1. Chapter 11 States of Matter; Liquids and Solids Dr. Peter Warburton peterw@mun.ca http://www.chem.mun.ca/zcourses/1011.php

2. © Peter Warburton 2008 All media copyright of their respective owners2 Gases Gases are compressible fluids. This behaviour arises because the gas molecules are in constant random motion through mostly empty space.

3. © Peter Warburton 2008 All media copyright of their respective owners3 Liquids Liquids are incompressible fluids. This behaviour arises because the molecules of the liquid are in constant random motion without much empty space to move around in.

4. © Peter Warburton 2008 All media copyright of their respective owners4 Solids Solids are incompressible and rigid. This behaviour arises because the molecules of the solid can only vibrate because they have almost no empty space to move into.

5. © Peter Warburton 2008 All media copyright of their respective owners5 Figure 11.2

6. © Peter Warburton 2008 All media copyright of their respective owners6 Ideal gas law We’ve treated gases in the past as behaving ideally, so that PV = nRT However, this gas law ASSUMES that molecules: 1)Have NO SIZE (no repulsive intermolecular forces) 2)DO NOT have attractive intermolecular forces

7. © Peter Warburton 2008 All media copyright of their respective owners7 All molecules have intermolecular forces (IMFs) Repulsive IMFs (real molecules have size) will limit the compressibility of a group of molecules. SQUEEZE!

8. © Peter Warburton 2008 All media copyright of their respective owners8 All molecules have intermolecular forces (IMFs) Attractive IMFs cannot be overcome at sufficiently low temperatures (molecules have low average kinetic energies). Lower T!

9. © Peter Warburton 2008 All media copyright of their respective owners9 Changes in state We can often change the physical state of a substance (called a phase transition) by changing the temperature and/or pressure by suitable amounts.

10. © Peter Warburton 2008 All media copyright of their respective owners10 Phase transitions From solid to liquid is fusion. Change in enthalpy (or heat) of fusion is  H fus From liquid to gas is vaporization. Change in enthalpy of vaporization is  H vap From solid to gas is sublimation. Change in enthalpy of sublimation is  H sub

11. © Peter Warburton 2008 All media copyright of their respective owners11 Phase transitions

12. © Peter Warburton 2008 All media copyright of their respective owners12 Enthalpy of phase transitions The enthalpy change of a phase transition tells us how much heat must be added to (or removed from) a substance in a given phase so it changes phase. Since ALL of the the energy is involved in the phase change, the temperature MUST REMAIN CONSTANT during the change.

13. © Peter Warburton 2008 All media copyright of their respective owners13 Enthalpy of phase transitions

14. © Peter Warburton 2008 All media copyright of their respective owners14 Vapour pressure Vapour pressure is the partial pressure (the part of the total pressure that comes from a given substance) of the vapour (gas) above the liquid (or solid) phase measured at EQUILIBRIUM at a GIVEN TEMPERATURE

15. © Peter Warburton 2008 All media copyright of their respective owners15 Equilibrium and vapour pressure Vapour pressure should be measured when the vapour pressure has STOPPED CHANGING. This equilibrium means that the rate of molecules leaving the liquid (or solid) phase is BALANCED EXACTLY by the rate of the molecules in the vapour joining the liquid (or solid) phase.

16. © Peter Warburton 2008 All media copyright of their respective owners16 Equilibrium and vapour pressure

17. © Peter Warburton 2008 All media copyright of their respective owners17 Temperature and vapour pressure The vapour pressure depends on how easily molecules can overcome attractive intermolecular forces that keep it in the liquid (or solid) phase. Higher temperatures mean each molecule, on average, has more kinetic energy that COULD ALLOW the molecule to “escape” from the attractive forces.

18. © Peter Warburton 2008 All media copyright of their respective owners18 Temperature and vapour pressure

19. © Peter Warburton 2008 All media copyright of their respective owners19 Temperature and vapour pressure Since all groups of molecules AT THE SAME TEMPERATURE have the SAME AVERAGE KINETIC ENERGY, then molecules that have WEAKER intermolecular forces are MORE VOLATILE (have GREATER vapour pressures at the GIVEN temperature) than molecules with STRONGER intermolecular forces.

20. © Peter Warburton 2008 All media copyright of their respective owners20 Temperature and vapour pressure In this Figure, the MOST VOLATILE (weakest IMFs) liquid is on the left, while the LEAST VOLATILE (strongest IMFs) is on the right.

21. © Peter Warburton 2008 All media copyright of their respective owners21 Boiling point The boiling point of a liquid is the temperature where the vapour pressure IS THE SAME AS the external pressure. Therefore, if the external pressure changes, the boiling point temperature ALSO changes.

22. © Peter Warburton 2008 All media copyright of their respective owners22 Normal boiling point The normal boiling point of a liquid is the temperature where the vapour pressure IS THE SAME AS an external pressure of exactly 1 ATMOSPHERE. 1 atm = 760 mmHg = 101.325 kPa

23. © Peter Warburton 2008 All media copyright of their respective owners23 Freezing (or melting) point The freezing point of a liquid is the temperature where at which the phase transition from liquid to solid occurs. Since melting (fusion) is the exact opposite transition (solid to liquid) the freezing point and melting point are IDENTICAL.

24. © Peter Warburton 2008 All media copyright of their respective owners24 Freezing and boiling points Since the normal freezing and normal boiling points of a PURE substance are fixed properties, measuring them is an easy and useful first step in identifying unknown compounds.

25. © Peter Warburton 2008 All media copyright of their respective owners25 Intermolecular forces The forces between molecules are electrostatic. They depend on charges (like charges repel, opposite charges attract), and the distance between the charges. Larger charges (like those found on ions) and smaller distances between molecules tend to lead to stronger forces.

