Bonding and Forces of
Attraction
(mostly Chapter 9)
Metallic Bonding
• In
metallic bonding instead of sharing electrons
between two atoms (covalent bonding), the valence
electrons are shared between all the atoms in a
positive metal crystal. These atoms are attracted to
the negatively charged 'cloud’ or ‘sea' of electrons.
• The movement or delocalization of the free
electrons means that metallic bonded materials have
good thermal and electrical conduction.
http://www.ider.herts.ac.uk/school/courseware/materials/bonding.html
Ionic Bonds
Ionic Bonds
Occur when the nonmetal
“steals” one or more electrons
away from a metal.
The nonmetal becomes a negative ion
while the “victim” metal becomes a
positive ion. The atoms are held
together by their opposite charges.
Ionic Bonds
Strength of crystal lattice depends on two
factors, size and charge transferred.
Smaller atoms have stronger ionic bonds.
Ex: NaF is stronger than NaCl
Atoms transferring more electrons are
stronger. Ex: MgCl2 is stronger than NaCl.
Covalent Bonding
The bonds of nature.
Covalent Bond
• Shared valence electrons
• Complete outer energy levels
• Molecule has 2 or more nonmetal atoms
covalently bond
– Carbohydrates, proteins, fats, DNA,
stupendous seven (H2, N2, O2, F2, Cl2, Br2, I2)
How do they form?
Distance is to
great
E - E – repulsive
P – E - attractive
Repul = Attract
P – P – repulsive
E – E - repulsive
Attractive forces balance the repulsive forces
Electronegativity
• Electronegativity is the
ability of atoms in a
molecule to attract
electrons to themselves.
• On the periodic chart,
electronegativity
increases as you go…
– …from left to right across
a row.
– …from the bottom to the
top of a column.
What types of bonds are they?
MgO, water, Calcium Carbide, Potassium Oxide,
Nitrogen trihydride
Electronegativity and Bond Type
• Ability to attract electrons in a chemical
bond—subtract for bond type.
• Ionic Bonds 1.7 or greater
• Polar Covalent Bonds < 1.7 and >0.2
• Pure or Nonpolar Covalent bonds <0.2
Dash Line Represents the Cut Off
Polar Covalent Bonds
• Electrons are shared, but not equally
• Non polar bonds – equal sharing, H2, O2,
Cl2
1 valence e
1 valence e
1 s1
1 s1
1 s2
Stable
Helium
configuration
Single Covalent bond –
shared electrons (also
known as a sigma bond σ)
1s1
1s1
1s2 2s2 2p4
H –1s2 - He
O - 1s2 2s2 2p6 - Ne
Lone pairs –
unshared elecrons
The molecules are more stable because they have
complete outer energy levels – Octet Rule
Bonding Capacity
Atom
Carbon
Nitrogen
Oxygen
Halogens
Hydrogen
Number of
Valence
Electrons
Number of Bonding
Bonding
Capacity
Electrons
Drawing Lewis Dot Structures
1. Count the valence electrons.
2. Predict the location of the atoms
a. Hydrogen is a terminal atom
b. The central atom has the smallest electronegativity.
3. Draw a pair of electrons between the central
atom and the surrounding atoms.
4. Use the remaining electrons to complete the
octets of each atom. If there are electrons left
over, place them on the central atom.
5. If the central atom does not have a complete
octet then try double or triple bonds.
a. If the atom has 1, 2, or 3 valence electrons, it
doesn’t require an octet.
STEP 1: count the total # of
valence e- for all atoms
involved in the bonding
Carbon: 1 carbon with
4 valence
electrons (1x4) = 4
CCl4
4+28
=32
Chlorine: 4 chlorine
with 7 valence
electrons (4x7) = 28
STEP 2–place the single atom in the
center and other atoms around it
evenly spaced
CCl4
4+28
=32
Cl
Cl C Cl
Cl
STEP 3: place the electrons in pairs
between the central atom
and each non-central atom
CCl4
4+28
=32
-8
=24
Cl
Cl C Cl
Cl
STEP 4: place the remaining
electrons around the noncentral atom until each has
8 electrons (H atoms have)
CCl4
only 2e4+28
=32
-8
=24
-24
=0
Cl
Cl
C
Cl
Cl
Step 5: If you run out of electrons
before the central atom has an
octet…form multiple bonds until
it does.
Hydrogen—1 electron
Carbon—4 electrons
Ex—HCN
Nitrogen—5 electrons
TOTAL—1+4+5=10 e-s
H:C:N
..
H:C:N:
..
