons Transferred
When electrons are transferred, the number of electrons lost and gained must be equal. To do
this more atoms are added to provide a balanced number of valence electrons. For the first
example below, beryllium, Be, needs to lose both outer electrons, but chlorine, Cl, only needs
one. To balance the transfer of electrons needed for the atoms to reach a stable noble gas
electron configuration another chlorine is added. The other examples demonstrate the
balanced transfer of electrons. Here are some examples using electron dot diagrams:
•
Be +
•
••
••
2+
•• 1–
•
•
•
•Cl• + •Cl•  Be + 2 • Cl••
••
••
••
•
•
••
1+
•• 2–
K + K + O  2 K + •• O••
••
••
•
•
•
Mg + Mg + Mg +
•
•
•
••
••
2+
•• 3–
N + N  3 Mg + 2 •• N••
•
•
••
In each of these reactions the number of valence electrons lost from each metallic element must
have a place to be gained by the nonmetal element. The total number of electrons lost then
gained is the least common multiple of the number of valence electrons (or you can just keep
track of the electrons and add elements and valence electrons until the number of electrons lost
equals the number of electrons gained).
For example, Identify the number and type of ions formed in a reaction of calcium and carbon.
Step 1: Determine the number of valence electrons for each element and whether the element
loses or gains electrons. Calcium has 2 valence electrons and is a metal so it will give up two
electrons. Carbon is the nonmetal and with 4 valence electrons it will gain 4 electrons (carbon
gains and loses electrons depending on what the other element does, in this case calcium
wants to give electrons so carbon will gain electrons).
Step 2: Determine how many extra elements need to be added so the number of valence
electrons gained equals the number lost. Carbon needs more than 2 electrons from calcium so
add another calcium. Two calciums will have 4 valence electrons to lose which equals the 4
valence electrons one carbon gains.
Step 3: Show the reaction:
•
•
2+
•• 4–
2Ca + C  2 Ca +•• C••
••
Electrostatic Attractions
When ions are created they can usually dissolve in water (we will discuss solubility in another
unit), or will form a solid by sticking together. The opposite charges of cations and anions will
Page 8 of 28
Chemistry
force the ions to attract. Electrostatic attraction is the name of the strong bonding force
holding the cations and anions together in an ionic compound.
Crystal Lattice
There are, of course, a multitude of cations and anions present in any reaction. So each cation
can attract several anions and each anion can attract several cations. The result is a repeating
and organized arrangement of cations and anions called a crystal lattice. A crystal lattice is a
very stable bonding for ions and a lattice bonding energy is associated with the repeated
arrangement of cations and ions.
The crystal lattice of sodium
chloride. Small sodium cations
and larger chlorine anions in a
repeating pattern held together by
electrostatic attractions.
http://commons.wikimedia.org/wiki/File:Sod
ium-chloride-unit-cell-3D-ionic.png
Chemical Formulae and Naming.
There is a specific method for using element symbols from the periodic table of elements to
represent compounds. For ionic compounds the cation is listed first and then the anion. The
number of cations and anions needed to balance the charge is listed as subscripts. Two
comments before showing some examples: 1) a 1 in chemistry is usually dropped, so that Na1+
equal Na+ or Mg1Cl2 equals MgCl2; and 2) since every solid ionic compound is composed of
many, many cations and anions, we say that the chemical formula of an ionic compound
represents the lowest whole number ratio of elements making up the compound.
Here are some examples of compounds that form from a pair of elements. The number of
elements needed to balance the reaction is not shown.
K + S  K2S
Ba + P  Ba3P2
See if you can determine these compounds using electron dot structures.
Sn + F  SnF4
The name of ionic compounds has a similarly straightforward methodology. Name the cation
with the same name as the element, then name the anion but add an -ide ending to create a
word other than the element so that a reader understands that a compound is being named.
For example sodium reacts with chlorine (the two elements) to form sodium chloride (NaCl).
Another example barium reacts with phosphorous (see example above) to form barium
phosphide (Ba3P2). Note that the number of cations and anions is not listed since every
chemist could figure out the number of cations and anions.
