CHM2S1-AIntermolecular Forces Dr R. L. Johnston

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CHM2S1-A Intermolecular Forces Dr R. L. Johnston

• Handout 2 The Importance of Intermolecular
  Forces
• III Intermolecular Forces in Action
• Consequences of Intermolecular Forces
• Anomalous Properties of Water
• The Hydrophobic Effect
• Protein Structure

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8. Consequences of Intermolecular Forces

• 8.1 Real Gases
• Ideal (or perfect) gas equation of state
• where R (the gas constant) 8.3145 J K?1 mol?1.
• Assumptions (1) atoms/molecules have no size
• (2) there are no interactions between the
  atoms/molecules
• Real (imperfect or non-ideal) gases dont obey
  this equation, due to the failure of both
  assumptions.

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p,V isotherms for an ideal gas
Note that an ideal gas can never liquify, however
low the temperature. The closest we get is He
for which TBP 4.2 K
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• van der Waals equation of state
• Introducing the molar volume, Vm V/n, this
  becomes
• van der Waals coefficients a, b gt 0.
• a measures strength of attractive interactions
  between molecules
• b measures volume of molecules
• These equations have the form peff.Veff nRT
  

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Ideal and Real (Non-Ideal) Gases
In real (non-ideal) gases, we allow for both
non-zero intermolecular forces and non-zero size
of molecules.
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• Pressure
• b reduction of available volume for molecules
  to move in, due to non-zero size of molecules.
  (Takes account of repulsive forces by modelling
  molecules as hard spheres). Less volume to move
  in ? more frequent collisions between molecules ?
  pressure increases.
• a attractive long range interactions between
  molecules lead to a decrease in the frequency and
  the force of collisions between molecules ?
  pressure decreases.
• Note
• at high T or high Vm, vdW equation ? perfect gas
  equation
• liquid and gas coexist when p 0 (when 2 terms
  in equation balance).

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p,V Isotherms for a van der Waals Gas
vdW gases can only liquify for T ? Tc
(independent of p).
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From the vdW equation, the following expressions
can be derived Tc 8a / 27Rb Vc 3b
pc a / 27b2 i.e. the lower a (or higher
b), the lower the temperature needs to be for
liquids to form. e.g. CO2 Tc (observed)
304 K Tc (predicted by vdW) 300 K.
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Comparison of van der Waals coefficients
Gas a / Pa m6 mol?2 b / 10?5 m3 mol?1 ? / kJ mol?1 Tb / K
He 0.004 2.370 0.1 4
Ar 0.138 3.219 1.2 87
Xe 0.431 5.105 2.1 165
H2 0.025 2.661 0.3 20
N2 0.143 3.913 0.9 77
CO2 0.369 4.267 2.0 (subl.) 195
CH4 0.231 4.278 1.3 112
C6H6 1.848 11.54 3.1 353
H2O 0.561 3.049 20.0 373

• More polarisable molecules behave in a less ideal
  manner, due to larger dispersion forces
  (reflected in a larger van der Waals a factor).
• Magnitude of b correlates to the size of the
  molecule.

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• There are a number of other, more accurate
  equations of state. Here, we will mention only
  one other.
• Virial equation of state
• B second virial coefficient
• C third virial coefficient
• B, C depend on T. B is more important than C
  (B/Vm ? C/Vm2).
• B has units of cm3 mol?1.

B (273 K) B (600 K)
Ar ?21.7 11.9
CO2 ?149.7 ?12.4
N2 ?10.5 21.7
Xe ?153.7 ?19.6
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• Gas Compressibility
• Ideal (perfect) gases Z 1
• Real gases
• v. low p Z ? 1 molecules far apart ? weak
  interactions ? behaves like perfect gas
• medium p Z lt 1 attractive forces dominate ?
  easier to compress
• high p Z gt 1 repulsive forces dominate ? harder
  to compress

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Gas Compression Factor (Z)
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• 8.2 Non-Ideal Solutions and Mixtures
• Ideal Solutions obey Raoults Law
• pA partial vapour pressure of A in liquid
  mixture
• pA vapour pressure of pure liquid A
• xA mole fraction A in liquid mixture.
• Total pressure, p
• In terms of chemical potentials (?) we can write
• Raoults law implies that all interactions A?A,
  B?B, A?B are the same (i.e. UAA UBB UAB).
• Note does not assume no interactions, but ?mixH
  0.
• Raoults law is obeyed well by mixtures of
  similar (shape and bonding) molecules e.g.
  benzene/toluene.

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• Non-Ideal Solutions strong deviations from
  ideality (positive or negative) shown by mixtures
  of dissimilar liquids e.g. CS2/acetone (UAA ?
  UBB ? UAB).
• UAB gt UAA , UBB ? ?mixH lt 0 (exothermic mixing)
• negative deviation
• UAB lt UAA , UBB ? ?mixH gt 0 (endothermic mixing)
• positive deviation
• In terms of chemical potential
• where aA is the activity ( effective mole
  fraction) of liquid A in the mixture.

