Chapter 10
Chemical
Bonding
Tro, 2nd ed.
Nonmetals
arefound
foundto
tothe
theleft
right
metalloids.
Metals are
ofofthethe
metalloids
LEWIS STRUCTURES OF
ATOMS
Metals form cations and nonmetals form anions to
attain a stable valence electron structure.
These rearrangements occur by losing, gaining or
sharing electrons.
The Lewis structure of an atom is a representation
that shows the valence electrons for that atom.
Valence electrons: the electrons that occupy the
outermost energy level of an atom.
Valence electrons are responsible for the electron
activity that occurs to form chemical bonds.
The Lewis structure of an atom
uses dots to show the valence
electrons of atoms.
B
2
1
2s 2p
The number of dots equals the number of s
and p electrons in the atom’s outermost shell.
Lewis Structures of the first 20 elements
Notice that C’s e- config is 2s22p2.
CHEMICAL BONDING
Atoms will do one of three “things” to get to a noble
gas electron configuration:
1. Take electrons from another atom
2. Give electrons to another atom
3. Share electrons with atom(s)
In choices 1 & 2: cause ions to form, then ionic
bonds
In choice 3: sharing electrons results in covalent
bonds
With the exception of hydrogen & helium, this
structure consists of eight electrons in the
outermost energy level (The Octet Rule)
The Ionic Bond
Cations (+) have given up e-s and anions
(-) have gained e-s, and now have
opposite electrical charges
Results in strong electrostatic force of
attraction
All cations and anions exist in crystal
lattices, defined geometric
structures with repeating 3-D
pattern
The 3s
electron
of sodium
transfers
to the 3p is
orbital
of chlorine.
The
force
holding
Na+ and
Cl- together
an ionic
bond.
A sodium ion (Na+) and a chloride ion (Cl-) are formed.
Lewis representation of sodium chloride formation.
2+ and two Cl- together are
2+
-) are
The
forces
holding
Mg
Two
3s
electrons
of
magnesium
the
A magnesium ion (Mg ) and twotransfer
chloride to
ions
(Cl3p
ionic bonds.
orbitals
of two chlorine atoms.
formed.
In
NaCl
theiscrystal
made up
each
of sodium
cubic crystals.
ion is
surrounded by six chloride ions.
In the crystal each chloride ion is
surrounded by six sodium ions.
The ratio of Na+ to Cl- is 1:1
There is no molecule of NaCl
The Ionic Bond
Using your Periodic Table, determine
the cation and anion each atom is
likely to form, then write the Lewis
structures and made the compounds.
Finish by writing the compound’s
chemical formula.
Practice: Al & F, Mg & O, Na & O, Na
& N, Al & O
A sodium ion is smaller than a sodium atom
because:
(1) the sodium atom has lost its outermost
electron.
(2) the 10 remaining electrons are now attracted
by 11 protons and are drawn closer to the
nucleus.
A chloride ion is larger than a chloride atom
because:
(1) the chlorine atom has gained an electron
and now has 18 electrons and 17 protons.
(2) The nuclear attraction on each electron
has decreased, allowing the chlorine to
expand.
Transition Metals form
cations – a little different
Transition metals lose their “s”
electrons first, because they are in
the highest principle energy level,
then they lose their “d” electrons.
Zn  Zn2+ + 2 e- Cu  Cu+ + 1e[Ar]4s23d10[Ar]3d10
[Ar]4s13d10[Ar]3d10
COVALENT BONDING
A covalent bond consists of a pair of
electrons shared between two atoms.
In the millions of chemical compounds that
exist, the covalent bond is the
predominant chemical bond.
Substances which covalently bond exist as
molecules.
Carbon dioxide bonds covalently.
It exists as individually bonded
covalent molecules containing
one carbon and two oxygen
atoms.
The term molecule is not used when
referring to ionic substances.
Instead they are called Formula Units.
Sodium chloride bonds ionically.
It consists of a large aggregate
of positive and negative ions. No
molecules of NaCl exist.
COVALENT BONDING
Nonmetal Atoms have deficiency of electrons in
outermost shell and want to gain electrons to get
full shell
Since two nonmetal atoms both want more electrons,
they will share electrons to get full shell
H has 1 e- and wants 2
Cl has 7 e-s and wants 8
Both satisfied if they share a pair of electrons
between them
Each contributes 1 e- to the pair and each gets to
share the 2 e-s in the pair
H has 2 e-s and Cl has 8 e-s and they are “HAPPY”
COVALENT BONDING
IMPORTANT: in giving e-s to be shared, atom
actually gains e-s
The number of e-s an atom contributes to be shared
is equal to the number of e-s it needs to have an
octet! (or a full shell)
A pair of shared e-s is called a covalent bond
1 pair of e-s between two atoms = single bond
2 pairs of e-s betwn two atoms = double bond
3 pairs “ “
“
“ “
= triple bond
LEWIS STRUCTURES OF
COMPOUNDS
In writing Lewis structures, the most
