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Intermolecular Forces

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Presentation on theme: "Intermolecular Forces"— Presentation transcript:

1 Intermolecular Forces
Chapter 11

2 Intermolecular Forces: Introduction
Forces between separate molecules and dissolved ions (not bonds) “Van der Waals Forces” 15% as strong as covalent or ionic bonds

3 Intermolecular Forces: Introduction
Low temperature – strong High temperature – kinetic energy of motion overcomes the IMF Boiling point is a good indicator Stronger IMF = higher boiling point Weaker IMF = lower boiling point

4 Ion-Ion Full to Full Charge Ion-Dipole Full to Partial Charge Hydrogen Bonds Partial to Partial Charge (H involved) Dipole-Dipole Partial to Partial Charge (No H) London Dispersion Forces Non-polar to Non-Polar

5 Predict what type of IMF would form between:
Br2 and I2 KCl and water Water and ammonia Two SO2 molecules NaCl ions in a crystal

6 Type 1: Ion-Ion Forces Full Charges to full charges
High melting points (ionic solids) Ex: Melting Point Na ~98 oC NaCl ~800 oC

7 Type 2: Ion-Dipole Forces
Full Charges to partial charges Very Strong

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10 Type 3: Hydrogen Bonds Stronger than dipole-dipole that do not have hydrogen (no inner electrons, strong + pull) Generally involves hydrogen and O, N or F R-H · · · · O-R R-H · · · · N-R R-H · · · · F-R

11 Miscibility

12 Miscibility “Like dissolves like”
Substances that can hydrogen bond dissolve in one another.

13 Glucose and other sugars

14 Water Beading

15 Ice

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17 Ice

18 DNA

19 DNA is TWO molecules that are hydrogen bonded (like a zipper)

20 Boiling Point Generally increases with increasing molar mass
H2O unusually high - H-bonding

21 Type 4: Dipole-Dipole Slightly weaker IMF
Involve d+ and d- charges other than those in hydrogen bonding

22 Draw Lewis Dot Structures to explain the following boiling points
MM (amu) BP (oC) CH3CH2CH CH3CHO CH3CN

23 Type 5: London Dispersion Forces
Very weak IMF Caused by temporary imbalances in electrons

24 London Forces: Inorganic Molecules
More electrons, more chance for temporary dipole Boiling Point Table Halogen Molar Mass BP(oC) Noble Gas Atomic Mass F2 (g) 38.0 -188 He 4.0 -268 Cl2 (g) 71.0 -35 Ne 20.2 -246 Br2 (l) 159.8 59 Ar 39.9 -186 I2 (s) 253.8 185 Kr 83.8 -152

25 Explain the differences in boiling point between Cl2 (-35oC) and Ar (-186oC)

26 London Forces: Organic Compounds
The longer the carbon chain, the higher the London Dispersion Forces (the higher the melting point and boiling point) Chainlike molecules greater London Forces than “bunched up” molecules (branched)

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28 Ex 1 Rank the following compounds in terms of increasing melting points: NaCl, CF4, CH3OH.

29 Ex 2 Separate the following compounds by whether they have dipole-dipole attractions (including H-bonding) or London Forces. Which should have the highest dipole-dipole attraction? Which should have the strongest London Force? Br2, Ne, HCl, N2, HF

30 Ex 3 Rank the following in order of increasing boiling point: BaCl2, H2, CO, HF, Kr

31 Ex 4 Rank the following in order of increasing boiling point: N2, KBr, O2, CH3CH2OH, HCN

32 Properties of Liquid Viscosity – resistance of a liquid to flow
Oil is more viscous than water Water has H-bonds (stronger) Oil has London forces (weaker, but there are many more of them, long carbon chain)

33 40 W oil 10 W oil

34 Miscibility

35 Surface Tension

36 Surface Tension

37 Capillary Action Water is attracted to the glass
Mercury more attracted to itself

38 Heating Curves Changes of state do not have a temperature change.
Melting/Freezing Boiling/Condensing A glass of soda with ice will stay at 0oC until all of the ice melts. Graph “flattens out” during changes of state

39 Heating Curves Steam heats up Temperature (oC) Boiling Water warms up
Melting Ice warms up Heat (Joules)

40 Heating Curve No phase change is occurring (heating ice, water, or steam): q = mCpDT Melting or boiling: q = mLf or q = mLv

41 Lf and Lv Latent Heat - heat for phase changes. No temperature change.
Lf –latent heat of freezing/melting Lv –latent heat of boiling/condensing

42 Heating Curves Temperature (oC) Use q = mLv Boiling Use q = mLf
Melting Use q = mCpDT Heat (calories)

43 Important Values Substance Cp Steam 2.01 J/goC Water 4.18 J/goC
Ice J/goC Latent Heat of fusion (water) Lf = J/g Latent Heat of vaporization(water) Lv = J/g

