Molecular Structure & Intermolecular Forces

1. Molecular Structure & Intermolecular Forces Saturday Study Session #2 3rd Class

2. Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

3. Draw the Lewis structure for CH2Cl2 OR

4. Draw the Lewis Structure for NO+

5. Draw the Lewis Structure for XeF2 • *Xe can have more than an octet of electrons!

6. From the Lewis Structure we can determine: Electron geometry Molecular geometry Hybrid Orbital Polarity Intermolecular bond

7. Molecular Shapes • The shape of a molecule plays an important role in its reactivity. • By noting the number of bonding and nonbonding electron pairs, we can easily predict the shape of the molecule.

8. Electron Domains • We can refer to the electron pairs as electron domains. • In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain. • The central atom in this molecule, A, has four electron domains.

9. Valence-Shell Electron-Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

10. If all electron domains are bonds the molecular shapes are like these below:

11. If some of the electron domains are unshared pairs of electrons then the molecular shapes are as indicated in this chart and the one on the following slide.

12. Let’s practice molecular geometry

13. Molecular Geometry answers • CH2Cl2 • NO+ • XeF2 • Tetrahedral • Linear • Linear, but its electron geometry is octahedral due to 3 unshared pairs of e-

14. Bond Angles 120o 45o

15. Bond Angles for molecules without lone pairs of electrons.

16. Nonbonding Pairs and Bond Angle • Nonbonding pairs are physically larger than bonding pairs. • Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

17. Multiple Bonds and Bond Angles • Double and triple bonds place greater electron density on one side of the central atom than do single bonds. • Therefore, they also affect bond angles.

18. Bond and Molecular POLARITY

19. Polar Covalent BondsIntramolecular • Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. • Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. • Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

20. Polar Covalent Bonds • \ The greater the difference in electronegativity, the more polar is the bond.

21. Polarity • Just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

22. Polarity of Molecules By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

23. Polarity

24. Molecular Polarity Polar Molecules Non-Polar Molecules Polar or nonpolar bonds Dipoles cancel Insoluble in Water Look for Symmetrical molecule • Must have some polar bonds (∆EN > 1.7) • Overall net dipole • Soluble in Water • Look for • Asymmetry • -OH, -NH2groups

25. Practice Molecular Polarity • Which molecule is more polar? 1. CS2 or SF2 2. BH3 or NH3 NONPOLAR POLAR POLAR NONPOLAR

26. Which is more polar? 3. Benzene POLAR OR NONPOLAR Glucose

27. Molecules Stick Together • All molecules have some attractive forces for each other. • Polar molecules have more types of attractive forces than do nonpolar molecules. • These attractive forces are called INTERMOLECULAR FORCES (IMF)

28. Inter vs. Intra molecular forces Inter (between molecules) Intra (inside molecules) Ionic bonds Covalent bonds Metallic bonds • London dispersion forces • Dipole-dipole forces • Hydrogen bonds

29. Dipole-Dipole forces

30. Hydrogen bonding

31. Which IMFs are present? 1. CF4 2. BF3 3. NH3 4. H2CS • London dispersionforces 2. London dispersion forces, Dipole – dipole forces 3. London dispersion forces, Hydrogen bonding 4. London dispersion forces, Dipole – dipole forces (in water, weak H bonding)

32. Effects of IMFs • States of matter • Phase changes • Melting points • Boiling points • Vapor pressure

33. States of Matter (Increasing) Molecular Interactions ARE Intermolecular Forces

34. Particles getting farther apart means they are overcoming intermolecular forces by adding energy.

35. Energy In Energy Out Must overcome IMFs IMFs cause particles to congregate

36. Effects of IMFs on properties • Greater IMFs = higher melting and boiling points and lower vapor pressure. • Greater Molar Mass = more electrons = greater IMFs • Volatile substances have high VP due to low IMFs • Greater IMFs = high ΔHvap

