Intermolecular Forces and

Liquids and Solids

1

• Why do some solids dissolve in water but others do not?

• Why are some substances are gaseous at room temperature, but others are liquid or solid?

• What gives metals the ability to conduct electricity, what makes non-metals brittle?

• The answers have to do with …

Intermolecular forcesIntermolecular forces

Questions

2

Why mixtures mixWhy mixtures mix• Consider a glass of wine. Why do alcohol, water,

& pigment mix together?• There must be attractive forces.

Intramolecular forces occur between atoms

Intermolecular forces occur between molecules

• The factors that determine solubility are the strength of IMFs and speed of molecules.

3

Intermolecular Forces

11.2

Intermolecular forces are attractive forces between molecules.

Intramolecular forces hold atoms together in a molecule.

Intermolecular vs Intramolecular

• 41 kJ to vaporize 1 mole of water (inter)

• 930 kJ to break all O-H bonds in 1 mole of water (intra)

Generally, intermolecular forces are much weaker than intramolecular forces.

“Measure” of intermolecular force

boiling point

melting point

Hvap

Hfus

Hsub

4

The strengths of intermolecular forces are generally weaker than either ionic or covalent bonds.

16 kJ/mol (to separate molecules)

431 kJ/mol (to break bond)

++-

-

5

6

Excess lone pairs

7

HydrogenHydrogen

Excess H atomsExcess H atoms

8

Each lone pair and H atoms has a partner

on another water molecule.

H bond

9

(Mr=34)b.p =431K

(Mr=38)b.p=85K

(Mr=36.5)b.p=188K

H bond

10

(Mr=44)b.p =231K (Mr=44)

b.p =294K, CH3COH, ethanal

(Mr=46)b.p =352K,

CH3CH2OH, ethanol

Types of Intermolecular Forces

3. Dipole-Dipole Forces

Attractive forces between polar molecules

Orientation of Polar Molecules in a Solid

11

Types of intermolecular forces (between neutral molecules):

Dipole-dipole forces: (polar molecules)

SO O.. ::

....

:

+

--

SO O.. ::

....

:

+

--

dipole-dipole attraction

What effect does this attraction have on the boiling point?13

Polar molecules have dipole-dipole attractions for

one another.

+HCl----- +HCl-

dipole-dipole attraction

14

Types of Intermolecular Forces4. Dispersion Forces – van der Walls forces/London forces (weakest)

Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules

15

Dispersion ForcesOccur between every compound and arise from the net attractive forcesamount molecules which is produced from induced charge imbalances

The larger the molecule the greater it’s Dispersion Forces are.

16

London forcesLondon forces• Non-polar molecules do not have dipoles like

polar molecules. How, then, can non-polar compounds form solids or liquids?

• London forces are named after Fritz London (also called van der Waal forces)

• London forces are due to small dipoles that exist in non-polar molecules

• Because electrons are moving around in atoms there will be instants when the charge around an atom is not symmetrical

• The resulting tiny dipoles cause attractions between atoms/molecules

17

London Dispersion Forces

• Non - polar molecules also exert forces on each other.

• Otherwise, no solids or liquids.• Electrons are not evenly distributed at

every instant in time.• Have an instantaneous dipole.• Induces a dipole in the atom next to it.• Induced dipole- induced dipole

interaction.

18

“electrons are shifted to overload one side of an atom or molecule”.

London dispersion forces: (instantaneous dipole moment)( also referred to as van der Waal’s forces)

+ +- -

attraction

London Dispersion Forces

19

London forcesLondon forces

Instantaneous dipole: Induced dipole:

Eventually electrons are situated so that tiny dipoles form

A dipole forms in one atom or molecule, inducing a

dipole in the other20

polarizability: the ease with which an atom or molecule can be distorted to have an instantaneous dipole.

In general big moleculesare more easily polarized

than little ones.

21

Halogen Boiling Pt (K)

Noble Gas Boiling Pt (K)

F2 85.1 He 4.6

Cl2 238.6 Ne 27.3

Br2 332.0 Ar 87.5

I2 457.6 Kr 120.9

Which one(s) of the above are most polarizable?Hint: look at the relative sizes.

22

Intermolecular Forces4. Dispersion Forces Continued

Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted.

Polarizability increases with:

• greater number of electrons

• more diffuse electron cloud

Dispersion forces usually increase with molar mass.

23

SO

O

What type(s) of intermolecular forces exist between each of the following molecules?

HBrHBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.

CH4

CH4 is nonpolar: dispersion forces.

SO2

SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

24

25

Solubility and Intermolecular Forces

• Like dissolves like– Polar solutes dissolve in polar solvents

– Nonpolar solutes dissolve in nonpolar solvents

• Molecules with similar intermolecular forces will mix freely

• Molecules that can form H bondings will have greater solubility in water.

26

Ionic Solute with Polar Solvent

27

The hydration enthalpy is a measure of the strength of the interaction of the water molecule w/ the ion.

• Hydration enthalpy is the energy change when one mole of gaseous ions is converted to one mole of hydrated ions.

• It increases w/ the charge on the ion and decreases w/ the size of the ions. Therefore, it’s greater for cations than the anions since they’re smaller in size and the water molecules can give a better packing because of its bent geometry.e.g hydration enthalpy is great for Al3+ .

28

Ionic Solute withNonpolar Solvent

29

Nonpolar Solute withNonpolar Solvent

30

Nonpolar Solute with Polar Solvent

31

32

33

Example:• Compare the solubilities of

a)ethanoic acid, ethanol and propane.

b)Propanone and HI d) ethanoic acid and pentanoic acid

c)

34

• Hydrocarbon chain disrupts the H bonding in

water.

35

• Take a look at the solubility rules on p.128.

