CHEMICAL BONDING
Matter is composed of either
(1) Metals
-
Atoms
-
Metallic Bonding
(2) Nonmetals
-
Molecules
-
Covalent Bonding
(3) Metals and Nonmetals
-
Ions
-
Ionic Bonding
Chemical bonding involves the valence electrons of atoms
2A-1 (of 15)
(2) MATTER COMPOSED OF METALS
1904 ARNOLD SOMMERFELD
Proposed that metal atoms release their valence electrons,
and share them between large numbers of metal atoms
2A-2 (of 15)
METALLIC BOND – The electrostatic attraction of the shared valence
electrons to the nuclei of the many bonding metal atoms
Metallic bonding forms crystalline networks
containing billions of metal ATOMS that are
strongly attracted together
2A-3 (of 15)
(3) MATTER COMPOSED OF NONMETALS
1916 GILBERT NEWTON LEWIS
Proposed that nonmetal atoms share valence electrons to
achieve the electron configurations of Noble Gases
Diatomic chlorine
:
:
: Cl
Cl :
:
:
:
:
:
:
: Cl – Cl :
2A-4 (of 15)
LEWIS STRUCTURE – A representation of chemical bonding using
electron dot notation
:
:
BONDING PAIRS: in red
LONE PAIRS:
2A-5 (of 15)
in green
:
:
: Cl Cl :
COVALENT BOND – The electrostatic attraction of the shared electrons to
the nuclei of the bonding nonmetal atoms
Covalent bonding forms individual units called
MOLECULES that are weakly attracted to each
other
2A-6 (of 15)
LEWIS STRUCTURES
To draw a proper Lewis Structure for a covalently bonded species:
1 – Add up the valence e-s for all of the atoms in the molecule or ion
2 – Draw a skeletal structure by using pairs of electrons to make bonds
3 – Complete octets (or duets for H) for all atoms, outer atoms first,
using the remaining valence e-s
4 – If octets are not produced, make the atoms that have octets share
more e- pairs with atoms that do not have octets
2A-7 (of 15)
Oxygen difluoride, OF2
6 + 7 + 7 = 20 valence e-s
F
O
2A-8 (of 15)
F
Nitrogen tribromide, NBr3
5 + 7 + 7 + 7 = 26 valence e-s
Br
N
Br
2A-9 (of 15)
Br
(1) MATTER COMPOSED OF METALS AND NONMETALS
1904 RICHARD ABEGG
Proposed that atoms gain or lose valence electrons to
achieve the electron configurations of Noble Gases
2A-10 (of 15)
Metal atoms easily lose valence e-s, forming positive ions
Al
1s22s22p63s23p1
Al3+
1s22s22p6
Fe
[Ar]4s23d6
Fe2+
[Ar]3d6
Fe3+
[Ar]3d5
Nonmetal atoms gain e-s to their valence shells, forming negative ions
O
1s22s22p4
O2-
1s22s22p6
Once these ions are formed, they are stable (or unreactive)
2A-11 (of 15)
IONIC BOND – The electrostatic attraction between positive metal ions and
negative nonmetal ions
Ionic bonding forms crystalline networks
containing billions of positive and negative
IONS that are strongly attracted together
2A-12 (of 15)
Sodium chloride
Na
.
..
Cl :
..
..
+
Na
:
Cl
..
:
-
A sodium chloride crystal is a symmetrical array of
sodium and chloride ions in a 1:1 ratio
EMPIRICAL FORMULA – The simplest whole number
ratio of ions of different elements in a compound
Empirical Formula: NaCl
2A-13 (of 15)
Calcium fluoride
Ca
.
..
F:
..
..
2+
Ca
.
:
F:
..
..
F:
..
-
..
:
F:
-
..
Empirical Formula: CaF2
2A-14 (of 15)
REPRESENTING IONIC BONDING WITH ELECTRON DOT NOTATION
K3N
.
K.
K.
K.
+
+
+
.
N:
.
..
K
K
2A-15 (of 15)
K
:
N:
..
3-
Fluorine, F2
7 + 7 = 14 valence e-s
F
F
SINGLE BOND – One shared pair of e-s between two atoms
2B-1 (of 15)
Oxygen, O2
6 + 6 = 12 valence e-s
O
O
DOUBLE BOND – Two shared pairs of e-s between two atoms
2B-2 (of 15)
Nitrogen, N2
5 + 5 = 10 valence e-s
N
N
TRIPLE BOND – Three shared pairs of e-s between two atoms
2B-3 (of 15)
BOND ORDER – The number of shared pairs of electrons
BOND ENERGY – The energy needed to break a bond
BOND LENGTH – The distance between the nuclei of the 2 bonding atoms
Bond Order
Bond Energy (kJ/mol)
Bond Length (nm)
2B-4 (of 15)
F2
O2
N2
1
154
2
495
0.121
3
941
0.142
0.110
H
H
Cl
Cl
I
I
S
S
P
P
Longest Bond Length?
