1. Lewis diagrams molecular geometry bond and molecular polarity IMFAs
Unit 7
Lewis diagrams
molecular geometry
bond and molecular polarity
IMFAs

2. Lewis dot diagrams
add up the total number of valence electrons for all atoms in the molecule
arrange the atoms to pair up the separate atoms’ single electrons as much as possible
confirm that:
the total number of electrons exactly matches the total valence electrons of the original atoms, and
each atom has an octet of electrons (8), except
H and He have a duet of electrons (2)

3. structural formulas
also called “Lewis structures” or “Lewis diagrams” (but not “Lewis dot structures”)
replace each shared pair of electrons with a solid line representing a covalent bond consisting of two shared electrons
continue to show the lone pairs of electrons (which are unshared)
double-check that the lone pairs plus bond pairs still add up to the correct total number of valence electrons

4. multiple bonds
additional bonds may need to be added to a Lewis structure if
single electrons remain
atoms do not have octets
in simple cases, you may be able to pair up single electrons on adjacent atoms to form additional bonds, e.g.
CO2 N2
C2H4

5. multiple bonds
in other cases, you cannot strictly keep electrons with their original atoms; the electrons are free to move elsewhere in the molecule as needed to complete octets, e.g.
carbon monoxide, CO
ozone, O3
in these cases, atoms may not form their “normal” number of bonds
but the total number of valence electrons must not change; they are just rearranged

6. multiple bonds computational approach
you can also calculate exactly how many bonds are in a molecule in the following way
add up the valence electrons that the atoms in the molecule actually have
separately add up the valence electrons those atoms need in order to have noble gas configurations
calculate the difference, need – have
that difference is the number of shared electrons the molecule must have
every 2 shared electrons make one bond

7. multiple bonds computational approach
O O
C O
have: = 12
have: = 10
need: = 16
need: = 16
4 shared e-
6 shared e-
thus 2 bonds
thus 3 bonds
after building the basic skeleton with bonds
add remaining electrons as needed to complete octets
double-check that the total number of electrons is exactly the number of valence electrons (“have”)

8. general hints for Lewis structures
if a given molecule can be drawn with both symmetrical and asymmetrical structures, the symmetrical one is more likely to be correct
central atoms are often
written first in the formula
the least electronegative element
the element that can form the most bonds
hydrogen and halogens
only form one bond, thus are terminal atoms
are generally interchangeable in molecules

9. exceptions to octet “rule”
most atoms have octets (8 valence electrons) when in molecules, but there are exceptions
group
number of electrons
number of bonds
examples
column 1
duet (2) 1
H2, LiH
column 2
quartet (4) 2
BeH2 , MgI2
column 3
sextet (6) 3
BH3 , AlCl3
columns 4-8
octet (8)
4 bonds
3 bonds + 1 lone pair
2 bonds + 2 lone pairs
1 bond + 3 lone pairs
CH4
NH3
H2O
HCl

10. molecular shapes: VSEPR model
valence shell electron-pair repulsion
groups of electrons naturally find positions as far apart from each other as possible
different molecular shapes result based on how many groups of electrons are present
each of the following counts as one “set” of electrons around the central atom
a lone pair
a single bond (2 shared e-)
a double or triple bond (4 or 6 shared e-)

11. VSEPR model—central atom with:
2 sets of e–
3 sets of e–
4 sets of e–
5 sets of e–
6 sets of e–
linear
trigonal planar
tetrahedral
trigonal
bipyramidal
octahedral
e.g. BeF2
e.g. BF3
e.g. CF4
e.g. XeF6
e.g. SF5

12. electron geometry vs. molecular shape
each set of electrons occupies a position around the central atom
the number of sets defines the electron geometry
but lone pairs are essentially transparent
even though they are invisible, lone pairs make their presence known by distorting the positions of the bonds around them (since lone pairs repel the electrons in the bonds)
this results in several related molecular shapes within each general class of electron geometry

13. tetrahedral electron geometry 4 electron sets
bonds
lone pairs
molecular shape
example
4 single
tetrahedral
CH4
3 single 1
triangular pyramid
NH3
2 single 2
bent (~109°)
H2O
1 single 3
linear
HCl

14. tetrahedral electron geometry

15. tetrahedral electron geometry

16. triangular planar electron geometry 3 electron sets
bonds
lone pairs
molecular shape
example
3 single
triangular planar
BH3
2 single + 1 double
CH2O
1 single + 1 double 1
bent (~120°) O3

17. linear electron geometry 2 electron sets
bonds
lone pairs
molecular shape
examples
2 single
linear
BeH2
2 double
CO2
1 single + 1 triple
HCN
in addition, any diatomic molecule must be linear (since any two points lie on a line)

18. triangular planar and linear electron geometry

19. bond polarity
two electrons shared between two atoms form a covalent bond
if those electrons are shared equally (or nearly equally), it is a non-polar covalent bond
if one atom attracts the electrons much more strongly than the other atom, it is a polar covalent bond
if one atom completely removes an electron from the other atom, the result is an ionic bond

