Autumn Lecture 4 (Intermolecular Forces)
Autumn Lecture 4 (Intermolecular Forces)
Autumn Lecture 4 (Intermolecular Forces)
Van der Waals (vdW) forces -evidence for van der Waals forces -temporary and induced dipoles -boiling points of alkanes
The hydrogen bond -formation of hydrogen bonds -effect of hydrogen bonding on b.p. of hydrides -important hydrogen bonded structures: ice, carboxylic acids, DNA Summary -three types of intermolecular forces -relative strength and influence on physical properties
Most bonds are between extremes: H d+ dC N polar covalent bonds electron density is not transferred between H atoms but is unequally shared H C CF3
H
ionic bond other extreme very different atoms so electrons completely transferred K+
Cl-
ionic bond
e.g. in H-F most electron density is on fluorine leaving H partially positive molecule is polar
F is more electronegative than H electronegativity the tendency of an atom to attract electron density towards itself in a chemical bond
The greater the difference in electronegativity between two elements, the greater the polarity of the bond
Cs+
+
P2O5 (difference = 1.4) completely covalent Li LiI (difference = 1.5) partly ionic, partly covalent Al2O3 (difference = 2.0) mostly ionic, slightly covalent CsF (difference = 3.3) completely ionic
F -
electron-poor carbon
d-
-computer modelling alows electron rich (d-) and electron poor (d+) regions of molecule to be represented using colours of the rainbow
d+
d+
delectron-rich carbon
Bond Polarity (9.1) Q: One bond is indicated (by ) in each of the following: Indicate the polarity of each bond in each case:
H3CI H3CB(CH3)2
OCH2
BrMgCH3
(use d+ and d- to show electron density) Which molecules have zero dipole moment?
Bond Polarity (9.1) Q: One bond is indicated (by ) in each of the following: Indicate the polarity of each bond in each case:
(use d+ and d- to show electron density) Which molecules have zero dipole moment?
Molecule Polarity (9.1) -bond polarity can cause the entire molecule to become possess a dipole moment (m) e.g.
H C
dichloromethane (m = 1.6 D)
difluoromethane (m = 2.3 D)
H C
Cl Cl
F F
Molecule Polarity (9.1) -bond polarity can cause the entire molecule to become possess a dipole moment (m) e.g.
H C
dichloromethane (m = 1.6 D)
difluoromethane (m = 2.3 D)
H C
Cl Cl
F F
-but sometimes dipoles within molecule cancel each other due to symmetry e.g.
m=0D
Cl C Cl
m=0D
F Cl F B F
Cl
Be
Cl Cl
-electrostatic attraction between opposite charges on neighbouring molecules (molecules arrange themselves so mutual attraction occurs lower energy)
b.p. 180 C
b.p. 174 C
o-dichlorobenzene is said to be polar BUT p-dichlorobenzene is non-polar as electron density is symmetrical (greater thermal energy needed to separate molecules bound by dipole-dipole forces)
b.p. 180 C
b.p. 174 C
o-dichlorobenzene is said to be polar BUT p-dichlorobenzene is non-polar as electron density is symmetrical dipole-dipole attraction explains higher boiling point of o-dichlorobenzene
first isomer has greater dipole-dipole attractions - more heat needed to separate molecules into vapour higher boiling point:
db.p. 61 C
d+
b.p. 47 C
polar
non-polar
d+
d-
H3C F
-for small molecules presence of a polar bond can change the boiling point by a hundred degrees!
van der Waals forces (9.3B) -BUT even non-polar molecules have some force between them -evidence: b.p. of noble gases and tetra-substituted methanes
CH4
-b.p. -162C
CF4
-129C
CCl4
77C
CBr4
190C
even non-polar
molecules attract each other slightly
bigger molecules > more electrons > more polarisable > bigger dipoles > higher boiling points: CH4 CF4 CCl4 CBr4
b.p.
-162C
-129C
77C
190C
- all atoms and molecules interact with London forces to some extent
-extent of van der Waals force depends on: -no. of electrons -surface area
-boiling points of alkanes depend mainly on intermolecular van der Waals forces
n-pentane (C5H12)
H
H C H H H H C C C H H H H C H H
2-methyl-butane (C5H12)
neopentane (C5H12)
-isomers have the same empirical formula (C5H12), same mass -extent of van der Waals forces depends on surface area
H C H H H H C C C H H H H C H H
-isomers have the same empirical formula (C5H12), same mass -extent of van der Waals forces depends on surface area
e.g. n-pentane has greatest area greatest v der W forces highest boiling point
-H atoms have strong partial positive charge (d+) when attached directly to 2nd period electronegative atom
-water molecules held in hydrogen bonded hexagonal lattice -open cage structure low density ice floats
-hydrogen bonding much stronger than other dipole-dipole interactions - drastically raises boiling points: (actual b.p. with hydrogen bonding)
no hydrogen bonding
(fig 9.10)
carboxylic acid
H2N
C H2
H2 C
H2N
C H2
H2 C
a-helix
Hydrogen Bonding in Proteins (9.8D) -when protein chains align a b-sheet structure results:
b-sheet
-two dimensional creased sheet like arrangement
-two strands of DNA hydrogen bond to and coil around each other forming a double helix:
(base pairs)
(base pairs)
liquid
gas
-permanent dipoles are stronger than induced dipoles (v der Ws) -hydrogen bonds are stronger than dipole-dipole interactions
Intermolecular interactions in solutions ion-dipole (15-40 kJ/mol): when an ion is in a polar solvent, solvent molecules surround the ion burying their oppositely charged parts against the ion solvation of the ion
hydrated cation
compare: ion-ion interaction (200-400 kJ/mol): strong attraction between oppositely fully charged ions
determine
strength of intermolecular forces present
determines
compounds
physical properties
e.g.
solubility boiling point
d+
H3C
dCl
Cl CH3
d-
d+
dO d+ H CH3 H d+ O d-