I2
biggest atoms
Shortest Bond Length?
H2
smallest atoms
Highest Bond Energy?
P2
most bonding electrons
Lowest Bond Energy?
I2
least bonding electrons, and
they are most shielded from the
nuclei
2B-5 (of 15)
Formaldehyde, CH2O
4 + 1 + 1 + 6 = 12 valence e-s
H
C
H
2B-6 (of 15)
O
Sulfate, SO426 + 4(6) + 2 = 32 valence e-s
2-
O
O
S
O
2B-7 (of 15)
O
NO35 + 3(6) + 1 = 24 valence e-s
-
O
N
O
O
-
↔
O
N
O
O
-
↔
O
N
O
O
RESONANCE – When more than one Lewis structure can be drawn for a
molecule or ion
RESONANCE STRUCTURES – The Lewis structures that can be drawn for
the molecule or ion
The bonding in the real nitrate ion is an average of its resonance structures
The average N-O bond order is (1+1+2) / 3 = 11/3
2B-8 (of 15)
1932 LINUS PAULING
Described how atomic orbitals are involved in covalent
bonding
VALENCE BOND THEORY – Two atoms share electrons by overlapping a
valence atomic orbital from each atom, creating a region of space
between the nuclei where the electrons reside
2B-9 (of 15)
H atom
H2 molecule
1s atomic orbital
with 1 valence e-
H atom
1s atomic orbital
with 1 valence e-
2 valence e-s in a
MOLECULAR ORBITAL
The attraction of the e-s in the molecular orbital to the 2 nuclei bonds the
atoms together
2B-10 (of 15)
ELECTRONEGATIVITY – A property developed by Pauling, measuring the
attraction of an atom for shared electrons
Metals – Low EN’s (the most active metals having the lowest EN’s)
Nonmetals – High EN’s (the most active nonmetals have the highest EN’s)
Atom with the highest EN?
Atom with the lowest EN?
F
Cs
(4.0)
2B-11 (of 15)
(0.7)
EN differences between atoms indicates their type of bonding
EN Difference
Bonding
0
Small (0.1 – 1.6)
Large (1.7 – 3.3)
Nonpolar Covalent
Polar Covalent
Ionic
2B-12 (of 15)
2 atoms with the same EN’s have an EN difference of 0
N–N
(EN of N = 3.0)
NONPOLAR COVALENT BOND – A bond between 2 atoms in which the
electrons are shared evenly
2B-13 (of 15)
2 atoms with close EN’s have an EN difference that is small
H – Br
(EN of H = 2.1, EN of Br = 2.8)
Dipole Moment Arrow
POLAR COVALENT BOND – A bond between 2 atoms in which the
electrons are shared unevenly
2B-14 (of 15)
2 atoms with extreme EN’s have an EN difference that is large
Na – Cl
(EN of Na = 0.8, EN of Cl = 3.0)
IONIC BOND – A bond between 2 atoms in which the electrons are
transferred, creating ions
2B-15 (of 15)
FORMAL CHARGE
While atoms that covalently bond are not charged, they can be given
charges based upon where the bonding electrons are assigned
FORMAL CHARGE – The charge given to an atom assuming one electron in
each bond is assigned to that atom
2C-1 (of 11)
F
F.
0
S
.
F
S.
0
.