H:C:::N:
Drawing Lewis Dot Structures
Draw Lewis Dot Structures for:
PH3
H2S
HCl
CCl4
SiH4
CH2Cl2
Draw Lewis Dot Structures
Cl2
NF3
CS2
BH3
CH4
SCl2
C2H6
BF3
Strength of Bonds
• Based on proximity – bond length
• Influenced by atom size and number of
shared electrons
– F2 is greater than Cl2 is greater than Br2
F2 1.43 x 10-10 m - single
O2 1.21 x 10-10 m - double
N2 1.10 x 10-10 m – triple
Sigma Bonds - σ
• Overlapping
electron orbitals
create a bonding
orbital – likely
hood of finding an
electron
• S overlaps with S
• S overlaps with P
• P overlaps with P
Pi Bonds--π
Pi bonds happen
after sigma bonds. While sigma bonds are head
to head bonds, pi bonds are side to side bonds of
the same two atoms. Every bond has one sigma
bond, but subsequent bonds between the same
two atoms must have a different way of
connecting so they use pi bonds!
Multiple Covalent Bonds – Double
and Triple Bonds
6 valence
electrons
6 valence
electrons
12 valence
electrons
1 sigma bond
Octet satisfied
More stable
and stronger
1 pi bond
5 valence
electrons
5 valence
electrons
10 valence
electrons
1 sigma bond
Octet satisfied
More stable
and stronger
2 pi bonds
Molecular Shapes
• The shape of a
molecule plays an
important role in its
reactivity.
• By noting the number
of bonding and
nonbonding electron
pairs we can easily
predict the shape of
the molecule.
What Determines the Shape of
a Molecule?
• Simply put, electron
pairs, whether they be
bonding or nonbonding,
repel each other.
• By assuming the electron
pairs are placed as far as
possible from each other,
we can predict the shape
of the molecule.
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
“The best
arrangement of a
given number of
electron domains is
the one that
minimizes the
repulsions among
them.”
Molecular Shape Chart
Formula
BeH2
BF3
CH4
NH3
H2O
Dot
Structure
Name of
Shape
Nonbonding
e- pairs
Bonding
electrons
Polarity
Hybridization
Polarity
Molecular Polarity
• Symmetrical molecules have equal charge
distribution so are nonpolar like tetrahedral CCl4
with 109.5o between chlorine atoms.
• Unshared electron pairs in molecules push the
shared pairs closer together, so NH3 (trigonal
pyramid) has 107o between hydrogens, and
water, H2O (bent), has 104.5o between hydrogen
atoms.
Hybridization
• The mixing of atomic
orbitals in an atom to
generate a new set of
bonding orbitals
(hybrids)
– Different shapes than
atomic orbitals
– Requires energy but
the energy is returned
during bond formation
2s
2p
2sp3
Hybrid Orbitals
• Consider beryllium:
– In its ground electronic
state, it would not be
able to form bonds
because it has no
singly-occupied orbitals.
Hybrid Orbitals
But if it absorbs the
small amount of
energy needed to
promote an electron
from the 2s to the 2p
orbital, it can form two
bonds.
This new orbital set is
hybridized with equal
energy—2sp orbitals
1s
2sp orbitals
2s
2p
2sp
2s
2p
2sp2
Group 2A
elements
make sp
hybridized
orbitals
Group 3A
elements
make sp2
hybridized
orbitals
Elements in group 4A have 4 valence electrons
and need four bonds to make an octet, therefore,
they will have sp3 hybridization.
2s2
2p2
2sp3
Questions
• Why do some solids dissolve in water
but others do not?
• Why are some substances gases at
room temperature, but others are liquid
or solid?
• What gives metals the ability to conduct
electricity, what makes non-metals
brittle?
• The answers have to do with …
Intermolecular forces
Intermolecular forces
There are 2 types of attraction in molecules:
intramolecular bonds & intermolecular forces
– intramolecular bonds attraction between
atoms in a molecule (covalent, ionic, and
metallic bonds)
– Intermolecular forces (IMF) the attraction
between molecules (hydrogen bonds, dipoledipole interactions, and London dispersion
forces
– These are listed from strongest to weakest in
each category.
Intermolecular Forces
Collectively, all intermolecular forces of
attraction are known as van der Waals
forces. They are not chemical bonds, but
simply attractions between two covalent
molecules.
For different substances of similar molar
mass, the weakest of these are London
dispersion forces. Carbon dioxide,
O=C=O, is a good example.
Dipole - Dipole attractions
• We have seen that polar
+ –
molecules have a separation of
charge.
• This happens in both ionic and
polar bonds (the greater the EN, H
Cl
the greater the dipoles)
• Molecules are attracted to each other in a
compound by these +δ and -δ opposite
charges
+ –
+ –
Polarity in Bonds vs. in Molecules
• NH3 and H2O are both
polar molecules with
polar bonds.
N-H: 3.04 - 2.20 = 0.84
O-H: 3.44 – 2.20 = 1.24
• CH4 and BeH2 are both
nonpolar molecules
with polar bonds.
C-H: 2.55 – 2.20 = 0.35
H-Be: 2.20 – 1.57 = 0.63
Electronegativity difference is
between 0.2 and 1.7 so bonds
in molecule are polar
Unshared electron pairs in trigonal
pyramid and bent molecules
make both molecules polar
Again, electronegativity difference is
between 0.2 and 1.7 so bonds in
molecule are polar
Symmetry of the tetrahedral and 3atom linear molecules make both
molecules nonpolar
London forces (temporary dipoles)
Since non-polar molecules do not have
dipoles, how can they attract each other to
form solids or liquids?