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Vapour Pressures of Solutions
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• 8.3 Other Consequences of IMFs
• Different phases adopted by various elements and
  compounds.
• Structures of solids and liquids.
• Liquid crystals unusual properties due to
  anisotropic intermolecular interaction (e.g.
  disk-like or cigar-shaped molecules).
• Transport properties (viscosity, thermal
  conductivity, diffusion).
• Properties of electrolyte solutions (solvated
  ions).
• Supramolecular chemistry (aggregation,
  self-ordering, molecular recognition, protein
  folding, drug-protein interactions, DNA ).

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• 8.4 Some Experimental Techniques for
  Investigating IMFs
• Molecular Beams study collisions and scattering
  between individual molecules.
• X-ray and neutron diffraction determine long
  range structures of crystalline solids and short
  range structure of liquids.
• Spectroscopy determine structures, binding
  energies and electronic, vibrational and
  rotational energies of loosely bound van der
  Waals molecules.
• Measurement of gas imperfection e.g. pV
  isotherms, Joule-Thomson effect, compressibility.
• Measurement of solution non-ideality deviations
  from Raoults law and Henrys law.
• Measurement of transport properties
• Atomic Force Microscopy direct measurement of
  intermolecular forces between surfaces and
  adsorbed molecules.

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9. Anomalous Properties of Water

• 9.1 Water
• Water is the most abundant liquid on Earth.
• But it is considered to be anomalous because
  it behaves differently from simple liquids (e.g.
  Ar).
• Differences are due to hydrogen bonding in water.
• The water molecule is small and compact, with two
  H atoms and two lone pairs arranged tetrahedrally
  around the O atom
• The dipole moment (?) of the isolated water
  molecule is 1.85 D.
• Water forms hydrogen bonds the O?H bonds act as
  H-bond donors and the O lone pairs act as H-bond
  donors.
• Each water molecule can take part in up to 4
  H-bonds.

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• In the gas-phase optimal H-bond bond strength
  between water molecules 23 kJ mol?1.
• H-bonding in condensed phases of water is
• cooperative (non-additive) the strength of
• H-bonding increases with increasing number
• of water molecules, as this increases the
• polarization of the O?H bonds.
• This shows up in the increase in the average
  dipole moment per water molecule, which increases
  from 1.85 D (isolated H2O) to 2.4?2.6 D (liquid
  H2O at 0?C).

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Phase Diagram for Water
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• 9.2 Ice (Solid Water)
• The structure of ice is based on tetrahedral
  coordination of the water molecules, which each
  take part in 4 H-bonds.
• There are a number of different solid ice phases.
  At 1 atm. the most stable form is hexagonal ice
  Ih.

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• 9.3 Liquid Water
• For water, the liquid is more dense than the
  solid (ice). This is in contrast to most
  liquids. Maximum density of liquid is at around
  4?C. Above 4?C, water behaves like other liquids
  expanding as it gets warmer.
• This is due to the disruption of the long-range
  ordered tetrahedral network in liquid water. The
  average number of nearest neighbours around each
  H2O molecule increases from 4 to approx. 4.4 on
  melting.
• There is a fluctuating network of H-bonds in
  liquid water.
• Higher densities are favoured by increasing van
  der Waals (D-D and dispersion) interactions,
  though H-bonding favours lower coordination and
  lower density. ? on melting, the H-bonding is
  weaker but the vdW bonding is stronger.
• Consequences ice-bergs burst water pipes in
  winter

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• Applying pressure to ice causes melting.
  According to the Clapeyron equation
• ? every 133 atm. of applied pressure, decreases
  the melting temperature of ice by 1 K.
• This may contribute to enabling ice skating!

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• Other properties of liquid water
• Liquid water is less compressible than ice.
  Compressibility decreases with T until 46?C.
• Liquid water has a high dielectric constant
• (because H-bonds are polarizable) so it is
• a good solvent for ions.
• H-bonding leads to higher cohesive energies than
• for similar-sized molecules (especially compared
  with
• H2X molecules from the same group) ? relatively
  high
• boiling and melting points (same true for HF and
  NH3).
• The extended H-bonded network in liquid water
  leads to rapid transfer of H and OH? ? changes
  of pH move rapidly through aqueous solutions.
• Water has a high enthalpy (40 kJ mol?1) and
  entropy of vaporization (109 JK?1 mol?1),
  indicating that the liquid still has quite a lot
  of the order (and cohesion) of the solid ? water
  has a very high liquid range (100 K). This is
  critical for life on Earth!

