Ch.2 Polar Bonds and Their Consequences. 2.1 Polar Covalent Bonds and Electronegativity. polar covalent bonds: electron distribution is unsymmetrical

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1 2.1 Polar ovalent Bonds and Electronegativity polar covalent bonds: electron distribution is unsymmetrical Ionic haracter δ+ δ- + - X Y X Y X Y symmetrical covalent bond polar covalent bond ionic bond

3 Elecrtonegativity (EN): the intrinsic ability of an atom to attract electrons in a covalent bond N F Li Be Si P S l Na Mg Br K a I s EN difference: (polar bond), >2.0 (ionic)

4 direction of polarity l δ- δ+ MgBr δ+ δ- organometallic compounds

5 inductive effects: shifting of electrons in σ-bond in response of the electronegativity of nearby atoms (atom's ability to polarize a bond) ; polarize bonds, chemical reactivity l electronegative atoms inductively withdraw electrons MgBr metals inductively donates electrons

7 2.2 Polar ovalent Bonds and Dipole Moment dipole moment (μ): net molecular polarity (vector sum) ; resulted from individual bond polarities and lone pair contributions ; polar substances dissolve in polar solvents, nonpolar substances soluble in nonpolar solvents µ = Q x r Q: charge at either end of dipole r: distance between the charges 1 debyes (D): x m For one positive charge and one negative charge seperated by 100 pm 100 pm B - A + µ = Q x r = (1.60 x ) x (100 x m)(1d / x m) = 4.80 D

8 δ- l δ+ µ = 1.87 D -l: 178 pm δ- l δ+ 0.2 electron excess 0.2 electron deficient For fully seperated chloromethane (if the -l bond were ionic, + l - ) µ = D x 1.78 = 8.54 D but the actual dipole moment is 1.87 D 1.87 x 100= 22% 8.54 therefore, the chlorine atom in chloromethane has an excess of about 0.2 electron, and the carbon has a deficiency of about 0.2 electron

9 N net N net nonbonding electron pair: stick out from positively charged nuclei

11 Table 2.1 dipole moments of some compounds Dipole moment (D) Nal 9.0 Dipole moment (D) N N l l =N=N 1.50 BF 3 0

12 symmetrical molecules: μ = 0 D

13 2.3 Formal harges related to bond polarity, dipole moment ; common for atoms that have an apparently abnormal number of bonds N formal positive charge formal negative charge Nitromethane

14 electron counts - covalent bond: each atom owns one electron ; single bond (1 e - ); double bond (2 e - ); triple bond (3 e - ) - non-bonding electrons: owned by the atom 4 valence electrons for isolated carbon atom 8 2 = 4 valence electrons 5 valence electrons N N for isolated nitrogen atom = 5 valence electrons

15 Formal harge = # of valence electrons in free atom _ # of valence electrons in bound atom = # of valence electrons _ half of bonding electrons _ # of nonbonding electrons

16 N Nitromethane nitrogen valence electrons nitrogen bonding electrons nitrogen nonbonding electrons oxygen valence electrons oxygen bonding electrons oxygen nonbonding electrons = 5 = 8 = 0 Formal harge = 5-8/2-0 = + 1 = 6 = 2 = 6 Formal harge = 6-2/2-6 = - 1 For neutral molecules: sum of the formal charges equal to zero

17 Table 2.2 A summary of formal charges on atoms N N N dipolar molecule: neutral overall but have +/- charges on individual atoms ; dipolar character of molecules chemical reactivity

18 Practice formal charge Dimethyl sulfoxide or 3 S 3 3 S 3 3 S 3 Diazomethane 2 N N 2 N N Acetonitrile 3 N 3 N Methyl isocyanide 3 N 3 N

19 2.4 Resonance h.2 Polar Bonds and Their onsequences For some molecules simple Lewis structure cannot describe the actual structure correctly For example, nitromethane has two equivalent N- bonds, but one Lewis structure can't represent it N N - Nitromethane Both oxygen-nitrogen bonds are 122 pm in length, midway between the length f a typical N- single bond (130 pm) and a typical N= double bond (116 pm)

20 resonance form: individual Lewis structures that represent one molecule

21 2.5 Rules for Resonance Forms 1. Individual resonance forms are imaginary, not real; The real structure is composite (or resonance hybrid) of different forms. 2. Resonance forms differ only in the placement of their π or nonbonding electrons; Neither the position nor the hybridization of any atom changes from one resonance form to another. N N - Nitromethane A curved arrow shows movement of two electrons.

22 Benzene (two resonance forms)

23 3. Different resonance forms of a substance don't have to be equivalent; The real structure is composite (or resonance hybrid) of different forms. 3 3 strong base 3 3 more stable form (negative charge on electronegative oxygen) For non-equivalent resonance forms, the actual structure of the resonance hybrid is close to the more stable form than the less stable form.

24 4. Resonance forms must be valid Lewis structures and obey normal rules of valency; The octet rule still applies electrons on carbon Acetate ion (two resonance forms) NT a valid resonance form 5. The resonance hybrid is more stable than any individual resonance form; The larger the number of resonance forms, the more stable a substance is.

25 2.6 Drawing Resonance Forms Any three atom grouping with a multiple bond has two resonance forms. X Y * Z X Y Z* *X Y Z * X Y Z multiple bond (double or triple bond) * = 0, 1, or 2 electrons Y X Z

30 Practice Resonance forms arbonate ion: 2-3

31 Practice Resonance forms allylic radical Pentadienyl radical A half-headed curved arrow shows movement of one electron.

