Section 6.5 “Polar Bonds and Intermolecular Forces”

1. Section 6.5“Polar Bonds and Intermolecular Forces” Ball-and-stick model Pre-AP Chemistry *Use after VSEPR Theory powerpoint

2. Important Information: • Chapter 6 Notes (5 blocks) • Chapter 6 Vocabulary (25 words) • Sec 6.5 Practice Problems – Set E p. 188, Set F p. 191,p. 202 (48a-c, 49a-d) p. 202 (45a-d,47a-e) • Activity: VSEPR Predicting Molecular Shapes due monday • 6.5 Study Guide due Monday • Chapter 6 TestTuesday • Bonus: Interpreting Graphics due Tuesday

3. 6.5 Practice problems Set E (p. 188) 1. Use the VSEPR theory to predict the molecular geometry of the following molecules: • HI • CBr4 • CH2Cl2

4. 6.5 Practice problems Set F (p. 191) 1. Use the VSEPR theory to predict the molecular geometries of the molecules whose Lewis structures are given below: • F S F • Cl P Cl • Cl

5. 6.5 Practice problems p. 202 48. Draw a Lewis structure for each of the following molecules, and then use the VSEPR theory to predict the molecular geometry of each. • SCl2 • PI3 • Cl2O • NH2Cl • SiCl3Br • ONCl

6. 6.5 Practice problems p. 202 49. Draw a Lewis structure for each of the following polyatomic ions, and then use VSEPR theory to predict the geometry of each: • NO3- • NH4+ • SO42- • ClO2-

7. Section 6.5Molecular Geometry • OBJECTIVES: • Negative particles (electrons) repel and move away from each other. VSEPR Theory • Multiple orbitals can combine to form hybrid orbitals. Hybridization • Weak forces exist between molecules. Dipole, Hydrogen bonding, London dispersion forces

8. Chapter 6 Section5 Molecular Geometry Hybridization • VSEPR theory is useful for predicting and explaining the shapes of molecules. • A step further must be taken to explain how the orbitals of an atom are rearranged when the atom forms covalent bonds. • For this purpose, we use the model of hybridization, which is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies.

9. Chapter 6 Section5 Molecular Geometry Hybridization • Take the simple example of methane, CH4. The carbon atom has four valence electrons, two in the 2s orbital and two in 2p orbitals. • Experiments have determined that a methane molecule is tetrahedral. How does carbon form four equivalent, tetrahedrally arranged, covalent bonds? • Recall that s and p orbitals have different shapes. To achieve four equivalent bonds, carbon’s 2s and three 2p orbitals hybridizeto form four new, identical orbitals called sp3orbitals. • The superscript 3 on the p indicates that there are three p orbitals included in the hybridization. The superscript 1 on the s is left out, like in a chemical formula.

10. Chapter 6 Section5 Molecular Geometry Hybridization • The four (s + p + p + p) hybrid orbitals in the sp3-hybridized methane molecule are equivalent: they all have the same energy, which is greater than that of the 2s orbital but less than that of the 2p orbitals. • Hybrid orbitals are orbitals of equal energy produced by the combination of two or more orbitals on the same atom. • Hybridization explains the bonding and geometry of many molecules.

11. Hybrid Orbitals Methane CH4 4 Single bonds Ethene C2H4 1 double bond, 2 single Ethyne C2H2 1 triple bond, 1 single sp3 sp sp2

12. Chapter 6 Section5 Molecular Geometry Geometry of Hybrid Orbitals

13. Section 6.5Polar Bonds and Molecules

14. Polar Bonds and Polar Molecules • OBJECTIVES: • Describe how electronegativity values determine the distribution of charge in a polar molecule. • Describe what happens to polar moleculeswhen they are placed between oppositely charged metal plates. • Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.

15. Bond Polarity • Covalent bonding means shared electrons • but, do they share equally? • Electrons are pulled, as in a tug-of-war, between the atoms nuclei • In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond

16. Bond Polarity • When two different atoms bond covalently, there is an unequal sharing • the more electronegative atom will have a stronger attraction, and will acquire a slightly negative charge • called a polar covalent bond, or simply polar bond.

17. Electronegativity? The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994

18. Table of Electronegativities Higher electronegativity

19. Bond Polarity • Refer to Table • Consider HCl H = electronegativity of 2.1 Cl = electronegativity of 3.0 • The electronegativity difference: 3.0 - 2.1 = 0.9 • the bond is polar • the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge

20. Bond Polarity #1 The more electronegative atom attracts electrons more strongly and gains a slightly negativecharge. The less electronegative atom has a slightly positivecharge. • Table shows how the electronegativity can also indicate the type of bond that tends to form

21. Bond Polarity • Only partial charges, much less than a true 1+ or 1- as in ionic bond #2Written as: H Cl • the positive and minus signs (with the lower case delta: ) denote partial charges. d+d- d+ andd-

22. Bond Polarity #3 Can also be shown: • the arrow points to the more electronegative atom. • Table shows how the electronegativity can also indicate the type of bond that tends to form H Cl

23. Electronegativity differences/Bond Types

24. Polar molecules • A polar bond tends to make the entire molecule “polar” • areas of “difference” • HCl has polar bonds, thus is a polar molecule. • A molecule that has two poles is called dipole, like HCl

25. Polar molecules • The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape • carbon dioxide has two polar bonds, and is linear = nonpolar molecule! • Bond polarities cancel (are in opposite directions)

26. Polar molecules • The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape • water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e- away from H = very polar!

