Molecular Bonds
(Putting Elements
Together)
Molar Mass
• Each atom has an atomic mass
• Molar mass is the atomic mass of all the
atoms in the molecule summed together
• For Example:
H2O = 2 x Atomic Mass of H +
1 x Atomic Mass of O
Counting Atoms in a Molecule
In the example, NH3, the subscript 3 only applies to the
hydrogen.
– Therefore: there is 1 N and 3 H in ammonia
In the example, 3Ca3(PO4)2, the number of atoms changes
due to the Coefficient in front of the molecule
The 3 is multiplied to the Ca, P and O
The subscript 2, multiplies the P and O
3Ca3(PO4)2
3 Ca3 ( P O4 )
This means that
there are 3 x 3 Ca,
3 x 2 P and 3 x (4 x 2) O
2
Bonds. . .
No, not that kind – bonds between
atoms to form molecules
It all depends upon the atom’s valence
(outer shell) electrons
These are the e- in the last Energy
Level (n = 1 through 7)
Figure these out using the Periodic Chart
and/or Lewis Dot Diagrams
The Roman Numerals Tell You How Many Valence Electrons for the
Primary or Representative Elements;
The Valence Electrons for the Transition Elements Vary
I
II
III IV V VI VII VIII
Group I is monovalent; II is divalent; III is
trivalent; IV is tetravalent; V is back to being
trivalent (since three e- openings); VI is
divalent; VII is monovalent and VIII has a
complete octet, so these seldom react or bond
Bond Types (In General):
• Pure or Non-Polar Covalent
Χ difference = 0 to 0.5 on the Pauling EN Scale
The pair of e- shared are done so equally
Two nonmetals bonded together
• Polar Covalent
A shared pair of e-, but not equally
χ difference = 0.5 to 1.6
Molecule has Partial + and – Charges
• Ionic Bonds
χ difference = 1.7 or higher to the maximum of 4.0
Metal bonded with a nonmetal
• Metallic Bonds are similar to Ionic Bonds
Metallic Bonds
• Two or more metals mixed are called
alloys
• Two major formats
– Interstitial and Substitutional
These bonds permit the
roaming of e- which creates
a sea of dissociated eCalled the Electron Sea
Model
Ionic Bonds
• These are the bonds between a metal and
a nonmetal
• The metal Ion is positively charged and
called a cation
• The nonmetal Ion is negatively charged
and called an anion
• The bonded molecule should be neutrally
charged when finished
Knowing where the metals and nonmetals are
on the table will make your life easier
Let’s take a moment to discuss
polyatomic ions. . .
• This is a molecule that acts as a cation or
anion
• For example:
NH4+ ammonium
ClO4- perchlorate
HCO3- bicarbonate
CrzO7-2 chromate
ClO3- chlorate
N3- azide
CN- cyanide
OH- hydroxide
NO3- nitrate
C2H3O2- acetate
• Don’t PANIC – I gave a list to you!
In an Ionic Bond – one or more electrons are lost or gained
by the atoms involved
This allows the atoms to have a complete valence shell –
following the octet rule
In an Ionic Compound – balance the
molecule using the criss-cross rule
Mg +2 +
Cl-1
Mg Cl2
The one is understood.
This applies even if using a polyatomic ion
NH4+ +
O-2
(NH4)2O
The parentheses are used to keep
the polyatomic together
Pb+4
+
CO3-2
Pb2 (CO3)4
Pb(CO3)2
and this can be simplified by
reducing the subscripts to
Naming Ionic Compounds is really simple:
1. Name the cation (metal) using its proper
name; if it is a polyatomic, do the same
2. Then, using the stem of the anion (nonmetal),
simply add the suffix “ide”
Zinc + Chlorine = Zinc Chloride
Iron + Oxygen = Iron Oxide
Lithium + Cyanide = Lithium Cyanide
Ammonium + Fluorine = Ammonium Fluoride
Cobalt + Phosphorous = Cobalt Phosphide
Transition Metals present an issue for
balancing and naming molecules since
they can have varying oxidation states
For example:
Manganese can be a +2 or +3
Iron can be a +2 or +3
Lead can be a +2, or even a +4
Copper is a +1 or +2
Gold is usually a +1 or +3
And Hydrogen is a +1 or a -1!
