Molecular Polarity & Intermolecular Forces

How to Predict Whether a Molecule is Polar or Nonpolar

– You already learned how to determine whether a BOND is polar or nonpolar using a Table of Electronegativity.

– You learned to determine and draw the correct molecular geometry of a molecule.

– You can put these two pieces together and determine whether the overall molecule is polar.

– Although you might think that a molecule which has at least one polar bond would automatically be a polar molecule, this would be an incorrect assumption.

– There are lots of nonpolar molecules which contain polar bonds. However, it is true that a molecule with no polar bonds and no lone pairs must be nonpolar.

– How is this possible? Because of the overall 3-D shape!

– If the molecule and the polar bonds are symmetrically arranged, they may cancel out.

– Thus, the molecule would be nonpolar with polar bonds.

– Let’s take a look:CCl4 CH3Cl NH3 PH3

H2S SiH4

– So not only are the individual bond dipoles important, but lone pairs of electrons affect the molecular polarity as well.

– Lone pairs of electrons pull electron density away from the central atom, so they have a dipole as well.

– We can draw an overall dipole for the entire molecule, called the dipole moment.

– This dipole moment is the vector addition of the individual bond dipoles as well as the pull of lone pairs on electron density.

– If you look at a table of molecular geometries, there are some generalizations about molecular polarity which can be made:

• Your Turn: Determine whether CO2, CF4 , CH2Cl2, and H2O are polar or nonpolar

Partial Ionic Character

• We often say that a molecule like HF has “partial” ionic character.

• Or we can say that an “ionic” compound like AlN has “partial” covalent character.

• This can actually be calculated.

– If a molecule is polar, it has a nonzero dipole moment. If a molecule is nonpolar, it has a dipole moment of zero.

– The dipole moment is defined as: μ= Qr• where μ is the Dipole moment in debye units (D),

Q is the charge in coulombs (C), and r is the distance between the charges.

– The higher the dipole moment, the more polar the molecule is or you could say that the bonds have more ionic character.

– We can use this equation to calculate a dipole moment, but as it is actually easy to experimentally measure a molecule’s dipole moment, we more often use the dipole moment to calculate the partial charge or %-ionic character of a molecule.

– HBr has a bond length of 141 pm and a dipole moment of 0.79 D. Given that the full charge of an electron is 1.60x10-19C and 1 D = 3.336x10-30Cm, what is the %-ionic character of the H-Br bond?

Intermolecular Forces

Intermolecular Forces– You already know that CO2 is a gas at room

temperature, while water is a liquid, and sucrose is a solid.

– Why? They are all molecular species.– What holds water molecules together in the

liquid phase at room temperature?– Or what determines what phase or state a

compound will exist in at room temperature?

– Let’s review the 3 states of matter:

Gas Liquid Solid

no fixed volume fixed volume fixed volume

no fixed shape no fixed shape fixed shape, so rigid

very low densities high densities high densities

density varies with T and P

not compressible not compressible

rapid, random motion

fluid, motion little motion

high kinetic motion

some kinetic energy

little kinetic energy

Intermolecular Forces– So what determines the state of a

compound at room temp?– The strength of the attractions

between separate molecules: the stronger the attraction, the more likel the compound will be a solid, the weaker the attractions, the more likely it will be a gas.

Intermolecular Forces– These attractive forces “glue”

solids or liquids together.– What are these forces?

Intermolecular Forces

– The forces that hold individual molecules together in the solid or liquid phases are called intermolecular forces.

– They are also called van der Waal forces (although there is also a repulsive van der Waal force)

– They are responsible for many physical properties, including MPt and BPt

Intermolecular Forces

– There are several types DEPENDING on the MOLECULAR POLARITY!

– Here are the 3 Main Types• Ion-Dipole (Ch 12)•Dipole-Dipole (and H-Bonding)•London Dispersion

Dipole Dipole Forces• Attractive force between POLAR molecules (have a

dipole moment)

• Electrostatic attraction of partial positive end of

molecule to partial negative end of another molecule

• In liquid or solid, molecules align themselves so are

attracted to several other molecules

• These dipole-dipole forces are much weaker than a

real covalent bond, about 3-4kJ/mol.

• So they may be broken with a low amount of energy,

so the solid melts, and the liquid evaporates!

431 kJ/mol 16 kJ/mol

Strength of Dipole-Dipole Forces

• There are 2 main factors in the strength

of dipole-dipole forces:

– Distance between molecules (the

closer they are, the stronger the

dipole-dipole forces)

– Molecular Polarity (the more polar, the

stronger the dipole-dipole force)

Hydrogen Bonding: A special case of dipole-dipole forces

• Some molecules have such a strong

dipole-dipole force, that this extra-strong

dipole force was given its own name:

Hydrogen Bonding

• It is NOT really a bond, it is just an extra

strong dipole-dipole force.

• It occurs under certain circumstances.

H-Bonding

• Occurs when a molecule is small and very

polar

• Occurs when have a N-H, O-H, or H-F bond

• The molecules have a very large dipole

moment, and they can get very close to one

another due to the small size of H, N, O, F

• H-Bonding is responsible for the very high

melting and boiling point of water

• H-Bonding is responsible for shape of

proteins and DNA

London Dispersion Forces

• But CO2 is a nonpolar molecule,

so how can it ever be a solid?

• There is another force, called

the London Dispersion Force.

• The London Force acts upon

ALL molecules, polar or not!

Loondon Dispersion Forces

• But London Forces are the ONLY

intermolecular force that operates

in nonpolar molecules.

• So how does it work when there

are no dipoles to create an

attraction?

• Instantaneous, induced dipoles!

London Dispersion Forces

• London Forces seem weak, but

they can be very important,

and many nonpolar molecules

are solids or liquids.

• Waxes, oil, gasoline are

examples.

London Dispersion Forces

• There are 2 factors in the

strength of London Force:

– Size and mass of molecule or

atom

– Shape of molecule

Effect of Size/Mass on London Dispersion Forces• The larger the size or mass of a molecule,

the more electrons it has.

• These electrons are also generally

further from the nucleus.

• So the electrons can be distorted or

pushed and pulled from one end of the

molecule more readily, creating a larger

temporary dipole.

Effect of Shape on London Dispersion Forces

• The closer molecules can align, the

stronger the London Force, and the

higher the melting and boiling points.

• Bulky molecules with “branching” can’t

get as close, and so have lower London

Forces and will have lower MPt and BPt

than an “unbranched” molecule of the

same mass.

Intermolecular Force Summary

• London Forces act upon all molecules, but are most

important for larger, heavier molecules

• Dipole-dipole forces are found in polar molecules and

depend on the size and polarity of the molecules

• H-bonding occurs when very small and very electronegative

N, O, or F atoms are present along with N-H, O-H, or H-F

bonds

Intermolecular Force Summary

• Although they are weak compared to covalent

bonds, these intermolecular forces greatly affect

the boiling and melting points of a compound

• The stronger the forces, the higher the melting and

boiling points

• The stronger the forces, the more likely a

compound will be a solid or liquid

• They also mean more deviations of a gas from the

Ideal Gas Law!