States of Matter

1. States of Matter • We have seen a general law to explain physical properties of gases. • There is no general law for the behaviour of condensed states • Molecules are tightly packed • Intermolecular forces cannot be ignored INTERMOLECULAR ATTRACTIONS Intermolecular forces cause compounds to exist in solid, liquid, or gas and affect properties such as the mp and bp of compounds as well as the solubility of one substance in another (later). Solid: not compressible nor does it flow (shape and volume). Liquid: not compressible but can flow (volume). Gas: compressible and can flow.

3. Solids-Ionic and Metallic Electrostatic forces - these forces occur between charged species and are responsible for the extremely high melting and boiling points of ionic compounds and metals. Metal ions interact with the sea of electrons that surround them. This attraction is strong (bp of calcium 1484C). Substances which bear full charges, anions and cations, are attracted very strongly (bp NaF 1695C) Bond strengths for these solids are normally >1000 kJ/mol

4. Liquids: Intermolecular ForcesIon-dipole Important in solvation of salts to yield solutions. Strength~40-600 kJ/mol

5. Dipole-dipole interactions Polar covalent molecules contain regions that are rich and poor in electrons. The electron-rich regions of one molecule can attract the electron-poor region of an adjacent molecule. Strengths of ~5-25kJ/mol

6. Hydrogen Bonds Hydrogen bonds are abnormally strong dipole-dipole attractions that involve molecules with -OH, -NH, or FH groups. H atoms are very small (r= 37 pm, they're smaller than any other atom but helium). When a bonded highly electronegative atom (oxygen, nitrogen, or fluorine) pulls electrons away from the hydrogen atom, the positive charge that results is tightly concentrated. The hydrogen is intensely attracted to small, electron-rich O, N, and F atoms on other molecules. Strength for water=19kJ/mol: liquid at STP!!

7. Biochemistry and hydrogen bonds

8. DNA

9. Dipole-induced dipole A polar molecule can also induce a temporary dipole in a non-polar molecule. The electron cloud around a non-polar molecule responds almost instantaneously to the presence of a dipole, so this "dipole-induced dipole" force isn't as orientation-dependent as the dipole-dipole interaction. Strength ~2-10kJ/mol

10. BP for alkanes Alkanes: non-polar. Increasing alkane chain length gives higher bp. There must be some other intermolecular force!

11. Instantaneous dipole-induced dipole Random instantaneous imbalances in the distribution of electrons in non-polar molecules creates instantaneous dipoles that can induce a dipole on a neighbouring non-polar molecule. The net force is attractive and sometimes called the London force (in honour of their discoverer). London forces always contribute to intermolecular attractions. Usually weaker than other forces. Explain intermolecular forces between nonpolar molecules or noblegas atoms. Molecules with large, diffuse electron clouds can have London forces that are as strong as other forces.(hexadecane is a solid at room temperature!).

12. Intermolecular Forces

13. Properties that are dependent on intermolecular forces shape and volume surface tension Tendency of liquids to reduce their exposed surface to the smallest possible area. The molecules within the liquid are attracted equally from all sides, but those near the surface experience unequal attractions and thus are drawn toward the centre of the liquid. Because of surface tension, various small insects are able to skate across the surface of a pond, objects of greater density than water can be made to float. A steel needle placed carefully on water, for example, can be supported by the surface tension.

14. Wetting: The spread of liquid across a surface to form a thin film. Water on glass: glass is a solid silicon oxide network. On clean glass H-bonding between water and the surface overcomes the energy needed to expand the surface area. On greasy glass the water will try to minimise contact with the non-polar surface. Beads form. To prevent this add a a surface-active agent called a surfactant that has a non-polar “tail” and a polar “head” to dissolve the grease. The surfactant reduces the surface tension of the water.

16. Viscocity Resistance to change of form of a liquid. Proportional to the strength of intermolecular forces and to the ability of the molecule to entangle (polymers) Evapouration and sublimation Molecules with sufficient speed (kinetic energy) can escape the attraction of neighbours and evapourate (or sublimate). Equilibrium: lg, sg

17. Vapour Pressure • In a closed container, the pressure exerted by evapourated molecules at equilibrium is called the vapour pressure (vp). • Equilibrium vapour pressure: rate of evapouration is equal to rate of condensation • The factors that affect the vapour pressure are • Temperature (see Maxwell’s distribution on previous slide) • Chemical composition (intermolecular forces!)

18. Vapour Pressure and Heat What happens to a solid substance when it is heated? The compound can simply get hotter or a phase change can occur. The transition from the solid phase to the liquid phase is an example of a phase change, which is often called melting. Boiling or vapourisation is an example of a phase change from the liquid to the gas phase. Phase changes can be expressed as enthalpy changes at constant temperatures (Claussius-Clapeyron equation).

19. Boiling points

20. The boiling point of a liquid can now be defined as the temperature at which the vapour pressure of the liquid is equal to the prevailing atmospheric pressure. • At 1atm pressure: Normal Boiling Point (100C for water) • At 1bar pressure: Standard Boiling Point (99.6C for water; 1bar=0.987atm)

21. Phase diagrams Show regions of P and T at which various phases are thermodynamically stable Triple point: three phases in equilibrium Vp curve for solid Vp curve for liquid Melting point line Critical point

22. Heating curves

23. Supercritical fluids Heating liquid in a closed vessel does not produce boiling. The vp (density) of the vapour rises with increasing T, as the liquid density decreases, until both are equal and a single phase exists (neither liq nor vap.).

24. Simple water phase diagram Note melting line for solid Skating!

25. Water Density and Ice Maximum density of pure water at a pressure of 1 atm 999.861 kg/m³ at 3.98°C (277.13 K). 931 kg/m3 Aquatic life at sub-zero T, iceberg etc.

26. Phase diagrams Water Carbon dioxide Triple pt: 6.11mbar, 273.16K Critical pt: 215bar, 647.3K Triple pt: 5.11bar, 218.8K Critical pt: 72bar, 304.2K

27. CO2 Dry ice fog-special effects Supercritical fluids Caffeine extraction from coffee beans Dry-cleaning Polymerisations Chromatography

28. Carbon allotropes