April 20, 2010

Solvents and solutes are either Polar or Non Polar

Polar substances have an unequal charge distribution (asymmetrical)

H2O = polar:

CH4 = Nonpolar

LIKE DISSOLVES LIKE!

Polar solutes dissolve in polar substances and non-polar solutes dissolved in non-polar substances

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INTERMOLECULAR BONDS

- bonds between molecules
- 3 types

1.) London Dispersion Force (L.D.F)

- Results from temporary electron dipoles
- Weakest intermolecular force
- Increases as the $ e- increases
- Occurs in any compound that has e- (ie: everything)

2.) Dipole Dipole

- Results from a permanent dipole in molecules
Polar molecules experience this force
Polarity depends how much elements want e- (electronegativity)
- Electronegativity increases to the right and up
- The strength of a dipole- dipole bond depends on the difference in electronegativity between the two atoms
- Only polar molecules experience this

Substance Boiling Point # of e-
N2 -196 degrees C 14
O2 -183 degrees C 16
NO - 152 degrees C 15
ICl 97 degrees C 70
Br2 59 degrees C 70

More electrons = higher the boiling point

3.) Hydrogen Bonds (H-Bonds)
-This is a special type of dipole- dipole bond between H and O, F, or N
- Any molecule that: H-F, H-O or H-N

Identify the substances with H-Bonds:
1) CH4
2) CH3OH
3) H2S
4) CH3-NH2
5) HCl
6) CH2-OH-OH2
/ / /
OH OH OH

Answer: Number 2, 4, and 6

Compare the boiling points of:
- Ethanol (C2H5OH)
- Ehtane (C2H6)
- Methanol (CH3OH)
- Methane (CH4)

> London Forces are the weakest intermolecular force and hydrogen bonds are the strongest.

> More electrons, the higher the boiling point.

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IONS IN SOLUTIONS
- The formation of a solution depends on the ability of the solute to dissolve in the solvent
- Solvation is the interaction between solutes and solvents
- Ionic solids (salts) are cyrstals made up of ions
- Molecular solids are crystals made up of neutral molecules
- Dissolving ionic solutions produces ions in a process called disssociation (remember?)

Ionizaiton is the break up of a neutral molecule into charged particles
Examples:
1) FeCl3 (s) ----->Fe 3+ (aq) + 3 Cl -1 (aq)
2) Ag2O (s) -----> 2 Ag + (aq) = O 2- (aq)

3) Na3PO4(s) -----> 3 Na + (aq) + PO4 (aq)

4) (NH4)2SO4 (s) -----> 2NH4 + (aq) + SO4 2- (aq)

Determining concentrations is relatively easy.

Examples:

What is the [Cl-] in a solution of 0.50 M AgCl3?

AgCl3 -----> Ag + (aq) + 3Cl- (aq)
Cl = 3x as many moles = (0.5 M) x 3 = 1.5 M

What is the [NO3-] in a solution of 0.82 M Fe(NO3)2?
Fe(NO3)2 -----> Fe 2+ (aq) + 2 NO 3- (aq)
(0.82 M) x 2 = 1.64 M

What is the [Cr2O7 2-] and [K+] when 3.5 g of K2Cr2O7 dissolved in 40 mL of water?
K2CrO7 -----> CrO7 2- (aq) + 2K + (aq)
3.5 g x 1 mol/294.2 g = 0.0