26. © Peter Warburton 2008 All media copyright of their respective owners26 Intermolecular forces At everyday conditions, the forces between molecules tend to be weakly attractive overall. These intermolecular forces are generally called van der Waals forces. However, vdW forces can be subdivided into two different groups.

27. © Peter Warburton 2008 All media copyright of their respective owners27 Dipole-dipole forces We’ve already seen that some molecules have a permanent molecular dipole. Since full charges are not involved in molecular dipoles, these dipole-dipole intermolecular interactions are relatively weak as compared to ionic bonds, where full charges are involved.

28. © Peter Warburton 2008 All media copyright of their respective owners28 Dipole-dipole forces The larger the dipole moment (the molecules are more polar), the stronger the IMFs tend to be.

29. © Peter Warburton 2008 All media copyright of their respective owners29 London (dispersion) forces Nonpolar molecules still interact with each other despite their lack of permanent dipoles. In a nonpolar molecule, “on average,” the electrons do not prefer one part of the molecule over the other. However, at any given instant, they might not be evenly distributed and so the molecule ends up having a “temporary dipole.”

30. © Peter Warburton 2008 All media copyright of their respective owners30 London (dispersion) forces When a molecule with a temporary dipole comes close to another molecule, the electrons of the second molecule will try to move away from the negative partial charge of the first molecule, leading to the second molecule having a temporary induced dipole.

31. © Peter Warburton 2008 All media copyright of their respective owners31 London (dispersion) forces London forces tend to be very weak because the partial charges tend to be small and fleeting. However, ALL chemical species have London forces between them. Variations in the strength of London forces depend on two factors.

32. © Peter Warburton 2008 All media copyright of their respective owners32 London (dispersion) forces Polarizability – is the ability for electrons to move freely within the molecule. The more freely electrons can move, the larger the induced dipole can be. Larger molecules and atoms are more polarizable, and generally have larger London forces.

33. © Peter Warburton 2008 All media copyright of their respective owners33 London (dispersion) forces Shape – The shape of a molecule plays a part in determining how the electrons can move in a molecule. More compact shapes are usually more symmetrical and allow less contact between molecules. They generally have smaller induced dipoles with weaker London forces. Stronger London forces Weaker London forces

34. © Peter Warburton 2008 All media copyright of their respective owners34 Hydrogen bonding A hydrogen bond is an attractive interaction between a hydrogen atom bonded to a very electronegative atom (O, N, and F), and an unshared electron pair on another electronegative atom.

35. © Peter Warburton 2008 All media copyright of their respective owners35 Hydrogen bonding Hydrogen bonds are really just a very special case of dipole-dipole forces. H-F, H-O, and H-N bonds are very polar (larger partial charges) Also, because the hydrogen is very small, it is possible for another molecule to approach it very closely (short distance). Hydrogen bonds are relatively strong intermolecular forces

36. © Peter Warburton 2008 All media copyright of their respective owners36 Hydrogen bonding (Figure 11.18) With London forces, boiling points will increase with molecular size (polarizability). We EXPECT boiling points to follow the trend CH 4 < SiH 4 < GeH 4 < SnH 4 and so on across the periodic table.

37. © Peter Warburton 2008 All media copyright of their respective owners37 Hydrogen bonding (Figure 11.18) This trend is generally true, except for NH 3, H 2 O, and HF because of the hydrogen bonds (stronger than London forces) that can occur.

38. © Peter Warburton 2008 All media copyright of their respective owners38 IMFs and bonding at a glance

39. © Peter Warburton 2008 All media copyright of their respective owners39 Problem 11.37 For each of the following substances, list the kinds on intermolecular forces expected: BF 3 CH 3 CHOHCH 3 HI

40. © Peter Warburton 2008 All media copyright of their respective owners40 Problem 11.43

41. © Peter Warburton 2008 All media copyright of their respective owners41 Classification of solids Ionic solids are formed by regular arrangements of cations (positive) and anions (negative). Because of the strong ionic bonds involved, the melting points are usually high, and the solids are brittle and hard. NaCl is an example.

42. © Peter Warburton 2008 All media copyright of their respective owners42 Classification of solids Molecular solids are separate molecules held together through the mainly weaker dispersion, dipole-dipole and hydrogen bond intermolecular forces. Weaker forces mean lower melting points and softer consistency. Also, since electrons can’t easily move from one molecule to another, the solids are nonconducting of electricity. Ice and sugar are examples.

43. © Peter Warburton 2008 All media copyright of their respective owners43 Classification of solids Covalent network solids are formed by a large array of atoms. These solids are formed from repeating units, but are almost better considered to be “one large molecule”. If the arrays are three-dimensional, the solids are rigid, and so the melting points are usually high, and the solids are hard but usually not brittle. Quartz and diamond are examples.

44. © Peter Warburton 2008 All media copyright of their respective owners44 Classification of solids Covalent network solids can also have structures of loosely held rigid sheets or chains. Graphite can be used as a lubricant because the sheets of carbon can easily slide past each other. Asbestos forms chains.

45. © Peter Warburton 2008 All media copyright of their respective owners45 Diamond and graphite

46. © Peter Warburton 2008 All media copyright of their respective owners46 Classification of solids Metallic solids are large arrays of metal nuclei found in an “electron sea”. Since the forces vary widely amongst metals we see widely variable melting points and hardnesses. However, since electrons can move from one atom to another, the solids are conducting of electricity and heat. Silver and iron are examples.

47. © Peter Warburton 2008 All media copyright of their respective owners47 Classification of solids

48. © Peter Warburton 2008 All media copyright of their respective owners48 Classification of solids