27. 6.5 Practice problems p. 202 45. For each of the following polar molecules, indicate the direction of the resulting dipole: • H - F • H - Cl • H - Br • H - I

28. Polar molecules • When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates.

29. 6.5 Practice problems p. 202 47. On the basis of individual bond polarity and orientation, determine whether each of the following molecules would be polar or nonpolar: • H20 • I2 • CF4 • NH3 • CO2

30. Chapter 6 Section5 Molecular Geometry Intermolecular Forces • The forces of attraction between molecules are known asintermolecular forces. • The boiling point of a liquid is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules. • Intermolecular forces vary in strength but are generally weaker than bonds between atoms within molecules, ions in ionic compounds, or metal atoms in solid metals. • Boiling points for ionic compounds and metals tend to be much higher than those for molecular substances: forces between molecules are weaker than those between metal atoms or ions.

31. Chapter 6 Comparing Ionic and Molecular Substances Section5 Molecular Geometry

32. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued • The strongest intermolecular forces exist between polar molecules. • Because of their uneven charge distribution, polar molecules have dipoles. A dipole is created by equal but opposite charges that are separated by a short distance. • The direction of a dipole is from the dipole’s positive pole to its negative pole.

33. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued • A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows:

34. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued • The negative region in one polar molecule attracts the positive region in adjacent molecules. So the molecules all attract each other from opposite sides. • Such forces of attraction between polar molecules are known as dipole-dipole forces. • Dipole-dipole forces act at short range, only between nearby molecules. • Dipole-dipole forces explain, for example the difference between the boiling points of iodine chloride, I–Cl (97°C), and bromine, Br–Br (59°C).

35. Chapter 6 Comparing Dipole-Dipole Forces Section5 Molecular Geometry

36. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued • A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons. • The result is a short-range intermolecular force that is somewhat weaker than the dipole-dipole force. • Induced dipoles account for the fact that a nonpolar molecule, oxygen, O2, is able to dissolve in water, a polar molecule.

37. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued • Some hydrogen-containing compounds have unusually high boiling points. This is explained by a particularly strong type of dipole-dipole force. • In compounds containing H–F, H–O, or H–N bonds, the large electronegativity differences between hydrogen atoms and the atoms they are bonded to make their bonds highly polar. • This gives the hydrogen atom a positive charge that is almost half as large as that of a bare proton.

38. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued • The small size of the hydrogen atom allows the atom to come very close to an unshared pair of electrons in an adjacent molecule. • The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule is known as hydrogen bonding.

39. Chapter 6 Section5 Molecular Geometry Intermolecular Forces • Hydrogen bonds are usually represented by dotted lines connecting the hydrogen-bonded hydrogen to the unshared electron pair of the electronegative atom to which it is attracted. • An excellent example of hydrogen bonding is that which occurs between water molecules. The strong hydrogen bonding between water molecules accounts for many of water’s characteristic properties.

40. Chapter 6 Hydrogen Bonding Section5 Molecular Geometry

41. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued London Dispersion Forces • Even noble gas atoms and nonpolar molecules can experience weak intermolecular attraction. • In any atom or molecule—polar or nonpolar—the electrons are in continuous motion. • As a result, at any instant the electron distribution may be uneven. A momentary uneven charge can create a positive pole at one end of an atom of molecule and a negative pole at the other.

42. Chapter 6 Section5 Molecular Geometry Intermolecular Forces, continued London Dispersion Forces, continued • This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between temporary dipoles. • The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces. • Fritz London first proposed their existence in 1930.

43. Intermolecular Attractions • They are what make solid and liquid molecular compounds possible. • Weaker than either ionic or covalent • The weakest are called van der Waal’s forces • Two kinds: • London dispersion forces • dipole interactions

44. Intermolecular Attractions #1.London dispersion forces - weakest force - caused by motion of e- - increases as # e- increases - Increases as atomic mass increases - only intermolecular force acting on noble gases and nonpolar molecules -Explains Halogens: start as gases; bromine is liquid; iodine is solid – all in Group 17

45. Dipole interactions #2. Dipole-Dipole Forces • Occurs when polar molecules are attracted to each other • positive region of one molecule attracts the negative region of another molecule. • Slightly stronger than London dispersion forces. • Opposites attract, but not completely hooked like in ionic solids. • responsible for high boiling points • includes Hydrogen Bonding

46. d+d- d+d- d+d- d+d- H F O O H F H O H #2. Dipole interactions • Slightly stronger than London dispersion forces. • Dipole-induced dipole = polar molecule attracts the e- in a nonpolar molecule

47. d+d- d+d- d+d- d+d- d+d- d+d- d+d- #2. Dipole Interactions d+d-

48. #3. Hydrogen bonding • …is the attractive force caused by hydrogen bonded to N, O, F, or Cl • N, O, F, and Cl are very electronegative, so this is a very strong dipole. • And, the hydrogen shares with the lone pair in the molecule next to it. • This is the strongest of the intermolecular forces.

49. #3. Hydrogen bonding defined: • When a hydrogen atom is: • a) covalently bonded to a highly electronegative atom, AND • b) is also weakly bonded to an unshared electron pairof a nearby highly electronegative atom. • The hydrogen is left very electron deficient (it only had 1 to start with!) thus it shares with something nearby • Hydrogen is also the ONLY element with no shielding for its nucleus when involved in a covalent bond!

50. d- d+ O d+ H d+ d- H H O H d+ Hydrogen Bonding(Shown in water) This hydrogen is bonded covalently to: 1) the highly negative oxygen, and 2) a nearby unshared pair.