Intermolecular Forces

1. Intermolecular Forces Mr. Claus Modified from Holt 2006

2. Mercury in Three States

3. Liquid Forces Lead to Surface Wetting and Capillary Action • Liquid particles can have cohesion, attraction for each other. • Liquid particles can also have adhesion, attraction for particles of solid surfaces. • The balance of cohesion and adhesion determines whether a liquid will wet a solid surface. • The forces of adhesion and cohesion will pull water up a narrow glass tube, called a capillary tube.

4. Intermolecular forces and surface tension • Below the surface of a liquid, the particles are pulled equally in all directions by cohesive forces. • However, surface particles are pulled only sideways and downward, so they have a net downward force. • It takes energy to oppose this net force and increase the surface area. • The tendency of liquids to decrease surface area to the smallest size possible is surface tension.

5. Changes of State • Most substances can undergo six changes of state: freezing, melting, evaporation, condensation, sublimation, and deposition. • Generally, adding energy to a substance will increase the substance’s temperature. • But after a certain point, adding more energy will cause a substance to experience a change of state instead of a temperature increase.

6. Changes of State

7. Liquids evaporate to a gas…. • Energy is required to separate liquid particles. They gain energy when they collide with each other. • If a particle gains a large amount of energy, it can leave the liquid’s surface and join gas particles. • Evaporation is the change of state from liquid to gas. Evaporation is an endothermic process. • Boiling point is the temperature and pressure at which a liquid and a gas are in equilibrium. • Water latent heat of vaporization (enthalpy of vaporization): 40.68 kJ/mole or 40,680 J/mol or 2,260 J/gram or 540 calories/gram

8. Gases condense to liquids • When gas particles no longer have enough energy to overcome the attractive forces between them, they go into the liquid state. • Condensation is the change of state from a gas to a liquid. Condensation is an exothermic process - Water latent heat of condensation (enthalpy of condensation): 40.68 kJ/mole or 40,680 J/mol or 2,260 J/gram or 540 calories/gram • Condensation can take place on a cool night, causing water vapor in the air to form dew on plants.

9. Solids melt into liquids • As a solid is heated, the particles vibrate faster and faster in their fixed positions. • At a certain temperature, some of the molecules have enough energy to break out of their fixed positions. • Melting is the change of state from solid to liquid. Melting is an endothermic process. • Melting point is the temperature and pressure at which a solid becomes a liquid.

10. Liquids freeze to solids • As a liquid is cooled, the movement of particles becomes slower and slower. • At a certain temperature,the particles are pulled together into the fixed positions of the solid state. • Freezing is the change of state from a liquid to a solid. Freezing is an exothermic process. • The Latent Heat of Fusion of Water is 334.7 J/gram or 6009 J/Mole, or 6.0 kJ/Mole, or 80 calories/gram – 1440 calories/mole • Freezing point is the temperature at which a substance freezes.

11. Solids sublime to gas • The particles in a solid are constantly vibrating. Some particles have higher energy than others. • Particles with high enough energy can escape from the solid. • Sublimation is the change of state from solid to gas. • Sublimation is an endothermic process.

12. Gas deposits to solid • Molecules in the gaseous state become part of the surface of a crystal. • When a substance changes state from a gas to a solid, the change is often called deposition. • Deposition is an exothermic process.

13. Comparing ionic to covalent compounds • It takes energy to overcome the forces holding particles together. • Thus, it takes energy to cause a substance to go from the liquid to the gaseous state. • The boiling point of a substance is therefore a good measure of the strength of the forces that hold the particles together. • Melting point also relates to attractive forces between particles.

14. Most covalent compounds melt at lower temperatures than ionic compounds do.

15. Oppositely Charged Ions Attract Each Other • Ionic substances generally have much higher forces of attraction than covalent substances. • For small ions, attractions between ions of opposite charge hold the ions tightly in a crystal lattice. • These attractions are overcome only by heating to very high temperatures.

16. Larger Vs Smaller ions • If the ions are larger, then the distances between them are larger and the forces are weaker. e.g., CsI • Thus, ionic compounds with small ions have high melting points. e.g., LiF

17. Intermolecular Forces – (between covalent compounds) • Intermolecular forces are the forces of attraction between molecules of covalent compounds. • Intermolecular forces include dipole-dipole forces and London dispersion forces.

