Chapter 12: Intermolecular Forces

Why do gases differ from liquids & solids?

Gases are tiny particles far apart with no attraction for each other (ideally) and they are moving rapidly in random directions
Gases obey a set of laws: Ideal Gas Laws

But liquids & solids don't have a set of "laws" because....
	
Liquids:  - condensed from gases
		   - not compressible therefore not much space
		   between molecules
		   -moving randomly but more slowly, are attracted
		   to each other!
Solids:   - ordered fixed in place particles
	         - close together
	         - strong forces of attraction!

Interparticle Forces
Interparticle forces of attraction between "particles" that affect physical state & physical behavior
Note:  particles can be ions, atoms or molecules
True chemical bonding forces are intra-particle forces: within a chemical substance
	They are atom to atom or ion to ion, not molecule to molecule:  ionic bonding, metallic bonding, covalent bonding, and (new to us) network covalent bonding
Interparticle forces are between particles, not within them
Special group of interparticle forces between molecules and atoms:
	- called intermolecular forces
	- affect behavior of covalent compounds

See table in your packet of handouts, also at end of this doc.

Table 12.2 Comparison of Bonding and Nonbonding (Intermolecular) Forces 
My version:
Force/Bonding*	Basis of Attraction	Energy, kU/mol	Example	Properties	
Ionic*	Cation – Anion	400 – 4000	NaCl	High MP, brittle solid	
Covalent*	Shared pair of e-s	150-1100	H-H	Low MP, soft when solid	
Metallic*	Delocalezed e-s	75 – 1000	Fe	Range of MPs, malleable	
Network Covalent*	Extended cov bonding	500+	Diamond	Very high MP, very hard solid	
Ion-Dipole	Opp charges	40 – 600	Na+--H2O	Salts dissolving in polar solvents	
Hydrogen-bonding**	H from one molecule to N,O,F of other molecule	10 – 40	See figure	Intermolecular Force	
Dipole-Dipole**	Partial opp charges	5 – 25	HCl – HCl	Intermolecular Force	
Ion – Induced Dipole	Opp partial charge	3 – 15	Fe2+---O2		
Dipole – Induced Dip	Partial charges	2 – 10	HCl –Cl2	Intermolecular Force	
London Forces**	Polarizable e- clouds	0.05- 40	F2 – F2	Intermolecular Force	
*Also the names for the four types of solid.
**True intermolecular forces in pure substances.
Network Covalent Bonding
Make special note of the network covalent solids:  
	- covalent bonds are extended throughout the crystal solid
	- diamond and SiC, also SiO2
Diamond is a face-centered cubic (fcc) crystal with 4 C's in "holes" in unit cell, which we see later in chapter
Graphite has unique structure (see diagram in the powerpoint file)

Ion-Dipole Attraction
Interparticle forces involved in dissolving ionic solids in water or other polar solvents:  
Ion-dipole attraction has to overcome ion-ion attraction in a solid’s crystal lattice
This is why some compounds are soluble and some are not
Called Hydration of an ion:  typically endothermic - takes energy to pull ions apart
Some energy is gained back as Heat of Hydration - polar water molecules orient themselves and surround the individual ions
More detail coming soon in Chapter 13

The Major InterMolecular (IM) Forces
The three major IM Forces are very weak to moderate forces of attraction between molecules (or atoms)
		-London (dispersion) forces (LDF)
		- Dipole-dipole attractions
		- Special case of enhanced dipole attraction called Hydrogen-bonding
		- Dipole-induced dipole attraction, between different types of molecules or atoms

London Dispersion Forces (LDF)
"Instantaneous dipole" causes neighboring electron clouds to also move to one side, inducing a dipole in them  
Leads to a small force of attraction between slight positive and slight negative ends of two different particles – these attractions are called London dispersion forces
Strength of LDF depends on:  size of electron cloud in atom or number of atoms in a molecule 
Polarizability increases down a group
Polarizability decreases left to right across a period
Look at the elemental halogens:  fluorine and chlorine are gases; bromine is liquid and iodine is solid because of size of e- cloud around each molecule

Figure 12.14 Dispersion forces among nonpolar molecules
Figure 12.15 Molar mass and boiling point
Figure 12.16  Molecular shape and boiling point
Figure 12.11 Polar molecules and dipole-dipole forces

Dipole-dipole Attraction
Dipoles involve polar molecules which are attracted to each other because of the slight positive and slight negative "poles" to the molecules
Compare molecules of F2, HF, HCl, HBr, HI
Boiling points:  F2<HCl<HBr<HI<HF

