Chemistry 281(01) Winter 2014

1. Chemistry 281(01) Winter 2014 CTH 27710:00-11:15 am Instructor: Dr. UpaliSiriwardane E-mail:  upali@latech.edu Office:  311 Carson Taylor Hall ; Phone: 318-257-4941; Office Hours:  MTW 8:00 am - 10:00 am; TR 8:30 - 9:30 am & 1:00-2:00 pm. January 14, 2014 Test 1 (Chapters 1&,2), February 6, 2014 Test 2 (Chapters 3 &4) February 25, 2014, Test 3 (Chapters 5 & 6), Comprehensive Final Make Up Exam: February 27, 2012 9:30-10:45 AM, CTH 311.

2. Molecular structure and bonding Lewis structures 2.1 The octet rule 2.2 Structure and bond properties 2.3 The VSEPR model Valence-bond theory 2.4 The hydrogen molecule 2.5 Homonuclear diatomic molecules 2.6 Polyatomic molecules Molecular orbital theory 2.7 An introduction to the theory 2.8 Homonuclear diatomic molecules 2.9 Heteronuclear diatomic 2.10 Bond properties

3. Lewis Theory of Bonding Octet RuleAll elements except hydrogen ( hydrogen have a duet of electrons) have octet of electrons once they from ions and covalent compounds.

4. Noble gas configuration The noble gases are noted for their chemical stability and existence as monatomic molecules. Except for helium, They share a common electron configuration that is very stable. This configuration has 8 valence-shell electrons. All other elements reacts to achieve Noble Gas Electron Configurations. valence e- He 2 Ne 8 Ar 8 Kr 8 Xe 8 Rn8

5. The octet rule • Atoms are most stable if they have a filled or empty outer layer of electrons. • Except for H and He, a filled layer contains 8 electrons - an octet. • Two atoms will gain or lose (ionic compounds) share (covalent compounds) Many atoms with fewer electrons will share(metallic compounds)

6. What changes take place during this process of achieving closed shells? a) sharing leads to covalent bonds and molecules b) gain/loss of electrons lead to ionic bond c) Sharing with many atoms lead to metallic bonds

7. Lewis Electron Dot symbols • Basic rules • Draw the atomic symbol. • Treat each side as a box that can hold up to two electrons. • Count the electrons in the valence shell. • Start filling box - don’t make pairs unless you need to. X

8. Lewis symbols Lewis symbols of second period elements Li Be B C N O F Ne

9. What is a Lewis Structure (electron-dot formula) of a Molecule? • A molecular formulas with dots around atomic symbols representing the valence electrons • All atoms will have eight (octet) of electrons (duet for H) if the molecule is to be stable.

10. F F Single covalent bonds H C H H H H H Do atoms (except H) have octets?

11. Lewis structures • This is a simple system to help keep track of electrons around atoms, ions and molecules - invented by G.N. Lewis. • If you know the number of electrons in the valence-shell of an atom, writing Lewis structures is easy. • Lewis structures are used primarily for s- and p-block elements.

12. How do you get the Lewis Structure from Molecular formula? • Add all valence electrons and get valence electron pairs • Pick the central atom: Largest atom normally or atom forming most bonds • Connect central atom to terminal atoms • Fill octet to all atoms (duet to hydrogen)

13. Lewis Structure of H2O

14. Types of electrons Bonding pairs Two electrons that are shared between two atoms. A covalent bond. Unshared (nonbonding ) pairs A pair of electrons that are not shared between two atoms. Lone pairs or nonbonding electrons. Unshared pair oo H Cl oo oo oo Bonding pair

15. Lewis Structure of H2O 2 bond pairs= 2 x 2 = 4 2 lone pairs = 2 x 2 = 4 Total 8 = 4 pairs Bond pairs: an electron pair shared by two atom in a bond. E.g. two pairs between O--H in water. Lone pair : an electron pair found solely on a single atom. E.g. two pairs found on the O atomat the top and the bottom.

16. Lewis Structure of H2S

17. Lewis Structure of CCl4

18. What is the Lewis Structure? • CO2 • NH3 (PH3) • PCl3 (PF3, NCl3)

19. O=C=O Lewis structure and multiple bonds This arrangement needs too many electrons. O C O How about making some double bonds? That works! = is a double bond, the same as 4 electrons

20. Multiple bonds So how do we know that multiple bonds really exist? The bond energies and lengths differ! Bond Bond Length Bond energy type order pm kJ/mol C C 1 154 347 C C 2 134 615 C C 3 120 812

21. Formal Charges Formal charge = valence electrons- assigned electrons • There are two possible Lewis structures for a molecule. Each has the same number of bonds. We can determine which is better by determining which has the least formal charge. It takes energy to get a separation of charge in the molecule • (as indicated by the formal charge) so the structure with the least formal charge should be lower in energy and thereby be the better Lewis structure

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

23. Electron counts" and formal charges in NH4+ and BF4- "

24. What is Resonance Structures? • Several Lewis structures that need to be drawn for molecules with double bonds • One Lewis structure alone would not describe the bond lengths of the real molecule. • E.g. CO32-, NO3-, NO2-, SO3

25. Resonance structures Sometimes we can have two or more equivalent Lewis structures for a molecule. O - S = O O = S - O They both - satisfy the octet rule - have the same number of bonds - have the same types of bonds Which is right?

