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Lewis dot diagramsadd 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)
structural formulasalso 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
multiple bondsadditional bonds may need to be
added to a Lewis structure ifsingle electrons remainatoms 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
multiple bondsin 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, COozone, 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
multiple bondscomputational approach
you can also calculate exactly how many bonds are in a molecule in the following wayadd up the valence electrons that the
atoms in the molecule actually haveseparately add up the valence electrons
those atoms need in order to have noble gas configurations
calculate the difference, need – havethat difference is the number of
shared electrons the molecule must have
every 2 shared electrons make one bond
multiple bondscomputational approach
O2
after building the basic skeleton with bondsadd remaining electrons as needed to
complete octetsdouble-check that the total number of
electrons is exactly the number of valence electrons (“have”)
have: 6 + 6 = 12need: 8 + 8 = 16
O O
4 shared e-
thus 2 bonds
CO
have: 4 + 6 = 10need: 8 + 8 = 16
C O
6 shared e-
thus 3 bonds
general hints for Lewis structuresif a given molecule can be drawn
with both symmetrical and asymmetrical structures, the symmetrical one is more likely to be correct
central atoms are oftenwritten first in the formulathe least electronegative elementthe element that can form the most
bondshydrogen and halogens
only form one bond, thus are terminal atoms
are generally interchangeable in molecules
exceptions to octet “rule”most atoms have octets (8 valence
electrons) when in molecules, but there are exceptionsgroup number of
electronsnumber of bonds example
s
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 bonds3 bonds + 1 lone
pair2 bonds + 2 lone
pairs1 bond + 3 lone
pairs
CH4
NH3
H2OHCl
molecular shapes: VSEPR modelvalence shell electron-pair repulsiongroups 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 atoma lone paira single bond (2 shared e-)a double or triple bond (4 or 6 shared e-)
VSEPR model—central atom with:
2 sets of e–
linear
e.g. BeF2
3 sets of e–
trigonal planar
e.g. BF3
4 sets of e–
tetrahedral
e.g. CF4
5 sets of e–
trigonalbipyramidal
e.g. SF5
6 sets of e–
octahedral
e.g. XeF6
electron geometry vs. molecular shapeeach set of electrons occupies a position
around the central atomthe number of sets defines the electron
geometrybut lone pairs are essentially transparenteven 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
tetrahedral electron geometry4 electron sets
bonds lone pairs molecular shape example
4 single 0 tetrahedral CH4
3 single 1 triangular pyramid NH3
2 single 2 bent (~109°) H2O
1 single 3 linear HCl
triangular planar electron geometry
3 electron setsbonds lone
pairsmolecular shape example
3 single 0 triangular planar BH3
2 single + 1 double
0 triangular planar CH2O
1 single + 1 double
1 bent (~120°) O3
linear electron geometry2 electron sets
bonds lone pairs
molecular shape examples
2 single 0 linear BeH2
2 double 0 linear CO2
1 single + 1 triple 0 linear HCN
in addition, any diatomic molecule must be linear (since any two points lie on a line)
bond polaritytwo electrons shared between two
atoms form a covalent bondif 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
bond polaritythe electronegativity difference
between the two atoms determines how polar a bond is
Cℓ2 HCℓ LiCℓ
bond type ΔEN, electronegativity difference
non-polar
polar
ionic
0.0 – 0.40.5 – 1.7
> 1.7
dipole moment is the actual measureable quantity related to bond polarity
the size of the dipole moment is affected byelectronegativity differencebond length
we will focus on ΔEN and a qualitative sense of bond polarity
bond polarity
molecular polaritythe overall polarity of a molecule depends
on the combined effect of the individual polar bondsindividual bonds polar
individual bonds polar
overall moleculenonpolar
overall moleculepolar
molecular polarity
what allows bond dipoles to cancel?geometric symmetry of the
moleculehaving 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
molecular polarityinherently
symmetrical shapes (if all surrounding atoms are the same)tetrahedraltriangular planarlinear
inherently asymmetrical shapesbenttriangular pyramideven symmetrical shapes become
asymmetrical if different terminal atoms are attached
IMFA: intermolecular forces of attraction
“bricks”— individual atoms, ions, or molecules of a solid
“mortar”— holds the separate pieces together(the IMFA)
types of IMFAstrongest
weakest
London forces
dipole-dipole attraction
hydrogen bond
metallic bond
ionic bond
covalent network
occurs between
non-polar molecules
polar molecules
ultra-polar molecules(those with H–F, H–O, or H–N bonds)
metal atoms
cations and anions (metals with non-metals in a salt)
atoms such as C, Si, & Ge (when in an extended grid or network)
van
der
Waa
ls f
orc
es
consequences of IMFAsmelting points and boiling points rise with
strength of IMFAincreasing molar mass
substances generally mix best with other substances having the same or similar IMFAs”like dissolves like”non-polar mixes well with non-polarpolar mixes well with polar(polar also mixes well with ultra-polar and
