Chemistry
Chemical Bonding
The Development of Atomic Models
1.
Dalton – solid, indivisible mass
2.
Thomson – “plum-pudding” model
-
Negatively charged e- (raisins) stuck in
positively charged proton dough
No neutrons
3.
Rutherford – electrons surrounding a
dense nucleus
4.
Bohr model – elctrons arranged in
symmetrical orbits around the nuclues
-
“planetary model”
Electrons in a given path have a fixed
energy level
5. Quantum mechanical model – modern
mathematical description of the atom
Sodium atom:
Energy level – region around the nucleus
where the electron is likely to be
moving.
An electron can jump from one level to
another by absorbing energy.
Quantum – the amount of energy
required to move an electron from its
present energy level to the next higher
one
“quantum leap”
Quantum mechanical model – uses
mathematical equations to describe the
location and energy of electrons in an
atom
 Developed by Erwin Schrodinger
 Electrons are not in definite paths
 Their location is described in terms of
probability of being in a certain region
 Electron cloud (ceiling fan)
 Conventionally, the border is drawn at
90% probability
Atomic orbital – region in space that an
electron is likely to be in
Electrons can be described by a
series of 4 quantum numbers.
1. Principle quantum number (n)
Describes the energy level
 Values of 1, 2, 3, 4, etc.

2. Azimuthal quantum number ( l )
Describes the shape of atomic orbitals
 Sublevels
 Values of 0 to n-1
 0 = s, 1 = p, 2 = d, 3 = f

s = spherical, p = peanut shape,
d&f = more complex shapes
d = “daisy”
f = “fancy”
So if n = 1, then l can be 0 (s) = 1 sublevel
n = 4, then l can be 0 (s), 1 (p), 2(d) , 3(f), = 4 sublevels
s orbitals – spherical
p orbitals – dumbbell-shaped
d orbitals – daisy-shaped
f orbitals – fancy shapes
3. Magnetic quantum number (ml)


Orientation of the orbital in space
Values of –l to +l
So s has 1 orbital
p has 3
d has 5
f has 7
4. Spin quantum number (ms)
Values of +½ and -½
 Each orbital can hold 2 electrons with
opposite spins
 Since spinning charged objects create a
magnetic field, the electrons must spin
opposite directions to minimize
repulsion

