George Mason University General Chemistry 211 Chapter 10 The Shapes Geometry of Molecules Acknowledgements Course Text:

2. Lewis Electron-Dot Symbols A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element Note that the group number indicates the number of valence electrons

3. Lewis Electron-Dot Formulas A Lewis electron-dot formula is an illustration used to represent the transfer of electrons during the formation of an ionic bond The Magnesium has two electrons to give, whereas the Fluorines have only one �vacancy� each Consequently, Magnesium can accommodate two Fluorine atoms

4. Lewis Structures The tendency of atoms in a molecule to have eight electrons (ns2np6) in their outer shell (two for hydrogen) is called the octet rule You can represent the formation of the covalent bond in H2 as follows: This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots

5. The Electron Probability Distribution for the H2 Molecule

6. Lewis Structures The shared electrons in H2 spend part of the time in the region around each atom In this sense, each atom in H2 has a helium (1s2) configuration

7. Lewis Structures The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar)

8. Lewis Structures Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures An electron pair is either a: bonding pair (shared between two atoms) lone pair (electron pair that is not shared) Hydrogen has no unbonded pairs Chlorine has 3 unbonded pairs

9. Lewis Structures Rules for obtaining Lewis electron dot formulas Calculate the number of valence electrons for the molecule (= group # for each atom), add the charge of anion and subtract the charge of a cation Write the skeleton structure of the molecule or ion Put atom with lowest group number and lowest electronegativity as the central atom Distribute electrons to atoms surrounding the central atom to satisfy the octet rule for each atom Distribute the remaining electrons as pairs to the central atom (or atoms) If the Central atom is deficient in electrons to complete the octet; move electron pairs from surrounding atoms to complete central atom valence electron needs

10. Practice Problem Write a lewis structure for CCl2F2 Step 1: Arrange Atoms (Carbon is �Central Atom� because is has the lowest group number and lowest electronegativity Step 2: Determine total number of valence electrons 1 x C(4) + 2 x Cl(7) + 2 x F(7) = 32 Step 3: Draw single bonds to central atom and subtract  2 e- for each single bond (4 x 2 = 8) ? 32 � 8 = 24 remaining Step 4: Distribute the 24 remaining electrons in pairs around surrounding atoms (3 electron pairs around each Fluoride atom)

11. Writing Lewis Dot Formulas The Lewis electron-dot formula of a covalent compound is a simple two-dimensional representation of the positions of electrons in a molecule Bonding electron pairs are indicated by either two dots or a dash In addition, these formulas show the positions of lone pairs of electrons

12. Writing Lewis Dot Formulas The following rules allow you to write electron-dot formulas even when the central atom does not follow the octet rule To illustrate, draw the structure of: Phosphorus Trichloride

13. Writing Lewis Dot Formulas Step 1: Total all valence electrons in the molecular formula. That is, total the group numbers of all the atoms in the formula  For a polyatomic anion, add the number of negative charges to this total For a polyatomic cation, subtract the number of positive charges from this total

14. Writing Lewis Dot Formulas Step 2: Arrange the atoms radially, with the least electronegative atom in the center. Place one pair of electrons between the central atom and each peripheral atom

15. Writing Lewis Dot Formulas Step 3: Distribute the remaining electrons to the peripheral atoms to satisfy the octet rule

16. Writing Lewis Dot Formulas Step 4: Distribute any remaining electrons (2) to the central atom. If the number of electrons on the central atom is less than the number of electrons required to complete the octet for that atom, use one or more electrons pairs from other atoms to form double or triple bonds

17. Exceptions to the Octet Rule Although many molecules obey the octet rule, there are exceptions where the central atom has more than eight electrons Generally, if a nonmetal is in the third period or greater it can accommodate as many as twelve electrons, if it is the central atom These elements have unfilled �d� subshells that can be used for bonding

18. Exceptions to the Octet Rule For example, the bonding in phosphorus pentafluoride, PF5, shows ten electrons surrounding the phosphorus

19. Exceptions to the Octet Rule In xenon tetrafluoride, XeF4, the xenon atom must accommodate two extra lone pairs

20. Delocalized Bonding: Resonance The structure of ozone, O3, can be represented by two different Lewis electron-dot formulas Experiments show, however, that both bonds are identical

21. Delocalized Bonding: Resonance According to Resonance Theory, these two equal bonds are represented as one pair of bonding electrons spread over the region of all three atoms  This is called delocalized bonding, in which a bonding pair of electrons is spread over a number of atoms

22. Resonance & Bond Order Recall (Chap 9) � Bond Order The number of electron pairs being shared by any pair of �Bonded Atoms� or The number of electron pairs divided by the number of bonded-atom pairs Ex. Ozone

23. Practice Problem In the following compounds, the Carbon atoms form a �single ring.� Draw a Lewis structure for each compound, identify cases for which �resonance� exists, and determine the C-C bond order(s). 

