Polarity of Molecules

Electronegativity The pull an atom has for the electrons it shares with another atom in a bond. Electronegativity is a periodic trend  As atomic radius increases and number of electron shells increases, the nucleus of an atom has less of a pull on its outermost electrons

Periodic Table with Electronegativies increases decreases

Polar Bond A polar covalent bond is when there is a partial separation of charge One atom pulls the electrons closer to itself and has a partial negative charge. The atom that has the electrons farther away has a partial positive charge

ModelElectronegati vity of central atom Electronegati vity of outer atom(s) Subtraction Total Type of Bond 3.0 0Non-polar Covalent N N

Two atoms sharing equally 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

Atoms sharing almost equally Electronegativities: H = 2.1 C = 2.5 The carbon pulls on the electrons slightly more, pulling them slightly towards the carbon Put the difference isn’t enough to create a polar bond This is a non-polar covalent bond CHH H H

Sharing unevenly Electronegativities: H = 2.1 C = 2.5 O = 3.5 The carbon-hydrogen difference isn’t great enough to create partial charges 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 COH H ++ -- COH H

Ionic Bonds Ionic bonds occur when the electronegativies of two atoms are so different that they can’t even share unevenly…one atom just takes them from the other

How to determine bond type Find the electronegativies of the two atoms in the bond Find the absolute value of the difference of their values  If the difference is 0.4 or less, it’s a non-polar covalent bond  If the difference is greater than 0.4 but less than 1.4, it’s a polar covalent bond  If the difference is greater than 1.4, it’s an ionic bond

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 3.5 – 3.0 = 0.5 polar 4.0 – 4.0 = 0.0 non-polar 2.5 – 3.0 = polar

Polar Bonds versus Polar Molecules Not every molecule with a polar bond is polar itself  If the polar bonds cancel out then the molecule is overall non-polar. The polar bonds cancel out. No net dipole The polar bonds do not cancel out. Net dipole

The Importance of VSEPR You must think about a molecule in 3-D (according to VSEPR theory) to determine if it is polar or not! Water drawn this way shows all the polar bonds canceling out. But water drawn in the correct VSEPR structure, bent, shows the polar bonds don’t cancel out! Net dipole H O H O H H

Let’s Practice Example: Is NH 3 a polar molecule?

Let’s Practice 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

Intermolecular Forces

Intra- versus Inter-molecular Forces So far this chapter has been discussing intramolecular forces  Intramolecular forces = forces within the molecule (chemical bonds) Now let’s talk about intermolecular forces  Intermolecular forces = forces between separate molecules

Breaking Intramolecular forces Breaking of intramolecular forces (within the molecule) is a chemical change  2 H 2 + O 2  2 H 2 O  Bonds are broken within the molecules and new bonds are formed to form new molecules

Breaking Intermolecular forces Breaking of intermolecular forces (between separate molecules) is a physical change  Breaking glass is breaking the intermolecular connections between the glass molecules to separate it into multiple pieces.  Boiling water is breaking the intermolecular forces in liquid water to allow the molecules to separate and be individual gas molecules.

London Dispersion Forces 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 Once the electrons have “ganged up” and created a partial separation of charges, the molecule is now temporarily polar. The positive area of one temporarily polar molecule can be attracted to the negative area of another molecule. ++ -- ++ --

Strength of London Dispersion Forces 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 are the weakest of the intermolecular forces because molecules can’t form it all the time.

Strength of London Dispersion Forces 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

Dipole Forces Polar molecules have permanent partial separation of charge. The positive area of one polar molecule can be attracted to the negative area of another molecule. ++ -- ++ --

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 Dipole forces are stronger than London Dispersion Forces

Hydrogen Bonding Hydrogen has 1 proton and 1 electron.  There are no “inner” electrons. It bonds with the only one it has. When that electron is shared unevenly (a polar bond) with another atom, the electron is farther from the hydrogen proton than usual.  This happens when Hydrogen bonds with Nitrogen, Oxygen or Fluorine This creates a very strong dipole (separation of charges) since there’s no other electrons around the hydrogen proton to counter-act the proton’s positive charge.

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 Bond N H H N H H Hydrogen bond

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. It takes energy to break IMF’s Energy is released when new IMF’s are formed

IMF’s and Changes in State Some IMF’s are broken to go from solid  liquid. All the rest are broken to go 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 melting point is much higher This is very important because we need liquid water to exist!

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

Solubility In order from something to be dissolved, the solute and solvent must break the IMF’s they form within itself They must then form new IMF’s with each other

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)

Oil & Water Water has London Dispersion, Dipole 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 Surface tension is the resistance of a liquid to spread out.  This is seen with water on a freshly waxed car The higher the IMF’s in the liquid, the more the molecules “stick” together. 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 form IMF’s and “stick” together. This means that the surface tension of water is lower when soap is added. The lower surface tension allows the water to spread over the dirty dishes.

What did you learn about soap?

Soap Inter-molecular forces Works based on Molecular Geometry Bonding types & Structures Determined by