section 5.4—polarity of molecules. two atoms sharing equally: draw n 2 n n each nitrogen atom has...
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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
Wisconsin online :intermolecular forces
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
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