Unit 7

Lewis diagramsmolecular geometry

bond and molecular polarityIMFAs

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

tetrahedral electron geometry

tetrahedral electron geometry

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)

triangular planar and linearelectron geometry

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 polarity

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)

IMFA: intermolecular forces of attraction

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)

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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

dipole-dipole attractionsδ+ δ-

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)

m.p. = ~1600°C

m.p. = 3550°C

C60buckminsterfullerine

“bucky ball”

summary of propertiesstrongest

weakest

London

dipole

hydrogen

metallic

ionic

network

strength

soft and brittle

strong, malleable, ductile

hard but brittle

extremely hard

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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

soaps and emulsifiers

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