IB Chemistry on Polarity, Hydrogen Bonding and Van Der Waals forces

Chemical Bonds

Ionic Bonds

Transfer of electron from metal to non metal

Metal donate e Non Metal accept e

Positive ion (cation)

Negative ion (anion)

Ionic compound

Covalent Bonds Sharing of electron bet non metal atoms

Equal sharing electrons

+ - electrostatic forces

attraction

4 0.4 0 Difference in electronegativity

2

EN - 0.9 EN – 3

Diff = (3 – 3 ) = 0

H

EN – 2.1

Diff = 3 – 0.9 = 2.1

Polar covalent Bonds Non Polar covalent Bonds

Unequal sharing electrons

Covalent Polar Non polar

CI CI

EN – 3

Covalent Non polar

CI

Covalent polar Ionic

EN – 3

Diff = (3 – 2.1 ) = 0.9

Na+ CI-

EN – 3

Shared electron cloud closer to O

Electronegativity

Electronegativity (EN) •Tendency of atom to attract/pull shared/bonding electron to itself •EN value higher – pull/attract electron higher (EN value from 0.7 – 4)

Electronegativity • EN increase up a Group • EN increase across a Period

H 2.2

Li Be B C N O F

CI

Br

I

1 1.6 2 2.6 3 3.4 4

Electronegativity values

N, O, F have high EN value

3.2

3

2.7

Molecule Diff in EN Polarity

H - F (4.0 – 2.2) = 1.8 Most polar

H - CI (3.2 – 2.2) = 1

H - Br (3.0 – 2.2) = 0.8

H - I (2.7 – 2.2) = 0.5 Least polar

Polarity

Shape Diff in EN

Symmetrical Asymmetrical

Bond polarity cancel out each other

Polar bonds – molecule NON POLAR

Bond polarity cancel out each other

Polar bonds – molecule POLAR

Lewis structure

VSEPR

Geometry

1

4 ECC

3 bond pair

1 lone pair

.. N

H H

H

Polarity

2

3

4

Polar

✓ ✗

✗

In presence of electric field

Separation of charges

Unequal distribution electron due to diff EN value

shared electron closer to O shared electron closer to F

Covalent Bonds

Polar covalent Bonds Non Polar covalent Bond

Equal sharing electron Unequal sharing electron

Sharing of electron

Formation electric dipole

Partial +/-

Dipole moment

towards O

Partial + ( δ+)

Partial – (δ−)

Net dipole moment

Molecule is polar (dipole)

Net Dipole moment

Measured in Debye

Turning force / Dipole moment = Force x Distance Polar covalent Bonds

+ - O

III

C δ+

δ-

Turning force – dipole moment

+ -

O

II

C

II

O

δ+

δ-

δ-

No Turning force – No dipole moment

✓

Molecule polar ✓

O

O

Polarity

Shape

Asymmetrical

Polar bond Polarity dont cancel (ASYMMETRICAL)

Net dipole moment Molecule POLAR

Polar bond Polarity cancel (SYMMETRICAL) NO net dipole moment Molecule NON POLAR

Shape

Symmetrical

Polar bonds

CI Polar bonds

δ-

δ+ δ+

δ+

δ-

δ-

δ-

δ-

δ-

δ-

Bond polarity

don’t cancel

Bond polarity

cancel

H

Net Dipole moment No Net Dipole moment

✗

Asymmetrical Symmetrical

δ- δ+

Polar bonds

Bond polarity

don’t cancel

Net Dipole moment

C O

Polar bond Polarity dont cancel (ASYMMETRICAL)

Net dipole moment Molecule POLAR

δ- δ-

Polar bonds

Bond polarity cancel

No Net Dipole moment

Polar bond Polarity cancel (SYMMETRICAL)

NO dipole moment Molecule NON POLAR

✗ ✗ ✓ ✓

I

Bonding Forces

Bonding Forces within molecule Bonding Forces bet molecule

Intermolecular force bet molecule (IMF) Ionic bond

Covalent bond

Metallic bond

Permanent dipole

Ions Molecules/NOT ions

Ion dipole forces

Polar Non Polar

Hydrogen bonding

Temporary dipole (instantaneously induced dipole) London dispersion forces

Van Der Waals’ Forces attraction

Polar molecule (dipole – dipole attraction)

_ _ _

Attraction bet ions with polar molecules

Dipole/dipole attraction Dipole/dipole attraction

(involving H atom attach to N,O F)

Hydrogen bonding (dipole – dipole attraction)

Forces bet molecule

Permanent dipole

Polar Non Polar



Polar molecule (dipole – dipole attraction)

Polar molecules due to diff in EN values

Dipole/dipole interaction

Electrostatic forces attraction bet molecules



H atom bond to electronegative atom, N, O, F

Partial H+ attracted to lone pair electron on N, O, F

Electrostatic force attraction bet molecules involve H

Intermolecular force bet molecule (IMF)