Transition metals
Page 9 of 28
Chemistry
Recall transition metals form multiple cations. Some important examples are
and Cu2+,
Fe2+ and Fe3+, Co2+ and Co3+, Pb2+ and Pb4+, Sn2+ and Sn4+, and Hg22+ and Hg2+ . To
differentiate between the compounds that can form using different ions a roman numeral is
used to name each of the cations. For example for Cu1+ and Cu2+ the two ionic compounds
with an oxygen anion, O2–, are Cu2O and CuO. The compounds are copper (I) oxide and
copper (II) oxide respectively.
Cu1+
Some metals have only one cation so they do not use roman numerals in their names: Ag
forms only Ag1+ , Zn forms only Zn2+, Al forms only Al3+ , Cd forms only Cd2+.
Below are examples of how to identifying the chemical formula of a compound from its
elements, how to change from a chemical formula to the name of the compound, and how to
write the chemical formula if given the name . For the last three examples, the oxidation
number of the metal would not be known from just the element, but you can change from
chemical formula to name and name to chemical formula
Elements
(metal / nonmetal)
Al
+
Cl
Chemical Formula of
the Compound
Ions
Compound Name
Al3+
&
Cl1–
AlCl3
aluminum chloride
Ca
+ O
Ca2+
&
O2–
CaO
calcium oxide
Na
+ N
Na1+
&
N3–
Na3N
sodium nitride
Cu
+ F
Cu1+
Cu2+
&
&
F1–
F1–
CuF
CuF2
copper (I) fluoride
copper (II) fluoride
Pb
+ S
Pb2+
Pb4+
&
&
S2–
S2–
PbS
PbS2
lead (II) sulfide
lead (IV) sulfide
Mo
+ O
Mo6+
Mo4+
&
&
O2–
O2–
MoO3
MoO2
molybdenum (VI) oxide
molybdenum (IV) oxide
It is necessary to find the ions of the compound before you can determine the chemical
formula or name. Furhermore, you may have noticed a pattern for changing from ion to
chemical formula called criss-cross. Using the charges of the cation and anion as the subscripts
in the chemical formula for the opposite ion, you can write the chemical formula of the
compound. This pattern works about 2/3rds of the time, but, as with the last two examples,
the criss-cross can give subscripts that are too large. All subscripts for chemical formulae must
be written in lowest whole number ratio.
Finally, the compound name can be broken up into the two ions to determine the chemical
formula. For example, calcium phosphide would be the ions Ca2+ and P3– which form Ca3P2.
Metallic Bonds
Metals as either elements or alloys, which are mixtures of metals, have distinct properties.
Metals are highly conductive of heat and electricity. They are dense, have high luster (shiny),
Page 10 of 28
Chemistry
metallic colors (copper, silver, and gold), can be pulled into wires (ductile), and can be beaten
into sheets (malleable). The special bonding of metals is responsible for many of these
properties. Metals bond by sharing their valence electrons and some of the d-orbital electrons
between cations of the metal. These shared, freely-moving electrons are called a “sea of
electrons” and is how metallic compounds and elements bond as liquids and solids.
Page 11 of 28
Chemistry
Covalent or Molecular Compounds
Most compounds form when elements share valence electrons to form a covalent bond.
Covalent compounds include organic compounds, which are compounds formed by living
matter (for example, DNA, sugars, vitamins, and proteins), medicines, plastics and polymers,
compounds that are gases and liquids at normal temperatures and pressures, acids, and many
others. Most covalent compounds are combinations of nonmetallic elements, and may have
hundreds even thousands of elements bound together. The smallest individual set of atoms
making up each compound that has the properties of that compound is called a molecule.
Sharing of Electrons
Ionic bonding takes place by creating ions with filled s and p-orbitals, and sharing of electrons
will take place until the atoms reach the same stable electron configuration. For example, each
chlorine atom has 7 valence electrons, by sharing 1 valence electron each chlorine will have 8
(sharing electrons does not take away electrons it only adds electrons). Just like ionic
compounds, the covalent compounds will have 8 valence electrons around each atom to create
a stable electron configuration.
••
••
•• ••
••Cl• + •Cl••  ••Cl••Cl••
••
••
•• ••
Diatomic Elements
A set of elements, which includes chlorine, form molecules of two atoms at normal
temperature and pressures: H2, O2, N2, F2, Cl2, Br2, & I2. This group of elements includes the
halogens, hydrogen, oxygen, and nitrogen and can be recognized because they are the only
elements that have names ending with -gen or -ine.