ionic)other physical properties such as
strength, conductivity, etc. are related to the type of IMFA
predicting melting points, boiling pointsstronger IMFAs cause higher m.p. and
higher b.p.when atoms/ions/molecules are more strongly
attracted to each other, temperature must be raised higher to overcome the greater attraction
more polar molecules have 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/heavier atoms (higher molar mass) have
more e–
larger e– clouds can be distorted (polarized) more by London or dipole forces, causing greater attraction
strategy to predict m.p. and b.p.first sort atoms/molecules into the six IMFA
categoriesthen sort those in each category from lightest
to heaviest
same IMFA: sort by molar mass
thus at room temperature: F2 (g)
Cℓ2 (g)
Br2 (ℓ)
I2 (s)
°C
–250
–200
–150
–100
–50
0
+50
+100
+150
ex: halogen familyall are non-polar (London
force) lowest to highest m.p. and
b.p. matches lightest to heaviest
–219.62F2
(38)
melt boil
–101.5Cℓ2
(71)
–7.2Br2
(160)
+113.7I2
(257)
–182.95F2
(38)
–34.04
+58.8
+184.4
Cℓ2
(71)
Br2
(160)
I2
(257)
same mass: sort by IMFA type
°C
–50
0
+50
+100
+150
ex: organic molecules
all are ~60 g/moldifferent types of
IMFA
–0.5 butane (non-polar)
+10.8 methyl ethyl ether (slightly polar)
+56.2 acetone (more polar)
+97.4 1-propanol (ultra-polar = H-bonds)
+198 ethylene glycol(can form twice as many H-bonds)
the stronger the IMFA, the higher the boiling point
isomers (and an isobar)
n- and neo pentane
glycerol and 1-propanol
1-propanol and methyl ethyl ketone
butane and 2-methylpropane
1-propanol and 2-propanol
details about each IMFAstrongest
weakest
London forces
dipole-dipole attraction
hydrogen bond
metallic bond
ionic bond
covalent network
London (or dispersion) forcesnon-polar molecules (or single atoms)
normally have no distinct + or – poleshow can they attract each other enough
to condense or freeze?they form temporary dipoleselectron clouds are slightly distorted by
neighboring moleculessort of like water sloshing in a shallow
pan
London dispersion forces in action
non-polar molecules, initially with uniform charge distribution
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
dipole-dipole attractionspolar molecules have permanent dipolesthe molecules’ partial charges (δ+, δ-)
attract the oppositely-charged parts of neighboring molecules
this produces stronger attraction than the temporary polarization of London forcestherefore polar molecules are more likely to be
liquid at a temperature where similar non-polar molecules are gases
hydrogen bonding (or ultra-dipole attractions)
H—F, H—O, and H—N bonds are more polar than other similar bondsthese atoms are very small, particularly HF, O, and N are the three most electronegative
elementsthese bonds therefore are 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
hydrogen bonding (or ultra-dipole attractions)
non-polar molecules(lower boiling points)
ultra-polar molecule(much higher boiling point)
hydrogen bonds (between molecules, not within them)
hydrogen bonding (or ultra-dipole attractions)
H H
O H H
O
H H
OH H
O
Beware!!These are not hydrogen bonds. They are normal covalent bonds between hydrogen and oxygen.These are hydrogen bonds. They are between separate molecules (not within a molecule).
metallic bondingstructure
nuclei arranged in a regular grid or matrix
“sea of electrons”—delocalized valence electrons free to move throughout grid
metallic “bond” is stronger than van der Waals attractions but generally is weaker than covalent bond since there are not specific e– pairs forming bonds
resulting propertiesshiny surfaceconductive (electrically and
thermally)strong, malleable, and ductile
alloy = mixture of metals
ionic bonding (salts)structure: orderly 3-D array
(crystal) of alternating + and – charges
made ofcations (metals from left side of periodic
table)anions (non-metals from right side of
periodic table)
propertieshard but brittle (why?)non-conductive when solidconductive when melted or dissolved
why are salts hard but brittle?
1. apply some force
2. layer breaks off and shifts
3. + repels + – repels –
4. shifted layer shatters away from rest of crystal
covalent networksstrong covalent bonds hold together
millions of atoms (or more) in a single strong particle
propertiesvery hard, very strongvery high melting temperaturesusually non-conductive (except graphite)
examplescarbon (two allotropes: diamond, graphite)pure silicon or pure germaniumSiO2 (quartz or sand)other synthetic combinations averaging 4 e–
per atom: SiC (silicon carbide), BN (boron nitride)
summary of propertiesstrongest
weakest
London
dipole
hydrogen
metallic
ionic
network
strength
soft and brittle
strong, malleable, ductile
hard but brittle
extremely hard
van
der
Waa
ls f
orc
es
m.p. & b.p.
low
medium to high
medium to high
very high
conductive?
no
very(delocalized e–)
if melted or dissolved(mobile ions)
usually not
soaps and emulsifiers
some molecules are not strictly polar or non-polar, but have both characteristics within the same molecule
non-polar
region
polar region
this kind of molecule can function as a bridge between molecules that otherwise would repel each other
oil
water
soap or
emulsifier