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Chemistry 20

Chemistry 20. Chapter 2. PowerPoint presentation by R. Schultz. robert.schultz@ei.educ.ab.ca. •. •. ‾. Na +. Cl. •. •. •. •. •. •. electron transfer. 2.1 Three-Dimensional Structures. Ionic Recall electron transfer (Chapter 1)

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Chemistry 20

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  1. Chemistry 20 Chapter 2 PowerPoint presentation by R. Schultz robert.schultz@ei.educ.ab.ca

  2. • • ‾ Na + Cl • • • • • • electron transfer 2.1 Three-Dimensional Structures • Ionic • Recall electron transfer (Chapter 1) • As ions form, they group – not in pairs – in stable crystals, of anions and cations in an arrangement so that each cation is surrounded by anions and the reverse • Crystal is called (ionic) crystal lattice • • + Na + Cl • • • • • •

  3. photo of NaCl crystal: note its cubic shape – matches theoretical model fig 2.1 A, page 48 2.1 Three-Dimensional Structures • NaCl crystal lattice: • In the crystal there are equal numbers of Na+ and Cl‾ fig 2.2, page 48

  4. 2.1 Three-Dimensional Structures • Formula is written NaCl – this is the formula unit – the simplest ratio of anions to cations in the lattice – same as formula from balancing charges • Different ionic compounds have different crystal shapes, 2 factors: • Relative ion size (see fig 2.4A, page 69) • Relative ion charge (see fig 2.4B, page 69)

  5. 2.1 Three-Dimensional Structures • There are no molecules of ionic compounds, the term is formula unit

  6. 2.1 Three-Dimensional Structures • Molecular • The smallest unit of a molecular compound is a molecule • Molecular structures can be determined using Lewis Diagrams along with the VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) Theory • Theory based on repulsion between valence electron pairs, both lone (LP) and bonding (BP) pairs

  7. 2.1 Three-Dimensional Structures • Valence electron pairs move as far apart as possible to minimize repulsion • Repulsion order: LP-LP > LP-BP > BP-BP • Shapes around central atoms can be determined using the following chart (an improvement of chart page 55 of your text)

  8. 2.1 Three-Dimensional Structures not necessary to state

  9. 2.1 Three-Dimensional Structures • Work through examples using chart as reference: CH4, NH3, H2O, C2H4, SO2, C2H2 • VSEPR Worksheet (BLM 2.1.8B) – “prelab” • Models tomorrow

  10. 2.1 Three-Dimensional Structures 0 LP’s,4 BP’s Each electron pair around C is a BP. They move to get as far apart as possible. Shape: tetrahedral Bond angle: 109º H •• •• •• H C H •• H behind page flat on page whenever there are 0 LP’s and 4 BP’s around central atom shape will be tetrahedral out of page

  11. •• •• •• N H H •• H 2.1 Three-Dimensional Structures 1 LP, 3 BP’s around Nall repel each other shape like tetrahedral, but one bond space taken up by LP shape name: pyramidal (book calls it “tetrahedral (trigonal pyramidal)”) bond angle approx 107º a little smaller than 109º since LP-BP repulsion greater than BP-BP whenever there is 1 LP and 3 BP’s around central atom, shape will be pyramidal

  12. •• •• •• O H •• H O H H 2.1 Three-Dimensional Structures 2 LP’s, 2 BP’s around Oall repel each other shape like tetrahedral, but 2 bond spaces taken up by LP’s shape name: bent or v-shaped (book calls it “tetrahedral (bent)”) bond angle approx 105º a little smaller than 109º since LP-BP repulsion greater than BP-BP – there are now 2 LP-BP sets whenever there are 2 LP’s and 2 BP’s around central atom, shape will be v-shaped or bent

  13. H •• •• •• •• •• H H C C •• H H H C = C 120º 120º 120º 120º H H 120º 120º flat on page! 2.1 Three-Dimensional Structures New rule: count multiple bonding pairs as if they were single bonding pairs 2 central atoms, each with 0 LP’s, 3 BP’s around themall repel each other shape name: trigonal planar bond angle 120º whenever there are 0 LP’s and 3 BP’s around central atom, shape will be trigonal planar

  14. •• •• •• •• •• •• •• •• S O O •• 2.1 Three-Dimensional Structures 1 LP, 2 BP’s around Sall repel each other shape like trigonal planar, but 1 bond space taken up by LP shape name: trigonal planar (bent) bond angle approx 120º whenever there is 1 LP and 2 BP’s around central atom, shape will be trigonal planar (bent)

  15. •• H •• •• •• •• C H C H C C H 2.1 Three-Dimensional Structures 0 LP’s, 2 BP’s around each Call repel each other shape name: linear bond angle approx 180º whenever there is 0 LP’s and 2 BP’s around central atom, shape will be linear