Transition Metals
• To determine the correct Roman Numeral
to place after the metal:
Roman Numeral = - (Charge # anion)(#anions)
(# cations)
This is needed because, for example,
iron chloride can be either FeCl2 or FeCl3;
or iron (II) chloride or iron (III) chloride
Therefore – Ionic Bonds are:
•
•
•
•
Metal
+
+ ion
cation
monatomic
(except NH4+)
left of steps
Nonmetal
- ion
anion
monatomic or
polyatomic
right of steps
Reactions are Exothermic
Form Crystal Lattice Structures
Covalent Compounds
• These can be monatomic or polyatomic
compounds
• It is a bond between two nonmetals
• They share a pair of electrons
• They can be subgrouped into polar or
nonpolar
• If a binary compound (2 atoms) – use
the same naming rules as in Ionic
Compounds
• If it has more than two atoms – need to
use the prefixes
Number
1
2
3
4
5
6
Prefix
Mono
Di
Tri
Tetra
Penta
Hexa
Number
7
8
9
10
11
12
Prefix
Hepta
Octa
Nona
Deca
Undeca
Dodeca
Naming Covalent Compounds
Process:
1. Prefix Indicating # + full name of first
nonmetal
2. Prefix Indicating # + root name of
second nonmetal + suffix “ide”
3. Watch for polyatomics and use their
proper names
For Example:
•
•
•
•
P4S10 becomes Tetraphosphorous Decasulfide
P2O5 Becomes Diphosphorous Pentaoxide
SF6 becomes Sulfur Hexafluoride
SiBr4 becomes Silicon Tetrabromide
Covalent Bonds can be Polar or Nonpolar
A nonpolar has no discernable
negative or positively charged sides
(EN difference is 0)
A polar covalent bond means one
side is negative and the other positive
Electronegativity
Percent Ionic
Bond
Difference
Character
Type
• 0.2
1%
Non-polar
• 0.4
4
Covalent
• 0.5
• -------------------------------------------------------------------------• 0.6
9
• 0.8
15
• 1.0
22
Polar
• 1.2
30
Covalent
• 1.4
39
• -------------------------------------------------------------------------• 1.6
47
Ionic if metal/nonmetal
• 1.8
55
Polar Cov. if non/nonmetal
• 2.0
63
• -------------------------------------------------------------------------• 2.2
70
• 2.4
76
Pure Ionic
• 2.6
82
• 2.8
86
• 3.0
89
• 3.2
92
•
•
•
Some elements are able to form more
than one oxyanion (polyatomic ions that
contain oxygen), each containing a
different number of oxygen atoms.
For example, chlorine can combine with
oxygen in four ways to form four different
oxyanions: ClO4-, ClO3-, ClO2-, and ClO(Note that in a family of oxyanions, the
charge remains the same; only the
number of oxygen atoms varies.)
The most common of the chlorine
oxyanions is chlorate, ClO3-. In fact, you
will generally find that the most common
of an element’s oxyanions has a name
with the form (root)ate.
• The anion with one more oxygen atom
than the (root)ate anion is named by
putting per- at the beginning of the root
and -ate at the end. For example, ClO4- is
perchlorate.
• The anion with one fewer oxygen atom
than the (root)ate anion is named with -ite
on the end of the root. ClO2- is chlorite.
• The anion with two less oxygen atoms
than the (root)ate anion is named by
putting hypo- at the beginning of the root
and -ite at the end. ClO- is hypochlorite.
Oxyanion Example
•
•
•
•
ClOClO2ClO3ClO4-
Hypochlorite
Chlorite
Chlorate
Perchlorate
• Some compounds have common
names as well as their scientific names
– you should learn these and others!
– NO
– H2O
– NH3
– CH4
– C4H10
nitrogen monoxide
nitric oxide
dihydrogen monoxide water
nitrogen trihydride
ammonia
carbon tetrahydride
methane
tetracarbon decahydride butane
Some atoms are Diatomic – KNOW THESE!
H2 N2 O2 F2 Cl2 Br2 and I2
and P is usually found as P4
while Sulfur is found as S8
Other elements will bond beyond the octet
rule – like PCl5, and the noble gas Xe
bonds with F in XeF6, XeF2, XeF4, and
XeO4 – and this is due to a thing called
“hypervalence” or “expanded octet”
Molecular Geometry
• The 3-Dimensional Shapes of Molecules depend
upon the valence e-’s of the atoms involved
• Valence Bond Theory and VSEPR Model both
use the same shapes
– Basically – they focus on covalent bonds with
the shared bonding pairs of electrons (BP)
– The assumption is made that the molecule will
adopt a geometry to minimize the repulsion
between e-’s
• The General Shapes:
Basic Geometry Bond Angles
•
•
•
•
Linear
Trigonal Planar
Tetrahedral
Trigonal Bipyramidal
• Octahedral
180o
120o
109.5o
90o and
120o
90o
Molecular Orbital Theory
MOT uses atomic orbitals (AO), e- λ’s and edensity regions to examine bonds
This is the end of Part I
• Next:
– Van der Waals and London Dispersion Forces
– Polarity
– Intermolecular Forces
– Lewis Dot Diagrams with Covalent Bonds
– Determining Molecular Structure
– Resonance Structures