Title: Lesson 7 Intermolecular Forces Learning Objectives: – Learn to identify and explain the three types of intermolecular forces: Van der Waals Permanent

Title: Lesson 7 Intermolecular Forces

Learning Objectives:

– Learn to identify and explain the three types of intermolecular forces:

• Van der Waals• Permanent dipole-dipole• Hydrogen bonds

– Understand and explain the effects of the above on melting/boiling points

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Use the VSEPR theory to deduce the shape of H3O+ and C2H4. For each species, draw the Lewis structure, name the shape, and state the value of the bond angle(s).

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Note Taking Do not copy the notes, re-express them Include diagrams

Van der Waals / Temporary Dipole-

Induced Dipole

Dipole-Dipole / Permanent Dipole Hydrogen Bonds

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Intermolecular Forces

The attractive forces between molecules

It is these that are partially broken during melting, and fully broken during boiling

Note: when molecular compounds melt/boil, the bonds in the molecule do not break, it is just the attractive forces between the molecules that break

Intermolecular Forces (imf)

These are weak electrostatic forces of attraction between neighbouring molecules.

They are much weaker than covalent, ionic or metallic bonding.

They influence ONLY the physical properties of molecules.

GIANT structures

SIMPLE molecules (eg H2O, H2, CH4 etc)

imf not applicable because NO separate MOLECULES exist

covalent (eg diamond), or ionic (eg NaCl) or

have high melting and boiling points

have much lower melting and boiling points

imf are applicable

metallic (eg Cu)

IMF influence PHYSICAL properties :

Melting points and boiling points

Solubility in water and other solvents

3D shapes of complex molecules such as DNA

Viscosity of liquids.

Density

etc etc

Boiling point variations are very good indicators of variations in IMF

Don’t forget !!!!Strong covalent bonds within molecules

Weak imf between molecules

are not broken when molecular substances are vaporized

are broken when molecular substances are vaporized

Boiling point variations suggest 3 types of imf :

1. London (Dispersion)

2. Dipole-dipole forces

3. Hydrogen bonds

For similar size molecules, imf strength INCREASES

Boiling point INCREASE IMF strength INCREASE

NOTE: Van der Waals’ forces is an umbrella term to cover both London dispersion and dipole-dipole attractions...

GROUP FORMULA& BPt /K

FORMULA& BPt /K

FORMULA& BPt /K

FORMULA& BPt /K

IVCH4

109SiH4

161GeH4

190SnH4

221

VNH3

240PH3

185AsH3

218SbH3

256

VIH2O373

H2S212

H2Se246

H2Te280

VII HF293

HCl188

HBr206

HI238

PERIOD 2 3 4 5

Noble Gases

He20

Ne27

Ar87

Kr121

NH3

240

H2O373

HF293

CH4 109

SiH4

161GeH4

190SnH4

221

He20

Ne27

Ar87

Kr121

PH3

185AsH3

218SbH3

256

H2S212

H2Se246

H2Te280

HCl188

HBr206

HI238

London Forces only Hydrogen bonds DP-DP +L forces

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London (dispersion) akaTemporary or instantaneous induced dipole forces

Non-polar molecules such as Cl2, have no permanent separation of charge (no permanent dipole)...

However, random electron movements create a small, temporary dipole

This induces a similar dipole in a neighbouring molecule

This creates a small attraction between them

These are weak and exist only for the tiniest fraction of a second

London (dispersion) forces are present in all molecules

Increase with molecular mass Decrease with the roundness of a molecule

Leading to momentary attraction between temporary dipoles

Consider non-polar molecules such as Ne, I2, CH4 etc

Electron cloud of the molecule is in constant random motion

Leading to momentary electron density imbalance.

Leading to a temporary dipole in the molecule

+ - + -

which induces a temporary dipole in neighbouring molecule.

which IS the van der Waal force

Van der Waal Forces

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e-

e-e-

e-

e-

e-

e-e-

e-

e-

e-e-

e-

e-

e-

e-e-

12 of 43 © Boardworks Ltd 2009

Van der Waals forces


Strength of van der Waals forces

The strength of van der Waals forces increases as molecular size increases.

Atomic radius increases down the group, so the outer electrons become further from the nucleus. They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce.

050

100150200

-50-100-150-200

bo

ilin

g p

oin

t (°

C)

Br2

This is illustrated by the boiling points of group 7 elements. F2 Cl2 I2

element


Strength of van der Waals forces

Straight chain alkanes can pack closer together than branched alkanes, creating more points of contact between molecules. This results in stronger van der Waals forces.

butane (C4H10)

boiling point = 272 K

2-methylpropane (C4H10)

boiling point = 261 K

The points of contact between molecules also affects the strength of van der Waals forces.

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Boiling points increase as number of electrons increases...

• More electrons mean a larger electron cloud density and this will induce a stronger attraction between molecules...


