Liquids

1. Liquids Polar bonds and dipoles Intermolecular forces Liquid properties Phase changes Evaporation, vapour pressure and boiling point Clausius-Clapeyron equation

2. Intermolecular forces • In the sequence gas → liquid → solid • Intermolecular attractions increase • Gases – essentially no interactions • Liquids – movement allowed • Solids – completely rigid

3. Polarity redux • Electronegativity differences between atoms creates polar bonds – the more electronegative atom attracts the electrons

4. Molecular dipole • Molecules are assemblies of several bonds • Molecular polarity depends on the orientation of the individual dipoles • If the dipoles cancel out, molecular is non-polar • If the dipoles don’t cancel, the molecule is polar

5. Symmetry and polarity

6. - + Studying bonds is an approximation • We can calculate the centers of gravity of the negative and positive charges in a molecule • If they do not coincide, the molecule is polar • These calculations are involved, so studying individual bonds is a good approximation

7. Dipole moments • The dipole moment is the charge x length of the dipole • An electron and proton separated by 0.1 nm (a typical bond length) • Where 1 D (Debye) = 3.336 x 10-30 Cm

8. Algorithm for predicting molecular polarity • Establish molecular skeleton • Draw Lewis dot structure • Count groups of charge around central atom • Establish electronic geometry using VSEPR • Determine molecular shape • Identify polar bonds and lone pairs • Inspect molecule: do polar bonds/lone pairs cancel out?

9. Percent ionic character • We have seen that we can calculate the dipole moment for a given charge separation • Comparison with experimental values permits estimation of “ionic character” • In HCl the experimental dipole moment is 1.03 D. • The theoretical dipole given the bond length of 0.127 nm is 6.09 D • Percent ionic character = 1.03/6.09 x 100 % = 16.9 %

10. May the force be with you • Covalent and ionic bonds are the intramolecular forces that hold the atoms in molecules together • Intermolecular forces hold the molecules together • Collectively, the intermolecular forces are called van der Waals forces • All arise from electrostatic interactions

12. Ion - dipole • Characteristic of interactions in solutions of ionic compounds in polar solvents • Negative ion with the positive dipole end • Positive ion with the negative dipole end

13. Dipole - dipole • Important attractive force in polar substances • Strength of the order of 3 – 4 kJ/mol (compared with 200 – 400 kJ/mol for covalent bonds)

14. Manifested in boiling points: • Nonpolar substances have much lower boiling points • Acetone (polar) 56ºC butane (nonpolar) -0.5ºC • Boiling point increases with dipole strength

15. London calling • Even molecules with no net dipole moment attract each other. • Electrons are not static but mobile: • Fluctuation creates dipole in one molecule which induces dipole in another molecule • Effect increases with atomic number – as atom becomes more polarizable • Boiling increases with atomic weight • Conventionally, dispersion forces are said to be weaker than other inter-molecular forces. For large molecules this is not really true. Large molecules are solids because of dispersion forces

16. Hydrogen bonds: the most important bond? • Key to life • Between H and O, N or F • Dipole-dipole bonds of unusual strength (up to 40 kJ/mol)

17. Hydrogen bonding • The ultimate expression of polarity • Small positive H atom exerts strong attraction on O atom • Other H-bonding molecules: HF, NH3 • H2O is the supreme example: two H atoms and two lone pairs per molecule

18. H2O has optimum combination of lone pairs and H atoms

19. H bonding generates three-dimensional network

20. Water: the miracle • All the properties of water that make it unique and life sustaining can be traced to hydrogen bonding • Density of ice lower than water • Anomalous high b.p. • High heat capacity • Universal solvent

21. Understanding the force • Predicting the forces acting between molecules means understanding the molecules • All molecules experience London forces, but only some will have dipole-dipole or hydrogen bonds. Where present, the latter will dominate

22. Properties of liquids depend on intramolecular forces • Water flows but syrup is sticky • Viscosity measures resistance to flow • Small non-polar molecules flow easily • Large or highly polar molecules flow less easily • Units of viscosity are kg/m-s

23. Surface tension? Take a tablet • Surface tension is the tendency of a liquid to resist spreading out • Arises from molecules at the surface experiencing inward pull • Walking on water: it’s no miracle, it’s surface tension • Surface tension is the energy required to increase the surface area of a liquid – units are J/m2

24. Cohesive and adhesive • Cohesive forces are the attractive forces between like molecules • Adhesive forces are the attractive forces between unlike molecules

25. Meniscus • Adhesive forces pull H2O molecules to maximize coverage • Cohesive forces between H2O molecules drag liquid up • Gravity pushes liquid down

26. Capillary action • Combined effects of cohesive, adhesive and gravitational forces cause liquid to rise towards edge of container • In very thin columns the effect of gravity is diminished and the liquid rises higher • Originally used as explanation (incorrect) for transport of water through plants (Osmosis is the cause)

27. Just a phase I’m going through • A phase change occurs when matter changes from one state to another • Solids can exhibit more than one phase which also undergo phase changes (gray tin to white tin)

28. Energetics of phase changes • In the series: solid → liquid → gas • Energy is required to break intermolecular forces • Distribution of molecules is more disordered (entropy) – greater disorder is more favourable

29. Roadmap of changes • More condensed to less condensed means heat absorption and entropy gain which are opposing

30. Phase changes involve “latent” heats • With matter in a single phase, heating the substance gives a T increase depending upon S.H. • At a phase change, two phases are in equilibrium and heat is absorbed to convert one into the other without a change in T. Hence the term “latent” heat – a term no longer in popular use.

31. Fusion versus vaporization • For all substances, the heat of vaporization is much larger than the heat of fusion • More bonds are broken in creating the vapour

32. Vapour pressure • Liquids do not turn into a vapour only at the boiling point • At any temperature, there is vapour in equilibrium with the liquid • A puddle of water on the sidewalk evaporates • A liquid develops a pressure in a manometer • The pressure exerted by the vapour in equilibrium with the liquid is the vapour pressure

33. Maxwell, Boltzmann and vapour pressure • Molecules exhibit a range of energies, which moves to higher energy as T increases • More molecules have sufficient energy to escape liquid as T increases • When the vapour pressure = atmospheric pressure, the liquid boils

34. Contrasting volatile and non-volatile liquids

35. Clausius – Clapeyron equation • The vapour pressure in equilibrium with a liquid obeys the following equation • Calculate ΔHvap from vapour pressure data • Calculate vapour pressure as f(T) given ΔHvapand one vapour pressure value