32 2.7 Acids and Bases: The BrØnsted-Lowry Definition BrØnsted-Lowry Definition: proton donor/acceptor Lewis Definition: electron pair donor/acceptor BrØnsted-Lowry Acid: proton ( + ) donor BrØnsted-Lowry Base: proton ( + ) acceptor

33 In general, -A + B -B + + A - Acid Base onjugate acid onjugate base Acid Base onjugate acid onjugate base N 2 N Acid Base onjugate acid onjugate base

34 2.8 Acid and Base Strength -A A - K eq = [ 3 + ][A - ] [A][ 2 ] in dilute solution [ 2 ] remains constant [ 2 ] = 55.6 M for pure water Acidity constant, K a K a = K eq [ 2 ] = [ 3 + ][A - ] [A] pk a = -logk a stronger acid: larger K a, smaller pk a weaker acid: smaller K a, larger pk a

35 Table 2.3 Relative strength of some common acids and their conjugate bases Acid pk a onjugate base stronger acid N 3 F N 3 l N F - N 3 - l - stronger base

36 -A + B -B + + A - Acid Base onjugate acid onjugate base - inverse relationship between the acid strength and its conjugate base strength - a strong acid loses + easily ; its conjugate base has little affinity for + - a weak acid loses + with difficulty ; its conjugate base has high affinity for + for example, l is a strong acid ; l - does not hold + tightly and is thus a weak base

37 2.9 Predicting Acid-Base Reactions from pk a Values An acid with a lower pk a will go to the conjugate acid with a higher pk a Acetic acid (pk a = 4.76) ydroxide ion Acetate ion Water (pk a = 15.74) stronger acid stronger base weaker base weaker acid

39 Practice Acidity Na + N 2 -? 3 2 Na + + N 3 Acetone (pk a ~ 19) Amonia (pk a ~ 35) + -? + - pk a = 25 pk a = N 2 pk a = 25 + N 3 pk a = 35

40 2.10 rganic Acids and rganic Bases organic acids are two kinds: - and =-- 3 pk a = Anion is stabilized by having negative charge on a highly electronegative atom pk a = 4.76 Anion is stabilized by having negative charge on a highly electronegative atom and by resonance

42 organic acids pk a = 19.3 Anion is stabilized resonance and having negative charge on a highly electronegative atom 3 3 pk a = 10 pk a = 9 3 pk a = 4.76

43 rganic bases: only one main kind- nitrogen atom -l + 3 N 3 N l - N Pyridine ( 3 2 ) 3 N Triethylamine xygen-containing compounds: act as bases with strong acids also act as acids (-, -) with bases 3 3 3

45 2.11 Acids and Bases: The Lewis Definition Lewis Acid: electron pair acceptor Lewis Base: electron pair donor filled orbital vacant orbital B Lewis base + A B A Lewis acid

46 Lewis Acid / Base Reaction l + + l - F F B F F + 3 F B 3 F 3 3 l l Al l 3 l + N 3 l Al 3 l 3 N 3 3

48 Lewis definition of acid/base is broader than BrØnsted-Lowry definition: Lewis acid include many other species other than + Lewis acids: neutral proton donors: 2 l Br N 3 2 S cations: Li + Mg 2+ Br + metal compounds: All 3 BF 3 Til 4 Fel 3 Znl 2 Snl 4

49 ommon Lewis Acids All 3 BF 3 tetrahedral complex (maximum 4 N) All 3 BBr 3

50 Lewis Base: nonbonding electron pair l + Lewis base + l -

51 most oxygen and nitrogen containing compounds are Lewis bases Lewis bases: l N 2 3 N S 3

52 Some compounds, such as alcohols and carboxylic acids, can act as both acids and bases. BrØnsted-Lowry Acid : proton ( + ) donor Lewis base : electron pair donor

53 S S 4 - BF BF 3 3 3

54 Some compounds, such as carboxylic acids, ester, and amides, have more than one atom with a lone pair of electrons and therefore react at more than one site. - Reaction normally occurs only once in such instances, and the more stable of the two possible protonation products is formed. For acetic acid, protonation occurs on the doubly bonded oxygen S 4 3 or 3 + S 4 - more stable (resonance stabilization) 3

55 2.12 Drawing hemical Structures condensed structure: - and - single bonds aren't shown condensed structure or 3 2 ( 3 ) 2

56 skeleton structure: line drawing of - bond - carbon atoms aren't usually shown - hydrogen atoms bonded to carbons aren't shown - atoms other than, are shown -, can be indicated for emphasis and clarity

57 N 3 N 3 Adrenaline

58 Molecular Models Buckminsterfullerene ( 60 ) Wire-frame model Ball-and-stick model Space-filling model

59 Work Alkaloids: Naturally ccurring Bases alkaloid: amines derived from natural sources N Methylamine (found in rotting fish) N N 3 N 3 Morphine Atropine Ephedrine

Ch 2 Polar Covalent Bonds

Ch 2 Polar Covalent Bonds h 2 Polar ovalent Bonds Two primary bond types: ovalent (shared e -1 s) and Ionic (transferred e -1 s) Ionic bonds can have covalent character, such as with Na:l. An e -1 pair on l -1 can fill the 3s orbital