Boiling points of alkanes

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Dipole-dipole forces akaPermanent dipole forces Different atoms have different electronegativities,

which means there will be variations in the electron charge density in different parts of a molecule

If a molecule is not symmetrical, the variation produces a dipole where a molecule has a positive and a negative end The end with high charge density is - The end with low charge density is +

Oppositely charged dipoles attract each other.

This is a relatively strong attractive force

If a molecule is symmetrical, variations in electron charge density cancel each other out and the molecule is non-polar... (Think of the tug of war example!)

-

+

- + -


Permanent dipole–dipole forces

If molecules contain bonds with a permanent dipole, the molecules may align so there is electrostatic attraction between the opposite charges on neighbouring molecules.

Permanent dipole–dipole forces (dotted lines) occur in hydrogen chloride (HCl) gas.

The permanent dipole–dipole forces are approximately one hundredth the strength of a covalent bond.


Permanent dipole–dipole or not?

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Melting and Boiling Points are stronger in Dipole-Dipole Strength can vary depending on the distance and relative

orientation of the dipoles. Generally stronger than London forces so energy needed to

separate bonds will be greater. NOTE: Weaker London forces also occur alongside

Dipole-Dipole forces...

NOTE! It’s important to compare substances with a similar molecular mass – otherwise the difference can be attributed to stronger London forces based on more electrons...

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Van der Waals’ forces is an umbrella term for:

I.e. van der Waals’ forces refers to all forces between molecules that do not involve electrostatic attractions between ions or bond formation.

Why do molecules such as CCl4, BF3 and BeCl2 NOT show dipole-dipole forces?

Individual bonds are polar eg δ+Be-Clδ-

but the molecules are NOT because

they are SYMMETRICAL

bond dipoles CANCEL

NON-POLAR molecule

Cl–Be–Clδ- δ-δ+

δ+

B

F

F Fδ-

δ-

δ-

3δ+

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Hydrogen bonds akaH-bonds The strongest type of intermolecular force

They occur between a nitrogen, oxygen or fluorine and a hydrogen that is bonded to a nitrogen, oxygen or fluorine

N, O and F are the three most electronegative elements, and all have lone-pairs when bonded

When H is bonded to N, O or F, the electrons in the bonded are strongly attracted to the N/O/F, leaving the H very positive

The lone pair on the N/O/F is strongly attracted to the positive hydrogen


What is hydrogen bonding?

When hydrogen bonds to nitrogen, oxygen or fluorine, a larger dipole occurs than in other polar bonds.

This is because these atoms are highly electronegative due to their high nuclear charge and small size. When these atoms bond to hydrogen, electrons are withdrawn from the H atom, making it slightly positive.

Hydrogen bonds are therefore particularly strong examples of permanent dipole–dipole forces.

The H atom is very small so the positive charge is more concentrated, making it easier to link with other molecules.


Hydrogen bonding

In molecules with OH or NH groups, a lone pair of electrons on nitrogen or oxygen is attracted to the slight positive charge on the hydrogen on a neighbouring molecule.

Hydrogen bonding makes the melting and boiling points of water higher than might be expected. It also means that alcohols have much higher boiling points than alkanes of a similar size.

hydrogen bond lone pair

directly bonded to a VERY electronegative atom

For these to occur you need:

1. A VERY electronegative atom with

ONLY F, O & N are sufficiently electronegative

Hydrogen Bonding

ONLY δ+H-F , δ+H-O or δ+H-N

and 2. a H atom

are appropriate.

an available lone pair of electrons,

- electronegative atom draws e- away from H (de-shields it) making it SLIGHTLY positive, δ+H

A hydrogen bond = the attraction between a lone pair on a N, O or F atom and a de-shielded H

atom in a δ+H-F, δ+H-N or δ+H-O bond

HF has approx. one H bond per molecule

In SOLID, H-bonds are permanent

H2O has approx. two H bond per molecule

In LIQUID, H-bonds are continuously breaking and reforming

In GAS, H-bonds are completely broken

δ+ H-F: δ+ H-F: δ+ H-F:

H

δ+ H-O:

δ+

..

H

δ+ H-O:

δ+

..

H

δ+ H-O:

δ+

..

Hence water’s unusually HIGH mpt and bpt

NH3 has approx. one H bond per molecule

H

δ+ H-N:

H

H

δ+ H-N:

H

H

δ+ H-N:

H

Decreasing strength of individual H bonds

δ+H-F: , δ+H-O: , δ+H-N:

because electronegativity decreases

but water forms TWO H-bonds per molecule

order of b pt is H2O >> HF > NH3

not HF > H2O > NH3

CH3CH2OH

CH3-C-CH3

H2S

will H-bond.

will not H-bond.

will not H-bond

Further examples :

-O-Hδ+ - - - :O-

O bonded to C, not H

H bonded to S which is NOT electronegative enough for H-bonds

O

H2O has approx. two H-bonds per molecule

H

δ+ H-O:

δ+

..