Ex. How many orbitals are in the
following?
A. 3p
B. 2s
C. 4f
D. 4p
E. 3d
F. 3rd energy level
How many electrons can be in each of the
above?
Are these possible?
n
l
ml
ms
1
0
0
+1/2
5
2
-5
-1/2
2
1
-1
+1/2
Electron Configuration – ways in which
electrons are arranged around nuclei of
atoms.
Rules that govern filling of
atomic orbitals:
1. Aufbau principle – electrons enter
orbitals of lowest energy first.
2. Pauli exclusion principle – An
atomic orbital can describe at
most two electrons. They must
have opposite spins.
3. Hund’s rule – When electrons
occupy orbital of equal energy,
one electron enters each orbital
until all orbitals contain one
electron with parallel spins.
1s
2s 3s 4s 5s 6s 7s
2p 3p 4p 5p 6p
3d 4d 5d
4f
Na = 11
1s22s22p63s1
Cd =48
1s22s22p63s23p64s23d104p65s24d10
Practice: Write the electron configuration
for the following elements:
Li
O
Sc
More practice: Identify each of the following
atoms on the basis of its electron configurations.
a) 1s22s22p6
neon
b) 1s22s22p63s1
sodium
c) [Kr] 5s24d2
zirconium
d) [Xe] 6s24f6
samarium
Ground state – lowest energy level for
an electron. (Normal, nonexcited state)
Exceptional Electron
Configurations:
Cr: expected: 1s22s22p63s23p64s23d4
actual: 1s22s22p63s23p64s13d5
Cu: expected: 1s22s22p63s23p64s23d9
actual: 1s22s22p63s23p64s13d10
Half-filled energy levels are more stable than
other partially filled energy levels. There
are other exceptions.
Light and Atomic Spectra
Electromagnetic radiation – a series of
energy waves that includes radio waves,
microwaves, visible light, infrared, and
ultraviolet light, X-rays, and gamma rays.
Parts of a wave:
Wavelength, l
crest
trough
Amplitude – height of the wave from the
origin to the crest.
Wavelength - l - distance between the crests
Frequency – f (or n) – the number of wave
cycles to pass a given point per unit of time.
The units of frequency are 1/s, s-1, or Hertz
(Hz)
c= lf
where c = speed of light = 3.00x108 m/s
or 3.00x1010 cm/s
As l increases, f decreases.
As wavelength increases, frequency decreases.
As wavelength decreases, frequency increases.
Ex. A certain wavelength of yellow light has a
frequency of 2.73x1016 s-1. Calculate its
wavelength. Convert to nm.
C = lf
l = c/f
l = (3.00x108 m/s )/(2.73x1016 s-1)
l = 1.10x10-8 m
Spectrum – series of colors produced when
sunlight is separated by being passed
through a prism.
ROY G. BIV
Red: longest wavelength, lowest frequency
Violet: shortest wavelength, highest
frequency
Atomic emission spectrum – series of lines
of colored light produced by passing light
emitted by an “excited atom through a
prism.” This can be used to identify the
element. The atomic emission spectrum
of hydrogen shows three series of lines.
The lines in the UV region (Lyman series)
represent electrons falling to n=1, lines in
the visible region (Balmer series)
represent electrons falling to n=2 and
lines in the IR region (Paschen series)
represent electrons falling to n=3.
Max Planck found that the energy emitted
or absorbed by a body changes only in
small discrete units he called quanta. He
determined that the amount of radiant
energy, E, absorbed or emitted by a body
is proportional to the frequency of the
radiation.
E=hf
E = energy (J) f = frequency
h = Planck’s constant, 6.626x10-34 Js
Einstein studied the photoelectric effect
whereby light of sufficient frequency
shining on a metal causes current to flow.
The amplitude of the radiation was not
important, the frequency was. This told
him that light must be in particles, each
having a given energy. Einstein proposed
that electromagnetic radiation can be
viewed as a stream of particles called
photons:
E=hf
Photoelectric Effect
light
Electron
(photons)
metal
Energy of a photon:
E = hf
Example: Calculate the energy of an individual
photon of yellow light having a frequency of
2.73x1016 s-1.
E=hf
E = (6.626x10-34 Js)(2.73x1016 s-1)
E = 1.81x10-17 J
Einstein’s special theory of relativity:
E=mc2
Matter and energy are different forms of
the same entity.
Going Further:
Louis deBroglie suggested that very small
particles like electrons might also display
wave particles and he came up with:
deBroglie’s equation:
l=h
mv
m = mass in kg
v = velocity in m/s
h = Planck’s constant, 6.626x10-34 Js
DeBroglie’s equation is used to find the
wavelength of a particle. It was
determined that matter behaves as
through it were moving in a wave. This
is important in small object such as
electrons but is negligible in larger
objects such as baseballs. Heavy objects
have very short wavelengths.
Example: Calculate the wavelength of an
electron traveling at 1.24x107 m/s. The mass of
an electron is 9.11x10-28g.
l=h
mv
9.11x10-28 g
1 kg = 9.11x10-31 kg
1000g
l=
(6.626x10-34 Js)
=
(9.11x10-31 kg)(1.24x107 m/s)
l = 5.87x10-11 m
End Going Further.
In the photoelectric effect, electrons
(called photoelectrons) are ejected by
metals (esp. alkali) when light of
sufficient frequency shines on them. Red
light won’t work. Photoelectric cells
convert light energy into electrical
energy. They are used in automatically
opening doors and security systems.
Heisenberg’s Uncertainty Principle –
it is impossible to determine
accurately both the momentum
and the position of an electron
simultaneously.
We detect motion by
electromagnetic radiation. This
interaction disturbs electrons.
Einstein and Heisenberg!
Mendeleev- arranged elements in
order of increasing atomic mass
Moseley- arranged elements in
order of increasing atomic number
Periodic Law
- When the elements are arranged in
order of increasing atomic number,
there is a periodic pattern in their
physical and chemical properties.
Be able to locate noble gases,
representative elements, transition
metals, inner transition metals.
Noble Gases
have completely filled shells of electrons
 similar electronic structures

