Unit 7

1. Unit 7 Lewis diagrams molecular geometry bond and molecular polarity IMFAs

2. Lewis dot diagrams • add up the total number of valence electrons for all atoms in the molecule • arrange the atoms to pair up the separate atoms’ single electrons as much as possible • confirm that: • the total number of electrons exactly matches the total valence electrons of the original atoms, and • each atom has an octet of electrons (8), except • H and He have a duet of electrons (2)

3. structural formulas • also called “Lewis structures” or “Lewis diagrams” (but not “Lewis dot structures”) • replace each shared pair of electrons with a solid line representing a covalent bond consisting of two shared electrons • continue to show the lone pairs of electrons (which are unshared) • double-check that the lone pairs plus bond pairs still add up to the correct total number of valence electrons

4. multiple bonds • additional bonds may need to be added to a Lewis structure if • single electrons remain • atoms do not have octets • in simple cases, you may be able to pair up single electrons on adjacent atoms to form additional bonds, e.g. • CO2 • N2 • C2H4

5. multiple bonds • in other cases, you cannot strictly keep electrons with their original atoms; the electrons are free to move elsewhere in the molecule as needed to complete octets, e.g. • carbon monoxide, CO • ozone, O3 • in these cases, atoms may not form their “normal” number of bonds • but the total number of valence electrons must not change; they are just rearranged

6. multiple bondscomputational approach • you can also calculate exactly how many bonds are in a molecule in the following way • add up the valence electrons that the atoms in the molecule actually have • separately add up the valence electrons those atoms need in order to have noble gas configurations • calculate the difference, need – have • that difference is the number of shared electrons the molecule must have • every 2 shared electrons make one bond

7. multiple bondscomputational approach O2 CO O O C O have: 6 + 6 = 12 have: 4 + 6 = 10 need: 8 + 8 = 16 need: 8 + 8 = 16 4 shared e- 6 shared e- thus 2 bonds thus 3 bonds • after building the basic skeleton with bonds • add remaining electrons as needed to complete octets • double-check that the total number of electrons is exactly the number of valence electrons (“have”)

8. general hints for Lewis structures • if a given molecule can be drawn with both symmetrical and asymmetrical structures, the symmetrical one is more likely to be correct • central atoms are often • written first in the formula • the least electronegative element • the element that can form the most bonds • hydrogen and halogens • only form one bond, thus are terminal atoms • are generally interchangeable in molecules

9. exceptions to octet “rule” • most atoms have octets (8 valence electrons) when in molecules, but there are exceptions

10. molecular shapes: VSEPR model • valence shell electron-pair repulsion • groups of electrons naturally find positions as far apart from each other as possible • different molecular shapes result based on how many groups of electrons are present • each of the following counts as one “set” of electrons around the central atom • a lone pair • a single bond (2 shared e-) • a double or triple bond (4 or 6 shared e-)

11. VSEPR model—central atom with: 2 sets of e– 3 sets of e– 4 sets of e– 5 sets of e– 6 sets of e– trigonal bipyramidal linear trigonal planar tetrahedral octahedral e.g. BeF2 e.g. BF3 e.g. CF4 e.g. XeF6 e.g. SF5

12. electron geometry vs. molecular shape • each set of electrons occupies a position around the central atom • the number of sets defines the electron geometry • but lone pairs are essentially transparent • even though they are invisible, lone pairs make their presence known by distorting the positions of the bonds around them (since lone pairs repel the electrons in the bonds) • this results in several related molecular shapes within each general class of electron geometry

13. tetrahedral electron geometry4 electron sets

14. tetrahedral electron geometry

15. tetrahedral electron geometry

16. triangular planar electron geometry3 electron sets

17. linear electron geometry2 electron sets • in addition, any diatomic molecule must be linear (since any two points lie on a line)

18. triangular planar and linearelectron geometry

19. bond polarity • two electrons shared between two atoms form a covalent bond • if those electrons are shared equally (or nearly equally), it is a non-polar covalent bond • if one atom attracts the electrons much more strongly than the other atom, it is a polar covalent bond • if one atom completely removes an electron from the other atom, the result is an ionic bond

20. bond polarity • the electronegativity difference between the two atoms determines how polar a bond is 0.0 – 0.4 0.5 – 1.7 > 1.7 Cℓ2 HCℓ LiCℓ

21. bond polarity • dipole moment is the actual measureable quantity related to bond polarity • the size of the dipole moment is affected by • electronegativity difference • bond length • we will focus on ΔEN and a qualitative sense of bond polarity

22. molecular polarity • the overall polarity of a molecule depends on the combined effect of the individual polar bonds individual bonds polar individual bonds polar overall molecule nonpolar overall molecule polar

23. molecular polarity • what allows bond dipoles to cancel? • geometric symmetry of the molecule • having identical terminal atoms (or atoms with the same electronegativity) • what prevents bond dipoles from canceling? • geometric asymmetry (due to lone pairs) • having different terminal atoms

24. molecular polarity

25. molecular polarity • inherently symmetrical shapes (if all surrounding atoms are the same) • tetrahedral • triangular planar • linear • inherently asymmetrical shapes • bent • triangular pyramid • even symmetrical shapes become asymmetrical if different terminal atoms are attached

26. IMFA: intermolecular forces of attraction “mortar”— holds the separate pieces together (i.e. the IMFA) “bricks”— individual atoms, ions, or molecules of a solid

27. IMFA: intermolecular forces of attraction

28. types of IMFA strongest occurs between covalent network atoms such as C, Si, & Ge (when in an extended grid or network) ionic bond cations and anions (metals with non-metals in a salt) metallic bond metal atoms hydrogen bond ultra-polar molecules (those with H–F, H–O, or H–N bonds) dipole-dipole attraction polar molecules van der Waals forces London forces non-polar molecules weakest