36

37

4-nitrophenolb.p=279°C

Forms H bond mainly w/ other 4-nitrophenol

molecules

2-nitrophenolb.p=216°C

Forms H bond mainly intra-molecularly

H bonding

38

Intramolecular H bonding in α-helix

A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions.

An amorphous solid does not possess a well-defined arrangement and long-range molecular order.

A unit cell is the basic repeating structural unit of a crystalline solid.

39

Properties of Ionic Compounds

1. Crystalline solids - a regular repeating arrangement of ions in the solid:

– Ions are strongly bonded together.– Structure is rigid.

2. High melting points• Coordination number- number of ions of

opposite charge surrounding it

40

- Page 198

Coordination Numbers:

Both the sodium and chlorine have 6

Both the cesium and chlorine have 8

Each titanium has 6, and each oxygen has 3

NaCl

CsCl

TiO2

41

Do they Conduct Electricity?

• Conducting electricity means allowing charges to move.

• In a solid, the ions are locked in place.• Ionic solids are insulators.• When melted, the ions can move around.3. Melted ionic compounds conduct.

– NaCl: must get to about 800 ºC.– Dissolved in water, they also conduct

(free to move in aqueous solutions)

42

- Page 198

The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

43

Ionic solids are brittle

+ - + -+- +-

+ - + -+- +-

Force

44

Ionic solids are brittle

+ - + -

+- +-+ - + -

+- +-

• Strong Repulsion breaks a crystal apart, due to similar ions being next to each other.

Force

45

46

Types of CrystalsCovalent (Network or macromolecular) Crystals – e.g.

Silica(SiO2),SiC, BN , Si.Stronger than IM forces.• Held together by covalent bonds• Hard, high melting point• Poor conductor of heat and electricity

diamondgraphitecarbon

atoms

47

48

49

This p bond overlap forms a huge p bonding network.

• Electrons are free to move through out these delocalized orbitals.

• The layers slide by each other.

50

51

52

SiC

• Take a look at the tables on p. 129, 133, and 134.

• Allotropes are different forms of an element that exist in the same physical state. e.g. C( graphite, fullerene, graphite), oxygen (O2 and O3)

53

The giant covalent structure of silicon dioxide

• There are three different crystal forms of silicon dioxide. The easiest one to remember and draw is based on the diamond structure.

• Crystalline silicon has the same structure as diamond. To turn it into silicon dioxide, all you need to do is to modify the silicon structure by including some oxygen atoms.

• Notice that each silicon atom is bridged to its neighbours by an oxygen atom. Don't forget that this is just a tiny part of a giant structure extending on all 3 dimensions.

54

Types of Crystals

Molecular Crystals• Formed from molecules• Held together by intermolecular forces• Soft, low melting point• Poor conductor of heat and electricity

55

Examples of Molecular solids

Iodine, I2• Iodine is a dark grey crystalline solid with a purple

vapour. M.Pt: 114°C. B.Pt: 184°C. It is very, very slightly soluble in water, but dissolves freely in organic solvents.

• Iodine is therefore a low melting point solid. The crystallinity suggests a regular packing of the molecules.

56

Examples of Molecular solids

Ice

• Ice is a good example of a hydrogen bonded solid.• There are lots of different ways that the water molecules

can be arranged in ice. This is one of them, but NOT the common one - I can't draw that in any way that makes sense! The one below is known as "cubic ice", or "ice Ic". It is based on the water molecules arranged in a diamond structure

57

Types of CrystalsMetallic Crystals – Typically weaker than covalent, but

can be in the low end of covalent• Lattice points occupied by metal atoms• Held together by metallic bonds• Soft to hard, low to high melting point• Good conductors of heat and electricity

Cross Section of a Metallic Crystal

nucleus &inner shell e-

mobile “sea”of e-

58

Metallic Bonds are…• How metal atoms are held together in the

solid.• Metals hold on to their valence electrons

very weakly.• Think of them as positive ions (cations)

floating in a sea of electrons:

+ + + ++ + + +

+ + + +59

How does electricity flow in metal?

• In metals, free electrons carry the current.

• In metal wire, all the atoms share the electrons in their outermost electron shell and these electrons are free to move anywhere within the wire.

60

Strength and Workability

• Malleability and ductility• Metals are described as malleable (can be

beaten into sheets) and ductile (can be pulled out into wires). This is because of the ability of the atoms to roll over each other into new positions without breaking the metallic bond.

• If a small stress is put onto the metal, the layers of atoms will start to roll over each other. If the stress is released again, they will fall back to their original positions. Under these circumstances, the metal is said to be elastic.

61

1) Ductility 2) Malleability

Due to the mobility of the valence electrons, metals have:

and

Notice that the ionic crystal breaks due to ion repulsion!

62

Malleable

+ + + ++ + + +

+ + + +

Force

63

Malleable

+ + + +

+ + + ++ + + +

• Mobile electrons allow atoms to slide by, sort of like ball bearings in oil.

Force

64

65

Strength and Workability

• If a larger stress is put on, the atoms roll over each other into a new position, and the metal is permanently changed.

66

An amorphous solid does not possess a well-defined arrangement and long-range molecular order.

A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing

Crystallinequartz (SiO2)

Non-crystallinequartz glass 67

Types of Crystals

68

Name type of solid Force(s) Melting Pt. (oC)

Boiling Pt. (oC)

Ne

molecular

-249

-246

H2S

molecular

-86

-61

H2O

molecular

0

100

Mercury

metallic

-39

357

W

metallic

3410

5660

CsCl

ionic

645

1290

MgO

ionic

2800

3600

Quartz (SiO2)

covalent network

1610

2230

Diamond (C)

covalent network

3550

4827

69

70

71

72