  16. 2.1 Three-Dimensional Structures • You must know the chart – it’s not just memory • know 3 lines; the rest can be worked out from that • Do worksheet BLM 2.1.8C

  17. 2.1 Three-Dimensional Structures • Polar Molecules • Chapter 1 – polar covalent bonds • Follow-up → polar molecules • Whether or not a molecule is polar affects its bonding and therefore its chemical and physical properties

  18. O H H non-polar polar 2.1 Three-Dimensional Structures • To be polar, molecule must have at least 1 polar covalent bond • Having 1 or more polar covalent bonds doesn’tnecessarily make molecule polar • Direction matters O = C = O

  19. 2.1 Three-Dimensional Structures e.g. CO2 e.g. COS Note: arrow directions are relative! e.g. H2O e.g. SO3 e.g. CH2O e.g. CH4 e.g. CH3Cl e.g. NH3

  20. H C C H F 2.1 Three-Dimensional Structures • Polarity of more complex molecules can be determined using the same methods • e.g. C2H5F • Do worksheet BLM 2.1.12B If F had been an H, dipole would’ve been in opposite direction and net dipoles from the 2 C’s in opposite directions, making it non-polar, but ………

  21. 2.1 Three-Dimensional Structures • Generalizations for later: • all hydrocarbons & halocarbons, CxHy, e.g. CH4, C3H8, C2F6 will be non-polar • all monosubstituted hydrocarbons & halocarbons CxHyX, e.g. C2H5Br, CH3OH, CHI3 will be polar • These will save you a lot of time later in chapter

  22. tetrahedral Fig 2.17B, page 60 2.1 Three-Dimensional Structures • Network Solids (also called Covalent Networks) covalent bonds – strongest chemical bonds • atoms bonded together by covalent bonds in a molecule made up of gigantic # of atoms • Examples: • Diamond, C(s) – Each C is covalently bonded to 4 other C’s in a arrangement 1 diamond is 1 gigantic molecule!

  23. trigonal planar fig 2.17A, page 60 2.1 Three-Dimensional Structures • Graphite, C(s) – Each C is covalently bonded to 3 other C’s in a arrangement forming 2-dimensional layers • Weaker bonds between layers • Why is diamond hard and brittle with a very high melting point (mp) while graphite is slippery with a somewhat lower mp? Each layer of crystal is 1 molecule

  24. 2.1 Three-Dimensional Structures • other allotropes of C(s): fullerenes, carbon nanotubes, graphene (Read pages 60, 61) • Silicon carbide, SiC(s) – structure like graphite, but mixture of Si and Chigh mp, hard, brittle like diamond – used in grindstones, carbide drill bits and saw blades • Silicon dioxide, SiO2(s) – main ingredient of sand – high mp, hard, brittle fig 2.19, page 61

  25. 2.2 Intermolecular Forces • Intramolecular Forces (intramolecular bonds) – bonds between atoms in a molecule or between atoms in a polyatomic ion, a.k.a. : • Intermolecular forces (intermolecular bonds) – bonds among molecules in molecular substances covalent bonds

  26. intermolecular forces O O intramolecular forces H H H H 2.2 Intermolecular Forces (polar) covalent bonds intermolecular forces are important in understanding the differing boiling points (bp’s) and mp’s of molecular compounds

  27. 2.2 Intermolecular Forces • Types of intermolecular forces: • dipole-dipole forces • hydrogen bonding • London (dispersion) forces • dipole-dipole forces (polar molecular compounds only) Also called van der Waals Forces

  28. δ- bond dipoles molecular dipole δ+ O O O O O O O O O O O O O H H H H H H H H H H H H H H H H H H H H H H H H H H 2.2 Intermolecular Forces • Recall polar molecules, e.g. H2O fig 2.21 page 63 “+” pole of 1 molecule attracts “–” poles of neighbours (much weaker than ionic bonds)

  29. 2.2 Intermolecular Forces • hydrogen bondinga special kind of dipole-dipole force that occurs when H is covalently bonded to a very electronegative atom (F, O, or N) • easy way to remember – instead of hydrogen bonding, think hydrogen FONding • strongest intermolecular bond in most situations • H has only 1 electron; when bonded to a very electronegative atom that electron is far

  30. •• O H •• •• •• H very +’ve •• •• hydrogen bond O O H H •• •• •• •• •• •• H H 2.2 Intermolecular Forces from the H, making it very +’ve. • The attraction of the +’ve H for electron pairs on the other atom is a hydrogen bond remember that actual shape around O is v-shaped leading to …

  31. 2.2 Intermolecular Forces Fig 2.23A, page 65 This accounts for the hexagonal configuration of snowflakes, for the fact that ice floats on water, and for water’s high surface tension Remember the pennies in the glass?