important consideration for forming a
stable compound is that the atoms attain
a noble gas configuration.
The most difficult part of writing Lewis
structures is determining the
arrangement of the atoms in a molecule or
an ion.
In simple molecules with more than two
atoms, one atom will be the central atom
surrounded by the other atoms.
Cl2O has two possible
arrangements.
The two chlorines can be bonded to each other.
Cl-Cl-O
The two chlorines can be bonded to oxygen.
Cl-O-Cl
Usually the single atom will be the central
atom. (also usually the “leftist or lowest” on
the Periodic Table)
Valence Electrons of Group A Elements
Atom
Group
Valence Electrons
Cl
VIIA
7
H
IA
1
C
IVA
4
N
VA
5
S
VIA
6
P
VA
5
I
VIIA
7
Procedures for Writing
Lewis Structures
Step 1. Obtain the total number of valence electrons to be
used in the structure by adding the number of valence
electrons in all the atoms in the molecule or ion.
If you are writing the structure of an ion, add one electron
for each negative charge or subtract one electron for
each positive charge on the ion.
Step 2. Write the skeletal arrangement of the atoms and
connect them with a single covalent bond (two dots or one
dash). Choose the “leftist or lowest” element as the central
atom. Arrange terminal atoms symmetrically around the
central atom.
Hydrogen, which contains only one bonding electron, can
form only one covalent bond.
Oxygen atoms usually have a maximum of two covalent
bonds (two single bonds, or one double bond).
Procedures for Writing
Lewis Structures
Step 3. Subtract two electrons for each single
bond you used in Step 2 from the total number
of electrons calculated in Step 1.
This gives you the net number of electrons
available for completing the structure by
adding lone pairs of electrons to the terminal
atoms until they have an octet. Any remaining
electrons become lone pairs on the central
atom.
Procedures for Writing
Lewis Structures
Step 4. Check that each atom is satisfied. If one
atom doesn’t have an octet, move lone pairs of
electrons in as bond pairs to make multiple covalent
bonds. Do this symmetrically.
Step 5. Check the total number of electrons in
the structure and make sure it matches the number
of valence electrons in step 1.
(Also learn the number of bonds an atom prefers to
make:
H and F always 1 bond and terminal atom; C mostly
4 (and usually a central atom); halogens mostly 1; O
and S mostly 2; N and P mostly 3)
Write the Lewis structure for H2O.
The total number of valence electrons is eight, two from the two
hydrogen atoms and six from the oxygen atom.
The two hydrogen atoms are connected to the oxygen atom
which is central. Write the skeletal structure:
:
H:O
H
or
H:O:H
Place two dots between the hydrogen and oxygen atoms to form the
covalent bonds.
Subtract the four electrons used from eight valence electrons to obtain four
electrons yet to be used around the oxygen. (Why not the H?)
H:O:
H
or
: :
: :
Distribute the four remaining electrons in lone
pairs around the oxygen atom. (Hydrogen atoms
cannot accommodate any more electrons. NEVER
have more than 1 bond to H or have lone pairs
around H.)
H:O:H
Thearrangements
shape of the molecule
not shown because
by the Lewis
These
are Lewisisstructures
each
structure.
atom
has a noble gas electron structure.
Covalent bonding with equal
sharing of electrons occurs in
diatomic molecules formed
from one element.
hydrogen
chlorine
iodine
A dash may replace a pair of dots that
represent a bond:
H-H
nitrogen
Practice Lewis
Structures
Look in packet for practice sheet and
work with one partner to draw the
Lewis structures on separate paper.
Bring them up to show on the document
camera.
Complex Lewis
Structures
Do the Lewis structures for the
following with a partner:
HCN, CH4, SO3, CH3OH, SF6, PCl3, NO2
Some will have EXCEPTIONS to the
Octet Rule.
Complex Lewis
Structures
Exceptions to Octet Rule:
Expanded valence shell: any central atom
with outermost e-s in period 3 or below
has d orbitals available for bonding and can
hold 10 or 12 e-s
Electron deficient or free radical
structures: have less than 8 e-s and will be
very reactive compounds
Complex Lewis
Structures
There are some molecules and polyatomic ions for
which no single Lewis structure consistent with all
characteristics and bonding information can be
written.
When more than one structure satisfies the rules,
we call them resonance structures.
Real molecule is a hybrid of all possible Lewis
structures.
Resonance stabilizes the molecule.
Try O3.
DRAWING LEWIS
STRUCTURES
Multiple Bonds: O2 and N2
Multiple Central Atoms: C2H6, N2H4,
C3H8, C6H6, CH3NH2, CH3COOH
Compounds Containing
Polyatomic Ions
A polyatomic ion is a stable group of
atoms that has either a positive or
