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Chemical Bonding and Molecular Structure

Chemical Bonding and Molecular Structure . Ionic vs. covalent bonding Molecular orbitals and the covalent bond Valence electron Lewis dot structures octet vs. non-octet resonance structures formal charges VSEPR (Valence shell electron pair repulsion)

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Chemical Bonding and Molecular Structure

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  1. Chemical Bonding and Molecular Structure • Ionic vs. covalent bonding • Molecular orbitals and the covalent bond • Valence electron Lewis dot structures • octet vs. non-octet • resonance structures • formal charges • VSEPR (Valence shell electron pair repulsion) • - predicting shapes of molecules • Bond properties • polarity, bond order, bond strength Bonding and Structure

  2. Chemical Bonding Problems and questions — • How is a molecule or polyatomic ion held together? • Why are atoms distributed at strange angles? • Why are molecules not flat? • Can we predict the structure? • How is structure related to chemical and physical properties? Bonding and Structure

  3. Forms of Chemical Bonds • There are 2 extreme forms of connecting or bonding atoms: • Ionic—complete transfer of electrons from one atom to another • Covalent—electrons shared between atoms Most bonds are somewhere in between. Bonding and Structure

  4. Ionic Bonds Ionic compounds - essentially complete electron transfer from an element of low (metal) to an element of high electron affinity (EA) (nonmetal) Na(s) + 1/2 Cl2(g)  Na+ + Cl-  NaCl (s) - primarily between metals (Groups 1A, 2A and transition metals) and nonmetals(esp O and halogens) - NON-DIRECTIONAL bonding via Coulomb (charge) interaction Bonding and Structure

  5. Br Br Covalent Bonding Covalent bond is the sharing of the VALENCE ELECTRONS of each atom in a bond Recall: Electrons are divided between core and valence electrons. ATOM core valence Na 1s2 2s2 2p6 3s1 [Ne] 3s1 Br [Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5 Bonding and Structure

  6. 8A 1A Valence Electrons 2A 3A 4A 5A 6A 7A Number of valence electrons is equal to the Group number. Bonding and Structure

  7. Covalent Bonding The bond arises from the mutual attraction of 2 nuclei for the same electrons. A covalent bond is a balance of attractive and repulsive forces. Bonding and Structure

  8. Polar and non-polar covalent bond • Fig. (a) In the nonpolar covalent bond present, there is a symmetrical distribution of electron density. (b) In the polar covalent bond present, electron density is displaced because of its electronegativity. Dipole moment, µ= e (esu) x d (angstrom) Greater the DM greater the polarity Bonding and Structure

  9. •• •• Cl H H Cl • • + • • •• •• Overlap of H (1s) and Cl (2p) Bond Formation A bond can result from a “head-to-head/ end-to-end” overlapof atomic orbitals on neighboring atoms. This type of overlap places bonding electrons in a MOLECULAR ORBITAL along the line between the two atoms and forms a SIGMA BOND (s). S-s, s-p and p-p orbitals form sigma bond. Bonding and Structure

  10. 6_H2pot.mov Sigma Bond Formation by Orbital Overlap Two s Atomic Orbitals (A.O.s) overlap to form an s (sigma) Molecular Orbital (M.O.) Bonding and Structure

  11. Sigma Bond Formation by Orbital Overlap Two s A.O.s overlap to from an s  M.O. Similarly, two p A.O.s can overlap end-on to from a p M.O. e.g. F2 Bonding and Structure

  12. Electron Distribution in Molecules • Electron distribution is depicted with Lewis electron dot structures • Electrons are distributed as: • shared or BOND PAIRS and • unshared or LONE PAIRS. G. N. Lewis 1875 - 1946 Bonding and Structure

  13. •• H Cl • • •• Unshared or lone pair (LP) shared or bond pair Bond and Lone Pairs • Electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. This is a LEWIS ELECTRON DOT structure. Bonding and Structure

