Section 5.4—Polarity of Molecules

Two atoms sharing equally: Draw N2

N NEach nitrogen atom has an electronegativity of 3.0

They pull evenly on the shared electrons

The electrons are not closer to one or the other of the atoms

This is a non-polar covalent bond.

All compounds that contain Non-polar bonds are NON-POLAR molecules.

Atoms sharing unequally: Draw H2S

Electronegativities: H = 2.1 sulfur = 2.5

The sulfur pulls on the electrons slightly more, pulling them slightly towards the sulfur.

This is a polar covalent bond

S H

H

Sharing unevenly: Draw CH2O

Electronegativities: H = 2.1 C = 2.5 O = 3.5

The carbon-hydrogen difference isn’t great enough to create partial charges : It’s actually a NON POLAR bond: **Exception to the rule

But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond

This is a polar covalent bond

C OH

H

Showing Partial ChargesThere are two ways to show the partial

separation of chargesUse of “” for “partial” Use of an arrow pointing towards the partial

negative atom with a “plus” tail at the partial positive atom. These are referred to as DIPOLES which are: separation of opposite charge!

C OH

H

+ -C OH

H

Let’s Practice

Example:If the bond

is polar, draw the polarity arrow

C – H

O—Cl

F—F

C—Cl

Let’s Practice

Example:If the bond

is polar, draw the polarity arrow

C – H

O—Cl

F—F

C—Cl

2.5 – 2.1 = 0.4 non-polar*exception

3.5 – 3.0 = 0.5 polar

4.0 – 4.0 = 0.0 non-polar

2.5 – 3.0 = - 0.5 polar

How do Dipoles Cancel?

Dipoles must move in equal but opposite directions in order for the forces to cancel The molecule is classified as NONPOLAR.

Polar Bonds versus Polar Molecules

Not every molecule with a polar bond is polar itselfIf the polar bonds form Dipoles that cancel out

then the molecule is overall non-polar.

The dipoles cancel out.No net dipole

The dipoles do not cancel out.

Net dipole

This one is hard to tell!

The Importance of VSEPR in Predicting Polarity.Shape is important. All molecules must be

drawn in the correct shape to see the proper canceling of dipoles to determine if its polar or nonpolar.

Water drawn this way shows all the dipoles canceling out.

But water drawn in the correct VSEPR structure, bent, shows the dipoles don’t cancel out!

Net dipoleH O H

O H H

Draw the molecule NH3

Example:Is NH3 a

polar molecule?

Example:Is NH3 a

polar molecule?

NH H

HElectronegativities:N = 3.0H = 2.1Difference = 0.9 Polar bonds

VSEPR shape = Trigonal pyramidal

Net dipole

Yes, NH3 is polar

Draw the molecule for dihydrogen monosulfide, H2S.

Is it polar or non-polar? What shape?

Net dipoleYes, H2S is polar

Is water polar or non-polar? What shape? Electronegativities:

S = 2.5H = 2.1Difference = .4 Polar bonds

VSEPR shape = bent

S H

H

Draw the molecule CO2

Example:Is CO2 a

polar molecule?

Draw the molecule of carbon dioxide, CO2

Example:Is CO2 a

polar molecule?

CO O

Electronegativities:C = 2.5O = 3.5Difference = 1.0 Polar bonds

VSEPR shape = linear

Dipole cancels

No CO2 is nonpolar

Draw the molecule for carbon tetrachloride, CCl4.

CH

H

H

H

Electronegativities:C = 2.5H = 2.1Difference = .4 NonPolar bonds

VSEPR shape = tetrahedral

No Net dipoleYes, CH4 is nonpolar

Let’s make this Simple:Nonpolar bonds= Nonpolar molecule.

Polar bonds with a lone pair on the central atom = most likely a polar molecule

Polar bonds & no lone pair on central atom & all terminal atoms are the same= nonpolar molecule.

If terminal atoms are different, its polar.

