Chapter 11: States of Matter & Intermolecular Forces

1. Chapter 11: States of Matter & Intermolecular Forces

2. Essential Question: Sections 1&2 How do particles interact with each other and how does this affect properties?

3. Warm Up • Recall the three common phases. • Think about the motion of particles in each of these phases- talk to people in your section. • Now let’s act them out! • What might cause substances to be in each of these phases?

4. States/Phases

5. Solids • Fixed, rigid position • Held tightly together • Various attractive forces hold them together • Often exist in crystalline form • Can be hard and brittle (like table salt) or soft (like potassium)

6. Liquids • Also held close together by attractive forces • Have enough energy to slide past each other • Viscosity (ability to flow) differs • More viscous = thicker (doesn’t easily flow) • Less viscous = thinner (easily flows) • Will a liquid wet a solid surface? Depends on two forces: • Cohesion: attraction to each other. • Adhesion: attraction to solid surfaces.

7. Liquid Water: Wax vs. Glass • Water on wax paper? • Droplets bead all over the surface • High cohesion for each other, low adhesion to wax • Water in a glass? • Meniscus! • Stronger adhesion forces pull water molecules up the sides • These water molecules pull other water molecules with them (very cohesive) • Water rises up until adhesive and cohesive forces are counterbalanced by gravity (capillary action)

8. Liquids: Surface Tension • Particles below the surface feel cohesion in all directions. • Particles at the surface only feel cohesion sideways and downward.

9. Liquids: Surface Tension Cont. • Takes work to pull particles to the surface (going against cohesive forces). • Surface tension: force that acts on the surface of liquids (attempts to minimize surface area).

10. Gases • Particles are extremely far apart • Attractive forces are hardly felt • Particles move quickly and almost independently • Considered fluids since they can move around freely

11. Phase Changes

12. Phase Changes • Remember what happens as energy is added to a substance? • Temperature increases UNTIL it begins to undergo a phase change; then it’s constant until the phase change is complete. • These temperatures are the melting & boiling points.

13. Evaporation • How can water evaporate without adding heat to make it boil? • Sometimes the energy of motion is enough to allow molecules to go from liquid to gas. • A water molecule hit by several molecules at the same time gains enough energy to leave the liquid’s surface. • Only happens at the surface! (Boiling causes vaporization throughout the entire sample.)

14. Additional Phase Change Info. • Condensation- why does it happen? • Gaseous molecules cool down enough that they no longer have enough energy to overcome cohesive forces. • Sublimation- think moth balls! • Just like vaporization, but more energy needed to go right from solid to gas. • Every phase change has energy involved. • Reverse processes have the SAME amount of energy involved. • Ex: melting a liquid absorbs 50kJ of energy; freezing the liquid back to a solid releases 50kJ of energy

15. Questions • Explain how adhesion plays a role in the formation of a meniscus. • Describe the motions of particles when water boils. • Explain the difference between evaporation and boiling.

16. Section 1 Review • Pg. 384 #2-5,8,11,12

17. Section 2: Warm Up Questions • Why do you think water molecules show cohesion? • What happens when oil and water are mixed? Why does this happen?

18. Intermolecular Forces • Ionic compounds are held together by attractive ionic forces between ions. • Two factors affect strength of ionic forces: • Size of charge on the ion. • Ex: NaCl vs. MgCl2; MgCl2 has stronger forces because Mg+2 has a larger positive charge than Na+. • Size of the ion itself (radius). • Ex: NaCl vs. KCl; NaCl has stronger forces because Na+ is smaller than K+. • Stronger forces = higher boiling/melting point!

19. Practice • Which ionic compound has stronger ionic forces? Which has the greater melting point? 1) Ba(NO3)2 vs. BaCl2 BaCl2 2) AlBr3 vs. CaBr2 AlBr3

20. Bonds vs. IFA The H and O atoms in each water molecule are held together by a covalent bond. Each water molecule is attracted to another water molecule because of the IFA felt between them.

21. Types of IFA • Three types: • London Dispersion • Dipole-Dipole • Hydrogen Bonding • In general, the strengths of each are: London Dispersion < Dipole-Dipole < H Bonding • Bonds between atoms are MUCH stronger than any of these intermolecular forces!