F
0
S naturally has 6 valence e-s , and now 6
 0
F naturally has 7 valence e-s , and now 7
 0
Quick way to determine formal charge:
(natural number of valence e-s – 1 e- per bond – each lone pair e-)
S:
F:
2C-2 (of 11)
6–2–4= 0
7–1–6= 0
C
O
C
O
-1
+1
C: 4 – 3 – 2 = -1
O: 6 – 3 – 2 = +1
Formal charges are used to determine the validity of a Lewis structure the most accurate Lewis structures are those with atoms that have formal
charges as close to 0 as possible
2C-3 (of 11)
thiocyanate, SCN6 + 4 + 5 + 1 = 16 valence e-s
S
S:
C:
N:
C
N
-
6–2–4= 0
4–4–0= 0
5 – 2 – 4 = -1
2C-4 (of 11)
↔
S
C
N
-
S: 6 – 3 – 2 = +1
C: 4 – 4 – 0 = 0
N: 5 – 1 – 6 = -2
↔
S
C
N
-
S: 6 – 1 – 6 = -1
C: 4 – 4 – 0 = 0
N: 5 – 3 – 2 = 0
thiocyanate, SCN6 + 4 + 5 + 1 = 16 valence e-s
S
S:
C:
N:
C
N
-
6–2–4= 0
4–4–0= 0
5 – 2 – 4 = -1
↔
S
C
N
-
S: 6 – 3 – 2 = +1
C: 4 – 4 – 0 = 0
N: 5 – 1 – 6 = -2
↔
S
C
N
S: 6 – 1 – 6 = -1
C: 4 – 4 – 0 = 0
N: 5 – 3 – 2 = 0
The best Lewis structures have
(1) formal charges for the most atoms as close to 0 as possible
(2) negative formal charges go on the atom with the greatest EN
2C-5 (of 11)
-
COVALENT COMPOUNDS THAT DO NOT OBEY THE OCTET RULE
(1) Molecules with hypovalent central atoms
(atoms with less than 4 valence electrons)
Covalent compounds with B and Be
BeH2
2 + 1 + 1 = 4 valence e-s
H
Be
2C-6 (of 11)
H
BF3
3 + 7 + 7 + 7 = 24 valence e-s
F
B
F
F
F
B
F
2C-7 (of 11)
F
NO!
F is too electronegative to share more
than 1 pair of e-s
(2) Molecules with hypervalent central atoms
(atoms that have empty d orbitals in their outer shell)
Nonmetal atoms in the 3rd, 4th, 5th, or 6th Periods
PF5
5 + 5(7) = 40 valence e-s
3s
F
F
F
P
F
2C-8 (of 11)
F
↑↓
___
3p
↑ ↑ ↑
___ ___ ___
3d
___ ___ ___ ___ ___
P can make 5 bonds using empty d
orbitals in its outer shell
ClF3
7 + 3(7) = 28 valence e-s
F
F
Cl
F
F
F
Cl
2C-9 (of 11)
F
Only 26 valence e-s
Sulfate, SO426 + 4(6) + 2 = 32 valence e-s
2-
O
O
S
S:
O
6 – 4 – 0 = +2
O: 6 – 1 – 6 = -1
O
Experimental data shows the S-O bonds are stronger than single bonds
Reducing the formal charge on atoms that can exceed the octet rule can
produce a more accurate Lewis structure
S must make 2 double bonds to reduce its formal charge to 0
2C-10 (of 11)
Sulfate, SO426 + 4(6) + 2 = 32 valence e-s
2-
O
O
S
O
O
+ 5 other
resonance structures
S: 6 – 6 – 0 = 0
O: 6 – 1 – 6 = -1
O: 6 – 2 – 4 = 0
Experimental data shows the S-O bonds are stronger than single bonds
Reducing the formal charge on atoms that can exceed the octet rule can
produce a more accurate Lewis structure
S must make 2 double bonds to reduce its formal charge to 0
2C-11 (of 11)
MOLECULAR SHAPE
VSEPR THEORY (Valence Shell Electron Pair Repulsion) – All atoms and
lone pairs attached to a central atom will spread out as far as possible to
minimize repulsion
A Lewis structure must be drawn to use the VSEPR Theory
2D-1 (of 15)
CO2
4 + 6 + 6 = 16 valence e-s
O
C
O
STERIC NUMBER (SN) – The sum of the bonded
atoms and lone pairs on a central atom
The steric number of carbon is 2 (SN = 2):
2 bonded atoms and 0 lone pairs
Linear
Bond angle is 180°
2D-2 (of 15)
O
C
O
BH3
3 + 1 + 1 + 1 = 6 valence e-s
H
B
H
H
SN = 3
3 bonded atoms and 0 lone pairs
H
Trigonal Planar
Bond angle is 120°
2D-3 (of 15)
B
H
H
SO2
6 + 6 + 6 = 18 valence e-s
O
S
O
SN = 3
2 bonded atoms and 1 lone pairs
Bent
Bond angle is 120°
2D-4 (of 15)
O
S
O
CH4
H
H
C
H
H
SN = 4
4 bonded atoms and no lone pairs
H
Tetrahedral
Bond angle is 109.5°
2D-5 (of 15)
C
H
H
H
NH3
H
N
H
H
SN = 4
3 bonded atoms and 1 lone pairs
Trigonal Pyramidal
Bond angle is 108°
2D-6 (of 15)
N
H
H
H
H2O
..