••
••
•• ••
Fluorine (and other halogens) ••F• + •F••  ••F •• F••
••
••
•• ••
Hydrogen , H2
H• + •H 
Oxygen
••
••
•• ••
•O• + •O•  O•• •• O
••
••
•• ••
He
B
G
e
Transition Metals
Halogens
Si
As
Sb Te
57
89
72
Po
104
••
••
••
••
Page 12 of 28
Noble Gases
Semimetals 
Alkaline Earth Metals
Alkali Metals
H
H •• H
Chemistry
•N• + •N•  N•• •• •• N
•
•
Nitrogen
Notice that hydrogen has only 2 valence electrons around each atom after sharing. This is
because only the first energy level is filled, which only requires two electrons.
Multiple Bonds
For oxygen and nitrogen, the atoms must share more than one electron to reach an octet for
each atom. When two atoms share two pairs of electron then a double bond forms; when two
atoms share three pairs of electrons then a triple bond forms. A single bond forms with only
one shared pair of electrons. Notice that for many simple examples of sharing the number of
lone electrons an atom has equals the number of shared electrons in the compound.
Unlike ionic compounds, most covalent compounds have more than two atoms. When three,
four, five, and more atoms combine they share electrons until each atom reaches an octet
(although there are exceptions).
Elements
Electron Dot Diagram
Molecule
Name of Compound
C&H
•
C + H
•
H
••
 H •• C •• H
••
H
CH4
methane
N&H
••
N + H
•
••
 H •• N •• H
••
H
NH3
ammonia
O&H
••
O + H
••
••
 H •• O •• H
••
H2O
water
C & Cl
•
••
•C• + •• Cl
•
••
••
•• Cl ••
•• •• ••
 •• Cl •• C •• Cl ••
•• •• ••
•• Cl••
••
CCl4
carbon tetrachloride
C&O
•
••
 C + O
•
••
CO2
carbon dioxide
P&F
••
••
•• •• ••
•P• + •• Cl  •• Cl •• P •• Cl ••
•
•• •• ••
•• Cl ••
••
PCl3
phosphorous trichloride
•
••
••
H  + C + N  H •• C •• •• ••
N
•
•
HCN
hydrogen cyanide
H & C &N
••
••
 O •• •• C •• •• O
••
••
Page 13 of 28
Chemistry
Bonding electrons can be shown as lines connecting atoms and lone pair electrons can be
shown as lines above an atom (http://commons.wikimedia.org/wiki/File:H2O.svg#file):
••
H •• O •• H =
••
Naming of Simple Covalent Compounds
Many covalent compounds have common names like water, ammonia, sugar, acetic acid,
vitamin C, hemoglobin, etc. but there is a systematic naming system. Name the first element
and name the second with an -ide ending just like ionic compounds, but also tell how many
atoms are in the molecule. Some examples are: carbon dioxide is CO2 and carbon monooxide
is CO, phosphorous trichloride is PCl3 and phosphorous pentachloride is PCl5; and nitrogen
dioxide is NO2 and dinitrogen tetroxide is N2O4. The numbering prefixes are given below.
1
2
3
4
5
6
7
8
9
10
mono-
di-
tri-
tetra-
penta-
hexa-
septa-
octa-
nona-
deca-
Use mono- only with the second atom never with the first.
Complex Molecules and Lewis Structures
Often molecules are complex and it is not easy to figure out the dot structure for the molecule.
There are a set of rules that allow you to determine the electron dot diagram for molecules.
Steps for Writing Lewis Structures of Covalent Compounds from the Chemical Formula
1. Write the atoms in the order given in the chemical formula, place the atom with the
lowest electronegativity in the center (usually the first element) and multiple atoms in a
chemical formula surrounding the central atom.
2. Determine and add the valence electrons of all atoms. If there are charges add one
electron for each negative charge and subtract one electron for each positive charge.
3. Attach all atoms with a bond.
4. Subtract 2e– for each bond from the total number of valence electrons (step 2).
5. Add pairs of nonbonding electrons to outer atoms. Continue adding pairs so that the
number of pairs on each outer atom are the same or nearly the same.
