Bonding

1. Bonding

2. What is a Bond? • A force that holds atoms together. • Why? • We will look at it in terms of energy. • Bond energy- the energy required to break a bond. • Why are compounds formed? • Because it gives the system the lowest energy.

3. Ionic Bonding • An atom with a low ionization energy reacts with an atom with high electron affinity. • A metal and a non metal • The electron moves. • Opposite charges hold the atoms together.

5. What about covalent compounds? • The electrons in each atom are attracted to the nucleus of the other. • The electrons repel each other, • The nuclei repel each other. • The reach a distance with the lowest possible energy. • The distance between is the bond length.

6. Energy 0 Internuclear Distance

7. Energy 0 Internuclear Distance

8. Energy 0 Internuclear Distance

9. Energy 0 Internuclear Distance

10. Energy 0 Bond Length Internuclear Distance

11. Energy Bond Energy 0 Internuclear Distance

12. Covalent Bonding • Electrons are shared by atoms. • polar covalent bonds: The electrons are not shared evenly. • One end is slightly positive, the other negative. • Indicated using small delta d.

13. Wait, you already know about polarity.

14. d+ d- H - F

15. d+ H - F d+ H - F d- d- d- H - F d+ d+ d- H - F H - F d+ d- d- H - F d+ d+ d+ H - F H - F d- d-

16. - + d+ H - F d+ H - F d- d- d- H - F d+ d+ d- H - F H - F d+ d- d- H - F d+ d+ d+ H - F H - F d- d-

17. Electronegativity • The ability of an electron to attract shared electrons to itself.

18. Electronegativity • Tends to increase left to right. • Decreases as you go down a group. • Most noble gases are inert, no electronegativity. • Difference in electronegativity between atoms tells us how polar the bond is.

19. Covalent Character decreases Ionic Character increases Polar Covalent Ionic Electronegativity difference Bond Type Zero Covalent Intermediate Large

20. Dipole Moments • A molecule with a center of negative charge and a center of positive charge is dipolar(two poles), or has a dipole moment. • Center of charge doesn’t have to be on an atom. • Will line up in the presence of an electric field.

21. - + d+ H - F d+ H - F d- d- d- H - F d+ d+ d- H - F H - F d+ d- d- H - F d+ d+ d+ H - F H - F d- d-

22. d+ d- H - F How It is drawn

23. Which Molecules Have Dipoles? • Any two atom molecule with a polar bond. • With three or more atoms there are two considerations. • There must be a polar bond. • Geometry can’t cancel it out.

24. Geometry and polarity • Three shapes will cancel them out. • Linear

25. Geometry and polarity • Three shapes will cancel them out. • Planar triangles 120º

26. Geometry and polarity • Three shapes will cancel them out. • Tetrahedral

27. Geometry and polarity • Others don’t cancel • Bent

28. Geometry and polarity • Others don’t cancel • Trigonal Pyramidal

29. Polar molecules • Must have polar bonds • Must not be symmetrical • Symmetrical shapes include • Linear • Trigonal planar • Tetrahedral • Trigonal bipyrimidal • Octahedral • Square planar

30. Ions • Atoms tend to react to form noble gas configuration. • Metals lose electrons to form cations • Nonmetals can share electrons in covalent bonds. • When two non-metals react.(more later) • Or they can gain electrons to form anions.

31. Ionic Compounds • We mean the solid crystal. • Ions align themselves to maximize attractions between opposite charges, and to minimize repulsion between like ions. • Can stabilize ions that would be unstable as a gas. • React to achieve noble gas configuration (to form an octet).

32. Size of ions • Ion size increases down a group. • Cations are smaller than the atoms they came from. • Anions are larger

33. Periodic Trends • Across the period nuclear charge increases so they get smaller. • Energy level changes between anions and cations. N-3 O-2 F-1 B+3 Li+1 C+4 Be+2

34. Size of Isoelectronic ions • Positive ions have more protons so they are smaller. N-3 O-2 F-1 Ne Na+1 Al+3 Mg+2

35. Bonding

36. What is a Model? • Explains how nature operates. • Derived from observations. • It simplifies them and categorizes the information. • A model must be sensible, but it has limitations.

37. Properties of a Model • A human inventions, not a blown up picture of nature. • Models can be wrong, because they are based on speculations and oversimplification. • Become more complicated with age. • You must understand the assumptions in the model, and look for weaknesses. • We learn more when the model is wrong than when it is right.Find the energy for this

38. Localized Electron Model • Simple model, easily applied. • A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Three Parts • Valence electrons using Lewis structures • Prediction of geometry using VSEPR • Description of the types of orbitals

39. VESPRNotes

40. Resonance Structures • We have assumed up to this point that there is one correct Lewis structure. • There are systems for which more than one Lewis structure is possible: • Different atomic linkages: Structural Isomers • Same atomic linkages, different bonding: Resonance

41. Resonance Structures (cont.) • The classic example: O3. Both structures are correct!

42. Resonance Structures (cont.) • In this example, O3 has two resonance structures: • Conceptually, we think of the bonding being an average of these two structures. • Electrons are delocalized between the oxygens such that on average the bond strength is equivalent to 1.5 O-O bonds.

43. Structural Isomers • What if different sets of atomic linkages can be used to construct correct LDSs: • Both are correct, but which is “more” correct?

44. Formal Charge • Formal Charge: Compare the nuclear charge (+Z) to the number of electrons (dividing bonding electron pairs by 2). Difference is known as the “formal charge”. #e- 7 6 7 7 6 7 Z+ 7 6 7 7 7 6 Formal C. 0 0 0 0 +1 -1 • Structure with less F. C. is more correct.

45. Formal Charge • Example: CO2 e- 6 4 6 6 4 6 7 4 5 Z+ 6 4 6 6 6 4 6 6 4 FC 0 0 0 0 +2 -2 -1 +2 -1 More Correct

46. Beyond the Octet Rule • There are numerous exceptions to the octet rule. • We’ll deal with three classes of violation here: • Sub-octet systems • Valence shell expansion • Odd-electron systems

47. Beyond the Octet Rule (cont.) • Some atoms (Be and B in particular) undergo bonding, but will form stable molecules that do not fulfill the octet rule. • Experiments demonstrate that the B-F bond strength is consistent with single bonds only.

48. Beyond the Octet Rule (cont.) • For third-row elements (“Period 3”), the energetic proximity of the d orbitals allows for the participation of these orbitals in bonding. • When this occurs, more than 8 electrons can surround a third-row element. • Example: ClF3 (a 28 e- system) F obey octet rule Cl has 10e-

49. Beyond the Octet Rule (cont.) • Finally, one can encounter odd electron systems where full pairs will not exist. • Example: Chlorine Dioxide. Unpaired electron

50. Summary • Remember the following: • C, N, O, and F almost always obey the octet rule. • B and Be are often sub-octet • Second row (Period 2) elements never exceed the octet rule • Third Row elements and beyond can use valence shell expansion to exceed the octet rule. • In the end, you have to practice…..a lot!