1. William L. Masterton Cecile N. Hurley http://academic.cengage.com/chemistry/masterton Chapter 7 Covalent Bonding Edward J. Neth University of Connecticut

2. Covalent Bonding NOT IN NOTES Recall that electrons in atoms are placed into atomic orbitals according to the Aufbau (fill 1 st energy level first), Pauli (2 e-s per orbital), and Hund’s Rules (1 e- in each first within a sublevel) In this section of the course, we will look at the location of electrons in molecules containing covalent bonds

3. Chemical Bonds: A Preview 1.Definitions chemical bond – an attraction strong enough to hold 2 atoms or ions together

4. 2. Hydrogen molecule Electron density – the area between 2 nuclei where the e-s are most likely to be found

5. Figure 7.1 – The Hydrogen Molecule

6. 7.1 Lewis Structures; The Octet Rule 1.Valence electrons - electrons in the highest principal energy level (outermost energy level of the atom) 2. Ionic bonds - attractive forces between positive and negative ions (due to e- transfer), holding them in solid crystals. 3. Covalent bonds - involve only nonmetals, one or more pairs of shared valence electrons between bonded atoms.

7. Vocab continued 4. Octet rule – main group A elements acquire a complete octet (8e-s) in their outershell (ns 2 np 6 ) during bonding. Transition metals do not follow octet rule. For hydrogen only 1 e- the duet rule applies 2e-s equal a full outer shell. 5. Isoelectronic – atoms/ions with the same number of electrons but different mass numbers

8. Lewis symbols for neutral atoms - 1A2A3A4A5A6A7A8A HBeBCNOFNe

9. Table 1.1

10. Electron Ownership An atom owns All lone electrons = Shown as lone pairs (dots) Half the number of bonding electrons A bond pair is shown as a line Multiple bonds Double bonds are two pairs (2 lines) Triple bonds are three pairs (3 lines)

11. The Octet Rule Main group elements seek to attain an octet of electrons Recall that an s 2 p 6 configuration is isoelectronic with a noble gas Closed electron shells Exceptions: The duet rule for H; Reduced octets (Be, B); and Expanded octets (N, P, etc.)

12. Drawing Lewis Structures 1.Count the number of valence electrons 2.Draw a skeleton structure for the species, joining the atoms by single bonds 3.Determine the number of valence electrons still available for distribution 4.Determine the number of valence electrons required to fill out an octet for each atom (except H) in the structure *see p. 169 of the text

13. Importance of Lewis Structures, bonding pairs and symbols Indicates number of and ways the atoms bond Shows the geometric structure of the molecule

14. Strategies: 1. H atoms almost always terminal atom 2. central atoms (usually only ONE present) 3. H bonded to O in alcohol and oxoacids 4. Molecules are clusters of atoms S = O - V

15. Examples of Lewis Structures OH -, H 2 O, NH 3, NH 4 +, C 2 H 4, C 2 H 2

16. More Examples:

17. Resonance Structures These are structures in which double bonds and/or triple bonds between atoms make for a structure that resonates between 2 or 3 simple structures. 1. Resonance forms are not different molecules 2. Resonance structures arise when two Lewis structures are equally possible 3. Only electrons can be shifted in resonance structures. Atoms cannot be moved.

18. Sulfur dioxide

19. Nitrate Ion; NO 3 -1

20. Benzene NOT IN NOTES

21. Example 7.3

22. Exceptions to the Octet Rule Electron deficient molecules Electron deficient atoms Be and B Odd electron species (free radicals) Example: NO

23. 1. Reduced Octets (Be and B) BH 3 BeF 2

24. 2. Expanded Octets elements that are capable of surrounding themselves with more than four pairs of electrons PCl 5, SF 6

25. Example 7.4 – Expanded Octets

26. 4. Radicals Examples:

27. 7.2 Molecular Geometry Diatomic molecules are the easiest to visualize in three dimensions HCl Cl 2 Diatomic molecules are linear

28. Figure 7.4 – Ideal Geometries There is a fundamental geometry that corresponds to the total number of electron pairs around the central atom: 2, 3, 4, 5 and 6 lineartrigonal planar tetrahedraltrigonal bipyramidal octahedral

29. Valence Shell Electron Pair Repulsion Theory The ideal geometry of a molecule is determined by the way the electron pairs orient themselves in space The orientation of electron pairs arises from electron repulsions The electron pairs spread out so as to minimize repulsion

30. The A-X-E Notation A = central atom X = terminal atom E = lone pair

31. Two electron pairs Linear Bond angles The bond angle in a linear molecule is always 180°

32. Three electron pairs Trigonal planar The electron pairs form an equilateral triangle around the central atom Bond angles are 120°

33. Four Electron Pairs Tetrahedral Bond angles are 109.5°

34. Bent and Pyramidal AX 2 E 2 AX 3 E

35. Five Electron Pairs Trigonal bipyramid Bond angles vary In the trigonal plane, 120° Between the plane and apexes, 90° Between the central atom and both apexes, 180° Example: PCl 5

36. Six Electron Pairs Octahedron or square bipyramid Bond angles vary 90° in and out of plane 180° between diametrically opposite atoms and the central atom Example: SF 6

37. Figure 7.5 - Molecular Geometry Summarized - 1

38. Figure 7.5 - Molecular Geometry Summarized, 2

39. Polarity - Bonds A polar bond has an asymmetric distribution of electrons X-X is nonpolar but X-Y is polar Polarity of a bond increases with increasing difference in electronegativity between the two atoms Bond is a dipole One end is (δ + ), while the other is (δ - )

40. Polarity - Molecules Molecules may also possess polarity Positive and negative poles Molecule is called a dipole Consider HF H is δ + while F is δ – Consider BeF 2 Be-F bond is polar BeF 2 is nonpolar molecule b/c it is symmetrical

41. Figure 7.11 - Polarity of Molecules

42. Valence Bond Theory Unpaired electrons from one atom pair with unpaired electrons from another atom and give rise to chemical bonds Simple extension of orbital diagrams

43. Figure 7.12 - Atomic Orbital Mathematics Two atomic orbitals produce two hybrid orbitals One s + one p  two sp

44. Table 7.4 - Hybrid Orbitals and Geometry

45. Hybrid Orbitals and Electron Occupancy Same rules we have seen before In an atom, an orbital holds two electrons In a molecule, an orbital also holds two electrons What electrons go into hybrid orbitals? Lone pairs One pair per bond Even for a double bond, only one pair goes into the hybrid orbital

46. Multiple bonds Sigma (σ) bonds Electron density is located between the nuclei One pair of each bond is called a sigma pair Pi bonds (π) Electron density is located above and below or in front of and in back of the nuclei One pair of a double bond is called pi (π) Two pairs of a triple bond are called pi (π)

47. Figure 7.13 - Ethylene and Acetylene

48. Hybrid Type 1.Draw the Lewis structure 2.Count the number of bonding or e- pair sites around the central atom * a “site” is a bond or a lone pair (double and triple bonds count as 1 site

49. Hybrid Type Hybrid type Example# of bonding sites