6. Continue adding electron pairs until you run out of valence electrons (step 4). If you
have octets around every outer electron and still have more electrons to add, then add
the rest of the electrons to the central atom(s).
7. All atoms should have octets. If there are too few electrons, then share nonbonding
outer electrons with the inner atom to make a double or triple bond.
8. Atoms in 3rd or more row can have expanded octets. Hydrogen will only form a single
bond without any outer electrons. Boron has only three pairs of bonding electrons
around it.
9. Finally, if there are still problems with octets, then move electron pairs if needed to
deficient outer atoms.
Page 14 of 28
Chemistry
Examples
Molecule
Step 1
Step 2
NO21–
SCl4
O N O
Cl
Cl S Cl
Cl
O = 6 e– X 2 = 12
N = 5 e–
(1–) = 1 e–
total e– = 18
S = 6 e–
Cl = 7 e– X 4 = 28
total e– = 34 e–
Step 3
O •• N •• O
Cl
••
Cl •• S •• Cl
••
Cl
Step 4
18 - 4 = 14
34 - 8 = 26
••
•• Cl ••
•• •• •• + 2e–
•• Cl •• S •• Cl •• ——
•• •• ••
•• Cl ••
••
Step 5 & Step 6
•• •• ••
•• O •• N •• O ••
••
••
Step 7
•• •• ••
•• •• ••
•• O •• N •• •• O •• or O •• •• N •• O ••
••
••
No change
Step 9
No change
There are five pairs of electrons around the
S, this is an expanded octet.
Molecular Shape
The shape of a molecule is important in determining its properties; for example, molecular
shape affects a molecule’s solubility in water. A theory called Valence Shell Electron Pair
Repulsion theory, VSEPR theory, is used to explain the different shapes that molecules adopt
and hybrid molecular orbitals (as apposed to the atomic orbitals from Unit 1) explain how
electrons can adopt different geometries.
VSEPR theory states that electron pairs of a central atom shift in the molecular geometry until
the electron-electron repulsions between both bonding and nonbonding electron pairs are as
small as possible. The table below shows the ideal geometries for most simple covalent
compounds.
Page 15 of 28
Chemistry
S
S = Symmetrical
AS = Asymmetrical
S
AS
S
AS
AS
S
AS
AS
S
S
AS
S
AS
S
Molecular Shapes from VSEPR Theory
http://commons.wikimedia.org/wiki/File:VSEPR_geometries.PNG#filelinks
In the table above, the steric number refers to the number of electron pairs, both bonding and
nonbonding, that surround the central atom. The number of nonbonding pairs called lone
pairs, is listed across the top of the table. When you have matched the Lewis structure with the
steric number and the number of lone pairs, you can find the geometry of the molecule and the
name of the molecular geometry.
Lewis Structures and the Corresponding Molecular Geometries
Page 16 of 28
Chemistry
Molecule
Complete Lewis Structure
H2O
••
H •• O •• H
••
Molecular Geometry
without
with
lone pair
lone pair
CCl4
••
•• Cl ••
•• •• ••
•• Cl •• C •• Cl ••
•• •• ••
•• Cl••
••
AsCl5
Page 17 of 28
Chemistry
Atomic orbitals, which are
spherical for s-orbitals and a
figure-8 shape for p-orbitals,
cannot explain the
tetrahedral shape of carbon
tetrachloride, CCl4, or the
other molecular shapes. The
shared valence electrons in
covalent compounds require
new wave equations to
formulate molecular orbitals
that explain the various
molecular shapes. Molecular
orbitals are hybrid orbitals,
which are combinations of
atomic orbital. Common
molecular orbitals are: 2 sp
orbitals from combining 1 sorbital and 1 p-orbital, three
sp2 orbitals result from a
combination of 1 s-orbital
and 2 p-orbitals, four sp3
orbitals where 1 s-orbital and
3 p-orbitals combine, five
sp3d orbitals form as 1 sorbital, 3 p-orbital, and 1 dorbital mix, and six sp3d2
orbitals, which is a
combination of 1 s-orbital, 3
p-orbital, and 2 d-orbitals.
