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

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unit 7

Unit 7

Lewis diagrams

molecular geometry

bond and molecular polarity


lewis dot diagrams
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)
structural formulas
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
multiple bonds
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
multiple bonds1
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
multiple bonds computational approach
multiple bondscomputational 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
multiple bonds computational approach1
multiple bondscomputational approach





have: 6 + 6 = 12

have: 4 + 6 = 10

need: 8 + 8 = 16

need: 8 + 8 = 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”)
general hints for lewis structures
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
exceptions to octet rule
exceptions to octet “rule”
  • most atoms have octets (8 valence electrons) when in molecules, but there are exceptions
molecular shapes vsepr model
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-)
vsepr model central atom with
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–




trigonal planar



e.g. BeF2

e.g. BF3

e.g. CF4

e.g. XeF6

e.g. SF5

electron geometry vs molecular shape
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
linear electron geometry 2 electron sets
linear electron geometry2 electron sets
  • in addition, any diatomic molecule must be linear (since any two points lie on a line)
bond polarity
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
bond polarity1
bond polarity
  • the electronegativity difference between the two atoms determines how polar a bond is

0.0 – 0.4

0.5 – 1.7

> 1.7




bond polarity2
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
molecular polarity
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


overall molecule


molecular polarity1
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
molecular polarity3
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
imfa i nter m olecular f orces of a ttraction
IMFA: intermolecular forces of attraction

“mortar”— holds the separate pieces together

(the IMFA)

“bricks”— individual atoms, ions, or molecules of a solid

types of imfa
types of IMFA


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)

dipole-dipole attraction

polar molecules

van der Waals forces

London forces

non-polar molecules


consequences of imfas
consequences of IMFAs
  • melting points and boiling points rise with
    • strength of IMFA
    • increasing molar mass
  • substances generally mix best with other substances having the same or similar IMFAs
    • ”like dissolves like”
    • non-polar mixes well with non-polar
    • polar 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 points
predicting melting points, boiling points
  • stronger 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 categories
    • then sort those in each category from lightest to heaviest
same imfa sort by molar mass
same IMFA: sort by molar mass



  • ex: halogen family
  • all are non-polar (London force)
  • lowest to highest m.p. and b.p. matches lightest to heaviest



  • thus at room temperature:
    • F2 (g)
    • Cℓ2 (g)
    • Br2 (ℓ)
    • I2 (s)

































same mass sort by imfa type
same mass: sort by IMFA type

ethylene glycol

(can form twice as many H-bonds)

  • ex: organic molecules
  • all are ~60 g/mol
  • different types of IMFA




1-propanol (ultra-polar = H-bonds)


acetone (more polar)



methyl ethyl ether (slightly polar)



butane (non-polar)


  • the stronger the IMFA, the higher the boiling point



isomers and an isobar
isomers (and an isobar)

butane and 2-methylpropane

glycerol and 1-propanol

n- and neo pentane

1-propanol and 2-propanol

1-propanol and methyl ethyl ketone

details about each imfa
details about each IMFA


covalent network

ionic bond

metallic bond

hydrogen bond

dipole-dipole attraction

London forces


london or dispersion forces
London (or dispersion) forces
  • non-polar molecules (or single atoms) normally have no distinct + or – poles
  • how can they attract each other enough to condense or freeze?
  • they form temporary dipoles
  • electron clouds are slightly distorted by neighboring molecules
    • sort of like water sloshing in a shallow pan
london dispersion forces in action
London dispersion forces in action

1. temporary polarization due to any random little disturbance



2. induced polarization caused by neighboring molecule

3. induced polarization spreads

4. induced polarization reverses

non-polar molecules, initially with uniform charge distribution

dipole dipole attractions
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
hydrogen bonding or ultra dipole attractions
hydrogen bonding (or ultra-dipole attractions)
  • H—F, H—O, and H—N bonds are more polar than other similar bonds
    • these atoms are very small, particularly H
    • F, O, and N are the three most electronegative elements
    • these 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 attractions1
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)

hydrogen bonding or ultra dipole attractions2
hydrogen bonding (or ultra-dipole attractions)


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 bonding
metallic bonding
  • structure
    • 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 properties
    • shiny surface
    • conductive (electrically and thermally)
    • strong, malleable, and ductile
  • alloy = mixture of metals
ionic bonding salts
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
    • hard but brittle (why?)
    • non-conductive when solid
    • conductive when melted or dissolved
why are salts hard but brittle
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 –

covalent networks
covalent networks
  • strong covalent bonds hold together millions of atoms (or more) in a single strong particle
  • properties
    • very hard, very strong
    • very high melting temperatures
    • usually non-conductive (except graphite)
  • 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)
m.p. = 3550°C



“bucky ball”

m.p. = ~1600°C

summary of properties
summary of properties



m.p. & b.p.



extremely hard

very high

usually not


hard but brittle

medium to high

if melted or dissolved

(mobile ions)

strong, malleable, ductile

medium to high


(delocalized e–)




soft and brittle



van der Waals forces



soaps and emulsifiers
soaps and emulsifiers

some molecules are not strictly polar or non-polar, but have both characteristics within the same molecule


polar region

soap or emulsifier


non-polar region

this kind of molecule can function as a bridge between molecules that otherwise would repel each other

soaps and emulsifiers1
soaps and emulsifiers

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