TOPIC: Intermolecular Forces How do particle diagrams of liquids & solids compare to those of gases?

Describe relative positions and motions of particles in each of 3 phases SOLID LIQUIDGAS

Why do some substances exist as gases, some as liquids, and some as solids at room temp?

Part of answer has to do with forces between separate molecules (called intermolecular forces)

Intermolecular forces between molecules. They are weaker. Intramolecular forces are between individual atoms (we will learn this later) Intramolecular forces Intermolecular forces

Intermolecular Forces-IMF Inter means “between” or “among” Intermolecular forces = forces between neighbouring compounds

All molecules have Dispersion forces (the regents calls these Van der Waals) 2 other types of forces (IMF): 1. Dipole-Dipole forces 2. Hydrogen bonds -if one of these are present, they are more important.

Most atoms don’t have a charge, unless they are ions, so we often refer to them as having partial charges and write them like this… This separation of Charge is responsible For the forces Between the molecules

1. Dispersion Forces (van der waals) 1. Dispersion Forces (van der waals): ● weakest IMF ● occur between nonpolar molecules ● Click here for animation (slide 4 of 13) Click here for animation NonnoNonpolar means no poles (+/-) Can’t tell one end of molecule from other end electrons are evenly distributed

instantaneous and momentary fluctuate results from motion of electrons induce if charge cloud not symmetrical will induce asymmetry in neighbor’s charge cloud!

4 categories of Nonpolar Molecules (you need to memorize) Noble Gas molecules:  He, Ne, Ar, Kr, Xe, Rn diatomics if both atoms are same: (7)  H 2, N 2, O 2, Cl 2, F 2, I 2, Br 2 Pure Hydrocarbons (C x H y ) :  CH 4, C 2 H 6, C 3 H 8 small symmetrical molecule  CO 2, CF 4, CCl 4

Dispersion Forces and Size Dispersion forces ↑ with molecule size larger the electron cloud, the greater the fluctuations in charge can be  Rn > Xe > Kr > Ar > Ne > He  I 2 > Br 2 > Cl 2 > F 2  C 8 H 18 > C 5 H 12 > C 3 H 8 > CH 4

Boiling point of N 2 is 77 K (-196˚C) IMF are very weak dispersion forces

2. Dipole-dipole forces 2. Dipole-dipole forces: intermediate IMF occur between polar molecules (they have a partial charge at each pole – one is typically much larger than the other) Click here for animation (slide 3 of 13) Click here for animation

Dipole-dipole Forces & Polar Molecules Molecule shows permanent separation of charge; has poles: one end partly (-) & one end partly (+)

Polar means molecule has poles: (+) & (-) geometry and electron distribution are not symmetrical Polar Molecules

3. Hydrogen bonds 3. Hydrogen bonds: strongest IMF occur between molecules that have: H-F H-O or H-N bonds ONLY

Hydrogen Bonding H-O N-H Occurs between molecules with H-F, H-O, or H-N bonds

Hydrogen Bonding Hydrogen bonding is extreme case of dipole-dipole bonding F, O, and N are all small and electronegative  strong electrons attraction  H has only 1 electron, so if being pulled away H proton is almost “naked” H end is always positive & F, O, or N end is always negative

Strength of Hydrogen Bonding Fluorine most electronegative element, so  H-F bonds are most polar and exhibit strongest hydrogen bonding  H-F > H-O > H-N (H-bonding…sound like FON to me!!!)

Hydrogen bonding: strongest IMF influences physical props a great deal H-F > H-O > H-N

Strongest Intermolecular Force Hydrogen Bonding Dipole-Dipole Dispersion

Indicate type of IMF for each molecule: NH 3 Ar N 2 HCl HF Ne O 2 HBr CH 3 NH 2 Hydrogen bonding Dispersion forces Dipole-dipole forces Hydrogen bonding Dispersion Dipole-dipole Hydrogen bonding

O H H O H H H-Bonding = strongest IMF much harder to “pull” molecules apart

C Dispersion Forces= weakest IMF much easier to “pull” molecules apart C H H H H H H H H

IMF vs. Physical Properties If IMF  then:  Boiling point   Melting point   Heat of Fusion   Heat of Vaporization  while:  Evaporation Rate  Change from solid to liquid w/o changing temp Change from liquid to gas w/o changing temp Rate at which conc. will go from liquid to gas

Why do some substances exist as gases, some as liquids, and some as solids at room temp? #1 reason = IMF #2 reason = temperature (avg. KE)

IMF vs. Temp IMF more important as temp is lowered  Low temperature – low evaporation rate  High temperature – high evaporation rate

Intermolecular forces determine phase “Competition” between strength of IMF & KE determines phase

 If IMF are strong, substance will be solid or liquid at room temp  Particles want to clump together  If IMF are weak, substance will be gas at room temp  Particles free to spread apart

It’s a balancing act! Intermolecular Forces weak Kinetic Energy High (fast) [this substance = a gas at room temperature]

Intermolecular Forces strong Kinetic Energy Low (slow) [this substance = a condensed phase (solid/liquid)]

REMEMBER… Temp = average KE If we change T we change KE Increase KE will help “pull” molecules apart (overcome IMF)