The 4 atomic orbitals
hybridize to become
4 molecular orbitals
The four
molecular
orbitals are
sp3
The 3 atomic orbitals
hybridize to become
3 molecular orbitals
The three
molecular
orbitals are
sp2
http://commons.wikimedia.org/wiki/File:S-p-Orbitals.svg & http://commons.wikimedia.org/wiki/File:Sp3Orbital.svg & http://commons.wikimedia.org/wiki/File:Sp2-Orbital.svg
Polar and Nonpolar Covalent Compounds
Many molecules dissolve in water and some do not. For example sugar, vinegar, and vitamin
C all dissolve in water, but wax, oil, and vitamin E do not dissolve well if at all. The reason
Page 18 of 28
Chemistry
that some substances dissolve in water and some do not is because some molecules are polar
and some are not polar.
Polar molecules have a shift in electron density in the molecular orbitals away from the center
of the molecule. This means that the electronegativities of the different atoms are unbalanced
and the electrons will be more likely to be found over the more electronegative parts of the
molecule (recall that electrons move randomly so they will move everywhere, but the pull of
the most electronegative nucleus keeps them near that atom more often). Water, H2O, itself is
polar. Each H has an electronegativity of 2.1 and oxygen, O, has an electronegativity of 3.5.
The electron density of the shared electrons is shifted towards oxygen since it is most
electronegative. The molecule H2O has a dipole moment, which is a shift of negative charge
towards one part of the molecule (this is not a transfer of electrons like ionic compounds but a
shift of electron density).
Water Molecule, H2O
Electron Density of the
molecule shifts to oxygen the
most electronegative atom
A Dipole Moment represents
a substantial shift in charge
with the negative end at the
arrowhead
Polar molecules must have two characteristics: 1) The difference in electronegativity (∆EN is
difference in electronegativity) between bonding atoms must be between 0.4 and 1.8. When
the electronegativity is greater than 1.8 the bonding is becoming ionic since the electron is
being transferred not shared. When the electronegativity is less than 0.4 then the electron
density in a bond is not shifting enough to create a dipole moment. 2) The molecule must be
asymmetrical so that the pull of electrons is not cancelled out by equal pulls in opposite
directions. For example, CCl4, carbon tetrachloride, is nonpolar because, while the bonds are
polar with ∆EN = 0.6 (C = 2.5 and Cl = 3.1), the molecule is tetrahedral and symmetrical so the
electron density moves towards each Cl in all directions of the molecule. Polar molecules
include: CHCl3, NH3, & COH2 and nonpolar molecules include: CH4, PH3, and SCl5. To figure
out why these are polar you must first draw the Lewis structure, determine the VSEPR
geometry, find the difference in electronegativity and finally decide whether both the
0.4≤∆EN≤1.8 and the geometry is symmetrical are true.
States of Matter.
Ionic compounds are solids. The electrostatic forces hold the ions locked into a solid state. But
covalent, or molecular, compounds can be gases, liquids, and solids. For a covalent compound
to be a solid the individual molecules must attract each other strongly or the molecular mass of
the substance must be high. The attraction between molecules is called intermolecular forces
and happens in three different ways with different amounts of attraction. When the electron
Page 19 of 28
Chemistry
density shifts in a molecule so that there is a positively charged end and negatively charged
end, then the molecules can be attracted to each other through the attractive force of the
positive and negative charges. These attractive forces are not bonds because they are 100 of
times weaker than the molecular and ionic bonds (sometimes referred to as intramolecular
forces)
Hydrogen bonds are the strongest intermolecular
bond (not inter- is a prefix meaning between).
Hydrogen bonds only occur with a select group of
atoms. When HF, HOR, or HNR2, where R is H or
more of the molecule, are present then H bonding
can occur with a O or N in a molecular compound
or with HF. The double helix structure of DNA is
due to the hydrogen bonds between the
nucleotides. Water also has hydrogen bonds that
account for many of its unique properties.
Hydrogen bonds in DNA
(the dotted lines)
http://commons.wikimedia.org/wiki/File:D
NA_chemical_structure.svg
Dipole-dipole forces are the next strongest intermolecular forces. Polar molecules have
shifted electron density that creates partially positive and partially negative charges at
opposite ends of the molecule. The molecules will attract each other with the positive end
attracting the negative end and visa versa. Unlike ionic compounds the molecules have strong
bonding forces in the molecule but the intermolecular forces are easily broken.
