Chemical Bonding

1. Chemical Bonding

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

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

4. Figure 7.1 – The Hydrogen Molecule

5. 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.

6. Vocab continued 4. Octet rule – main group A elements acquire a complete octet (8e-s) in their outershell (ns2np6) 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

7. Lewis symbols for neutral atoms - 1A 2A 3A 4A 5A 6A 7A 8A H Be B C N O F Ne

8. 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)

9. The Octet Rule • Main group elements seek to attain an octet of electrons • Recall that an s2p6 configuration is isoelectronic with a noble gas • Closed electron shells (filled with e-) • Exceptions: H • The duet rule for H; Reduced octets (Be, B); and Expanded octets (N, P, etc.)

10. Drawing Lewis Structures • Count the number of valence electrons • Draw a skeleton structure for the species, joining the atoms by single bonds • Determine the number of valence electrons still available for distribution • Determine the number of valence electrons required to fill out an octet for each atom (except H) in the structure

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

12. 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

13. Examples of Lewis Structures • OH-, H2O, NH3, NH4+, C2H4, C2H2

14. More Examples:

15. 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.

16. Sulfur dioxide

17. Nitrate Ion; NO3-1

18. Benzene NOT IN NOTES

19. Example 7.3

20. Exceptions to the Octet Rule • Electron deficient molecules • Electron deficient atoms Be and B • Odd electron species (free radicals) • Example: NO, NO2

21. 1. Reduced Octets (Be and B) BH3 BeF2

22. 2. Expanded Octets • elements that are capable of surrounding themselves with more than four pairs of electrons • PCl5, SF6

23. Example 7.4 – Expanded Octets

24. Radicals • A free radical is a molecule with an odd number of electrons. Free radicals do not have a completed octet and often undergo vigorous redox reactions.

25. Molecular Geometry • Diatomic molecules are the easiest to visualize in three dimensions • HCl • Cl2 • Diatomic molecules are linear

26. 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 linear trigonal planar tetrahedral trigonal bipyramidal octahedral

27. 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

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

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

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

31. Four Electron Pairs • Tetrahedral • Bond angles are 109.5°

32. Bent and Pyramidal AX2E2 AX3E

33. 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: PCl5

34. 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: SF6

35. Figure 7.5 - Molecular Geometry Summarized - 1

36. Figure 7.5 - Molecular Geometry Summarized, 2

37. Polarity - Bonds • A polar bond has an non-symmetric 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 then it has: • One end is (δ+), while the other is (δ-)

38. Polarity - Molecules • A polar molecule has an overall non-symmetric distribution of electrons • Molecules may also possess polarity • Positive and negative poles • Molecule is called a dipole • Consider HF • H is δ+ while F is δ– • Consider BeF2 • Be-F bond is polar • BeF2 is nonpolar molecule b/c it is symmetrical

39. Quick Questions for P/NP molecules: 1. Are there lone pair electrons on the central atom? • Yes = polar (except AX4E2 and AX2E3) • No = go to next question 2. Are all the terminal atoms of the same element? No = polar(if the EN diff. is greater than 0.45) Yes = go to next question 3. Is the molecule symmetrical? Yes = nonpolar No = polar (if the EN diff. is greater than 0.45)

40. Figure 7.11 - Polarity of Molecules

41. 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

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

43. 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

44. 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)

45. Figure 7.13 - Ethylene and Acetylene

46. Hybrid Type • Draw the Lewis structure • 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

47. Hybrid Type Hybrid type Example # of bonding sites sp BeF2 2 sp2 trigonal planar 3 sp3 tetrahedral(bent/pyra) 4 sp3d trigonal bipyramidal 5 sp3d2 octahedral 6