29. IMFAs – Trends and Characteristics • melting points and boiling points • Solubility • Conductivity (type specific) • Strength (type specific)

30. Melting Points and Boiling Points • stronger IMFAs cause higher m.p. and b.p. • atoms and molecules that are heavier and/or larger generally have higher m.p. and higher b.p. • larger e– clouds can be distorted (polarized) more by London or dipole forces, causing greater attraction • CH3CH2CH2CH2CH3 > CH3CH2CH3 • Polarity (for dipole dipole and H-bonds) more polar=higher b.p, m.p. • HCl > HI

31. Melting Points and Boiling Points • strategy to predict m.p. and b.p. • first sort atoms/molecules into the six IMFA categories • then sort within each IMFA category from lightest to heaviest (or least polar to most polar)

32. same IMFA: sort by molar mass melt boil • ex: halogen family • all are non-polar (London force) • lowest to highest m.p. and b.p. matches lightest to heaviest I2 (257) • thus at room temperature: • F2 (g) • Cℓ2 (g) • Br2 (ℓ) • I2 (s) +184.4 +150 I2 (257) +113.7 +100 Br2 (160) +58.8 +50 Br2 (160) –7.2 0 Cℓ2 (71) –34.04 –50 –100 Cℓ2 (71) –101.5 F2 (38) –150 –182.95 –200 F2 (38) –219.62 –250 °C

33. same mass: sort by IMFA type ethylene glycol (can form twice as many H-bonds) • ex: organic molecules • all are ~60 g/mol • different types of IMFA +198 +150 +100 1-propanol (ultra-polar = H-bonds) +97.4 acetone (more polar) +56.2 +50 methyl ethyl ether (slightly polar) +10.8 0 butane (non-polar) –0.5 • the stronger the IMFA, the higher the boiling point –50 °C

34. Solubility • substances generally mix best with others of similar IMFAs • ”like dissolves like” • non-polar mixes well with non-polar • polar mixes well with polar (and ultra-polar / ionic)

35. other physical properties • strength, conductivity, etc. are related to the type of IMFA

36. details about each IMFA strongest covalent network ionic bond metallic bond hydrogen bond dipole-dipole attraction London forces weakest

37. London (or dispersion) forces • non-polar molecules (or single atoms) normally have no distinct + or – poles • electron clouds are slightly distorted by neighboring molecules • sort of like water sloshing in a shallow pan • form temporary dipoles • Low MP and BP • soluble

38. London dispersion forces in action 1. temporary polarization due to any random little disturbance δ+ δ- 2. induced polarization caused by neighboring molecule 3. induced polarization spreads 4. induced polarization reverses non-polar molecules, initially with uniform charge distribution

39. dipole-dipole attractions • polar molecules have permanent dipoles • the molecules’ partial charges (δ+, δ-) attract the oppositely-charged parts of neighboring molecules • this produces stronger attraction than the temporary polarization of London forces • therefore polar molecules are more likely to be liquid at a temperature where similar non-polar molecules are gases

40. hydrogen bonding (or ultra-dipole attractions) • H—F, H—O, and H—N bonds are more polar than other similar bonds • Very small atoms, particularly H • F, O, and N are the most electronegative elements • particularly polar • molecules containing these bonds have much higher m.p. and b.p than otherwise expected for non-polar or polar molecules of similar mass • the geological and biological systems of earth would be completely different if water molecules did not H-bond to each other

41. hydrogen bonding (or ultra-dipole attractions) ultra-polar molecule (much higher boiling point) hydrogen bonds (between molecules, not within them) non-polar molecules (lower boiling points)

42. hydrogen bonding (or ultra-dipole attractions) Beware!! These are not hydrogen bonds. They are normal covalent bonds between hydrogen and oxygen. O O O O H H H H H H H H These are hydrogen bonds. They are between separate molecules (not within a molecule).

43. metallic bonding • structure • nuclei arranged in a regular grid or matrix • “sea of electrons”—delocalized valence electrons free to move throughout grid • resulting properties • conductive (electrically and thermally) • strong, malleable, and ductile shiny • Form alloys = mixture of metals • Bronze = copper and tin • Brass = copper and zinc • Steel = iron and carbon metallic “bond” is generally weaker than covalent bond since there are not specific e– pairs forming bonds

44. ionic bonding (salts) • structure: orderly 3-D array (crystal) of alternating + and – charges • made of • cations (metals from left side of periodic table) • anions (non-metals from right side of periodic table) • properties • non-conductive when solid • conductive when melted or dissolved • hard but brittle (why?)…

45. why are salts hard but brittle? 1. apply some force 2. layer breaks off and shifts 4. shifted layer shatters away from rest of crystal 3. + repels + – repels –

46. covalent networks • strong covalent bonds hold together millions of atoms (or more) in a single strong particle • properties • very high melting temperatures • usually non-conductive (except graphite) • very hard, very strong • examples • carbon (two allotropes: diamond, graphite) • pure silicon or pure germanium • SiO2 (quartz or sand) • other synthetic combinations averaging 4 e– per atom: • SiC (silicon carbide), BN (boron nitride)

47. m.p. = 3550°C C60 buckminsterfullerine “bucky ball” m.p. = ~1600°C

48. summary of properties strongest strength m.p. & b.p. conductive? network extremely hard very high usually not ionic hard but brittle medium to high if melted or dissolved (mobile ions) strong, malleable, ductile medium to high very (delocalized e–) metallic hydrogen dipole soft and brittle low no van der Waals forces London weakest

49. Metallic London Ionic ---------- Hydrogen Metallic Ionic Metallic Network Network Metallic Network Metallic Hydrogen Ionic Covalent

50. D F H B G J C I E A