◦
◦
◦
◦
◦
He
Ne
Ar
Kr
etc.
1s2
1s2 2s2 2p6
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s2 4p6
Representative Elements
elements in A groups
on periodic chart
 representative because they best
represent what we
know about
elemental structure
& periodicity

d - Transition Elements
 elements in B groups on periodic chart
 metals
 have d electrons
 transition from metals
to nonmetals
f - Transition Elements
 inner transition metals
Electron configuration using the
periodic table:
Columns are called groups or
families. Rows are called periods or
series.
Shorthand electron configuration:
3s
11
3p
Na Ne 
12
Mg Ne 
13
Al
Ne
14
Si
You do it!

Configuration
Ne 3s
2
Ne 3s
2
1
Ne 3s 3p
1

3d
4s
19 K Ar 

20
Ca Ar 

Sc Ar  

Ti Ar   

V Ar    

Cr Ar      

21
22
23
24
4p
Configuration
Ar  4s1
Ar  4s 2
Ar  4s 2 3d1
Ar  4s 2 3d 2
Ar  4s 2 3d 3
Ar  4s1 3d 5
There is an extra measure of stability associated
with half - filled and completely filled orbitals.
3d
25 Mn Ar      
4s

Fe Ar      

Co Ar      

28
Ni Ar      

29
Cu You do it!
26
27
4p
Configurat ion
Ar  4s2 3d5
Ar  4s2 3d6
Ar  4s2 3d7
Ar  4s2 3d8
Periodic Trends in atomic size:
covalent atomic radius - half the
distance between the nuclei of atoms
in a homonuclear diatomic molecule
(like Cl2)
Trends:

Atomic size increases going down a group
because electrons are added to higher energy
levels that are farther from the nucleus. It
decreases going across a period because as
each proton is added to the nucleus an electron
is being added to the same energy level. This
“shell” of electrons is pulled closer in towards
the nucleus.
Size changes little in the
transition metals because the
electrons being added are core
electrons.
Zeff = effective nuclear charge
-actual pull of the nucleus on the
valence electrons.
Zeff = Zactual - effect of e- repulsions
Trends: 
Increases from H to He
Decreases from He to Li because 1s
electrons shield 2s electrons
Increases from Li to Be
Decreases from Be to B because 2s shield 2p
Increases from B to C to N
Decreases from N to O because of
repulsion due to doubly occupied
orbitals
Increases from O to F to Ne
Decreases from Ne to Na because
1s,2s & 2p shield 3s
Know exceptions to Zeff trends and
reasons for these.
Ionization energy (IE)
- energy required to remove the highest
energy electron from a gaseous atom
Li(g) + energy  Li+(g) + e- depends on Zeff and size
Trends: 
Ionization energy decreases down a a
group because the valence electrons are
farther from the nucleus and are thus held
less tightly. Ionization energy increases across
a period because atomic size decreases and
valence electrons are held more tightly. Zeff
increases.
1st ionization energy = energy required to remove
the first electron
Al + energy  Al+ + e2nd ionization energy = energy required to remove
the second electron
Al+ + energy  Al2+ + e3rd ionization energy = energy required to remove
the third electron
Al2+ + energy  Al3+ + e-
For Na, the 1st ionization energy is
fairly low but the 2nd would be high.
1stIE
2ndIE
3rdIE
4thIE
IA
Na
496
4562
6912
9540
IIA
Mg
738
1451
7733
10,550
IIIA
Al
578
1817
2745
11,580
IVA
Si
786
1577
3232
4356
Know exceptions to ionization
energy trends and reasons for them.
Ionic Size
Anions are larger than the atoms
from which they were formed. Know
why!!!
Cations are smaller than the atoms
from which they were formed. Know
why!!!!
Sizes of Ions
Related to
Positions of
the Elements
in the Periodic
Table
Isoelectronic ions- a group of ions
with the same number of electrons
The one with the highest atomic
number is the smallest in size (More
protons pulling on the same # of
electrons).
Na+, Mg2+, Ne, F-, O2-, N3- are isoelectronic.
They all have 10 electrons. Mg2+ is the
smallest because it has 12 protons pulling on
10 electrons. (The protons win the “tug of
war”) N3- is the largest because it has 7
protons and 10 electrons. (The electrons win
the “tug of war”)
Electron Affinity
-energy change that occurs when an
electron is added to a gaseous atom
-usually exothermic
Cl(g) + e-  Cl-(g)+ energy
Trends: 
exceptions)
(but many
Electron Affinity vs Atomic Number
300
Be
4
200
Mg
12
Ca
20
100
He
2
0
-100
-200
-300
-400
H
1
Li
3
B
5
C
6
Ne
10
N
7
O
8
F
9
Na
11
Al
13
Ar
18
P
Si 15
14
K
19
S
16
Cl
17
Electronegativity
-relative tendency of an atom to
attract shared electrons to itself
Trends: 
FONCl (Phone Call)
Elements with high electronegativity
(nonmetals ) tend to gain electrons
to form anions.
Elements with low electronegativities
(metals) often lose electrons to form
cations.
Valence electrons- electrons in the
highest occupied energy level of an
atom. Valence electrons are the only
electrons involved in the formation of
chemical bonds.
Electron dot structures for atoms:
-each dot represents a valence electron
p
p
X
s
p
Examples:
N
N
O
O
Xe
Xe
Al
Al
Na
Na
I
I
Si
Si
One of the major “driving forces” in
nature is the tendency to go to lower
energy. Atoms lose, gain or share
electrons to become lower in energy
and thus more stable.
Metals lose electrons easily to become positively charged
cations. They will usually lose their valence electrons to
achieve a noble gas electron configuration.
Na
 Na+ + e[Ne]3s1 [Ne]
Al
 Al3+ + 3e[Ne]3s23p1
[Ne]
Some transition metals lose their highest
energy level s and p electrons but still have d
electrons remaining. Their electron
configuration is not quite that of a noble gas
but is still stable. It is called a pseudo-noble
gas electron configuration. For example, zinc
loses its two electrons in 4s but keeps the ten
electrons in 3d.
Transition metals always lose their
highest numerical energy level electrons
first. Transition metals in the 4th period
lose their 4s and 4p electrons before
losing any from 3d. Metals in groups 3, 4,
& 5 do this also.
Example:
Fe
[Ar]4s23d6
Fe2+
[Ar]3d6
Fe3+
[Ar]3d5
Nonmetals tend to gain electrons to become
stable and form negatively charged anions.
They achieve a noble gas electron
configuration.
Example:
Cl + e-  [ Cl ][Ne]3s23p5 [Ne]3s23p6
N + 3e-  [ N ]31s22s22p3
1s22s22p6
Ionic Bonding- the attraction of oppositely
charged ions (cations and anions)
When the electronegativity difference
between two elements is large, the elements
are likely to form a compound by ionic
bonding (transfer of electrons). The farther
apart across the periodic table two Group A
elements are, the more ionic their bonding
will be.
We can use Lewis dot formulas to represent the
formation of ionic compounds.
Na + Cl  Na+[ Cl ]- or NaCl
Mg:
Mg:
Mg:
N