negative charge and behaves as a
single unit in many chemical
reactions.
Practice: NH4+, SO32-, NO2-, NO3-, I3-,
A scale of relative
electronegativities
wasacross
developed
by
Electronegativity
generally
increases left to right
a period.
Linus Pauling. Electronegativity decreases down a group
for representative elements.
Metals are low in EN and nonmetals are high.
ELECTRONEGATIVITY
Electronegativity: The relative attraction
that an atom has for a pair of shared
electrons in a covalent bond.
If the two atoms that constitute a covalent
bond are identical then there is equal
sharing of electrons.
This is called nonpolar covalent bonding.
Ionic bonding and nonpolar covalent bonding
represent two extremes.
ELECTRONEGATIVITY
If the two atoms that constitute a covalent
bond are not identical then there is
unequal sharing of electrons.
This is called polar covalent bonding.
One atom assumes a partial positive charge
and the other atom assumes a partial
negative charge.
This charge difference is a result of the unequal
attractions the atoms have for their shared
electron pair.
ELECTRONEGATIVITY
H-H are the same atom, and have the same
“greediness,” so the two atoms are forced
to share equally.
F-F same - forced to share equally.
If two atoms have diff EN, the one with
higher EN will “take” the e-s in the pair
more often than the other atom.
H-F are not the same atoms, and are not
equal in greediness, F is far greedier,
takes the e-s more than half the time.
Partial positive
Partial negative
charge on hydrogen.
charge on fluorine.
Polar Covalent Bonding in HF
+
:
:
H F
-
Fluorine has a greater attraction
for
the
The
shared
electron
Shared
electron
pair.pair is
shared electron pair than hydrogen.
closer to fluorine than to
hydrogen.
Types of Covalent
Bonding
The polarity of a bond is determined by the
difference in electronegativity values of the
atoms forming the bond.
If the electronegativity difference between two
bonded atoms is greater than 1.9 to 2.0, the bond
will be more ionic than covalent.
If the electronegativity difference is greater than
2, the bond is strongly ionic.
If the electronegativity difference is less than 1.9
but greater than 0.5, the bond is polar covalent.
If the electronegativity differences is 0.5 or less,
the bond in nonpolar covalent.
Types of Covalent
Bonding
Estimate whether a bond is polar cov, mostly pure or
nonpolar cov or ionic by finding the absolute value of
the difference in EN between the two atoms in the
bond.
__________________________________________
| | | | | | | | | | | | | |
0 .2 .4 .6 .8 1 1.2 1.4 1.6 1.8 2.0 2.2 2.4 2.6
nonpolar | polar cov
| mostly ionic
covalent
Practice: H-Cl, C-Cl, C-O and C=O, N-Cl, Ca-N
H-Cl C-Cl C-O C=O N-Cl Ca-N
0.9
0.5
1.0
1.0
0.0
2.0
pol cov “
“
“
nonpol ionic
Molecular Geometry
(bent)
= bent
One more geometry is trigonal pyramidal.
Some Geometric Figures
Linear
2 atoms on opposite sides of central
atom
180° bond angles
180°
Trigonal Planar
3 atoms form a triangle around the
central atom
Planar
120° bond angles
Tetrahedral
4 surrounding atoms form a
tetrahedron around the
central atom
109.5° bond angles
120°
109.5°
Some Geometric Figures
Trigonal Pyramidal
3 atoms form a triangular
pyramid beneath the central
atom
Not planar
~109° bond angles
Derivative of tetrahedral
geometry
The Valence Shell
Electron Pair Repulsion
(VSEPR) Model
The VSEPR model is based on the idea that electron
pairs will repel each other electrically and will
seek to minimize this repulsion.
To accomplish this minimization, the electron pairs
will be arranged as far apart as possible around a
central atom.
The 3-dimensional arrangement of the atoms within
a molecule determines molecular interactions
(physical properties and chemical reactions).
BeCl2 is a molecule with only
two pairs of electrons around
beryllium, its central atom.
Its electrons are arranged 180o
apart for maximum separation.
LINEAR EP
ARRANGEMENT
& MOLECULAR
GEOMETRY
BF3 is a molecule with three pairs of
electrons around boron, its central
atom.
o apart for
Its
electrons
are
arranged
120
This arrangement of atoms is called
maximum
separation.
trigonal planar.
CH4 is a molecule with four pairs of
electrons around carbon, its central
atom.
An
obvious since
choicethe
for molecule
its atomic is
arrangement
However,
3o angle between its atoms with all of
is
a
90
dimensional the molecular structure is
its atoms in a single plane.
o
tetrahedral with a bond angle of 109.5 .
Ball and stick models of methane, CH4, and carbon
tetrachloride, CCl4.
Tetrahedral Shapes H
Tetrahedral
4 areas of electrons around
the central atom
109.5° bond angles
All Bonding = tetrahedral
3 Bonding + 1 Lone Pair =
trigonal pyramid
2 Bonding + 2 Lone Pair = bent