  14. This observation is called the OCTET RULE Rules of Lewis Structures • For Groups 1A-4A (Li - C), • no. of BOND PAIRS = group number • No. of valence electrons of an atom = Group number • For Groups 5A-7A (N - F), • no. of BOND PAIRS = 8 - group No. • Except for H • (and atoms of 3rd and higher periods), • #Bond Pairs + #Lone Pairs = 4 Bonding and Structure

  15. Building a Dot Structure 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. In ammonia, N is central Ammonia, NH3 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons or 4 pairs Bonding and Structure

  16. •• H H N H Building a Dot Structure 3. Form a sigma bond between the central atom and surrounding atoms. 4. Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while each H shares 1 pair. Bonding and Structure

  17. Sulfite ion, SO32- Step 1. Central atom = S Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 6 + 18 + 2 = 26 e- or 13 pairs Step 3. Form sigma bonds 10 pairs of electrons are left. Bonding and Structure

  18. Thanks to all Bonding and Structure

  19. •• O • • • • •• •• O S O • • • • •• •• •• Sulfite ion, SO32- (2) Remaining pairs become lone pairs, first on outside atoms then on central atom. Each atom is surrounded by an octet of electrons. NOTE - must add formal charges (O-, S+) for complete dot diagram Bonding and Structure

  20. Carbon Dioxide, CO2 1. Central atom = __C____ 2. Valence electrons = _16_ or _8_ pairs 3. Form sigma bonds. This leaves __6__ pairs. 4. Place lone pairs on outer atoms. Bonding and Structure

  21. Carbon Dioxide, CO2 (2) 4. Place lone pairs on outer atoms. 5. To give C an octet, form DOUBLE BONDS between C and O. The second bonding pair forms a pi (p)bond. Bonding and Structure

  22. H2CO SO3 C2F4 Double and even triple bonds are commonly observed for C, N, P, O, and S Bonding and Structure

  23. Sulfur Dioxide, SO2 1. Central atom = S 2. Valence electrons = 6 + 2*6 = 18 electrons or 9 pairs 3. Form pi () bond so that S has an octet — note that there are two ways of doing this. Bonding and Structure

  24. Equivalent structures called: Sulfur Dioxide, SO2 RESONANCE STRUCTURES The proper Lewis structure is a HYBRID of the two. A BETTER representation of SO2 is made by forming 2 double bonds Each atom has - OCTET - formal charge = 0 O = S = O Bonding and Structure

  25. Urea (NH2)2CO 1. Number of valence electrons = 24 e- 2. Draw sigma bonds. Leaves 24 - 14 = 10 e- pairs. 3. Complete C atom octet with double bond. 4. Place remaining electron pairs on oxygen and nitrogen atoms. Bonding and Structure

  26. BF3 SF4 Violations of the Octet Rule elements of higher periods. Usually occurs with: Boron Bonding and Structure

  27. Boron Trifluoride • Central atom = B • Valence electrons = 3 + 3*7 = 24 or electron pairs = 12 • Assemble dot structure The B atom has a share in only 6 electrons (or 3 pairs). B atom in many molecules is electron deficient. Bonding and Structure

  28. Sulfur Tetrafluoride, SF4 • Central atom = S • Valence electrons = 6 + 4*7 = 34 e- or 17 pairs. • Form sigma bonds and distribute electron pairs. 5 pairs around the S atom. A common occurrence outside the 2nd period. Bonding and Structure

  29. Explanation of the failure of octet rule Sidgwick’s rule of maximum covalency • It is not essential to have 8 electron surrounding in an atom • Maximum covalency depends on its period Bonding and Structure

  30. Sugden’s view of singlet linkages • Octet rule never violated • In PCl5, 3 Cl atom linked by normal covalent bond and • others 2 by sharing only 1 electron called singlet linkage Bonding and Structure

  31. Formal Charges • Formal charge is the charge calculated for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs.

  32. .. : O .. N H O .. : : O Nitric acid Formal charge of H • We will calculate the formal charge for each atom in this Lewis structure. ..