26. Resonance structures of SO2 They both are! O - S = O O = S - O O S O This results in an average of 1.5 bonds between each S and O.

27. Resonance structures of CO32- ion

28. Resonance structures of NO3- ion

29. Resonance structures of SO3

30. Resonance structures of NO2- ion

31. Resonance structures of C6H6 • Benzene, C6H6, is another example of a compound for which resonance structure must be written. • All of the bonds are the same length. or

32. Exceptions to the octet rule Not all compounds obey the octet rule. • Three types of exceptions • Species with more than eight electrons around an atom. • Species with fewer than eight electrons around an atom. • Species with an odd total number of electrons.

33. Atoms with more than eight electrons • Except for species that contain hydrogen, this is the most common type of exception. • For elements in the third period and beyond, the dorbitals can become involved in bonding. Examples • 5 electron pairs around P in PF5 • 5 electron pairs around S in SF4 • 6 electron pairs around S in SF6

34. O O S O O O || O S O - - || O An example: SO42- 1. Write a possible arrangement. 2. Total the electrons. 6 from S, 4 x 6 from O add 2 for charge total = 32 3. Spread the electrons around.

35. Atoms with fewer than eight electrons Beryllium and boron will both form compounds where they have less than 8 electrons around them.

36. H | :N - H | H F H | | F - B - N – H | | F H F | B | F F - + Atoms with fewer than eight electrons Electron deficient. Species other than hydrogen and helium that have fewer than 8 valence electrons. They are typically very reactive species.

37. What is VSEPR Theory Valence Shell Electron Pair Repulsion This theory assumes that the molecular structure is determined by the lone pair and bond pair electron repulsion around the central atom

38. What Geometry is Possible around Central Atom? • What is Electronic or Basic Structure? • Arrangement of electron pairs around the central atom is called the electronic or basic structure • What is Molecular Structure? • Arrangement of atoms around the central atom is called the molecular structure

39. Possible Molecular Geometry • Linear (180) • Trigonal Planar (120) • T-shape (90, 180) • Tetrahedral (109) • Square palnar ( 90, 180) • Sea-saw (90, 120, 180) • Trigonalbipyramid (90, 120, 180) • Octahedral (90, 180)

40. Molecular Structure from VSEPRTheory • H2O • Bent or angular • NH3 • Pyramidal • CO2 • Linear

41. Molecular Structure from VSEPR Theory • SF6 • Octahedral • PCl5 • Trigonalbipyramidal • XeF4 • Square planar

42. What is a Polar Molecule? • Molecules with unbalanced electrical charges • Molecules with a dipole moment • Molecules without a dipole moment are called non-polar molecules

43. How do you a Pick Polar Molecule? • Get the molecular structure from VSEPR theory • From c (electronegativity) difference of bonds see whether they are polar-covalent. • If the molecule have polar-covalent bond, check whether they cancel from a symmetric arrangement. • If not molecule is polar

44. Which Molecules are Polar • H2O • Bent or angular, polar-covalent bonds, asymmetric molecule-polar • NH3 • Pyramidal, polar-covalent bonds, asymmetric molecule-polar • CO2 • Linear, polar-covalent bonds, symmetric molecule-polar

45. What is hybridization? Mixing of atomic orbitals on the central atoms valence shell (highest n orbitals) Bonding: s p d sp, sp2, sp3, sp3d, sp3d2 Px dx2- y2 Py Pz dz2

46. What is hybridization? Mixing of atomic orbitals on the central atom Bonding a hybrid orbital could over lap with another ()atomic orbital or () hybrid orbital of another atom to make a covalent bond. possible hybridizations: sp, sp2, sp3, sp3d, sp3d2

47. What is Valence Bond Theory • Describes bonding in molecule using atomic orbital • orbital of one atom occupy the same region with a orbital from another atom • total number of electrons in both orbital is equal to two Be Cl2

48. sp2 and sp3 Hybridization BF3

49. What are p and s bonds s bonds single bond resulting from head to head overlap of atomic orbital p bond double and triple bond resulting from lateral or side way overlap of atomic orbitals

50. How do you tell the hybridization of a central atom? • Get the Lewis structure of the molecule • Look at the number of electron pairs on the central atom. Note: double, triple bonds are counted as single electron pairs. • Follow the following chart