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Comparison of boiling points of group 16 and
group 18 hydrides
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10. The Hydrophobic Effect

• 10.1 Definitions
• Hydrophobic Effect The low solubility of
  hydrocarbons and other non-polar molecules in
  water and their increased tendency to aggregate.
• Hydrophobic Interaction Enhanced effective
  attractions between hydrocarbon molecules etc.,
  when in water.
• Simple enthalpy explanation immiscibility (lack
  of solubility) of solute B in solvent A occurs
  when the A-B interactions are weaker than the A-A
  and B-B interactions (UAB lt UAA, UBB).
• This might be expected to be the case for B
  hydrocarbon (quite strong dispersion forces
  between long chain hydrocarbons) and A water
  (strong H-bonds), with A-B interactions being
  primarily dipole-induced dipole in nature
  (relatively weak).
• BUT this does not explain why solubility of oil
  in water, as a function of T, goes through a
  minimum at T ? 25?C. (Normally expect solubility
  ? as T ?).

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• 10.2 Origin of the Hydrophobic Effect
• Note the overall enthalpy of interaction of a
  non-polar solute with water is not particularly
  unfavourable (?H ? 0) because the non-polar
  molecules induce cage-like ordering of the first
  shell of water molecules, strengthening their
  H-bonding.
• The origin of the hydrophobic effect is mostly
  entropic.
• The ordering of the shell of water molecules
  around the hydrocarbon solute (so as to minimise
  dangling H-bonds), causes a significant
  decrease in the entropy of the water (?S lt 0).
• Typically, the total change in entropy in
  dissolving small hydrocarbon molecules in water
  (at 298 K)
• ?S ? ?100 J K?1 mol?1
• For T lt 25?C, entropy term dominates and becomes
  more unfavourable with increasing T ? solubility
  decreases as T rises.
• For T ? 25?C, the water cages start to break up
  (weakening H-bonds) so ?H, ?S increase ?
  solubility increases as T rises (enthalpy starts
  to dominate).

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• 10.3 Clathrates single hydrocarbon or other
  non-polar molecules (even small ones, such as CH4
  and CO2) surrounded by a polyhedral cage of water
  molecules.
• At high P, low T, these clathrates can
  precipitate out as solids.
• Examples CH4-H2O clathrates in oil pipelines.
• CO2-H2O clathrates in deep ocean sites.

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• 10.4 Micelles (examples of colloids)
  pseudo-spherical clusters of surfactant molecules
  consisting of hydrophilic heads (polar or
  charged groups) and hydrophobic tails
  (hydrocarbon chains) dispersed in water.
• Hydrophobic tails aggregate together (dispersion
  forces) this also minimizes the unfavourable
  hydrophobic entropy effect on the solvent
  (water). Centre of micelle is oil-like.
• Hydrophilic heads form a close-packed shell and
  have strong intermolecular interactions with the
  water molecules.
• Sizes range from 100s (charged heads) to 1000s
  of molecules.
• Used to solubilize hydrocarbons in aqueous
  solution
• e.g. detergents, drug carriers, organic
  synthesis, petroleum recovery.
• Analogous to biological membranes.

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Clathrates and Micelles
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11. Protein Structure

• 11.1 Proteins natural polymers polypeptides
  chains of amino acids (H2NCHRCO2H) joined by
  peptide links ?CO-NH?.
• Protein folding the folding up of the
  polypeptide chains under the influence of
  intermolecular forces.
• Note the chemical function of the protein is
  dependent on its 3D structure which depends on
  its folding.
• Primary structure sequence of amino acids.
• Secondary structure coiling into ?-helices or
  folding into ?-sheets, due to N?H?OC
  hydrogen-bonding between peptide groups (which
  are close in the sequence).
• Tertiary structure folding of the polypeptide
  chain by forming interactions (e.g. covalent
  disulfide ?S?S? links, ionic interactions,
  H-bonds) between side chains (R) of amino acids
  which are relatively far apart in the sequence.
  In aqueous solution, the hydrophobic effect may
  also be important.

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• Quaternary structure aggregation of more than
  one polypeptide chain due to similar interactions
  to those responsible for tertiary structure.
• Protein aggregation
• sometimes beneficial (e.g. for the function of
  haemoglobin an aggregate of 4 polypeptide
  chains)
• sometimes harmful (e.g. in protein misfolding
  diseases such as BSE, CJD).

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Secondary Protein Structure
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Secondary and Tertiary Protein Structure
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• 11.2 The hydrophobic effect in protein folding
• Globular proteins in aqueous solution have
  pseudo-spherical shapes
• cores rich in hydrophobic residues (amino acids
  with non-polar alkyl or aryl side-chains, R)
• outer shell rich in hydrophilic residues (polar
  side chains).
• Protein folding is partially driven by the
  hydrophobic effect burying ? ½ of hydrophobic
  residues reduces the unfavourable decrease in
  entropy of the surrounding water molecules.