Non Polar molecule (Induced dipole attraction)

Random movement /distribution of electron

Instantaneous negative charge on atom

Induced a temporary dipole separation


Non polar molecules

Polar molecules Polar molecules

Forces bet molecule

Molecules


H atom bond to electronegative atom, N, O, F

Partial H+ attracted to lone pair electron on N, O, F

Electrostatic force attraction bet molecules involve H

Permanent dipole

Polar Non Polar



Polar molecule

Polar molecules due to diff in EN values

Dipole/dipole interaction



Molecules

Non Polar molecule (Induced dipole attraction)

Random movement /distribution of electron

Instantaneous negative charge on atom

Induced a temporary dipole separation

Instantaneous dipole in one atom

induce a dipole in its neighbor


Non polar molecules Polar molecules

Polar molecules

Requirement for H bonding

•H atom bonded to N, O, F

•N, O, F - highly electronegative/ small size

•Attract electron close to itself – Polarised H+

•N---H, O—H, F—H bonds VERY POLAR

•Very polar H+ attract closely to lone pair on N, O, F

N ---- H O ---- H F ---- H δ- δ- δ- δ+ δ-+ δ+

Types of forces/Bonding

Factors affecting VDF forces Intermolecular force bet molecule (IMF)

Interaction Strength Boiling Point/C

Covalent Strongest High

Ionic Strong High

Ion dipole Strong HIgh

Dipole- dipole (H bond)

Moderate High

Dipole – dipole Weak Low

Temporary induce dipole (London dispersion)

Weakest Low

Dipole – dipole attraction Induced – dipole attraction London dispersion forces

Polar Non Polar

All molecules have London dispersion forces due to uneven distribution of electron cloud

- - -

- -

- -

-

-

-

- -

- - -

δ+ δ- London dispersion forces

RMM/Size Surface Area

London dispersion forces


N N

F F RMM – 38

RMM – 28

Size/ RMM increase

Number electrons increase

Temporary induced dipole increase

Van Der Waals bet molecule increase

RMM same

Surface area increase


Van Der Waals bet molecules increase

RMM – 72

RMM – 72

Pentane (C5H12)

Factor affecting b/p of molecules

Temporary dipole attraction London dispersion force

Permanent dipole attraction Dipole/dipole attraction

Hydrogen bonding

Factors affecting VDF forces Factor affecting b/p of molecules

RMM/Size Surface Area

N N

F F RMM – 38

RMM – 28

Size/ RMM increase

Number electrons increase


Van Der Waals bet molecule increase

RMM same

Surface area increase


Van Der Waals bet molecules increase

RMM – 72

RMM – 72

Pentane (C5H12)

Temporary dipole attraction London dispersion force

Permanent dipole attraction Dipole/dipole attraction

Hydrogen bonding

H2 N2 CI2 H2O

RMM 2 28 71 18

Boiling Point/C

-252 -196 -34 100

Forces London force

London force

London force

London force Dipole/dipole Hydrogen bonding

- -

- -

- -

H2

London forces

N2

London forces

CI2

London forces

H2O

London forces

Dipole/dipole

Hydrogen bond

RMM increase - London force/VDF increase – boiling point increase

- -

- -

- - -

- - -

- -

- - -

-

Hydrogen bonding Hydrogen bonding

RMM lowest - boiling point highest - due to hydrogen bonding

HCI HBr HI HF

RMM 36.5 81 128 20

Boiling Point/C

-85 -66 -35 19.5

Forces London force/VDF

London force/VDF

London force/VDF

London force/VDF Dipole/dipole Hydrogen bond

RMM increase - London force/VDF increase – boiling point increase

Which liquid has higher boiling point?

H H

H

H H

H

DNA

Br Br

Br

I I

I

Hydrogen bonding

RMM lowest - boiling point highest - due to hydrogen bonding

Br2 ICI

RMM 162 162

B/p/C 58 97

Forces London force/VDF

London force/VDF

Dipole/dipole


- - -

- - -

Br2

London forces bet molecules

Br Br Br Br I I I CI CI CI + + + - - -

ICI London forces + Dipole –dipole attraction

Hydrogen Bond bet

nitrogenous base

CH3CH2OH CH3CH2CH2OH CH3COOH C2H5-O-C2H5

RMM 46 60 60 74

Boiling Point/C

78 97 118 34

Forces London H2 bond

London H2 bond

London H2 bond

London

RMM - London force – boiling point


Stronger Hydrogen bond – boiling point

CH3COOH boiling point higher

C=O (carbonyl) – electron withdrawing gp

withdraw electron from OH gp

O-H gp more polarised

stronger H2 bond

Hydrogen

bond

Hydrogen

bond ✕

C - O - H

= o

C3H8 CH3CHO CH3CH2OH

RMM 44 44 46

Boiling Point/C

-42 20.2 118

Forces London

London Dipole/dipole


H2 bond

RMM highest

No Hydrogen Bond


.......

