TOPIC 4 CHEMICAL BONDING AND STRUCTURE

1. TOPIC 4CHEMICAL BONDING AND STRUCTURE 4.3 COVALENT STRUCTURES

2. ESSENTIAL IDEA Lewis (electron dot) structures show the electron domains in the valence shell and are used to predict molecular shape. NATURE OF SCIENCE (1.10) Scientists use models as representatives of the real world – the development of the model of molecular shapes (VSEPR) to explain observable properties.

3. THEORY OF KNOWLEDGE Does the need for resonance structures decrease the value or validity of Lewis (electron dot) theory? What criteria do we use in assessing the validity of a scientific theory?

4. UNDERSTANDING/KEY IDEA 4.3.A Lewis (electron dot) structures show all the valence electrons in a covalently bonded species.

5. UNDERSTANDING/KEY IDEA 4.3.B The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons.

6. APPLICATION/SKILLS Be able to deduce the Lewis (electron dot) structures of molecules and ions showing all valence electrons for up to four electron pairs on each atom.

7. Lewis Dot Structures • Lewis Dot structures are used to represent the valence electrons of atoms in covalent molecules • Dots are used to represent only the valence electrons. • Dots are written between symbols to represent bonding electrons 7

8. Writing Dot Structures Writing Dot structures is a process: • Determine the number of valence electrons each atom contributes to the structure • The number of valence electrons can usually be determined by the column in which the atom resides in the periodic table 8

9. Add up the total number of valence electrons Adjust for charge if it is a poly atomic ion Add electrons for negative charges Reduce electrons for positive charges Example SO32- 1 S = 6 e 3 0 = 6x3 = 18 e (2-) charge = 2 e --------- Total = 26 e Writing Dot Structures 9

10. Make the atom that is fewest in number the central atom. Distribute the electrons so that all atoms have 8 electrons. Use double or triple pairs if you are short of electrons If you have extra electrons put them on the central atom Electron Dot Structures 10

11. Electron Dot Structures Example 2: SO3 • 1 S = 6 e • 3 O = 6x3 = 18 e • no charge = 0 e --------- Total = 24 e Note: a double bond is necessary to give all atoms 8 electrons 11

12. Lewis Dot Structure for SO3 The diagram below shows the dot structure for sulfur trioxide. The bonding electrons are in shown in red and lone pairs are shown in blue. 12

13. Example 3: NH4+ 1 N = 5 e- 4 H = 4x1 = 4 e- (+) charge = -1 e- --------- Total = 8 e- Note: Hydrogen atoms only need 2 e- rather than 8 e- Electron Dot StructuresNH4+ 13

14. Example: Carbon Dioxide C 4 e-O 6 e- x 2 O’s = 12 e- Total: 16 valence electrons 1. Central atom = 2. Valence electrons = 3. Form bonds. This leaves 12 electrons (6 pairs). 4.Place lone pairs on outer atoms. • Check to see that all atoms have 8 electrons around it • except for H, which can have 2.

15. Carbon Dioxide, CO2 C 4 e- How many are in the drawing? O 6 e- X 2 O’s = 12 e-Total: 16 valence electrons • There are too many electrons in our drawing. We need to take away 4 electrons and then form DOUBLE BONDS between C and O. • Instead of sharing only 1 pair, a double bond shares 2 pairs. • Take away the lone pairs from the carbon. So one pair is taken away from each oxygen atom and replaced with another bond.

16. LEWIS STRUCTURE RULES • Add up the total number of valence electrons in the molecule. • Draw the skeletal structure. • Use a line between each element to symbolize an electron pair. • Distribute the remaining electrons around the elements in pairs to form octets. (Hydrogen can only ever have 2 electrons.) • If you do not have enough to form octets, make double or triple bonds. • Ions must have square brackets around them with the charge notated in the top right hand corner. • To be a correct Lewis structure, ALL electrons must be shown.

17. UNDERSTANDING/KEY IDEA 4.3.C Some atoms, like Be and B, might form stable compounds with incomplete octets of electrons.

18. BF3 SF4 Violations of the Octet Rule Violations of the octet rule usually occur with B and elements of higher periods. Some common examples include: Be, B, P, S, and Xe. Be: 4 B: 6 P: 8 OR 10 S: 8, 10, OR 12 Xe: 8, 10, OR 12

19. OCTET RULE EXCEPTIONS • Hydrogen will never have more than 2 electrons. • Be and B have less than 8 electrons. • Some elements like S and P can have expanded octets which hold more than 8 electrons. • Coordinate covalent bonds are formed when both electrons originate from the same atom. • An arrow is used to denote the direction in a coordinate covalent bond showing the atom from which both electrons originated. • This is common in double and triple bonds.

20. GUIDANCE Coordinate covalent bonds should be covered.

21. Coordinate Covalent Bonds • Coordinate covalent bonds occur when one atom donates both of the electrons that are shared between two atoms • Coordinate covalent bonds are also called Dative Bonds

22. UNDERSTANDING/KEY IDEA 4.3.D Resonance structures occur when there is more than one possible position for a double bond in a molecule.

23. RESONANCE • Resonance is a concept used to describe the structures when there are multiple ways to depict the same molecule. • If you can put a double bond in more than one position, you will be expected to draw the resonance structures. • The electrons are actually delocalized in the areas of the double bonds and are spread out equally among all bonding positions. • Bond strength and length are in between that of single and double bonds.

24. Resonance structures allow us to depict all the possible positions of the double bonds. • The true structure, however, is an intermediate form known as a resonance hybrid. • Double arrows are placed between all resonance structures. Ref: myweb.astate.edu

25. APPLICATION/SKILLS Be able to deduce resonance structures. Examples include but are not limited to C6H6, CO32- and O3.

