Attention BMC Students,
There will be no chemistry seminars on October 23rd (Thursday) and
October 24th (Friday).
For those who miss the seminar, an optional make-up seminar
will be given on
Oct. 27, Monday, 12-14 (Dept. Obstetrics and Gynecology,
Lecture Hall))
Chemical bonds 2
2014
Which of the following drawings is more likely to represent
an ionic compound, and which a covalent compound?
ionic
covalent
Each of the pictures (a)–(d) represents one of the following
substances at 25°C: sodium, chlorine, iodine, sodium chloride.
Which picture corresponds to which substance?
iodine
sodium
NaCl
chlorine
Which of the following alkaline earth oxides has the
larger lattice energy? Explain.
A
B
The magnitude of the lattice energy depends directly on the charge on the ions and
inversely on the distance between ions (that is, on the radii of the ions).
In this instance, all the ions in both drowings are doubly charged, either M2+ or O2-, so
only the size of the ions is important.
Q1Q2
Ek
d
B has the larger lattice energy as the ions in B are smaller.
.
Blue:
AB2
(SrF2)
Green:
AB4
(CBr4)
Red:
AB or AB2
(PbS or PbS2)
VALENCE ELECTRONS
Electrons in the outermost
shell of the atom
BONDING ELECTRONS
Number of valence
electrons shared with other
atoms
NON-BONDING ELECTRONS
Valence electrons, not
shared with other
atoms
Drawing Lewis Structures
• Sum the valence electrons from all atoms. Add one for each
negative charge and subtract one for each positive charge.
• Draw a skeleton structure with atoms attached by single bonds.
The central atom is the least electronegative one (except H).
• Complete the octets of atoms bound to the central atom.
• Place extra electrons on the central atom.
• If the central atom doesn’t have an octet, try forming multiple
bonds.
Give the electron-dot structures (Lewis structures) for
the following species:
(a) Ammonium ion
(c) IF3
(b) HCN
:
(d) Phosphorus
pentachloride
(e)
Sulfur
tetrafluoride
Resonance Structures
Resonance structures are attempts to represent a real
structure that is a mix between several extreme
possibilities.
Some molecules are not well described by Lewis Structures.
Typically, structures with multiple bonds can have similar
structures with the multiple bonds between different
pairs of atoms
Draw as many electron-dot resonance structures as possible for (a)
O3, (b) sulfur dioxide, and (c) the nitrite ion giving the octet of all
atoms.
• What if different sets of atomic linkages can be
used to construct Lewis Dot Structures?
Cl O Cl
Cl O
Cl
• Which is correct?
Cl Cl O
Cl Cl O
Formal charge (FC)
• Formal charge is the difference between the number of valence
electrons in the free atom and the number assigned to that atom
in the Lewis structure.
• Mathematically: FC = X – (Y + Z/2)
where X = the number of valence electrons in the free
atom (Group Number)
Y = the number of unshared electrons
Z = the number of shared electrons
• The most stable structure has:
the lowest formal charge on each atom,
the most negative formal charge on the most electronegative
atoms.
Formal Charge
Cl O Cl
Cl Cl O
#of e- (Y + Z/2)
7
6
7
7
6
7
#of e- in free atom (X)
7
6
7
7
7
6
F. C.
0
0
0
0 +1 -1
The most stable structure has:
the lowest formal charge on each atom,
the most negative formal charge on the most electronegative atoms.
Formal charge
O: 6 - 4½ - 4 = 0
O: 6 – 2 ½ - 6 = -1
N: 5 - 6 ½ -2 = 0
N: 5 – 8  ½ - 0 = + 1
More favorable
VALENCE SHELL ELECTRON PAIR
REPULSION: VSEPR
Pairs of electrons repel each other
Shape of molecule is governed by pairs of valence
electrons being as far apart as possible.
The VSEPR Model
• To apply VSEPR theory:
1. Draw the Lewis structure.
2. Determine the electron domain geometry.
3. Determine the molecular geometry.
Electron domain geometry: arrangement of electron
pairs around the central atom.
Molecular geometry: arrangement of atoms around the
central atom.
This does not include lone electron pairs. The molecular
shape differs from the electron shape if lone pairs are
present
Five basic shapes
Two charge clouds,
no lone pair
Three charge clouds,
no lone pair
Four charge clouds,
no lone pair
Five charge
clouds, no
lone pair
Six charge clouds, no
lone pair
Number of
Bonds
Number of
Lone Pairs
Number of
Charge Clouds
Geometry
Example
Number of
Bonds
Number of
Lone Pairs
Number of
Charge Clouds
Geometry
Example
Number of
Bonds
Number of
Lone Pairs
Number of
Charge Clouds
Geometry
Example
Number of
Bonds
Number of
Lone Pairs
Number of
Charge Clouds
Geometry
Example
Number of
Bonds
Number of
Lone Pairs
Number of
Charge Clouds
Geometry
Example
(a) Trigonal planar
(b) Trigonal bypiramidal
(d) Octahedral
(c) Linear
Draw the Lewis structure and predict the shape of the following
molecules:
(a) CO2
(b) H2CO
linear
trigonal planar
: :
(d) H2S
(c) CBr4
H-S-H
bent
tetrahedral
(f) PCl5
(e) PCl3
trigonal pyramidal
trigonal bypyramidal
What shape and bond angle do you expect for each of the following molecules?
(a) H2O
(b) BF3
(c) CO2
(d) SF6
(e) CCl4
(a) O: 4 Charge Clouds, bent, 104.5o
(b) B: 3 CC, trigonal planar, 120o
(c) C: 2 Charge Clouds,linear, 180o
(d) S: 6 CC, octahedral, 90o
(e) C: 4 CC, tetrahedral, 109o
Molecular Structure and
Covalent Bonding Theories
 bonds
 bond
Hybridization
• The valence shell orbitals (atomic orbitals)
commonly combine to change their character in
order to obtain a lower energy ‘mixed’ orbital set
for bonding in a particular geometry
• Hybridization – process by which atomic orbitals
combine to form a set of ‘mixed’ orbitals of lower
energy when bonding covalently
– The ‘mixed’ orbitals are called hybrid orbitals
Valence Bond(VB) Theory
sp3 Hybridization
sp2 Hybridization
sp Hybridization
Identify each of the following sets of hybrid orbitals:
What is the hybridization of the central atom in:
(a) AlCl3
(b) SiH4
(c) PCl5
sp2
sp3
sp3d
(d) [Cr(H2O)6]3+
sp
(e) H2Be
(f) ClO 4
sp3
sp3d2
Molecular Geometry and Polarity
• The polarity can be determine once the geometry is
known
• A polar bond is created if the atoms sharing the
electron pair have different electronegativities
– HCl and the associated dipole moment. This molecule is
polar. For diatomics, determination of polarity is easy.
What if the molecule has two or more atoms? All the
dipole have to be summed. If the sum equals zero, the
molecule has no dipole.
Molecular Geometry and
Polarity
Which of the following molecules are polar?
(a) H2O
(b) SO2
(c) CO2
(d) CH4
Intermolecular Forces
Which substance in each of the following pairs would you expect to
have the lower boiling point. Explain.
a) C6H6 or C6Cl6
b) H2C=O or CH3OH
c) He or Kr
d) H2Se or H2O
a) C6H6
smaller mass
weaker London forces
c) He
smaller mass
b) H2C=O
no H-bonding
d) H2Se
no H-bonding