18. Dipole-Dipole Forces Affect Melting and Boiling Points • Dipole-dipole forces are interactions between polar molecules. • When molecules are very polar, the dipole-dipole forces are very significant. • The more polar the molecules are, the higher the boiling point of the substance.

19. Hydrogen bonds • A hydrogen bond is a dipole-dipole force occurring when a hydrogen atom that is bonded to a highly electronegative atom of one molecule is attracted to two unshared electrons of another molecule. • In general, compounds with hydrogen bonding have higher boiling points than comparable compounds.

20. Hydrogen Bonding, cont’d • As the electronegativity difference of the hydrogen halides increases, the boiling point increases. • The boiling points increase somewhat from HCl to HBr to HI but increase a lot more for HF due to the hydrogen bonding between HF molecules.

21. Hydrogen Bonds Form with Electronegative Atoms • Strong hydrogen bonds can form with a hydrogen atom that is covalently bonded to very electronegative atoms in the upper-right part of the periodic table: nitrogen, oxygen, and fluorine. • The combination of the large electronegativity difference (high polarity) and hydrogen’s small size accounts for the strength of the hydrogen bond. Hydrogen Bonds Are Strong Dipole-Dipole Forces

22. Hydrogen bonding in water – six sided snow flakes? – crystaline water is less dense?

23. Hydrogen Bonding Explains Water’s Unique Properties • Each water molecule forms multiple hydrogen bonds, so the intermolecular forces in water are strong. • The angle between the two H atoms is 104.5°. • When water forms ice, the ice crystals have large amounts of open space. Thus, ice has a low density. • Water is unusual in that its liquid form is denser than its solid form.

24. London Dispersion Forces • A substance with weak attractive forces will be a gas because there is not enough attractive force to hold molecules together as a liquid or a solid. • However, many nonpolar substances are liquids. • What forces of attraction hold together nonpolar molecules and atoms? • London dispersion forces are the dipole-dipole force resulting from the uneven distribution of electrons and the creation of temporary dipoles.

25. London Dispersion Forces • As molar mass increases, so does the number of electrons in a molecule. • London dispersion forces are roughly proportional to the number of electrons present. • Thus, the strength of London dispersion forces between nonpolar particles increases as the molar mass of the particles increases.

26. London Dispersion Forces Result from Temporary Dipoles • The electrons in atoms can move about in orbitals and from one side of an atom to the other. • When the electrons move toward one side of an atom or molecule, that side becomes momentarily negative and the other side becomes momentarily positive. • When the positive side of a momentarily charged molecule moves near another molecule, it can attract the electrons in the other molecule.

27. Temporary Dipoles

28. Particle Size and Shape Also Play a Role • Dipole-dipole forces are generally stronger than London dispersion forces. • However, both of these forces between molecules are usually much weaker than ionic forces in crystals. • When the forces between ions are spread out over large distances, as with large ions or oddly shaped ions that do not fit close together, they do not have as great of an effect.

29. Enthalpy, Entropy, and Changes of State • Enthalpy is the total energy of a system. • Entropy measures a system’s disorder. • The energy added during melting or removed during freezing is called the enthalpy of fusion. • Particle motion is more random in the liquid state, so as a solid melts, the entropy of its particles increases. This increase is the entropy of fusion.

30. Enthalpy, Entropy, and Changes of State • As a liquid evaporates, a lot of energy is needed to separate the particles. This energy is the enthalpy of vaporization. • Particle motion is much more random in a gas than in a liquid. A substance’s entropy of vaporization is much larger than its entropy of fusion.

31. Enthalpy and Entropy Changes for Melting and Evaporation • Enthalpy and entropy change as energy in the form of heat is added to a substance. • The energy added as ice melts at its melting point is the molar enthalpy of fusion (∆Hfus). • ∆Hfus is the difference in enthalpy between solid and liquid water at 273.15 K. • ∆Hfus = H(liquid at melting point) H(solid at melting point)

33. Enthalpy and Entropy Changes for Melting and Evaporation, continued

34. Enthalpy and Entropy Changes for Melting and Evaporation, continued • The energy added as liquid evaporates at its boiling point is the molar enthalpy of vaporization (∆Hvap). • ∆Hvap is the difference in enthalpy between liquid and gaseous water at 373.15 K. • ∆Hvap = H(vapor at boiling point) H(liquid at boiling point) • Because intermolecular forces are not significant in the gaseous state, most substances have similar values for molar entropy of vaporization,∆Svap.

