H-H bond length Two e s shared by two Hs: covalent bonding. Coulomb attraction: Stronger attraction for e Fractional charge A dipole

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1 8 Bonding: General Concepts Types of chemical bonds Covalent bonding Ex. 2 E (kj/mol) Repulsions of nucleus and e s r r (nm) Zero interaction at long distance - bond length Two e s shared by two s: covalent bonding Ionic bonding Ex. a + : ionic forces operate Coulomb attraction: charges 9 Q1Q ( J nm) r Polar covalent bond Ex. Unequal sharing of e Charge distribution occurs E 2 tronger attraction for e ractional charge A dipole 1

2 Electronegativity The ability of an atom in a molecule to attract shared e s Linus Pauling ( ) BE A B Expected BE = (BEA-B ) actual (BEA-B ) AA BE 2 expected BB The greater E difference, the greater value Define E(A) E(B) = Assign E() = 4.0 luorine atom with the highest E E for all elements can be determined E difference for A B Bond type 0 covalent polar covalent large ionic (A +, B ) Increasing E A B igher E Decreasing E periodic table 2

3 Bond polarity and dipole moment A dipole has a dipole moment ( ) Will orient in an electric field = Q R distance (m) charge (c) Unit: debye (D) 1D = cm or polyatomic molecules net dipole The center of all positive atoms and the center of all negative atoms do not merge C : net dipole = zero Ions: Electron configurations and sizes Ionic compounds In general A M + non-metal on-metal: gains e Metal: loses e main group metal reach noble gas configuration Ex. Ca: [Ar]4s 2 Ca 2+ : [Ar] : [e]2s 2 2p 4 2 : [e]2s 2 2p 6 = [e] 3

4 izes of ions determines the structure and stability of ionic compounds egative ions: larger than parent (gain e ) Positive ions: smaller than parent (lose e ) Trend: Down a group larger Ex. Li + Be (pm) a + 95 The energy M + (g) + X (g) M X(s) ionization E Li(g) Li + (g) Born-aber electron affinity cycle (g) (g) ormation of an ionic compound energy depends on many factors difficult to predict qualitatively energy released = lattice E Q Q2 Lattice E = k( 1 ) r verall Ex. Li(s) + 1/2 2 (g) Li(s) Unit: lattice 1/2 BE of 2 kj/mol energy = 78.9 At 298 K LE theo = 966 4

5 Partial ionic character of covalent bonds In the gas phase: no true ionic bond Polar covalent bond percent ionic character measured dipole moment 00% calculated dipole moment X + Y Ex. a(g): 75% In general: > 50% ionic solid Ambiguity: 4 +, a 2 4 covalent bonded group Definition of ionic compound conducts electric current when melted 5

6 The covalent bond: A model Ex. C 4 C + 4 requires 1652 kj C 4 is more stable than the discreet atoms by 1652 kj Consider forces exist between C and 4 bonding Consider localized C bonds 1652 To break each bond = = 413 kj 4 Ex. C 3 C requires 1578 kj 3 C bonds, 1 C bond (413) = 339 kj/mol trength of C bond Assumption: bond strength does not vary with structures The covalent bond model Considering electrons are localized between two atoms 6

7 Covalent bond energies and chemical reactions act 1 C 4 (g) C 3 (g) + (g) 435 kj/mol C 3 (g) C 2 (g) + (g) 453 C 2 (g) C(g) + (g) 425 C (g) C(g) + (g) 339 total = 1652 Average: 413 BEs are structurally dependent Average BEs are used Bond types e s shared single 2 double 4 triple 6 Ex. BE (kj/mol) bond length (Å) C C C=C C C

8 Bond energy and enthalpy Ex. 2 (g) + 2 (g) 2(g) 2 2 2, 2 identical Bond E difference = = D (bonds broken) D (bonds formed) Bond E (+) = D - + D - 2D - = (565) = 544 kj Calculated from f : = 2( 271) = 542 kj f o () The localized electron bonding model (Valence bonding model) A covalent bond is formed between two atoms sharing electrons using atomic orbitals Pairs of e s localized on an atom: lone pairs Pairs of e s localized between atoms: bonding pairs 8