26. BENZENE www.pixmule.com

27. CARBONATE www.archives.evergreen.edu

28. OZONE www.chemwiki.ucdavis.edu

29. UNDERSTANDING/KEY IDEA 4.3.E Shapes of species are determined by the repulsion of electron pairs according to the VSEPR theory.

30. APPLICATION/SKILLS Be able to use the VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three and four electron domains.

31. Video VSEPR THEORY Valence Shell Electron Pair Repulsion theory. States that in a small molecule, the pairs of valence electrons are arranged as far apart from each other as possible. So far we have dealt with structural formulas which only show the types of atoms, bonds and lone pairs of electrons. They do not show the shape of the molecule.

32. 12 BASIC SHAPES • LINEAR – two atoms bonded to the central atom, no lone pairs of electrons on the central atom. • BENT- two atoms bonded to the central atom with one or two lone pairs of electrons on the central atom.

33. TRIGONAL PLANAR(flat triangle) – three atoms bonded to the central atom, no lone pairs of electrons on the central atom. • TRIGONAL PYRAMIDAL – three atoms bonded to the central atom, one lone pair of electrons on the central atom.

34. TETRAHEDRAL – four atoms bonded to the central atom, no lone pairs of electrons on the central atom. • TRIGONAL BIPYRAMIDAL – five atoms bonded to the central atom (octet rule exception)

35. SEE SAW (also called unsymmetrical tetrahedron) – derivative of the trigonal bipyramidal with one lone pair of electrons. • Note that on the three trigonal bipyramidal shape derivatives, the shapes come from pulling off atoms from the flat triangle, not the top and bottom.

36. T-SHAPED – also a derivative of the trigonal bipyramidal shape with 2 lone pairs of electrons on the central atom. • LINEAR – derivative of the trigonal bipyramidal shape with three lone pairs of electrons on the central atom. • This linear is different from the other linear as it has 3 pairs of lone electrons on the central atom.

37. OCTAHEDRAL – six atoms bonded to the central atom (another octet rule exception). • SQUARE PYRAMIDAL – derivative of the octahedral with one lone pair of electrons on the top or bottom. • SQUARE PLANAR – derivative of the octahedral shape with two lone pairs of electrons – one on top and one on bottom.

38. APPLICATION/SKILLS Be able to predict molecular polarity from bond polarity and molecular geometry.

39. MOLECULAR POLARITY • The polarity of a molecule depends upon: • The polar bonds it contains. • The shape of the molecule. • If the bonds are equally polar and arranged symmetrically, then they cancel each other out and are non-polar. • If the molecule contains bonds of different polarities or the bonds are not arranged symmetrically, then the molecule will be polar. • You can usually tell by the shape and lone pairs of electrons if the molecule is polar or not.

40. APPLICATION/SKILLS Be able to predict the bond angle from molecular geometry and presence of non-bonding pairs of electrons.

41. ELECTRON DOMAINS • Double and triple bonded electron pairs are orientated together and behave as a single unit known as an electron domain. • Lone pairs also count as electron domains so the 12 VSEPR shapes are narrowed down to 5 basic shapes. • For the electron domain shapes, you will need to know bond angles and hybridization.

42. LINEAR • A linear molecule has two electron domains. • The angle is 180 degrees and it has “sp” hybridization. • Non polar

43. TRIGONAL PLANAR • A trigonal planar molecule has 3 electron domains. • It has angles of 120 degrees and “sp2” hybridization. Non polar • The bent molecule can also have 3 electron domains. Polar

44. TETRAHEDRAL • A tetrahedral molecule has four electron domains. • It has angles of 109.5 degrees and “sp3” hybridization. Non polar • Trigonal pyramidal and “bent” with 2 lone pairs also have four electron domains. Polar

45. TRIGONAL BIPYRAMIDAL • A trigonal bipyramidal molecule has 5 electron domains. • It has angles of 90 and 120 degrees and “dsp3” hybridization. • Non polar (see saw and t-shaped are polar)

46. OCTAHEDRAL • An octahedral molecule has 6 electron domains. • It has angles of 90 degrees and “d2sp3” hybridization. Non polar (square pyramidal – polar)

47. GUIDANCE Allotropes of carbon (diamond, graphite, graphene and C60 buckminsterfullerene) should be covered.

48. ALLOTROPES OF CARBON • Some covalent structures are crystalline in nature like ionic lattices; however, they are linked together with covalent bonds. • The crystal is a single molecule with a regular repeating pattern of covalent bonds. • It is referred to as a giant molecular structure. • Allotropes are different forms of an element in the same physical state. • Different bonding within the structures gives rise to different physical properties. • Carbon has four allotropes.

49. GRAPHITE • Graphite • Each C atom is sp2 hybridized covalently bonded to 3 others forming hexagons in parallel layers with bond angles of 120 degrees. • The layers are held together by van der Waal’s forces so they can slide over each other. • Density is 2.26 g/cm3 • Contains one non-bonded, delocalized electron per atom so graphite conducts electricity due to the movement of these electrons. • Not a good heat conductor • Very high melting point, most stable allotrope • Non lustrous – grey solid • Used as a lubricant and in pencils

50. DIAMOND • Diamond • Each C atom is sp3 hybridized covalently bonded to 4 others tetrahedrally in a regular repeating pattern with bond angles of 109.5 degrees. • It is the hardest known natural substance. • Density is 3.51 g/cm3 • All electrons are bonded so it does not conduct electricity. • Does conduct heat better than metals. • Very high melting point, brittle • Lustrous crystal • Used in jewelry and tools