35. Enthalpy and Entropy Changes for Melting and Evaporation, continued

36. Gibbs Energy and State Changes • The relative values of H and S determine whether any process, including a state change, will take place. • A change in Gibbs energy is defined as: • ∆G = ∆H  T∆S • A process is spontaneous if ∆G is negative. • If ∆G is positive, then a process will not take place unless there is an outside source of energy.

38. Gibbs Energy and State Changes, continued • If ∆G is zero, then the system is at equilibrium. • At equilibrium, the forward and reverse processes are happening at the same rate. • For example, when solid ice and liquid water are at equilibrium, ice melts at the same rate that water freezes.

39. Gibbs Energy and State Changes, continuedEnthalpy and Entropy Determine State • You can confirm that ice melts only at 273.15 K by calculating ∆G at temperatures just above and below the melting point of ice. • Below the melting point of ice, at 273.00 K: • ∆G = ∆H  T∆S • = +6009 J/mol (273.00 K  +22.00 J/mol•K) • = +6009 J/mol 6006 J/mol = +3 J/mol • ∆G is positive, so melting will not occur at 273.00 K.

40. Enthalpy and Entropy Determine State, continued • Above the melting point of ice, at 273.30 K: • ∆G = ∆H  T∆S • = 6009 J/mol (273.30 K  −22.00 J/mol•K) • = 6009 J/mol 6013 J/mol = +4 J/mol • ∆G is positive, so water will not freeze at 273.30 K

41. Determining Melting and Boiling Points • For a system at the melting point, a solid and a liquid are in equilibrium, so ∆G is zero. • Thus, ∆H = T∆S. • Rearranging the equation, you get the following equation for the melting point of a substance:

42. Determining Melting and Boiling Points • For a system at the boiling point, a gas and a liquid are in equilibrium. • Rearranging the equation as before, you get the following equation for boiling point : • Note that when ∆Hvap > T∆Svap, the liquid state is favored, and when ∆Hvap < T∆Svap, the gas state is preferred

43. Calculating Melting and Boiling Points • Sample Problem A • The enthalpy of fusion of mercury is 11.42 J/g, and the molar entropy of fusion is 9.79 J/mol•K. The enthalpy of vaporization at the boiling point is 294.7 J/g, and the molar entropy of vaporization is 93.8 J/mol•K. Calculate the melting point and the boiling point.

44. Calculating Melting and Boiling Points • Sample Problem A Solution • First calculate the molar enthalpy of fusion and molar enthalpy of vaporization, which have units of J/mol. • Use the molar mass of mercury, 200.59 g/mol, to convert from J/g to J/mol: • ∆Hfus = 11.42 J/g  200.59 g/mol = 2291 J/mol • ∆Hvap = 294.7 J/g  200.59 g/mol = 59 110 J/mol

45. Sample Problem A Solution, continued • Set up the equations for determining Tmp and Tbp.

46. Phase Diagrams • A substance’s state depends on temperature and that pressure affects state changes. • A phase diagram is a graph that shows the relationship between the physical state of a substance and the temperature and pressure of the substance.

47. Phase Diagrams, continued • A phase diagram has three lines: • One line shows the liquid-gas equilibrium. • Another line shows the liquid-solid equilibrium • A third line shows the solid-gas equilibrium. • The three lines meet at the triple point, the temperature and pressure at which the three states of a substance coexist at equilibrium.

48. Phase Diagram for Water

49. Phase Diagrams Relate State, Temperature, and Pressure • For any given point (x, y) on the phase diagram for water, you can see in what state water will be. • For example, at the coordinates 363 K (x = 90°C) and standard pressure (y = 101.3 kPa), the point falls in the region labeled “Liquid.”

50. Gas-Liquid Equilibrium • Line AD shows liquid-vapor equilibrium. • As you follow line AD upwards, the vapor pressure is increasing, so the density of the vapor increases and the density of the liquid decreases. • At a temperature and pressure called the critical point the liquid and vapor phases are identical. • Above this point, the substance is called a supercritical fluid.