Chapter 11: States of Ma0er & Intermolecular Forces

Transcription

1 Chapter 11: States of Ma0er & Intermolecular Forces

2 Essen:al Ques:on: Sec:ons 1&2 How do par)cles interact with each other and how does this affect proper)es?

3 States

4 Solids Fixed, rigid posi)on Held )ghtly together O;en exist in crystalline form Can be hard and (like table salt) or so; (like potassium) Various forces hold solids together Ionic bonding Covalent bonding Metallic bonding

5 Liquids Also held close together by forces. Have enough energy to slide past each other. Liquids have three important Viscosity: resistance to flow More viscous = thicker (doesn t easily flow) Less viscous = thinner (easily flows) Cohesion: a@rac)on of liquid par)cles to each other. Adhesion: a@rac)on of liquid par)cles to solid surfaces.

6 Applica)ons of Cohesion & Adhesion Will a liquid wet a solid surface (spread over it) or not? Ex: What does spilled water on your table look like? Why? Water in a graduated cylinder? Meniscus at the surface! Why does this occur? Water s cohesive & adhesive proper)es are cri)cal for life! For example, plants use these proper)es to transport water from roots to leaves.

7 Liquids: Surface Tension In your own words, what is surface tension? How would you describe it to someone who has no idea what it is? Par)cles below the surface feel cohesion in all direc)ons. Par)cles at the surface only feel cohesion sideways and downward. This creates tension at the surface.

8 Liquids: Surface Tension Cont. Takes work to pull par)cles to the surface (going against cohesive forces). Surface tension: force that acts on the surface of liquids to minimize surface area). This tension creates a film or layer at the surface of the liquid. Imagine a piece of clear Saran Wrap pulled over the surface of the liquid. Demonstra)ons- string & boat.

9 Gases Par)cles are extremely far apart forces are hardly felt Par)cles move quickly and almost independently Considered fluids since they can move around freely

10 Phase Changes

11 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 un)l the phase change is complete. These temperatures are the mel)ng & boiling points.

12 Hea)ng/Cooling Curve

13 Evapora)on How can water evaporate without adding heat to make it change from a liquid to a gas? Some)mes the energy of mo)on is enough to allow molecules to go from liquid to gas. A water molecule hit by several molecules at the same )me gains enough energy to leave the liquid s surface. Only happens at the surface! (Boiling causes vaporiza)on throughout the en)re sample- that s why you see bubbles.)

14 Addi)onal Phase Change Info. Condensa)on- why does it happen? Gaseous molecules cool down enough that they no longer have enough energy to overcome cohesive forces. Sublima)on- think moth balls! Just like vaporiza)on, 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. One direc)on will absorb energy, the other will release energy.

15 Introduc)on Video The following mini- lesson will introduce the idea of intermolecular forces and how it effects liquids: v=bqqjpcdmip8

16 Sec)on 2: Intermolecular Forces Difference between bonds (intramolecular forces) and intermolecular forces: The H and O atoms in each water molecule are held together by a covalent bond. Each water molecule is a@racted to another water molecule because of the IFA felt between them.

17 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!

18 London Dispersion Arise from random shi;s in electrons in substances. Temporary (instantaneous) dipoles cause temporary par)al charges.

19 These temporary dipoles can induce other temporary dipoles in molecules nearby. Par)al charges molecules, un)l they disappear. All substances exhibit this type of intermolecular force (even ionic compounds) because all substances have electrons!

20 LDF Con)nued LDF are most no)ceable 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.

21 Dipole- Dipole IFA exhibited in polar covalent compounds. Polar = permanent par)al charges. Electronega)vity differences! Par)al posi)ve charges and par)al nega)ve charges a@ract in different molecules and hold them together. Stronger par)al charges = stronger dipole- dipole forces.

22 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 par)al charges to develop. If a molecule has one of these electronega)ve atoms bonded to a H, it will exhibit hydrogen bonding.