22. London Dispersion • Arise from random shifts in electrons in substances. • Temporary (instantaneous) dipoles cause temporary partial charges.

23. These temporary dipoles can induce other temporary dipoles in molecules nearby. • Partial charges attract molecules, until they disappear. • All substances exhibit this type of intermolecular force (even ionic compounds) because all substances have electrons!

24. LDF Continued • LDF are most noticeable in nonpolar substances because this is the only intermolecular force holding molecules together. • LDF increases in strength with more electrons/as molecules get bigger. • A very large, nonpolar molecule can have very strong LDF.

25. Dipole-Dipole • IFA exhibited in polar covalent compounds. • Polar = permanent partial charges. • Electronegativity differences! • Partial positive charges and partial negative charges attract in different molecules and hold them together. • Stronger partial charges = stronger dipole-dipole forces.

26. Hydrogen Bonding • Some dipole-dipole forces are especially strong and have a class of their own. • Hydrogen bonding: O, F, or N bonded to a H atom causes very strong partial charges to develop. • If a molecule has one of these electronegative atoms bonded to a H, it will exhibit hydrogen bonding.

27. Practice • Identify the type of intermolecular forces each of the following substances exhibits. Then rank them in order of weakest to strongest: CH3OH, F2, NaCl, CH3Br. • CH3OH = H bonding; F2= LDF; NaCl = ionic forces; CH3Br = dipole-dipole • F2 < CH3Br < CH3OH < NaCl

28. IFA & BP/MP • The strength and type of IFA that a substance exhibits directly impacts the melting and boiling points of the substance. • This also determines the phase at room temperature. • Stronger IFA = more E needed to break these forces of attraction = higher MP & BP. • Solids at room temperature have stronger IFA than liquids, and liquids at room temperature have stronger IFA than gases.

29. Practice • Rank the following substances in order of decreasing boiling point: Ar, H2O, Cl2. • H2O, Cl2, Ar

30. Section 2 Review • Pg. 392 #1,2,5-7,10,11

31. Lesson Essential Question: Section 3 • How can the spontaneity of a phase change or a reaction be determined?

32. Section 3: Energy & State Changes • Thermodynamics: studies the effect of heat, work, and energy on a system. • Will be examined more closely in AP chemistry. • Enthalpy and entropy are two important properties. • Enthalpy vs. entropy • Enthalpy: total energy of a system, H. • Depends upon heat energy & kinetic energy of particles. • Entropy: disorder in a system, S. • Not a form of energy! • Both influence whether a reaction will occur (is favorable) or not.

33. Enthalpy • Do substances prefer to be at higher or lower energy? • Lower energy is favorable: -∆H. • Less internal energy = greater stability. • This is why energy must be added to get substances to melt and vaporize; the particles have greater kinetic energy!

34. Entropy • More entropy (+∆S) is favorable. • Nature has a tendency to favor disorder. • Does your room get messier over time or does your room get cleaner over time? • You have to do work to organize/clean your room because it tends to get messy over time!

35. Entropy • Food coloring demonstration. • Particles also like chaos/spreading out! • Ex: diffusion occurs because particles have a natural tendency to spread out and become more disordered. It would take work to reverse this process and force the particles to come back together. • A gas in a container that is opened will float out into the air. Why? • More space in the air to spread out = more favorable disorder!

36. What Can ∆H & ∆S Tell Us? • Based on the signs of ∆H and ∆S, we can more easily determine whether a process will occur (spontaneous) or not. • We said that lower energy is favorable, -∆H, and higher entropy is favorable, +∆S. • If a process releases energy and at the same time increases its entropy, what can you conclude about the likelihood of that process occurring? • Very likely because it is favorable in terms of both enthalpy and entropy!

37. What Can Both ∆H & ∆S Tell Us? • What can you also say about the opposite: if a process absorbs energy and at the same time decreases its entropy, will the process occur? • Think about the signs for ∆H and ∆S! • +∆H and -∆S. • Very unlikely because it is unfavorable in terms of both enthalpy and entropy! • But what about processes that are favorable in terms of one and not the other? In other words when both signs of ∆H and ∆S are the same?