δ–
δ+
δ–
δ+
δ+
δ–
δ+
δ–
These diatomic (two atom) molecules have dipole-dipole attractions. Because of the
high electronegativity of one of the atoms the electron density has shifted to create a
partial negative charge, δ–, (on top on the first molecule) and partial positive charge, δ+.
Intermolecular bonding occurs through attraction of the positive and negative charges.
http://commons.wikimedia.org/wiki/File:Polare_atombindung.png
The weakest intermolecular forces are London Forces, also called Van der Waals forces. This
intermolecular force is due to temporary dipoles. Since electrons are in constant motion each
nonpolar molecules will form positive and negative end over time (the bonding electrons will
randomly show up all on one side). A nearby nonpolar molecule will adjust its electron
density because negative charges repel and positive charges attract electrons. This induces a
Page 20 of 28
Chemistry
dipole in a neighboring molecule so the two have an attraction. This is the mechanism for
intermolecular attraction between nonpolar molecules.
Nonpolar molecules have evenly distributed electron density as shown by the
first four molecules. But random motion can create a temporary dipole, the first
molecule to the right of the arrow. The molecules near this temporary dipole react
to the change in electron density and create their own dipoles.
They have induced dipoles.
http://de.wikipedia.org/w/index.php?title=Datei:Unpolare_atombindung.png&filetimestamp=20070704231926
London forces are stronger when the valence electrons are dispersed widely in large shells.
This is because the electrons will create temporary dipoles that are more difficult to remove
when the electrons are so distant from the nucleus and so dispersed in the large electron cloud.
The most prominent example of the change in the intermolecular forces due to atomic size is
the halogen family. Fluorine, F2, and chlorine, Cl2 are both gases. Like all halogens they are
nonpolar because the difference in electronegativity is zero. As gases the molecules of the
diatomic molecules do not attract each other and the molecules of gas move randomly and
distribute themselves throughout the volume of their container. Bromine, Br2, is a liquid; the
only other liquid element at typical temperatures and pressures is mercury, Hg. Bromine has
electrons filling the 4th energy level. The molecules of bromine have some attraction for each
other because the substance is a liquid, in which the molecules move randomly, but they slide
between molecules with small attractive forces holding the particles from separating to
become a gas. Finally, iodine, I2, is a solid. The temporary dipole is longer-lived with the
outer electrons in the 5th energy level and the molecules have longer, stronger attractions.
(Some of the changes in the state of matter—from gas to liquid to solid—are due to the larger
molecular weight, but temporary dipole changes have the most prominent role).
Differences in Attractive Forces
When molecules have strong intermolecular forces they are most likely to be solids. As solids
every molecule is locked in place by the strong attractions. As the intermolecular forces are
reduced, the molecules may be liquids where the forces keep the particles near to each other,
but the forces are not strong enough to stop random motion. Finally, weak intermolecular
forces are present in gases that cannot maintain any hold on each other so they move
randomly and in their own path filling the volume of their container. So water, H2O, is a
liquid despite its small molecular mass, while CH4, methane, is a gas (the natural gas used in
stoves and water heaters).
Remember that molecular weight has something to do with whether molecules are solids.
Waxes and your skin are essentially nonpolar, but they are solid, because they have large
numbers of atoms in each molecule (wax is C15 H31 CO2 C30 H61) so the molecules are massive
and have long, long molecular orbitals so induced and temporary dipoles are long lasting.
Page 21 of 28
Chemistry
Properties Affected by Intermolecular Forces.
Intermolecular forces affect the physical changes that substances undergo and their state of
matter at standard temperature and pressures. State of matter is determined by the type of
intermolecular forces. The stronger the intermolecular force the more likely it is a solid or a
liquid (generally, only molecules with high molecular masses will be a solid). Boiling point
and melting point (the temperature at which a substance boils or melts) is also affected by
intermolecular forces, with higher temperatures related to stronger intermolecular forces.
