Section 5.4—Polarity of Molecules

Two atoms sharing equally: Draw N 2 N N Each 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 H 2 S 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 SH H

Sharing unevenly: Draw CH 2 O 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 COH H

Showing Partial Charges There are two ways to show the partial separation of charges  Use 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! COH H ++ -- COH 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 = 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 itself  If 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 dipole H O H O H H

Draw the molecule NH 3 Example: Is NH 3 a polar molecule?

Example: Is NH 3 a polar molecule? NHH H Electronegativities: N = 3.0 H = 2.1 Difference = 0.9 Polar bonds VSEPR shape = Trigonal pyramidal Net dipole Yes, NH 3 is polar

Draw the molecule for dihydrogen monosulfide, H 2 S. Is it polar or non-polar? What shape?

Net dipole Yes, H 2 S is polar Is water polar or non-polar? What shape? Electronegativities: S = 2.5 H = 2.1 Difference =.4 Polar bonds VSEPR shape = bent SH H

Draw the molecule CO 2 Example: Is CO 2 a polar molecule?

Draw the molecule of carbon dioxide, CO 2 Example: Is CO 2 a polar molecule? COO Electronegativities: C = 2.5 O = 3.5 Difference = 1.0 Polar bonds VSEPR shape = linear Dipole cancels No CO 2 is nonpolar

Draw the molecule for carbon tetrachloride, CCl 4. C H H H H Electronegativities: C = 2.5 H = 2.1 Difference =.4 NonPolar bonds VSEPR shape = tetrahedral No Net dipole Yes, CH 4 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 H 2 + O 2  2 H 2 O  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 forces Breaking of intermolecular forces (between separate molecules) is a physical change  Breaking glass & Boiling water are examples  Example: H 2 O (l)  2 H 2 O (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 molecules Polar 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 strongest Rank 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: soft Strong Sheets of carbon rings but weak forces holding the sheets together

NETWORK COVALENT BONDS 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 online.com/Objects/ViewObject.aspx?ID= GCH 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 Forces  London Dispersion Forces  Dipole interactions  Hydrogen bonding The number and strength of the intermolecular forces affect the properties of the substance. Energy is needed to break IMF’s Energy 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 molecule In general small molecules have low melting and boiling points Based on it’s size, water should be a gas under normal conditions However, 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 Viscosity Viscosity 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?

Solubility Solute: 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 occur  An exception to this is if more energy is added somehow (such as heating)

Like Dissolves Like Polar solvents dissolve polar solutes Nonpolar solvents dissolve nonpolar solutes Polar 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 waterwater 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 tail Soap

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 online.com/Objects/ViewObject.aspx?ID= GCH online.com/Objects/ViewObject.aspx?ID= GCH6804 Wisconsin online :intermolecular forces