20. ΔEN, electronegativity difference
bond polarity
the electronegativity difference between the two atoms determines how polar a bond is
bond type
ΔEN, electronegativity difference
non-polar
polar
ionic
0.0 – 0.4
0.5 – 1.7
> 1.7
Cℓ2
HCℓ
LiCℓ

21. bond polarity
dipole moment is the actual measureable quantity related to bond polarity
the size of the dipole moment is affected by
electronegativity difference
bond length
we will focus on ΔEN and a qualitative sense of bond polarity

22. molecular polarity
the overall polarity of a molecule depends on the combined effect of the individual polar bonds
individual bonds polar
individual bonds polar
overall molecule
nonpolar
overall molecule
polar

23. molecular polarity what allows bond dipoles to cancel?
geometric symmetry of the molecule
having identical terminal atoms (or atoms with the same electronegativity)
what prevents bond dipoles from canceling?
geometric asymmetry (due to lone pairs)
having different terminal atoms

24. molecular polarity

25. molecular polarity
inherently symmetrical shapes (if all surrounding atoms are the same)
tetrahedral
triangular planar
linear
inherently asymmetrical shapes
bent
triangular pyramid
even symmetrical shapes become asymmetrical if different terminal atoms are attached

26. IMFA: intermolecular forces of attraction
“mortar”— holds the separate pieces together
(i.e. the IMFA)
“bricks”— individual atoms, ions, or molecules of a solid

27. IMFA: intermolecular forces of attraction

28. types of IMFA covalent network atoms such as C, Si, & Ge ionic bond
strongest
occurs between
covalent network
atoms such as C, Si, & Ge
(when in an extended grid or network)
ionic bond
cations and anions
(metals with non-metals in a salt)
metallic bond
metal atoms
hydrogen bond
ultra-polar molecules
(those with H–F, H–O, or H–N bonds)
van der Waals forces
dipole-dipole attraction
polar molecules
London forces
non-polar molecules
weakest

29. IMFAs – Trends and Characteristics
melting points and boiling points
Solubility
Conductivity (type specific)
Strength (type specific)

30. Melting Points and Boiling Points
stronger IMFAs cause higher m.p. and b.p.
atoms and molecules that are heavier and/or larger generally have higher m.p. and higher b.p.
larger e– clouds can be distorted (polarized) more by London or dipole forces, causing greater attraction
CH3CH2CH2CH2CH3 > CH3CH2CH3
Polarity (for dipole dipole and H-bonds) more polar=higher b.p, m.p.
HCl > HI

31. Melting Points and Boiling Points
strategy to predict m.p. and b.p.
first sort atoms/molecules into the six IMFA categories
then sort within each IMFA category from lightest to heaviest (or least polar to most polar)

32. same IMFA: sort by molar mass
melt
boil
ex: halogen family
all are non-polar (London force)
lowest to highest m.p. and b.p. matches lightest to heaviest I2
(257)
+184.4 °C
–250
–200
–150
–100
–50
+50
+100
+150 I2
(257)
+113.7
Br2
(160)
+58.8
thus at room temperature:
F2 (g)
Cℓ2 (g)
Br2 (ℓ)
I2 (s)
Br2
(160)
–7.2
Cℓ2
(71)
–34.04
Cℓ2
(71)
–101.5 F2
(38)
–182.95 F2
(38)
–219.62

33. same mass: sort by IMFA type°C
–50
+50
+100
+150
+198
ethylene glycol
(can form twice as many H-bonds)
ex: organic molecules
all are ~60 g/mol
different types of IMFA
+97.4
1-propanol (ultra-polar = H-bonds)
+56.2
acetone (more polar)
+10.8
methyl ethyl ether (slightly polar)
–0.5
butane (non-polar)
the stronger the IMFA, the higher the boiling point

34. Solubility substances generally mix best with others of similar IMFAs
”like dissolves like”
non-polar mixes well with non-polar
polar mixes well with polar (and ultra-polar / ionic)

35. other physical properties
strength, conductivity, etc. are related to the type of IMFA

36. details about each IMFA
strongest
covalent network
ionic bond
metallic bond
hydrogen bond
dipole-dipole attraction
London forces
weakest

37. London (or dispersion) forces
non-polar molecules (or single atoms) normally have no distinct + or – poles
electron clouds are slightly distorted by neighboring molecules
sort of like water sloshing in a shallow pan
form temporary dipoles
Low MP and BP
soluble

38. London dispersion forces in action
1. temporary polarization due to any random little disturbance δ+ δ-
2. induced polarization caused by neighboring molecule
3. induced polarization spreads
4. induced polarization reverses
non-polar molecules, initially with uniform charge distribution