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

24 Ionic Forces Ionic compounds are held together by ionic forces between ions. Although this is a BOND, we can also look at it as an intermolecular force, as it is responsible for holding ionic compounds together. Two factors affect strength of ionic forces: Size of charge on the ion. Ex: NaCl vs. MgCl 2 ; MgCl 2 has stronger forces because Mg +2 has a larger posi)ve 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/mel)ng point!

25 Lesson Essen)al Ques)on: Sec)on 3 How can the spontaneity of a phase change or a chemical reac)on be determined?

26 Sec)on 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 proper)es. Enthalpy vs. entropy Enthalpy: total energy of a system, H. Depends upon heat energy & kine)c energy of par)cles. Entropy: disorder in a system, S. Not a form of energy! Both influence whether a reac)on will occur (is favorable) or not.

27 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 par)cles have greater kine)c energy!

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

29 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 )me 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!

30 What Can Both H & S Tell Us? What can you also say about the opposite: if a process absorbs energy and at the same )me 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?

31 Spontaneity & Gibbs Free Energy There must be another thermodynamic quan)ty 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 nega)ve (the system gives off free energy), the process is spontaneous- it will occur. If G is posi)ve (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 reac)on rates are the same).

32 Spontaneity & Gibbs Free Energy G = H - T S Use H fusion for fusion (mel)ng) or freezing; use H vaporiza)on for vaporiza)on or condensa)on. 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.

33 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 posi)ve, at high temperatures it is likely that the reac)on will be spontaneous. Even if S is nega)ve, at low temperatures it is likely that the reac)on will be spontaneous.

34 Gibbs Free Energy Determines States The sign of G indicates whether the phase will change or not at a given temperature. Ex: Will ice melt at K? Just under 0 C. H fusion: 6,009 J/mol S fusion: J/(mol K) Signs of H and S? Mel)ng requires energy be taken in (+ H = +6,0009 J/mol). Changing from solid to liquid means more entropy (+ S = 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 ß Ice will not melt!

35 Summary

36 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 mel)ng and boiling points: G = H - T S 0 = H - T S T S = H T = H/ S T mp = H fus / S fus & T bp = H vap / S vap

37 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.

38 Lesson Essen)al Ques)ons: Sec)on 4 What happens to the par)cles of a substance as it undergoes a phase change? How do temperature and pressure affect phases?

39 Sec)on 4: Phase Equilibrium Equilibrium = interchange of par)cles In this case, between phases. Dynamic equilibrium: par)cles constantly move between phases, but there is no change in the amount of par)cles 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.

40 Vapor Pressure & Temperature When water goes between the liquid and gas phase in a closed container, the gas creates pressure. Gas molecules are constantly moving; when they strike the walls of the container, P is exerted. This is vapor pressure! BP can be defined as the T at which the vapor pressure equals the external pressure. If the pressures are equal then there s no greater pressure to push down on the liquid and keep the par)cles in the liquid phase.

41 Vapor Pressure & Temperature What happens to the vapor pressure as T increases? The vapor pressure increases! But why? Higher T = par)cles have greater KE; more KE means par)cles are more able to enter the gas phase. More par)cles in the gas phase = more P exerted because there are more collisions with the wall of the container.

42 Phase Diagrams Show the rela)onship between phase, T, and P. Phase diagrams are unique for every substance! Lines separate phases; points that are on the line exhibit. equilibrium between phases. Triple point: only T and P where all three phases are in equilibrium. Cri)cal point: point at which liquid & vapor phases become one- they are iden)cal. This is called a supercri)cal fluid: liquid & vapor pressures cannot be dis)nguished.

43 Solid & Liquid Densi)es Nega)ve slope of the solid/liquid line = solid less dense than liquid. Posi)ve slope = solid more dense than liquid (much more common). Water is unique- it has a nega)ve solid/liquid slope.