Section 5.5—Intermolecular Forces

Intramolecular Forces- versus Inter-molecular Forces

So far this chapter has been discussing “Intramolecular Forces”Intramolecular forces = forces

within the molecule AKA:chemical bonds

Breaking Intramolecular forces

Breaking of intramolecular forces (within the molecule) is a chemical change Example: 2 H2 + O2 2 H2O

Bonds are broken within the molecules and new bonds are formed to form new molecules

Requires a larger amount of energy to break than an intermolecular force

Inter-molecular Forces

Intermolecular forces = forces between separate molecules

Breaking Intermolecular forcesBreaking of intermolecular forces

(between separate molecules) is a physical changeBreaking glass & Boiling water are examplesExample: H2O(l) 2 H2O(g)

Does not require as much energy to break compared to an intramolecular force

London Dispersion Forces (LDF): are the primary force between nonpolar molecules but are found in all molecules!

All molecules have electrons.

Electrons move around the nuclei. They could momentarily all “gang up” on one side

This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another.

+ Positively charged nucleus - Negatively charged electron

+-

-

-

-

Electrons are fairly evenly dispersed.

+--

- -As electrons move, they “gang up” on one side.

+

-

London Dispersion Forces (LDF)

Once the electrons have “ganged up” and created a temporary dipole, the molecule is now temporarily polar.

The positive area of one temporarily polar molecule can be attracted to the negative area of another molecule.

+ - + -

London Dispersion Forces (LDF)

Strength of London Dispersion Forces (LDF)

Electrons can gang-up and cause a non-polar molecule to be temporarily polar

The electrons will move again, returning the molecule back to non-polar

The polarity was temporary, therefore the molecule cannot always form LDF.

London Dispersion Forces:the weakest of the intermolecular forces because molecules can’t form it all the time, only temporarily

Strength of London Dispersion Forces (LDF)

Larger molecules have more electrons

The more electrons that gang-up, the larger the partial negative charge.

The larger the molecule, the stronger the London Dispersion Forces

Larger molecules have stronger London Dispersion Forces than smaller molecules.

All molecules have electrons…all molecules can have London Dispersion Forces

London Forces explain why Chlorine is a gas, Bromine is a liquid and Iodine is a solid!Chlorine Gas = 34 e- Bromine Liquid = 70 e-

Iodine Solid = 106 e-

GREATER # ELECTRONS, STRONGER FORCES

Dipole- Dipole Forces: primary force between polar moleculesPolar molecules have permanent

permanent dipoles.The positive area of one polar molecule

can be attracted to the negative area of another molecule.

The partial positive & negative poles are shown as + and -

Strength of Dipole Forces

Polar molecules always have a partial separation of charge.

Polar molecules always have the ability to form attractions with opposite charges

In general, Dipole forces are stronger than London Dispersion Forces

+ - + -

Dipole-Dipole Forces

Hydrogen Bonding

A special dipole force between a hydrogen atom of 1 molecule and a F, O, or N of another molecule.

(ET fon home)A very strong dipole forms since F,

O, and N are all very small, highly electronegative atoms.

Hydrogen Bond

N

H H

N

H H

Hydrogen bond

Strength of Hydrogen Bond

Hydrogen has no inner electrons to counter-act the proton’s charge

It’s an extreme example of polar bonding with the hydrogen having a large positive charge.

This very positively-charged hydrogen is highly attracted to a lone pair of electrons on another atom.

This is the strongest of all the intermolecular forces.

Hydrogen Bonds

• The ladder rungs in a DNA molecule are hydrogen bonds between the base pairs, (AT and GC).