H — C — H

H
Tetrahedral Derivatives

H — N — H

H

H — O — H

Ammonia, NH3, has four electron
pairs around nitrogen.
The arrangement
of electron pairs
around nitrogen
is tetrahedral.
NH3 has one
lone pair of
electrons.
The NH3 molecule is
trigonal pyramidal.
Water has four electron pairs
around oxygen.
The arrangement
of electron pairs
around oxygen is
tetrahedral.
H2O has two
lone pairs of
The H2O molecule
electrons.
is bent.
The VSEPR Model
Summary: electron pair arrangement depends upon # of bonded
atoms and lone pairs around the central atom. Lone pairs exert
more repulsion that bond pairs. Count # of bonded atoms (B) and
# of LPs (E).
EP Arrangements: 2 = Linear; 3 = Trigonal Planar; 4 = Tetrahedral
Within these electron pair arrangements, the molecular geometry is
based only on “seeing” the atoms. I call the molecular geometry
the family members and the electron pair arrangement is the
electron pair family the members are in.
Electron Pair Families and their Molecular Geometry members:
Linear: AB2, linear only
Trigonal Planar: AB3 trigonal planar; AB2E bent
Tetrahedral: AB4 tetrahedral; AB3E trigonal pyramidal, and AB2E2
bent
VSEPR Practice
Use practice sheets in packet to fill in
Electron Pair Arrangement and
Molecular Geometry and Bond Angles.
Dipole Moments
A dipole is a molecule with positively and negatively
charged ends
Polar covalent bonds or molecules have one end slightly
positive, +; and the other slightly negative, (not “full” charges, come from nonsymmetrical electron
distribution)
Dipole Moment is a measure of the size of the polarity.
(We are NOT going to worry about the Debye unit or
actual numbers for dipole moment, just whether a
molecule has a dipole or not.)
Polarity of Molecules
For a molecule to be polar it must
- have polar bonds
electronegativity difference - theory
bond dipole moments – measured values
- have an unsymmetrical shape
vector addition
Polarity affects the intermolecular
forces of attraction
DIPOLE
NO DIPOLE

:OCO:
O
H

H
Polar covalent bonds
and unsymmetrical
shape cause molecule
to be polar
Polar covalent bonds,
but nonpolar molecule,
because vectors cancel
Adding Dipole Moments
Table
10.3
64