Data for IM Forces
			F2	HF	HCl	HBr	HI
Tot # e-s	18	10	18	36	54
MM		38	20	36.5	81	128
DEN		0	1.8	1.0	0.8	0.5
Dip Mom	0	1.4	1.1	0.8	0.4
% Disp	100	Low   81.4	94.5	99.5
% Dipole	0	High	18.6	5.5	0.5
BP, K	85 	291	188	206	238
DHvap 	6.86	High	16	18	20

Enhanced Dipole-Dipole or  Hydrogen-Bonding?
Dipole forces are decreasing down the hydrohalogen group because EN is decreasing
WHY is HF so very different in boiling point?
	HF represents a special case of dipole-dipole attraction called Hydrogen-bonding 
Occurs when H is bonded to a highly EN atom that is also very small:  H to F, O or N
Size of EN atom is important also, allows H to get very close, as seen with radius in pm below:
  N-70      O-73     F-72     Cl-100 (no H-bond)

Sample problem 12.3

Figure 12.13 Hydrogen bonding and boiling point

Dipole-Induced Dipole: between two different compounds’ particles
Dipole-induced dipole forces account for limited solubility of oxygen in water
Ability to do this = function of  polarizability of molecule
Compare H2 to I2:  bigger molecule polarizes, soluble in water, which is demonstrated by the much greater solubility of I2 in water

Sample problem 12.4

Practice with Intermolecular Forces:  explain the forces behind this data:

1.  Butane (CH3CH2CH2CH3) melts at -138oC and boils at 0.5oC, while acetone (CH3C=OCH3) melts at -95oc and boils at +56oC, yet both weigh 58 g/mol.  Draw Lewis structures and explain the differences in MPs and BPs.

2.   Guess BP order for CCl4, N2, Cl2, ClNO (chlorine-nitrogen-oxygen). 

Figure 12.17 (modified – see powerpoint and redraw it here, including Metallic Bonding which I added to this figure from the text)

Practice
See handout (on next page of these notes)
Also chapter problems: 2, 29, 31, 33, 37, 39, 43, 45, 
and 4th ed. #119:  What forces are overcome when the following events occur: (a) NaCl dissolves in water, (b) krypton boils, (c) water boils, (d) CO2 sublimes? 

QUESTIONS TO ASK IN PREDICTING THE KINDS OF
 INTERPARTICLE FORCES THAT WILL BE PRESENT IN A SOLID OR A LIQUID

Start at the top with the first question, “Is it metallic?”.  When you can answer yes, you are done.  If the answer is no, keep going down the list.  The “it” refers to whatever substance you are working with.

Question			If yes, this force is present*	MP & BP		Examples (MP, BP)

Is it metallic?			Metallic bonding		High			Iron (1555, 3000)
    (ONLY metal present)

Is it ionic?			Ionic bonding			High			NaCl(804, not defined)
    (cation & anion present)

Is it network 			Network covalent bonding	High			Diamond (3550, not def)
covalent compound?									SiC, SiO2
    (very few of these)

Is the substance molecular?
    (covalent bonds present)

In the molecule, is H		Hydrogen bonding		Medium		Water (0.100)
attached by a covalent
bond to F, O or N?

Does the molecule		Dipole-dipole attraction	Low			HCl (-114, -85)
have a dipole moment?

Is it a molecule		Only London Forces		Very low		Hydrogen (-257, -253)
with no dipole moment?								Iodine (114, 183)

Does the substance consist	Only London Forces		Extremely low		Neon (-249, -246)
of atoms with no covalent
bonds between them?

Remember – London forces are present in all liquids and solids.

Practice with these, supposing all to be in liquid or solid phase:  methane, ethanol, sucrose (look up structure), NaOH, SiC, F2O, Cl2O, octane, radon, uranium, hydrobromic acid.