H–O:
H
SN = 4
2 bonded atoms and 2 lone pairs
Bent
Bond angle is 105°
2D-7 (of 15)
O
H
H
PF5
F
F
P
F
F
SN = 5
5 bonded atoms and no lone pairs
F
Trigonal Bipyramidal
3 Equatorial F’s in a
plane, 120° apart
2 Axial F’s 180° apart,
90° from the plane
2D-8 (of 15)
F
F
F
P
F
F
SF4
:F:
F
S
F
:F:
SN = 5
4 bonded atoms and 1 lone pair
← 2 close 90º interactions
3 close 90º interactions →
← most stable configuration
e- pair in
equatorial position
2D-9 (of 15)
e- pair in
axial position
SF4
:F:
F
S
:F:
F
SN = 5
4 bonded atoms and 1 lone pair
See-Saw
e- pairs always go in
equatorial positions to
minimize repulsion
2D-10 (of 15)
F
F
F
P
F
ClF3
F
Cl
F
F
SN = 5
3 bonded atoms and 2 lone pairs
F
T-Shape
Cl
F
2D-11 (of 15)
F
XeF2
F
Xe
F
SN = 5
2 bonded atoms and 3 lone pairs
F
Linear
Xe
F
2D-12 (of 15)
SF6
F
F
F
S
F
F
SN = 6
6 bonded atoms and no lone pairs
F
F
Octahedral
90º and 180º
F
F
S
F
2D-13 (of 15)
F
F
IF5
F
F
F
I
F
F
SN = 6
5 bonded atoms and 1 lone pair
F
Square Pyramidal
2D-14 (of 15)
F
F
I
F
F
XeF4
F
F
Xe
F
SN = 6
4 bonded atoms and 2 lone pairs
F
Square Planar
2D-15 (of 15)
F
F
Xe
F
F
MOLECULAR POLARITY
A BOND is polar if it has a positive end and a negative end
A MOLECULE is polar if it has a positive end and a negative end
To determine if a molecule is polar or nonpolar:
1) Draw the correct Lewis structure
2) Draw its correct shape
3) Use EN’s to determine if the BONDS in the molecule are polar or
nonpolar
4) For the polar bonds, label the positive and negative ends with δ+ and δ5) If a line can be drawn separating all δ+’s from all δ-’s, the molecule is
polar, if not its nonpolar
2E-1 (of 13)
..
H–O:
H
δ+ H
δO δ-
H
δ+
2E-2 (of 13)
EN’s: O = 3.5, H = 2.1
3.5 – 2.1 = 1.4  the O-H BONDS are polar
All of the δ+’s can be separated from all of the δ-’s, 
the H2O MOLECULE is polar
H
N
H
H
EN’s: N = 3.0, H = 2.1
δδ- N δδ+ H H
δ+
2E-3 (of 13)
3.0 – 2.1 = 0.9  the N-H BONDS are polar
H δ+
All of the δ+’s can be separated from all of the δ-’s, 
the NH3 MOLECULE is polar
F
F
C
F
F
δF
δ+
δ+ C δ+
δ- F δ+ F δF
δ-
2E-4 (of 13)
EN’s: C = 2.5, F = 4.0
4.0 – 2.5 = 1.5  the C-F BONDS are polar
All of the δ+’s cannot be separated from all of the δ-’s,
 the CF4 MOLECULE is nonpolar
A more exact way to determine if a molecule is polar or nonpolar:
1) Draw the correct Lewis structure
2) Draw its correct shape
3) Use EN’s to determine if the BONDS in the molecule are polar or
nonpolar
4) For the polar bonds, draw a DIPOLE MOMENT ARROW pointing toward
the negative end of the bond
5) If the dipole moments are symmetrical the molecule is NONPOLAR
2E-5 (of 13)
Dipole moments of equal magnitude are symmetrical if:
1) there are 2 dipole moments that are linear
Y
2E-6 (of 13)
X
Y
Dipole moments of equal magnitude are symmetrical if:
2) there are 3 dipole moments that are trigonal planar
Y
X
Y
2E-7 (of 13)
Y
Dipole moments of equal magnitude are symmetrical if:
3) there are 4 dipole moments that are tetrahedral
Y
Y
Y
2E-8 (of 13)
X
Y
O
O
C
C
2E-9 (of 13)
O
O
Symmetrical dipole moments
 the CO2 MOLECULE is nonpolar
..
H–O:
H
O
H
H
2E-10 (of 13)
Assymmetrical dipole moments
 the H2O molecule is POLAR
H
N
H
H
N
H
H
2E-11 (of 13)
H
Assymmetrical dipole moments
 the NH3 molecule is POLAR
F
F
C
F
F
F
C
F
F
2E-12 (of 13)
Symmetrical dipole moments
 the CF4 molecule is NONPOLAR
F
Cl
F
C
F
F
Cl
C
F
F
2E-13 (of 13)
F
Assymmetrical dipole moments
because the C-Cl dipole moment is
smaller than the C-F dipole moments
 the CClF3 molecule is POLAR