Vapor Pressure, which is the pressure of the gas molecules of a pure substance that evaporate
from a pure substance at a temperature below the boiling point, is higher for weaker
intermolecular forces.Polyatomic Ions and Their Ionic Compounds
Some molecular compounds are most stable when they have one or more extra electrons
added or removed these are polyatomic ions. Most polyatomic ions are negative and one
common one is positive, NH4+, ammonium. These ions will react with elements or will trade
places with another ion to create ionic compounds with other ions. Here is a list of common
polyatomic ions.
Some Common Polyatomic Ions
1+
1–
1–
2–
3–
NH4+
ammonium
OH–
hydroxide
C2H3O2– or
CH3COO–
acetate
CO32–
carbonate
PO43–
phosphate
NO3–
nitrate
ClO–
hypochlorite
SO42–
sulfate
NO2–
nitrite
ClO2=
chlorite
SO32sulfite
HCO3–
hydrogen
carbonate
ClO3–
chlorate
CrO42–
chromate
CN–
cyanide
ClO4–
perchlorate
C2O42–
oxalate
Polyatomic ions used just like ions from elements, except parentheses are used to keep the
molecule together when you have to have more than one polyatomic ion. For example,
potassium carbonate has the ions K+ and CO32– which leads to a chemical formula of K2CO3.
Or, ammonium sulfide has the ions NH4+ and S2–and the chemical formula (NH3)2S.
Examples of Chemical Formulae of Compounds with Polyatomic Ions
Name of Ionic Compound
Ions
Chemical Formula
calcium chlorite
Ca2+
ClO2–
Ca(ClO2)2
aluminum carbonate
Al3+
CO32–
Al2(CO3)3
Page 22 of 28
Chemistry
iron (III) nitrite
Fe3+
Chemical Formula
Na2SO4
Mg(ClO)2
(NH4)2CrO4
Pt(SO4)2
NO2–
Ions
Na+
Mg2+
NH4+
Pt4+
Fe(NO2)3
Name of Ionic Compound
SO42–
sodium sulfate
ClO–
magnesium hypochlorite
CrO42–
ammonium chromate
SO42–
platinum (IV) sulfate
This should all be reminiscent of writing chemical formulae for elements and for determining
the name. All the polyatomic ions have endings different from the -ide used for ionic
compounds of elements (sodium chloride). This makes writing the ion for polyatomic ions as
simple as looking on a table to match the name of the ion with its chemical formula and
charge. The hardest problem is like the last one, where the cation is a transition metal with
unknown charge. But the key is to always balance charge with the number of cations versus
anions, or when given an unknown cation charge balance the number of electrons by choosing
the appropriate charge.
Summary
Compounds are pure substances that are composed of more than one element. Ionic
compounds are held together by electrostatic attractions between cations and anions. Most
cations are metal elements, which have electron configurations with three or fewer valence
electrons, and most anions are nonmetal elements, which have five to seven valence electrons.
The charge of cations is positive and the number of electrons lost and the charge of anions is
negative and the number of electrons gained. The number of electrons gained or lost leads to a
stable, noble gas electron configuration. After electrons are transferred the ions combine to
form a crystal lattice. The number of cations and anions is not always one-to-one in the
chemical formula, but must be in a ratio that makes the number of electrons lost and gained be
equal. This chemical formula of an ionic compound shows the ratio of elements using
subscripts to denote the number of elements.
Another type of compound is covalent compounds or molecular compounds. These
compounds are composed of molecules, which are individual groups of atoms bonded with
covalent bonds. The molecules form when electrons are shared between elements that are
usually nonmetals. Elements in covalent compounds share electrons until they have an octet
of electrons around them. This may require double bonds or triple bonds that have 2 pairs of
electrons or 3 pairs of electrons shared between two elements. Molecules can be either polar or
nonpolar and the shape of the molecule is one factor that determines this property. The shape
of a molecule can be determined by using VSEPR theory, which holds that both bonding and
nonbonding pairs of electrons will tend to move as far apart as possible because of the
electron-electron repulsions. If a shape is symmetrical then the molecule is nonpolar. If the
shape of a molecule is asymmetrical and the molecule has bonds that have electronegativity
differences of 1.8 to 0.4 then the molecule will be polar.