H

δ+ H-O:

δ+

..

H

δ+ H-O:

δ+

..

Hence water’s unusually HIGH melting point (0ºC)

for such SMALL molecules, water molecules are DIFFICULT (require a lot of added energy) to separate

and boiling point (100ºC)

when compared to other molecules of similar size / mass

eg H2S (a heavier molecule!) is a GAS at room temperature because it does not hydrogen bond

Hydrogen Bonding and the Unusual Physical Properties of Water

0

50

100

150

200

250

300

350

400

2 3 4 5Period

BPt (/K)

Noble gases

Group IV hydrides

Gp V hydrides

Gp VII hydrides

Gp VI hydrides

B Pt’s of NH3, H2O and HF are UNUSUALLY high imf UNUSUALLY STRONG

HYDROGEN BONDS

DIPOLE-DIPOLE FORCES but

increasing VW forces control b pt

LONDON FORCES ONLY


Hydrogen bonding and boiling points


Boiling points of the hydrogen halides

The boiling point of hydrogen fluoride is much higher than that of other hydrogen halides, due to fluorine’s high electronegativity.

02040

-20-40-60-80

-100b

oil

ing

po

int

(°C

)

HF HCl HBr HI

The means that hydrogen bonding between molecules of hydrogen fluoride is much stronger than the permanent dipole–dipole forces between molecules of other hydrogen halides. More energy is therefore required to separate the molecules of hydrogen fluoride.


Permanent dipole–dipole forces

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Comparing Boiling Points... If we compare these different forms of C2H6O

(known as isomers) we can see how the presence of H-bonds affects the boiling point...

O is bonded to C so no H-bond

O is bonded to H so H-bond present

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H2O the anomaly... Water has 2 hydrogen atoms and 2 pairs of lone pairs on the

oxygen Hence, it can form 4 hydrogen bonds with neighbouring

water molecules. Liquid water has fewer bonds, but ice uses up 4 bonds which

results in a tetrahedral shape that is fixed and open... So ice is less dense than water as it expands... Usually solids form closely packed particles and become

more dense...

4 H-Bonds circled

Structure of ice with 4 H-Bonds

Also because of it’s strong H-bonding, ice has an unusually LOW DENSITY compared to water

In ICE, the maximum number of H-bonds are operative

molecules are held apart in a tetrahedral arrangement of covalent and H-bonds

much empty space between molecules

larger volume than the same mass of water

LOWER DENSITY than water

expansion of water during freezing can burst pipes

and ice floats on water

Also because of it’s strong H-bonding, water has an unusually H IGH SURFACE TENSION

Hydrogen bonds “pull” molecules at the surface inwards creating a “skin-like” effect on the surface of water

This allows insects like the water strider to “walk on water”!

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Summary of IMF

Covalent substance have a lower M.P and B.P compared to ionic because the energy needed to overcome the IMF is lower than the energy needed to break electrostatic attractions in an ionic lattice...

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Solubility Non polar substances can dissolve in non polar solvents

by formation of London dispersion forces between solute and solvent. E.g. non polar Br2 can dissolve in non polar paraffin oil (a hydrocarbon)

Polar substances can dissolve in polar solvents e.g. water. Dipole interactions and H-Bonding are responsible. E.g. HCl, glucose (C6H12O6) and ethanol (C2H5OH) are polar substances that can dissolve

Note: Larger molecules where only a small part is polar will be less soluble as the non polar parts will not disassociate in water.

TIP! Think like for like...

Polar substances have a low solubility in non polar because the Dipole-Dipole forces keep them together and not interacting with the solvent...

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Electrical Conductivity

Covalent compounds do not conduct electricity (no ions) An exception is HCl dissolved in water – it’s ions H+ and Cl-

disassociate in water. Giant covalent molecules e.g. Graphite and Graphene are

conductors (mobile electrons). Fullerene and Silicon are semi conductors.

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Effects of intermolecular forces

Intermolecular forces play an important role in the properties of compounds including:

Melting/boiling point: Stronger intermolecular forces higher m.p./b.p.

Volatility: Stronger intermolecular forces lower volatility

Solubility: like dissolves in like Polar solutes dissolve best in polar solvents Non-polar solutes dissolve best in non-polar solvents

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Summary of Physical Properties

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Looking into intermolecular forces

Complete the activity here to research and model intermolecular forces

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Summary

Three types of intermolecular force, from strongest to weakest:

Hydrogen bonds Between N/O/F and H attached to N/O/F

Dipole-dipole Between permanent dipoles on asymmetric molecules

London (Dispersion) Between instantaneous dipoles formed on any

molecule/atom

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Title: Lesson 7 Intermolecular Forces Learning Objectives: – Learn to identify and explain the three types of intermolecular forces: Van der Waals Permanent