37. MC Question 1 In which of the following processes are covalent bonds broken? A) I2(s) → I2(g) B) CO2(s) → CO2(g) C) NaCl(s) → NaCl(l) D) C(diamond) → C(g) E) Fe(s) → Fe(l)

38. Question 1 Answer • Correct answer is D • Diamond is a covalently bonded network crystal. In order to form a gas the covalent bonds must be broken. • A and B the molecules remain intact, only IMF are “broken” • C is held together with ionic bonds • E is held together with metallic bonds

39. MC Question 2 The structural isomers C2H5OH and CH3OCH3 would be expected to have the same values for which of the following? (Assume ideal behavior.) A) Gaseous densities at the same temperature and pressure B) Vapor pressures at the same temperature C) Boiling points D) Melting points E) Heats of vaporization

40. Question 2 Answer • Correct answer is A • Density is a function of mass and volume. Isomers have the same molecular mass, same volume can be assumed. • All other answers the values change according to differences in IMFs. • C2H5OH has hydrogen bonding but CH3OCH3 does not.

41. MC Question 3 X: CH3–CH2–CH2–CH2–CH3 Y: CH3–CH2–CH2–CH2–OH Z: HO–CH2–CH2–CH2–OH Based on concepts of polarity and hydrogen bonding, which of the following sequences correctly lists the compounds above in the order of their increasing solubility in water? A) Z < Y < X B) Y < Z < X C) Y < X < Z D) X < Z < Y E) X < Y < Z

42. Question 3 Answer • Correct answer is E • The pure hydrocarbon butane X, is the least polar, thus has the lowest solubility in water. • The presence of an –OH group on butanol Y, makes it more soluble than butane, but less soluble than the 1,3-propanediol • Z, that contains two –OH groups. Aren’t you glad you don’t have to name all of them?

43. MC Question 4 Hydrogen Halide Normal Boiling Points, °C HF +19 HCl – 85 HBr – 67 HI – 35 The relatively high boiling point of HF can be correctly explained by which of the following? A) HF gas is more ideal. B) HF is the strongest acid. C) HF molecules have a smaller dipole moment. D) HF is much less soluble in water. E) HF molecules tend to form hydrogen bonds.

44. Question 4 Answer • Correct answer is E • Hydrogen bonding occurs when a H is bound to a “highly electronegative atom” (F, N or O) • The bonded H is attracted to an unshared electron pair or another highly electronegative atom on a neighboring molecule.

45. MC Question 5 Which of the following gases deviates most from ideal behavior? (Ideal gases assume no interparticle attractions) A) SO2 B) Ne C) CH4 D) N2 E) H2

46. Question 5 Answer • Correct answer is A • Deviations occur due to molecular volume (larger molecules have more mass as well) and attractive forces. • The more electrons present, the more polarizable a molecule, thus the greater the London dispersion (induced dipole-induced dipole) attractive forces become. • SO2has a higher molecular mass, more electrons and is more polarizable than the other answer choices.

47. MC Question 6 Molecular iodine would be most soluble in:  A) waterB) carbon tetrachloride C) vinegar (acetic acid and water)D) vodka (ethanol and water)E) equally soluble in all four

48. Question 6 Answer • Correct answer is B • Molecular iodine is a nonpolar molecule. • Carbon tetrachloride is the only nonpolar solvent listed.

49. FR Question 1 Explain each of the following in terms of the electronic structure and/or bonding of the compounds involved. At ordinary conditions, HF (normal boiling point = 20°C) is a liquid, whereas HCl (normal b.p. = -114°C) is a gas.

50. FRQ 1 Answer • HF exhibits hydrogen bonding but HCl does not. Both molecules have dispersion forces (HCl slightly greater than HF) but the hydrogen bonds are stronger in HF (F very highly electronegative) and require more energy to overcome to allow HF molecules to leave the liquid state and enter the gaseous state.