9 Lewis structures Describe the arrangement of valence electrons G.. Lewis ( ) observed: In most stable compounds the atoms achieve noble gas electron configuration Molecules with covalent bonds Duet rule, e : e: ctet rule C,,, : second row nonmetals e Ex Total: 8 e s 9

10 Ex. C 2 :C: Total: 16 C Does not obey octet rule Try and error C = C Ex. C ne extra charge = 10 e C ot good C ot good C ot good C Good C 10

11 Ex e isoelectronic with C Ex. 3 8 e C 4 C C 8 e = 10 e C 4 32 e C C Exceptions to the octet rule Less than 8 B nly 6 e ow about B act: B is very reactive toward electron rich molecules + B B 11

12 More than 8 12 e assical explanation use of low lying empty d orbitals (3d for sulfur) P (5 7) = 40 e P 10 e Allow the central atom to expand its valency Ex. Xe (3 6) = 26 e Xe obeys Be (2 7) = 16 e Be nly 4 e 12

13 Resonance Ex. 3 Predict: act: a short double bond two long single bonds Equal length (between a single and a double) A new theory must be formulated Actually there are three possible structures: We can consider the real structure as the avg. of the three a resonance hybrid Any single structure can not represent the real structure Each Lewis structure is called resonance structure (exists only on paper) Represented as double headed arrow (not ) 13

14 nly allow electrons to move (nuclei stay at the same position) Must be proper Lewis structures Ex. Carbon can not have five bonds dd-electron molecule itric oxide = 11 ot good A radical with unpaired e *Keep the number of unpaired e the lowest (6) = 17 C C C C ot a viable resonance structure 14

15 ormal charge Q: tructure of 4 2 (32 e ) 2 What about 2 2 or 10 e 12 e ow to decide? Method: Estimate the charge xidation state: 2 +6 The concept of formal charge ver estimated nly useful for bookkeeping Compare: number of e in a neutral atom A number of e in a bonded atom B When A = B no charge A > B (+) A < B ( ) 15

16 Define ormal charge = A B Assign lone pair e s to the atom completely Divide shared e s equally Ex. 2 or shared e s A = 6 B = 6 + 2/2 = 7 lone pair e s ormal charge on = 6 7 = 1 or : A = 6 B = 0 + 8/2 = 4 shared e s lone pair e s ormal charge on = 6 4 = verall charge = (+2) + 4() 2 ingle bonded : Double bonded : 6 [4 + 4/2] = 0 : 6 (10/2) = ok 16

17 (12/2) = 0 There are resonance structures Another school of thought: Keeping the octet rule is more important A better structure due to minimum formal charge ote: they are all resonance structures 2 etc. Guidelines or 2nd row elements: never exceed octet rule ormal charges as low as possible egative charge on more electronegative atoms B ot good 17

18 Molecular structures: The VEPR model To predict approximate 3-D structure A simple model: Valence shell electron pair repulsion (VEPR) model (useful for nonmetals) The main postulation: minimizing electron pair repulsions Bonding and nonbonding pairs 2 Be : linear (two pairs) As far away as possible B 3 B 120 o Trigonal planar (three pairs) C 4 3 C o Tetrahedral (four pairs) Cf. 90 o C quare planar (not good) Based on the position of atoms: trigonal pyramid 2 A bent structure (four pairs) 18

19 Bond angles lone pairs o o C 107 o Explanation lone pair is closer to the nucleus exerts large repulsion towards the other electron pairs squeezed ive pairs: P 5 P Trigonal bipyramidal ix pairs: P 6 P ctahedral 19

20 Ex. Xe 4 Xe ix pairs: octahedral arrangement Xe as far away as possible quare planar shape Ex. I 3 I I I ive pairs: trigonal bipyramidal arrangement I I I 90 o, 90 o, 180 o I I I 90, 90, 120 I I I 120, 120, 120 Best (no 90 o repulsions) linear 20

21 Treat double or triple bonds as one effective pair Ex. 2 3 pairs Bent structure In fact -- ~120 o with little distortion long pair is satisfied with 120 o 21