  32. 2.2 Intermolecular Forces • London (Dispersion) Forces • Many molecular compounds have neither dipole-dipole forces nor hydrogen bonding – they are non-polar and there is no charge-charge interaction, therefore……. • According to theory there would be no attraction between these molecules and their only possible state would be gas (at any temperature)

  33. wall  2.2 Intermolecular Forces • London Dispersion Forces (LDF) are present in ALL molecular compounds • some molecular compounds have LDF along with the other forces • illustration: charged balloon will be attracted to the uncharged wall because charges are induced in the wall as shown • motion of electrons in atom canproduce temporary dipoles fig 2.25, page 68

  34. 2.2 Intermolecular Forces • the temporary dipoles induce temporary dipoles in neighbouring molecules • attraction between these temporary dipoles and the induced dipoles is the source of London Dispersion Forces • Factors affecting strength of LDF: • total number of electrons in the molecule • shape of the molecule* and their neighbouring molecules

  35. 2.2 Intermolecular Forces • Read about the importance of hydrogen bonding in the DNA double helix and in myoglobin in muscle tissue • page 69-70

  36. 2.3 Relating Structures and Properties • State of Matter (mp and bp) page 72

  37. 2.3 Relating Structures and Properties • raising temperature is raising Ek of particles • the stronger the appropriate bonding, the higher the mp and bp of the substance • appropriate bonding: network covalents – covalent bondingionics – ionic bondingmetals – metallic bondingmoleculars – LDF, dipole-dipole, H-bonding(note: even though moleculars have covalent bonding this has nothing to do with mp and bp)

  38. 2.3 Relating Structures and Properties • Relative bond strengths: covalent > ionic > metallic > hydrogen > dipole-dipole > LDF(LDF varies greatly) strongest weakest

  39. 2.3 Relating Structures and Properties high mp & bp; metallic bonding is strong – no differentiation high mp & bp; ionic bonding is strong – no differentiation chart, page 72 low mp & bp; intermolecular forces are weak - differentiation

  40. 2.3 Relating Structures and Properties • Hydrogen bonding has a great affect on boiling points: Why do bp’s increase as you go from left to right? effect of H bonding How come group 14 always is lowest? Where water would be expected to be without H bonding fig 2.33, page 74 10 e‾ 18 e‾ 36 e‾ 54 e‾

  41. 2.3 Relating Structures and Properties • Examples: • Predict which will have the higher boiling point in each pair: • a) C2H6 or C3H8 both molecular: intermolecular forces only LDF 18 e‾ 26 e‾ DDF no no H bonds no no C3H8 will have higher bp due to stronger LDF

  42. 2.3 Relating Structures and Properties • b) CH3F and CH3OH LDF 18 e‾ 18 e‾DDF yes yesH bonds no yes Both molecular: intermolecular forces only CH3OH will have the higher bp because of hydrogen bonding Do worksheet BLM 2.3.3B

  43. 2.3 Relating Structures and Properties • Extension: Rank the following substances in order of  bp. State relevant bond types for each • CH3OH, SiC, PH3, LiF, CH3F, Mg Step 1: Identify moleculars – they will be lowest M M M LDF 18 e‾ no 18 e‾ no 18 e‾ no DDF yes no no no yes noH bonds yes no no no no no Other covalent ionic metallic (network) Answer: PH3, CH3F, CH3OH, Mg, LiF, SiC

  44. 2.3 Relating Structures and Properties • Note: on your test the substances will always be in pairs as on your previous worksheet • unlike the previous worksheet, not everything will necessarily be molecular – i.e. you may get a molecular with an ionic, an ionic with a covalent network, 2 moleculars; any grouping is possible

  45. 2.3 Relating Structures and Properties • Why are metals malleable according to the metallic bonding model? (possible question on next test) • Malleable means workable – bends without breaking; dents without shattering • metallic bonding model: an array of cations in a sea of valence electrons • because bonds are between cations and sea of electrons rather than between cations and cations, cations may be moved in the sea of electrons without breaking bonds

  46. 2.3 Relating Structures and Properties • Ionic compounds are not malleable – they shear along lines as diagram indicates fig 2.36, page 75

  47. magnified section of wire e- e- e- e- + e- e- 2.3 Relating Structures and Properties • Why do metals conduct electricity according to the metallic bonding model? (possible question on next test) +

  48. 2.3 Relating Structures and Properties • electricity is the flow of electrons • metallic bonding model: an array of cations in a sea of valence electrons • since valence electrons are loosely held to individual nuclei, they are delocalized – as a new electron enters the wire, another electron must leave to keep wire neutral

  49. 2.3 Relating Structures and Properties • Conductivity (others) • solid ionic: page 76

  50. 2.3 Relating Structures and Properties • aqueous ionic: page 77

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