N
Mg2+ [ N ]3Mg2+
or Mg3N2
Mg2+ [ N ]3-
Properties of Ionic Compounds:
They are usually crystalline solids with high
melting points (>400oC)
Their molten compounds and aqueous
solutions conduct electricity well because
they contain mobile charged particles.
Metals
Metals form metallic solids that consist of
positively charged metal cations in a “sea” of
loosely held valence electrons. This
arrangement allows metals to have their
unique properties.
Metals are ductile (can be pulled into a wire)
and malleable (can be hammered into a thin
sheet) because the valence electrons act as
“grease”, allowing the cations to slide past
each other without colliding with each other
and shattering. When ionic compounds such
as NaCl are hammered, like-charged ions
collide causing repulsion and the crystal
shatters.
Metals can conduct electricity easily.
Electricity is a flow of electrons. As electricity
(electrons) enters one end of a piece of
metal, an equal number of electrons exit the
other end.
Alloyssolutions of solids in solids
Substitutional alloy- atoms of one metal are
substituted for atoms of a similar-sized metal
in a metallic crystal.
Ex. brass, sterling silver, pewter
Interstitial alloy- smaller metal atoms fit into
holes in the crystal structure of a metal with
larger atoms
Ex. steel (carbon in iron)
Amalgam- alloy which contains
mercury
Covalent Bonding
Hydrogen and nonmetals of Groups 4,5,6 & 7
often become stable and gain noble gas
electron configurations by sharing electrons
to form covalent bonds. Atoms will usually
share electrons to follow the octet rule (eight
electrons, like most noble gases) or the duet
rule (2 electrons, like helium).
When atoms share one pair of electrons to
form a covalent bond, it is called a single
covalent bond. The electrons shared between
the atoms are a “shared pair”. A dash can be
used instead of two dots to represent the
shared pair. Any other electrons on the
atoms are “unshared pairs” or “lone pairs”.
Ex.
H2
H-H
Cl2
Cl-Cl
HCl
H-Cl
Atoms must sometimes share more than one
pair of electrons to become stable. When
two pair of electrons are shared between two
atoms, it is called a double bond. If three pair
are shared, it is a triple bond.
Ex. O2
N2
O=O
NN
Rules for Writing Lewis Structures
(electron dot structures):
(Use pencil!)
1.Add up the valence electrons from all the
atoms. Don’t worry about keeping track of
which electrons come from which atoms. If
you are working with an ion, you must add or
subtract electrons to equal the charge.
2.Use a pair of electrons to form a bond
between each pair of bound atoms.
3.Arrange the remaining electrons to satisfy
the duet rule for hydrogen and the octet rule
for everything else.
4.If necessary, change bonds to double or
triple.
5.Remember, we cannot create or destroy
electrons!
H2 O
8 electrons
H-O-H
H-O-H
NH3
8 electrons
H-N-H
H
NH4+
9-1 = 8 electrons
H
+
H-N-H
H
CO2
16 electrons
O-C-O
This used 20 electrons! BAD!!!
CO2
O=C=O
CCl4
32 electrons
CCl4
32 electrons
Cl
Cl-C-Cl
Cl
CN9 + 1 = 10 electrons
C-N
BAD!
CN9 + 1 = 10 electrons
CN
SO4232 electrons
O
2O S O
O
CO3224 electrons
O
O
C O
O
O
C O
2-
2-
Coordinate Covalent Bond- Bond in which
both electrons came from the same atom.
This bond is not really any different than any
other single bond.
Exceptions to the Octet Rule:
A few compounds are stable with less
than an octet. They include beryllium or
boron. These electron deficient
compounds are very reactive.
BF3
BeCl2
24 electrons
16 electrons
F
F-B-F
Cl-Be-Cl
Elements in the third period and below
can exceed the octet rule. They can place
extra electrons in empty d orbitals. Elements
in the second period can not exceed the
octet rule because there is no 2d orbital for
the extra electrons to go into. If it is
necessary to exceed the octet rule, place the