Page 23 of 28
Chemistry
Polar molecules have intermolecular forces called dipole-dipole attractions. These
intermolecular forces hold molecules next to each other and result in the substance being a
solid or liquid (compared to a nonpolar compound of the same molecular mass). Nonpolar
molecules also exhibit a weak intermolecular attraction called Van der Waal’s forces or
London forces. But these weak forces do not hold molecules together strongly so the
compounds are gases or easily boiled or melted.
Polyatomic ions are molecules with added or removed electrons which makes them anions or
cations respectively. They act like other ions, but the chemical formula of polyatomic ions is
listed in a table or needs to be memorized.
Page 24 of 28
Chemistry
Metallic
compoun
ds are
combinati
ons of
metal
elements
as
cations
with the
valence
electrons
free to
move
among
the
cations.
These
shared
valence
electrons
form a
bond for
metals
called a
“sea of
electrons
”, which
is
responsib
le for
many of
the
propertie
s of
metals.Ty
pe of
Bond
Metal or
nonmetal
elements
Share or
Transfer
electrons
Melting Point
& Boiling
Point (high or
low)
State of
Matter
Ionic
metal +
nonmetal
Transfer
high
solid
Covalent
nonmetal +
nonmetal
Share
low usually
gas, liquid and
solid (oxygen,
water, sugar)
Page 25 of 28
Chemistry
Metallic
compoun
ds are
combinati
ons of
metal
elements
as
cations
with the
valence
electrons
free to
move
among
the
cations.
These
shared
valence
electrons
form a
bond for
metals
called a
“sea of
electrons
”, which
is
responsib
le for
many of
the
propertie
s of
metals.Ty
pe of
Bond
Metallic
Metal or
nonmetal
elements
Share or
Transfer
electrons
Melting Point
& Boiling
Point (high or
low)
State of
Matter
metal + metal
Share ; a “sea
of electrons”
high but
mercury is a
liquid
solid (one
liquid)
Page 26 of 28
Chemistry
Type of
Bond
Hard or Soft
(brittle vs.
malleable)
Dissolves in
Water
Conducts
Electricity
Other Notes
hard
Often but not
always (like
chalk)
yes if it
dissolves,
yes if it is a
liquid
not as a solid
The positive ion,
cation, combines with
a negative ion, anion,
to create a neutral
compound.
Minerals are ionic.
Covalent
soft usually
Nonpolar
covalent
compounds
do not, but
Polar covalent
compounds do
dissolve.
no
The body’s soft tissue
is nearly all covalent
compounds.
Organic matter is
covalent.
Metallic
soft (malleable,
beaten into
sheets & ductile,
pulled into a
wire)
no
Yes, as a solid,
liquid, or gas
(but it doesn’t
dissolve)
Metals don’t form
compounds, they form
mixtures, called alloys,
or they are elements.
Ionic
Chemical Bonds
2. Biological, chemical, and physical properties of matter result from the ability of atoms to
form bonds from electrostatic forces between electrons and protons and between atoms and
molecules. As a basis for understanding this concept:
a. Students know atoms combine to form molecules by sharing electrons to form covalent or
metallic bonds or by exchanging electrons to form ionic bonds.
b. Students know chemical bonds between atoms in molecules such as H2, CH4, NH3,
H2CCH2, N2, Cl2, and many large biological molecules are covalent.
c. Students know salt crystals, such as NaCl, are repeating patterns of positive and negative
ions held together by electrostatic attraction.
d. Students know the atoms and molecules in liquids move in a random pattern relative to
one another because the intermolecular forces are too weak to hold the atoms or molecules
in a solid form.
e. Students know how to draw Lewis dot structures.
f. * Students know how to predict the shape of simple molecules and their polarity from
Lewis dot structures.
g. * Students know how electronegativity and ionization energy relate to bond formation.
Page 27 of 28
Chemistry
h. * Students know how to identify solids and liquids held together by Van der Waals forces
or hydrogen bonding and relate these forces to volatility and boiling/melting point
temperatures.
Starred standards are non-tested standards on the California Standards Test
Contributed by Kenneth Pringle
Edited by Kathleen Duhl
Formatted and Wiki Contribution by Christine Mytko
Page 28 of 28