44 Heating Curves: Example 1
How much energy must be removed to cool grams of water at 20.0oC to make ice at –10.0oC?

45 Heating Curves: Example 1
Temperature (oC) Melting (q=mLf) Water cools (q=mCpDT) Ice cools (q=mCpDT) Heat

46 Heating Curves: Example 1
Cooling the water q = mCpDT = (100 g)(4.18 J/goC)(0oC-20oC) q = J (8.36 kJ) (ignore the negative sign for now) Freezing the water q = mLf = (100.0 g)(334.7 J/g)= kJ 3. Cooling the ice down to –10.0oC q = mCpDT = (100 g)(2.09 J/goC)(-10oC-0oC) q = 2.09 kJ (we will ignore the negative sign for now) 8.36 kJ kJ kJ= kJ

47 Ex 2 How much energy is needed to convert 18.0 grams of ice at -25oC to steam at 125oC? ANS: 56.0 kJ

48 Ex 3 How much energy must be used to convert grams of water from steam at 110.0oC to ice at -25.0oC ? (309 kJ)

49 Vapor Pressure Pressure of a gas above a liquid caused by that liquid
Temperature – measure of the average kinetic energy of molecules At any given moment, some molecules have enough energy to escape EX: Even cold water will evaporate

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52 Volatility Volatile – liquids that evaporate easily
Acetone Often weak intermolecular forces Boiling point – point at which the vapor pressure of a liquid = vapor pressure of the atmosphere Normal Boiling Point – vapor pressure = 1 atm Steam pressure cookers – Forces water to boil at a higher temperature High Altitude – water boils at a lower temperature

53 Four Types of Solids Molecular Solids (single molecules)
Covalent Network Solids (one large molecule) Ionic Solids Metallic Solids

54 1. Molecular Solids Held together by IMF Ice Plastics

55 2. Covalent Network Solids
Basically one big molecule Held together by covalent bonds Diamond Graphite Quartz (SiO2)

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57 3. Ionic Solids Held together by electrostatic attraction (Ion:Ion)
Usually crystalline (unit cells)

58 4. Metallic Solids Atoms share electrons very freely
Positive nuclei in a “sea of electrons” Electrons held loosely Conducts electricity Photoelectric effect Malleable and ductile

59 Rank by boiling point (low to high)
Rank by boiling point (low to high). Below each, tell me which IMF is important: CH3OH Cl2 N2 CH3Cl CH3CH2CH2CH3 CH3CH2CH2OH Draw Lewis Dots for these CH3OCH3 CH3CH2CH3

60 How much energy must be removed to convert 100.0 grams of steam at 120.0oC to ice at -5.00oC?

61 (100. 0g)(2. 01 J/gK)(20. 0oC) = 4. 02 kJ (100. 0g)(2259. 4 J/g) = 225
(100.0g)(2.01 J/gK)(20.0oC) = 4.02 kJ (100.0g)( J/g) = kJ (100.0g)(4.18 J/gK)(100.0oC) = 41.8 kJ (100.0g)(334.7 J/g) = kJ (100.0g)(2.09 J/gK)(5.0oC) = 1.05 kJ kJ

62 The heat of vaporization of ammonia is 23. 35 kJ/mol
The heat of vaporization of ammonia is kJ/mol. How many grams of ammonia must evaporate to freeze grams of water initially at 10.00oC? Assume the ice stays at 0oC when frozen, and is used to cool a cup of too hot tea. (Ans: 27.4 g)

63 Water qcool (100)(4. 18)(10) = 4. 18 kJ qfreeze (100)(334. 7) = 33
Water qcool (100)(4.18)(10) = 4.18 kJ qfreeze (100)(334.7) = kJ kJ Ammonia qv = mLv m =qv/Lv m = kJ/23.35 kJ/mol = 1.61 mol mass = (1.61 mol)(17.0 g/mol) = 27.4 g

64 a) H-bonding, (b)London (c) Ion-dipole (d) dipole-dipole
a) H-bonding, (b)London (c) Ion-dipole (d) dipole-dipole. Ion-dipole and h-bonding are stronger 10. a) Solids = attractive forces (IMF) win Liquids = Balance Gases = Kinetic energy wins b) Increasing T increase KE, eventually overcoming IMF c) High pressure forces gas molecules clsoe together and IMF’s can win

65 a) Distance greater in liquid state
b) More movement, more volume, lower density 14. Overall, net forces are attractive 16.a) CH3OH has h-bonding, CH3SH does not b) Xe is heavier, greater London Forces c) Cl2 more polarizable than Kr d) Acetone has dipole-dipole forces 18. a) True b) False c) False d) True 20. a) Br2 b) C5H11SH c) CH3CH2CH2Cl

66 22. Propyl alcohol is longer and more polarizable 24
22. Propyl alcohol is longer and more polarizable 24. a) HF has hydrogen bonds, HCl dipole/dipole b) CHBr3 higher molar mass, more dispersion c) ICl has dipole-dipole, Br2 only dispersion 26.a) Dispersion, C8H18 higher boiling point b) C3H8(dispersion) CH3OCH3 (dip-dip) c) HOOH (h-bonding) HSSH (dip-dip) d) NH2NH2 (h-bonding) CH3CH3 (dispersion) 32. H2NNH2, HOOH, H2O can all h-bond

67 a) Exo b) Endo c) Endo d) Exo
g CCl2F2 kJ

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