  33. .. : O .. N H O .. : : O Nitric acid Formal charge of H • Hydrogen shares 2 electrons with oxygen. • Assign 1 electron to H and 1 to O. • A neutral hydrogen atom has 1 electron. • Therefore, the formal charge of H in nitric acid is 0. ..

  34. .. : O .. N H O .. : : O Nitric acid Formal charge of O • Oxygen has 4 electrons in covalent bonds. • Assign 2 of these 4 electrons to O. • Oxygen has 2 unshared pairs. Assign all 4 of these electrons to O. • Therefore, the total number of electrons assigned to O is 2 + 4 = 6. ..

  35. .. : O .. N H O .. : : O Nitric acid Formal charge of O • Electron count of O is 6. • A neutral oxygen has 6 electrons. • Therefore, the formal charge of O is 0. ..

  36. .. : O .. N H O .. : : O Nitric acid Formal charge of O • Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons). • A neutral oxygen has 6 electrons. • Therefore, the formal charge of O is 0. ..

  37. .. : O .. N H O .. : : O Nitric acid Formal charge of O • Electron count of O is 7 (6 electrons from unshared pairs + half of 2 bonded electrons). • A neutral oxygen has 6 electrons. • Therefore, the formal charge of O is -1. ..

  38. .. : O .. N H O .. : : O Nitric acid Formal charge of N • Electron count of N is 4 (half of 8 electrons in covalent bonds). • A neutral nitrogen has 5 electrons. • Therefore, the formal charge of N is +1. – ..

  39. .. : O .. N H O .. : : O Nitric acid Formal charges • A Lewis structure is not complete unless formal charges (if any) are shown. + – ..

  40. Formal Charge An arithmetic formula for calculating formal charge. Formal charge = group numberin periodic table number ofbonds number ofunshared electrons – –

  41. .. 1 H : : F + .. .. : : N B F F H H .. .. : : F H .. 4 "Electron counts" and formal charges in NH4+ and BF4- 7 – 4

  42. Resonance

  43. Resonance In chemistry, resonance or mesomerism is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. A molecule or ion with such delocalized electrons is represented by several contributing structures (also called resonance structures or canonical forms).

  44. General characteristics of resonance • Molecules and ions with resonance (also called mesomerism) have the following basic characteristics: • They can be represented by several correct Lewis formulas, called "contributing structures", "resonance structures“ or "canonical forms". However, the real structure is not a rapid interconversion of contributing structures. Several Lewis structures are used together, because none of them exactly represents the actual structure. To represent the intermediate, a resonance hybrid is used instead. • The contributing structures are not isomers. They differ only in the position of electrons, not in the position of nuclei. Bonding and Structure

  45. General characteristics of resonance • Each Lewis formula must have the same number of valence electrons (and thus the same total charge), and the same number of unpaired electrons, if any. • Bonds that have different bond orders in different contributing structures do not have typical bond lengths. Measurements reveal intermediate bond lengths. • The real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be. Bonding and Structure

  46. Contributing structures of some ion Bonding and Structure

  47. H H .. .. .. + – : : H C O N O H C O N O .. .. .. .. .. H H Resonance Structures of Methyl Nitrite • same atomic positions • differ in electron positions more stable Lewis structure less stable Lewis structure

  48. H H .. .. .. + – : : H C O N O H C O N O .. .. .. .. .. H H Resonance Structures of Methyl Nitrite • same atomic positions • differ in electron positions more stable Lewis structure less stable Lewis structure

  49. Why Write Resonance Structures? • Electrons in molecules are often delocalizedbetween two or more atoms. • Electrons in a single Lewis structure are assigned to specific atoms-a single Lewis structure is insufficient to show electron delocalization. • Composite of resonance forms more accurately depicts electron distribution.

  50. + – •• •• O O O •• •• •• •• Example • Ozone (O3) • Lewis structure of ozone shows one double bond and one single bond Expect: one short bond and one long bond Reality: bonds are of equal length (128 pm)

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