extra electrons on the central atom.
Ex.
PCl5
Cl
Cl
Cl
P
Cl Cl
SF6
F
F
F
S
F
F
F
I3-
22 electrons
[ I-I-I
]
More practice:
NF3
F
F-N-F
OF2
20 electrons
F-O-F
KrF4
36 electrons
F
F- Kr - F
F
BeH2
4 electrons
H - Be - H
SO3
2-
26 electrons
O
O-S-O
NO
3
24 electrons
O
O-N-O
H2 O 2
14 electrons
H-O-O-H
Resonance occurs when more than one valid
Lewis structure can be written for a
molecule. The actual structure is an average
of all of the resonance structures.
Ex. NO3-
O
O-N-O
O
O-N-O
O
O-N-O
In nitrate, the experimental bond length
is in-between that of a single bond and a
double bond. It acts like a 1 1/3 bond.
Ex. Benzene, C6H6
VSEPR
Lewis structures can be used to
determine the shapes of molecules.
Their shapes will tell us a lot about
their chemical behavior.
The valence shell electron pair repulsion (VSEPR)
theory tells us that valence electrons on the
central atom repel each other. They are arranged
as far apart as possible around the central atom so
that repulsions among them are as small as
possible. When we are using VSEPR to determine
molecular shape, we are really looking for regions
of electron density. Double and triple bonds
count the same as single bonds in determining
molecular shape.
In CO2, there are only two regions of
electron density (effective electron pairs)
around the central atom. These regions
arrange themselves as far apart as possible,
making the bond angle 180o and the
molecular shape linear.
O=C=O
In CO32-, there are three effective electron
pairs around the central atom. The bond
angle will be 120o and the shape will be
trigonal planar.
2O
C
O
O
In CH4, there are four effective electron pairs. We
might expect the bond angle to be 90o. Actually, since
molecules are three-dimensional, the electron pairs
are 109.5o apart (further than 90o) and take a
tetrahedral arrangement.
H
H
C
H
H
In NH3, there are four effective electron pairs. Three
are shared but one is unshared. Unshared pairs of
electrons take up more space than shared pair because
they are pulled closer to the nucleus. The presence of
the unshared pair distorts the other bond angles,
making them less than 109.5o and the shape is called
trigonal pyramidal. (The bond angles in ammonia are
about 107o.)
H - N - H
H
In H2O, there are four effective electron pairs,
also. Two are shared and two are unshared.
Since the unshared pairs repel more than
shared pair, the bond angle is less than 109.5o
(actually 104.5o for H2O) and the shape is
bent.
H–O–H
In PF5, there are five effective pairs, all shared. The
bond angles are 90o and 120o and the shape is
called trigonal bipyramidal. It is like two trigonal
pyramids with their bases touching.
F F
F-P-F
F
In SF6, there are six effective pairs, all
shared. The bond angles are 90o and the
shape is called octahedral. It is like two
square pyramids with their bases
touching.
Molecules that exceed the octet rule and
have unshared electrons can have more
complex shapes such as T-shaped, see-saw, and
square pyramidal.
Practice Determining Molecular
Shape:
H2 S
H-S-H
bent,
<109.5o
CCl4
tetrahedral,
109.5o
Cl
Cl-C-Cl
Cl
NH4+
tetrahedral,
109.5o
H
H-N-H
H
BF3
trigonal
planar,
120o
F
B
F
F
NO2-
O-N=O
bent,
<120o
PF6-
F F F
P
F
F
F
octahedral,
90o
SbCl5
Cl
Cl Cl
Sb
Cl Cl
trigonal
bipyramidal,
90o and 120o
SO2
O-S=O
bent, <120o
POLAR AND NONPOLAR
MOLECULES
Covalent bonds within a molecule can be
polar (electrons are shared unequally) or
nonpolar (electrons are shared equally). We
can predict polar bonds by looking up the
electronegativity values of each element in
the bond and subtracting the smaller value
from the larger value to determine the
electronegativity difference.
If the electronegativity difference (EN)
is zero, the bond is nonpolar covalent. If
the EN is between 0 and 2.0, we can
predict that the bond is polar covalent.