38. Spontaneity & Gibbs Free Energy • There must be another thermodynamic quantity that affects the spontaneity of processes besides just ∆H and ∆S. • ∆G = ∆H -T∆S • G stands for Gibbs free energy: energy in a system that’s available for work. • If ∆G is negative (the system gives off free energy), the process is spontaneous- it will occur. • If ∆G is positive (the system takes in free energy), the process is nonspontaneous- it will not occur. • If G equals zero, the process is at equilibrium (forward and reverse reaction rates are the same).

39. Spontaneity & Gibbs Free Energy • ∆G = ∆H -T∆S • Use ∆Hfusion for fusion (melting) or freezing; use ∆Hvaporization for vaporization or condensation. • Number is the same, sign changes (+ if E taken in, - if E released). • Same is true for ∆S. • + if disorder of new phase increases, - if disorder of new phase decreases. • T is temperature in Kelvin, K.

40. Spontaneity & Gibbs Free Energy • ∆G = ∆H -T∆S • You can also make general conclusions about spontaneity even if one term is favorable and the other is not. • Depends on T! • Even if H is positive, at high temperatures it is likely that the reaction will be spontaneous. • Even if S is negative, at low temperatures it is likely that the reaction will be spontaneous.

41. Gibbs Free Energy Determines Phases • The sign of ∆G indicates whether the phase will change or not at a given temperature. • Ex: Will ice melt at 273.00K? • ∆H fusion: 6,009 J/mol • ∆S fusion: 22.00 J/(mol·K) • Signs of ∆H and ∆S? • Melting requires energy be taken in (+∆H = +6,0009 J/mol). • Changing from solid to liquid means more entropy (+∆S = +22.00 J/mol·K). • ∆G = 6,009 J/mol – (273.00K) x (22.00 J/mol·K) • ∆G = 6,009 J/mol – 6,006 J/mol • ∆G = + 3 J/mol Ice will not melt! Just under 0°C.

42. Summary

43. Another Sample Problem • Will water freeze at 273.30K? • ∆H fusion: 6,009 J/mol • ∆S fusion: 22.00 J/(mol·K) • Signs of ∆H and ∆S? • Freezing means energy will be released (-∆H = -6,0009 J/mol). • Changing from liquid to solid means less entropy (-∆S = -22.00 J/mol·K). • ∆G = + 4 J/mol Water will not freeze! • Practice: pg. 367 # 3,12,14-16 Just over 0°C.

44. ∆G During Phase Changes • At phase changes, phases are in equilibrium with each other (both phases are present). • Therefore, ∆G = 0. • Can rearrange formula to find melting and boiling points: ∆G = ∆H -T∆S 0 = ∆H -T∆S T∆S = ∆H • T = ∆H/∆S Tmp= ∆Hfus/∆Sfus & Tbp= ∆Hvap/ ∆Svap

45. Sample Problem #1 For mercury, the enthalpy of fusion is 2,291J/mol, and the entropy of fusion is 9.79J/molK. The enthalpy of vaporization is 59,110 J/mol, and the entropy of vaporization is 93.8 J/molK. Calculate the MP and BP of mercury. Tmp= (2,291J/mol)/(9.79J/molK) = 234K Tbp= (59,110J/mol)/(93.8J/molK) = 630K

46. Sample Problem #2 For ammonia, the enthalpy of fusion is 5.66kJ/mol, and the entropy of fusion is 29.0J/molK. The enthalpy of vaporization is 23.33kJ/mol, and the entropy of vaporization is 97.2J/molK. What are the MP and BP? MP = 195K BP = 240K

47. A Note About Pressure • So far we have only looked at the effect of T on phase changes (besides G, H, and S). • P also has an effect, but only when gases are involved. • Gases are P dependent because they are compressible. • Liquids and solids are not compressible, so not affected by P during phase changes. • This will be examined more closely in AP chemistry.

48. Section 3 Review • Pg. 398 # 3-7

49. Lesson Essential Questions: Section 4 • What happens to the particles of a substance as it undergoes a phase change? • How do temperature and pressure affect phases?

50. Section 4: Phase Equilibrium • Equilibrium = interchange of particles • In this case, between phases. • Dynamic equilibrium: particles constantly move between phases, but there is no change in the amount of particles in each phase. • Moving at the same rate! • At 0°C, both liquid water and solid ice will be present; molecules will change between the two phases at equal rates.