• London forces! (named after Fritz London)
– London forces exist in non-polar molecules
– Since electrons move in atoms, there are
instants when the electrons shift to one side,
causing that side to be slightly negative. This
induces the surrounding molecules to do the
same.
– This results in the opposite side becoming
slightly positive.
London forces
Instantaneous dipole:
Induced dipole:
Eventually electrons A dipole forms in one atom
are situated so that or molecule, inducing a
tiny dipoles form
dipole in the other
Polar molecules already have a separation
of charge, due to either unshared electron
pairs, or a change of one of the atoms
around the central atom.
Nonpolar molecules have even charge
distribution because they are symmetrical.
Hydrogen bonding
Hydrogen bonding is a special type of dipole dipole attraction that is very strong
– It occurs only when hydrogen is bonded to
either N, O, or F
– The high electronegativity difference of N-H,
O-H, and H-F bonds cause these to be about
5x stronger than normal dipole-dipole forces
– Compounds containing these bonds are
important in biological systems (hold strands
of DNA together; cause shape of enzymes)
– An exceptionally important example is water!
In conclusion:
• If two molecules have roughly the same molar
mass, determine which has the strongest
attractions by drawing the electron dot structure.
• Those with the highest boiling point, are
attracted to each other the strongest, so for
boiling point:
hydrogen bonded > dipole-dipole > nonpolar
To Help You Review!
• Read through the questions that follow to
review this unit. These questions will help
you on your test. The answers follow!
Testing concepts
1. Which attractions are stronger: intermolecular or
intramolecular?
2. How much stronger is a covalent bond than a
dipole-dipole interaction?
3. What evidence is there that nonpolar molecules
attract each other?
4. Which chemical in table 10.4, p454, has the weakest
intermolecular forces? Which has the strongest?
How can you tell?
5. Suggest some ways that the dipoles in London
forces are different from the dipoles in dipole-dipole
attractions.
6. A) Which would have a lower boiling point: O2 or F2?
Explain. B) Which would have a lower boiling point:
NO or O2? Explain.
7. Which would you expect to have the higher melting
point (or boiling point): C8H18 or C4H10? Explain.
8. What two factors causes hydrogen bonds to be so
much stronger than typical dipole-dipole bonds?
9. So far we have discussed 3 kinds of intermolecular
forces: dipole-dipole, hydrogen bonding, and
London forces. What kind(s) of intermolecular
forces are present in the following substances:
a) NH3, b) SF6, c) PCl3, d) HBr, e) CO2
(hint: consider molecular shape/polarity)
Challenge: Ethanol (CH3CH2OH) and dimethyl ether
(CH3OCH3) have the same formula (C2H6O).
Ethanol boils at 78 C, whereas dimethyl ether boils
at -24 C. Explain why the boiling point of the ether
is so much lower than the boiling point of ethanol.
Challenge: try answering the question on the next slide.
H – bonding and boiling point
Boiling point
Predicted and actual boiling points
100
Group 4
Group 5
50
0
-50
Group 6
-100
-150
-200
Group 7
2
3
4
5
Period
• Why does BP as period , why are some BP
high at period 2?
Testing concepts
1. Intramolecular are stronger (COVALENT).
2. A covalent bond is 100x stronger.
3. The molecules gather together as liquids or solids
at low temperatures.
4. Based on boiling points, CH4 (-188) has the
weakest forces, H2O has the strongest (100).
• London forces
– Are present in all compounds
– Can occur between atoms or molecules
– Are due to electron movement not to EN
– Are transient in nature (dipole-dipole are more
permanent).
– London forces are weaker
Testing concepts
6. A) F2 would be lower because it is smaller.
Larger atoms/molecules can have their
electron clouds more easily deformed and
thus have stronger London attractions and
higher melting/boiling points.
B) O2 because it has only London forces. NO
has a small EN, giving it small dipoles.
7. C8H18 would have the higher melting/boiling
point. This is a result of the many more sites
available for London forces to form.
8. 1) a large EN, 2) the small sizes of atoms.
Testing concepts
9. a) NH3: Hydrogen bonding (H + N), London.
b) SF6: London only (octahedral).
c) PCl3: trigonal pyramid. Dipole-dipole,
London.
d) HBr: EN=2.8-2.1. Dipole-dipole, London.
e) CO2: London only (it is symmetrical)
Challenge: In ethanol, H and O are bonded (the
large EN results in H-bonding). In dimethyl
ether the O is bonded to C (a smaller EN
results in a dipole-dipole attraction rather
than hydrogen bonding).
Testing concepts
Boiling points increase down a group (as
period increases) for two reasons: 1) EN
tends to increase and 2) size increases. A
larger size means greater London forces.
Boiling points are very high for H2O, HF, and
NH3 because these are hydrogen bonds
(high EN), creating large intermolecular
forces
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