If the EN is 2.0 or greater, the bond is
usually considered to be ionic.
There are no strict dividing lines
between covalent and ionic bonds.
Physical properties are used to help
determine whether something is
covalent or ionic. Some covalent
substances are so polar that they ionize
partially or completely in water.
Polarity of a bond or molecule can be
represented by arrows or lowercase delta ()
symbols to show a partial charge.
Ex.
H F
+ This shows that the hydrogen end of the
molecule is more positive (less
electronegative) and the fluorine end is more
negative (more electronegative). A charge
difference such as this is called a dipole.
An entire molecule is polar if it has
polar bonds that do not cancel. This
happens in nonsymmetrical molecules.
Molecules with all nonpolar bonds such as H2,
O2, and Cl2, are always nonpolar. This means
that there are no positive and negative ends
on the molecule.
Heteronuclear diatomic molecules such
as HCl, BrCl, or HF are always polar
because the polar bonds can not cancel
each other out.
Linear triatomic molecules are polar only if
the two outside legs are different. CO2 is
nonpolar because the two C=O bonds cancel
each other out.
Bent molecules are polar.
Trigonal planar molecules are nonpolar if the
three legs are the same.
Trigonal pyramidal molecules are polar.
Tetrahedral molecules are nonpolar if the four
legs are all the same.
Trigonal bipyramidal and octahedral molecules are
nonpolar if all legs are the same element.
Practice: Polar or Nonpolar?
NH3
1st: Draw dot structure
2nd: Determine shape
3rd: Symmetrical = Nonpolar
Nonsymmetrical = Polar
Polar
H
H-N-H
SO2
O=S-O
Polar
H2 O
H-O-H
Polar
BF3
F
F-B-F
Nonpolar
CH4
Nonpolar
H
H-C-H
H
INTERMOLECULAR
FORCES
Intramolecular bondingsharing electrons
Intermolecular bonding- interactions
between particles (atoms, molecules
or ions)
Changes in state are due to changes
in intermolecular bonding, not
intramolecular bonding.
Dipole-dipole attraction- attraction of
polar molecules for each other.
(negative-positive)
*approx. 1% as strong as covalent or
ionic bonds.
*molecules orient themselves to
minimize repulsion and maximize
attractions
Hydrogen bonding- unusually strong dipoledipole attractions involving hydrogen atoms
which are covalently bonded to a very
electronegative element and a very
electronegative atom (F,O,N only) with
unshared electrons.
Two reasons for strength of hydrogen bonds:
1. small size of H atom allows closeness
2. great polarity
Substances with much H bonding have high
boiling points compared to similar substances.
Ex. H2O, NH3, HF
London dispersion forces (LDFs)- relatively weak
forces (usually) that exist between noble gas
atoms and nonpolar molecules. LDFs also exist in
compounds that have dipole-dipole and/or
hydrogen bonding. LDFs may be the most
important force in large molecules of these types.
LDFs occur because of momentary electron
imbalance (temporary dipole) which can induce
the same to occur in adjacent molecules.
This force is often very weak, thus the low
freezing point of noble gases. The freezing
point of noble gases increases going down the
group because heavier atoms have more
electrons and an increased chance of
temporary dipoles. They also have a lower
velocity and have more opportunity for
attractions. This causes London dispersion
forces to increase going down a group
on the periodic table.
Covalent network solidGroup 4 substances such as diamond, Si, SiC,
and Ge form extensive covalent bonds and
result in giant molecules. They have an atom
at each lattice point and are held together by
covalent bonds. These substances have the
strongest attractions.
General trends in strength of attraction:
(weakest)
LDF
dipole-dipole
hydrogen-bonding
metallic bonding
ionic bonding
covalent network bonding
(strongest)