Rank the forces of attraction in order of weakest to strongestRank the Intramolecular Forces: Ionic, Covalent, and Metallic

Rank the Intermolecular Forces: Dipole, London Dispersion, Hydrogen bonding

Rank ALL the Forces:

Covalent< Metallic < Ionic

London Dispersion forces< Dipole- Dipole forces< Hydrogen bonding

London Dispersion forces< Dipole- Dipole forces< Hydrogen bonding <Covalent< Metallic< Ionic

Bond Energy of Bonding Types

Carbon Allotropes: Diamond vs Graphite

Diamond: Hard Tetrahedral-Special : Network Covalent Bonds

Graphite: softStrong Sheets of carbon rings but weak forces holding the sheets together

NETWORK COVALENT BONDS

http://www.youtube.com/watch?v=fuinLNKkknI

•special covalent compounds•compounds that contain only carbon (diamond, graphite) or silicon compounds (silicon dioxide- quartz)•super strong bonding•super high melting points

Tutorial: must be in Mozilla

http://www.wisc-online.com/Objects/ViewObject.aspx?ID=GCH6804

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Section 5.6—Intermolecular Forces & Properties

IMF’s and Properties

IMF’s are Intermolecular ForcesLondon Dispersion ForcesDipole interactionsHydrogen bonding

The number and strength of the intermolecular forces affect the properties of the substance.

Energy is needed to break IMF’sEnergy is released when new IMF’s are

formed

IMF’s and Changes in State

IMF’s are broken to go from solid liquid. and from liquid gas.

Breaking IMF’s requires energy.

The stronger the IMF’s, the more energy is required to melt, evaporate or boil.

The stronger the IMF’s are, the higher the melting and boiling point

Water

Water is a very small moleculeIn general small molecules have low melting and

boiling pointsBased on it’s size, water should be a gas under

normal conditionsHowever, because water is polar and can form

dipole interactions and hydrogen bonding, it’s boiling point is much higher

This is very important because we need liquid water to exist!

Boiling Point of Polar Molecules

IMF’s and ViscosityViscosity is the resistance to flow

Molasses is much more viscous than water

Larger molecules and molecules with high IMF’s become inter-twined and “stick” together more

The more the molecules “stick” together, the higher the viscosity

An increase in temperature will help break the IMF’s and make a substance less viscous

What is More Viscous? Molasses or Water?

SolubilitySolute: the

substance that

is dissolved

Solvent: the

substance that

is doing the

dissolving

Solubility

- +

- +

- + - +- +

Solvent, water (polar)

+

-

- + Solute, sugar (polar)

Water particles break some intermolecular forces with other water molecules (to allow them to spread out) and begin to form new ones with the sugar molecules.

Solubility

Solvent, water (polar)

+

-

- + Solute, sugar (polar)

As new IMF’s are formed, the solvent “carries off” the solute—this is “dissolving”

- +

- +

- +- + - +

Solubility

If the energy needed to break old IMF’s is much greater than the energy released when the new ones are formed, the process won’t occurAn exception to this is if more energy is added

somehow (such as heating)

Like Dissolves Like

Polar solvents dissolve polar solutesNonpolar solvents dissolve nonpolar

solutesPolar solvents can also dissolve ionic

compounds because of the charged ends of both

Oil & Water

Water has London Dispersion, Dipole forces and hydrogen bonding. That takes a lot of energy to break

Water can only form London Dispersion with the oil. That doesn’t release much energy

Much more energy is required to break apart the water than is released when water and oil combine.

Water is polar and can hydrogen bond, Oil is non-polar.

Therefore, oil and water don’t mix!

Surface Tension of Water

metal paper clip on water water forms “beads”

Surface Tension

Surface tension is the resistance of a liquid to spread out.This is seen with water on a freshly waxed car

Due to higher IMF’s in the liquid, the more the molecules “stick” together, the less they want to spread out.

The higher the IMF’s, the higher the surface tension.

Soap & Water

Soap has a polar head with a non-polar tail

The polar portion can interact with water (polar) and the non-polar portion can interact with the dirt and grease (non-polar).

Polar head

Non-polar tailSoap

Soap & Water

The soap surrounds the “dirt” and the outside of the this Micelle can interact with the water.

The water now doesn’t “see” the non-polar dirt.

Dirt

Soap & Surface Tension

The soap disturbs the water molecules’ ability to “stick” together to form IMF’s

Soap lowers the surface tension of water This allows the water to spread over the

dirty dishes.

Tutorial: must be in Mozilla

http://www.wisc-online.com/Objects/ViewObject.aspx?ID=GCH6804

Wisconsin online :intermolecular forces