39. dipole-dipole attractions
polar molecules have permanent dipoles
the molecules’ partial charges (δ+, δ-) attract the oppositely-charged parts of neighboring molecules
this produces stronger attraction than the temporary polarization of London forces
therefore polar molecules are more likely to be liquid at a temperature where similar non-polar molecules are gases

40. dipole-dipole attractionsδ+ δ-

41. hydrogen bonding (or ultra-dipole attractions)
H—F, H—O, and H—N bonds are more polar than other similar bonds
Very small atoms, particularly H
F, O, and N are the most electronegative elements
particularly polar
molecules containing these bonds have much higher m.p. and b.p than otherwise expected for non-polar or polar molecules of similar mass
the geological and biological systems of earth would be completely different if water molecules did not H-bond to each other

42. hydrogen bonding (or ultra-dipole attractions)
ultra-polar molecule
(much higher boiling point)
hydrogen bonds
(between molecules,
not within them)
non-polar molecules
(lower boiling points)

43. hydrogen bonding (or ultra-dipole attractions)
Beware!! H O H O
These are not hydrogen bonds. They are normal covalent bonds between hydrogen and oxygen. H O H O
These are hydrogen bonds. They are between separate molecules (not within a molecule).

44. metallic bonding structure resulting properties
nuclei arranged in a regular grid or matrix
“sea of electrons”—delocalized valence electrons free to move throughout grid
resulting properties
conductive (electrically and thermally)
strong, malleable, and ductile shiny
Form alloys = mixture of metals
Bronze = copper and tin
Brass = copper and zinc
Steel = iron and carbon
metallic “bond” is generally weaker than covalent bond since there are not specific e– pairs forming bonds

45. ionic bonding (salts)
structure: orderly 3-D array (crystal) of alternating + and – charges
made of
cations (metals from left side of periodic table)
anions (non-metals from right side of periodic table)
properties
non-conductive when solid
conductive when melted or dissolved
hard but brittle (why?)…

46. why are salts hard but brittle?
1. apply some force
2. layer breaks off and shifts
4. shifted layer shatters away from rest of crystal
3. + repels +
– repels –

47. covalent networks
strong covalent bonds hold together millions of atoms (or more) in a single strong particle
properties
very high melting temperatures
usually non-conductive (except graphite)
very hard, very strong
examples
carbon (two allotropes: diamond, graphite)
pure silicon or pure germanium
SiO2 (quartz or sand)
other synthetic combinations averaging 4 e– per atom:
SiC (silicon carbide), BN (boron nitride)

48. buckminsterfullerine
m.p. = 3550°C
C60
buckminsterfullerine
“bucky ball”
m.p. = ~1600°C

49. summary of properties network ionic metallic hydrogen dipole London
strongest
strength
m.p. & b.p.
conductive?
network
extremely hard
very high
usually not
ionic
hard but brittle
medium to high
if melted or dissolved
(mobile ions)
strong, malleable, ductile
medium to high
metallic
very
(delocalized e–)
hydrogen
van der Waals forces
dipole
soft and brittle
low no
London
weakest

50. Metallic London Ionic ---------- Hydrogen
Network
Metallic
Hydrogen
Ionic
Covalent

51. D F H B G J C I E A

52. To do’s: Objective: Prepare for test tomorrow. Review HW
Characterize each of the 6 IMFA
Practice questions
Independent practice
Practice tests/homework
Ask questions

53. ionic Metallic Hydrogen bonded Dipole-dipole London ForcesMP BP
Conductivity
Solubility
Strength
Very strong, hard
Covalent network
mostly
Insoluble
Not conductive
C, Si, & Ge
Conductive in solution or molten state
(mobilized ions)
ionic
Soluble in polar solvents
Hard brittle
Metal + nonmetal
Good conductor of heat and electricity
mostly insoluble
Strong, malleable, ductile
Metallic
Elements/alloys left side of periodic table
Hydrogen bonded
Soluble in polar solvents
Soft,
brittle
Not conductive
H bonded to O, N, F
Dipole-dipole
Soluble in polar solvents
Not conductive
Soft, brittle
Polar molecules
London Forces
Not conductive
Soluble in non-polar solvents
Soft, brittle
non-polar molecules

54. soap or emulsifier soaps and emulsifiers oil water
some molecules are not strictly polar or non-polar, but have both characteristics within the same molecule
oil
polar region
soap or emulsifier
non-polar region
water
this kind of molecule can function as a bridge between molecules that otherwise would repel each other

55. soaps and emulsifiers
with a soap or emulsifier present to surround it, a drop of non-polar oil can mix into polar water

56. Practice Problems…

57. Draw 3 examples of molecules with polar bonds that overall are not considered to be polar molecules.

58. Draw the Lewis structure for AsH3.
Determine the shape and overall polarity.

59. Which of the following is most likely to dissolve NaCl?
BH3
CH4
NH3
BeH2

60. Draw 3 isomers of C5H12

61. Review for test tomorrow:
Quiz 2
HW # 13, 14, 16,17
Practice test
Review questions from slides