PROPERTIES OF LIQUIDS & INTERPARTICLE FORCES:  

Why would a metal object with higher density than water float on water?  
Why can we fill a glass of water above its rim? 
Surface tension is related to strength of attractive forces in liquid: the stronger the attractive forces the greater the surface tension
Surface tension is the energy required to increase surface area by a unit amount; units are J/m2

Figure 12.18 The molecular basis of surface tension

Table 12.3 Surface tension and forces between particles
Figure 12.19 Shape of water or mercury meniscus in glass

Properties of Liquids
Capillary action: rising of a liquid through a narrow space against the force of gravity
Viscosity: resistance to flow, units in Newton-seconds/m2

Table 12.4 Viscosity of water at several temperatures

Why water is special:
Water molecules are 80% H-bonded at normal conditions
Molecules are so close together that you cannot tell which H's belong to which O in each molecule
This is important to life on earth (& possibly elsewhere)
Ice floats on liquid water because the solid (ice) is less dense than the liquid (good for fishies)
Ice forms a crystal structure in tetrahedral arrangement
Hydrogen-bonding also accounts for other physical properties:
	Lower weight alcohols are very soluble in water because of the -OH functional group
	Great solvent properties because water is so polar
	Very high specific heat as noted back in Chapter 6
	High surface tension

Figure 12.20 The H-bonding ability of the water molecule
Figure 12.21 The hexagonal structure of ice

SOLIDS & CRYSTAL STRUCTURES
SOLIDS:  fixed particles that cannot move with velocity, but do vibrate and rotate in position, so they do have KE
Generally have long-range order - crystals have well-defined regular shapes, or if short-range order they are amorphous - no regular shape, like asphalt, wax, glass
Crystal structure includes the four types of solids
	ionic (all cation-anion units)
	metallic (Cu, Zn, U, etc.)
	molecular/atomic (ice, I2, etc.)
	network covalent (diamond, SiC, SiO2)

General Properties of the Four Types of Crystalline Solids
1.  Ionic (KNO3, MgO): high MP/BP; some water-solb, brittle, conduct only when molten or aqueous
2.  Molecular/Atomic (C10H8, I2): low MP/BP; more solb in nonpolar; nonconductors
3.  Network Covalent (Cdiamond, SiC, SiO2):  very high MP/BP; insolb, brittle, non- or semi-conductor
4.  Metallic (Cu, Fe, U):  wide range of MPs; insolb, malleable, ductile, elec conductor 

Figure 12.22 The striking beauty of crystalline solids

Crystal Structures
Crystals have a crystal lattice arrangement of which smallest pieces are unit cell in 3-D, containing > one formula unit
Seven basic types: cubic, tetragonal, orthorhombic, monoclinic, hexagonal, rhombohedral, triclinic
See packet of handouts and Dry Lab VI(?) in Lab Manual for “Crystal Structures and Characteristics” 
Simplest are the cubic, of which there are three types
	Simple cubic (sc): metals and ionic cmpds
	Body-centered cubic (bcc): metals
	Face-centered cubic (fcc): metals and ionic cmpds

Figure 12.23 The crystal lattice and the unit cell

Cell length and cell volume
See figure 12.28 for derivation of cell length based on which cubic structure makes up the unit cell.  For metals cell length is:  sc length = 2r; bcc length = 4r/(3)1/2; fcc length = 4r/(2)1/2
Cell length determination is different for ionic compounds, which are simple cubic or face-centered cubic:  sc length = 2(r+R)/(3)1/2; fcc length = 2(r+R)
Volume of any cube = (length)3 

Figure 12.24 The three cubic unit cells  (in three parts)

Crystal Structures
See handouts and practice problems
Also chapter problems: 61, 64, 67, 73, 75 
	4th ed. #98:  Polonium is a rare radioactive metal that is the only element with a crystal structure based on the simple cubic unit cell.  If its density is 9.142 g/cm3, calculate an atomic radius for a polonium atom.
	4th ed. #101:  Tantalum, with D = 16.634 g/cm3, has a bcc structure with an edge length of 3.3058 Angstroms.  Use its molar mass and this data to prove Avogadro’s number.

Figure on Phase Changes (similar to figure 12.2):  melting/freezing, vaporizing/condensing, sublimating/depositing

Vapor Pressure
Evaporation/vaporization: small fraction of molecules have high enough velocity to escape force of attraction at surface
RATE OF EVAPORATION:  will increase with increasing T, since fraction of molecules with escape vel will increase
In a closed system, a dynamic equilibrium will be reached: 
	Rate of evaporation = rate of condensation
	Vapor pressure:  vapor molecules exert a partial pressure called vapor pressure

Figure 12.4 Liquid-gas equilibrium
Figure 12.5 The effect of temperature on the distribution of molecular speed in a liquid

Vapor Pressure Practice
If 1.00 L of water is placed in 2.30x104 L closed room, will all the water evaporate?  Given D = 0.997 g/mL, Vapor Pressure = 23.8 torr
(Hint: Water will evaporate until room is at 23.8 torr partial pressure of water vapor.)
Do your work here:

Vapor Pressure vs. Temperature
Boiling point:  occurs when you see bubbles of gas forming in the liquid and coming to surface
Any pure liquid remains at constant T while boiling, since this is a change of state
Definition:
	BP is the Temperature at which VP = barometric P  
Why does water boil at 100oC in Fairfield and at 95oC in Denver? 
	Write your answer here:

Figure 12.6 and Figure 12.7

You practice drawing and labelling a generic VP curve

Vapor pressure curves:

Initial "phase diagrams" incorporated into P/T diagrams that will include solid phase later 
Why can NH3 be condensed from gas to liquid at Room T by compression, but N2 can't?