..........

London forces London forces

+

Dipole/dipole

London forces

+

Dipole/dipole

+

Hydrogen Bond

✕

Hydrogen

bond

Why 2 Nitrophenol has lower b/p than 4 nitrophenol?

Molecule symmetrical


No net dipole moment

Molecule NON POLAR

C6H5NO3 ( 2 nitrophenol)

C6H5NO3

( 4 nitrophenol)

RMM 139 139

Boiling Point/C

216 279

Forces London Intramolecular

H2 bond

London Intermolecular

H2 bond

Which NCI3 is polar but BCI3 non polar?

...

…

Intramolecular

H2 bond Non polar

…

…

✕

More intramolecular H2 bond

Lack intermolecular H2 bond

Intermolecular

H2 bond ✓

✓

More intermolecular H2 bond

Lack intramolecular H2 bond

BCI3

NCI3

RMM 117 120

Boiling Point/C

13 71

Forces London


…...

Non polar

...

δ-

δ- δ-

δ-+

polar

polar

….....

Molecule asymmetrical

Bond polarity does not cancel

NET dipole moment

Molecule POLAR

✓ Dipole dipole

Trans isomer – CI opposite side


NO Net dipole moment / NON POLAR

Intermolecular forces weaker

Molecule in linear shape

Able to pack closely together

Greater surface area for interaction

Intermolecular forces stronger

Polarity for Cis/Trans Geometrical Isomers

Molecule in kink/bend shape

Unable to pack closely together

Less surface area for interaction

Intermolecular forces weaker

Cis dichloroethene

Trans dichloroethene

RMM 96 96

Melting Point/C

-80 -50

Boiling Point/C

60 48

...

δ- Cis

Dipole dipole

Trans

... ...

...

Cis dichloroethene


Melting Point/C

-80 -50

Kink/ bend shape

Cis dichloroethene


Boiling Point/C

60 48

Greater attraction

Cis isomer – CI same side

Bond polarity does not cancel

Net dipole moment /POLAR

Intermolecular forces stronger

δ+ δ+

δ-

δ- δ-

... ...

Cis Trans

Non polar

Non polar

Polar

Polar

Cis Trans

…......

…......

✓ ✓

RMM increases

Number electron

VDF increases

H2O

•2 hydrogen atoms

•2 lone pairs on oxygen

4 Hydrogen bonding available

Evidence for Hydrogen Bonding

Boiling Point group 4 Hydrides

Group 4, 5, 6, 7 Hydrides

Group 4 – SiH4, GeH4, SnH4

Group 5 – PH3, AsH3, SbH3

Group 6 – H2S, H2Se, H2Te

Group 7 – HCI, HBr, HI

High boiling point for

NH3, HF, H2O due to hydrogen bonding

H2O HF

HF

•1 hydrogen atom

•3 lone pairs on fluorine

Lack hydrogen atoms for hydrogen bond

NH3

NH3

•3 hydrogen atoms

•1 lone pairs on ammonia

Lack lone pairs for hydrogen bond

CH3-C=O CH3-C-CH3 CH3C-O-H CH3-O-CH3 CH3-N-H CH3-C-N-H

I II II I II

H O O H O

Which of the following molecules are polar/non polar?

ICI BCI3 CH2CI2 SF6 NF3 CO2

δ+ δ- δ+ δ-

δ-

δ-

δ+

δ+

δ+ δ+

δ-

δ- δ-

δ-

δ-

δ-

δ-

δ- δ-

δ-

δ-

δ- δ-

Polar bond Polar Bond Polar Bond Polar Bond Polar Bond Polar Bond

Bond Polarity Bond Polarity Bond Polarity Bond Polarity Bond Polarity Bond Polarity

exist cancel exist cancel exist cancel

(Asymmetric) (Symmetrical) (Asymmetrical) (Symmetrical) (Asymmetrical) (Symmetrical)

✓ ✓ ✓

Which of the following molecules have hydrogen bonding?

CH3CHO CH3COCH3 CH3COOH CH3OCH3 CH3NH2 CH3CONH2

Hydrogen NOT No Hydrogen Bond Hydrogen attach No Hydrogen Bond Hydrogen attach Hydrogen attach

attach to N,O,F to N, O, F to N, O, F to N, O, F

✕ ✕ ✓ ✓ ✓ ✕

✕ ✕ ✕

Education

IB Chemistry on Polarity, Hydrogen Bonding and Van Der Waals forces