Relative Humidity
NOT IN TEXT:  
Relative humidity as reported by weather forecasters:  
%water evap:  actual partial P/equil vapor P * 100
If actual is 12.8 and VP for given T is 21.1, relative humidity is 61%

VP and DHvap
DHvap is related to VP and T thru the Clausius-Clapeyron equation
ln P = (-DHvap/RT) + C 	(where R = 8.314 J/mol-K, T in K)
If plotted on a graph, 
slope=(ln p2 – ln p1)/(1/T2 – 1/T1)= -DHvap/R
	Rearranges to Clausius-Clapeyron Equation (next slide)

The Clausius-Clapeyron Equation:  MEMORIZE THIS
	ln (P2/P1) = -DHvap (1/T2   -  1/T1)
			R

	Alternately, if you don’t want to use the negative sign: ln (P2/P1) = DHvap/R(1/T1-1/T2)

Clausius-Clapeyron equation examples:

Look at Sample Problem 12.2 in text.

My Example:  hexane has DHvap = 30.1 kJ/mol and at 25.0oC, VP = 148 torr.  What will VP be at 50.0oC?  
	Do your work here:

Practice:
Chapter problems: 17 & 18
17: A liquid has DHovap of 35.5 kJ/mol and a BP of 122oC at 1.00 atm.  What is its VP at 113oC?
18: What is the DHovap of a liquid that has a VP of 641 torr at 85.2oC and a BP of 95.6oC at 1.00 atm? 

Figure 12.8 from 4th ed. not in principles: Iodine subliming – going directly from solid to vapor

Simple Phase Diagrams
Are P vs. T diagrams showing three phases for pure elements or compounds, incorporates the VP curve
Critical Point is where liquid and gas cannot be distinguished from each other
Triple Point is where solid, liquid and gas phases meet and all three are present
Example:  For water the Triple Point is 0.01oC and 4.58 torr
For CO2 Triple Point is at -56.7oC and 5.1 atm

Figure 12.8 in Principles 2nd ed.:  Phase Diagrams for CO2 and H2O

You must draw and label phase diagrams based on data given to you and determine the physical state of a state of a substance from its placement on a phase diagram.

Practice: 20 & 22

See Table at end of document
SUMMARY OF INTERPARTICLE FORCES

INTERPARTICLE									STRENGTH		PREDICTED
FORCE			PARTICLES		DESCRIPTION			kJ/mol		MP/BP		EXAMPLES

Covalent			atoms			sharing of e-s				strong		high/high		H2O (0,100)
      Bonding						between two atoms		

Network Covalent		atoms			sharing of e-s				very strong	very high		diamond (35530,?)
       Bonding						betw two atoms extended	

Metallic Bonding		atoms			sharing of e-s in a sea			strong		high			iron (1555, 3000)
							between all atoms

Ionic Bonding			cations &		electrostatic attraction			400-4000	high			NaCl (804, 1413)
				anions			between opposite charges

Hydrogen Bonding*		polar molecules	attraction between H on		15-40		medium		H2O
				with H bonded to	one molecule to F, O or N
				F, O or N		on other molecules

Dipole-Dipole			polar molecules	attraction between slightly		5-25		low/medium		(CH3)2C=O (-95, 56)
      Attraction						opposite charges

London Dispersion		all atoms, ions		attraction between induced		0.05-40	low/medium		I2 (114, 183)
      Forces			and molecules		dipole opposite charges

Ion-Dipole			ions and polar		attraction between 			40-600		na			NaCl(aq)
     Attraction			molecules		opposite charges

Dipole-Induced		polar and nonpolar	attraction between dipole		2-10		na			O2 in H2O
    Dipole Attraction		molecules		and induced dipole

*Not really bonding, just an unfortunate choice for naming this intermolecular force of attraction.   Italics indicate the three Intermolecular Forces.