24. Practice Problem

25. Practice Problem

26. Formal Charge & Lewis Structures In certain instances, more than one feasible Lewis structure can be illustrated for a moleculeFor example, H, C and N The concept of �formal charge� can help discern which structure is the most likely Formal Charge: An atom�s formal charge is: Total number of valence electrons Minus all unshared electrons Minus � of its shared electrons Formal Charges must sum to actual charge of species: Zero Charge for a Molecule Ionic Charge for an Ion

27. Formal Charge & Lewis Structures When you can write several Lewis structures, choose the one having the least formal charge

28. Formal Charge & Lewis Structures

29. Formal Charge & Lewis Structures

30. Resonance/Formal Charge � Nitrate Ion

31. Resonance/Formal Charge � Cyanate Ion

32. Formal Charge vs Oxidation No �Formal Charge� is used to examine resonance hybrid structures , whereas �Oxidation Number� is used to monitor �REDOX� reactions Formal Charge - Bonding electrons are assigned equally to the atoms as if the bonding were �Nonpolar� covalent, i.e., each atom has half the electrons making up the bond Formal Charge = valence e- � (unbonded e- + � bonding e-) Oxidation Number - Bonding electrons are transferred completely to the more electronegative atom, as if the bonding were �Ionic� Oxidation No. = valence e- � (unbonded e- + bonding e-)

33. Formal Charge vs Oxidation No

34. The Valence-Shell Electron Pair Repulsion Model (VSEPR) The Valence-Shell Electron Pair Repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible Molecular geometry � The shape of a molecule is determined by the positions of atomic nuclei relative to each other, i.e., angular arrangement Central Atom Place atom with �Lower Group Number� in center(N in NF3 needs more electrons to complete octet) If atoms have same group number (SO3 or ClF3), place the atom with the �Higher Period Number� in the center (Sulfur & Chlorine)

35. VSEPR Model of Molecular Shapes The following rules and figures will help discern electron pair arrangements Select the Central Atom (Least Electronegative Atom) Draw the Lewis structure Determine how many bonding electron pairs are around the central atom. Determine the number of non-bonding electron pairs Count a multiple bond as �one pair� Arrange the electron pairs as far apart as possible to minimize electron repulsions Note the number of bonding and lone pairs

36. VSEPR Model of Molecular Shapes To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom Molecular Notation: A � The Central Atom (Least Electronegative  atom) X � The Ligands (Bonding Pairs) a � The Number of Ligands E � Non-Bonding Electron Pairs b � The Number of Non-Bonding Electron Pairs Double & Triple Bonds count as a �single� electron pair The Geometric arrangement is determined by: sum (a + b)

37. VSEPR Model of Molecular Shapes

38. VSEPR Model of Molecular Shapes

39. Arrangement of Electron Pairs About an Atom: Basic Shapes

40. Arrangement of Electron Pairs About an Atom: Basic Shapes

41. Arrangement of Electron Pairs About an Atom: Basic Shapes

42. Arrangement of Electron Pairs About an Atom: Basic Shapes

43. Arrangement of Electron Pairs About an Atom: Basic Shapes

44. Electron Pair Arrangement

45. Electron Pair Arrangement

46. Linear Geometry Two electron pairs (linear arrangement) Double bonds count as a �single electron pair� ? 2 bonding pairs 0 non-bonding pairs AXaEb = a + b = 2 + 0 = 2 (Linear) Thus, according to the VSEPR model, the bonds are arranged linearly (bond angle = 180o) Molecular shape of carbon dioxide is linear

47. Trigonal Planar Geometry Three electron pairs on Central atom  The three groups of electron pairs are arranged in a trigonal plane. Thus, the molecular shape of COCl2 is trigonal planar. The Bond angle is 120o

48. Trigonal Planar Geometry Effect of Double Bonds Bond angles deviate from ideal angles when surrounding atoms and electron groups are not identical A double bond has greater electron density and repels two single bonds more strongly than they repel each other

49. Trigonal Planar Geometry Effect of Lone Pairs The molecular shape is defined only by the positions of the nuclei When one of the three electron pairs in a trigonal planar molecule is a lone (non-bonding) pair, it is held by only one nucleus It is less confined and exerts a stronger repulsive force than a bonding pair This results in a decrease in the angle between the bonding pairs

50. Trigonal Planar Geometry Three electron pairs (Effect of �Lone� pairs) (trigonal planar arrangement) Ozone has two bonding electron pairs about the central oxygen (double bond counts as 1 pair) There is one non-bonding lone pair These groups have a: Trigonal Planar arrangement AXaEb (a + b = 2 + 1 = 3) Since one of the groups is a lone pair, the molecular geometry is described as bent or angular

51. Tetrahedral Geometry Four electron pairs (Tetrahedral Arrangement) Four electron pairs about the central atom lead to three different molecular geometries a + b = 4 + 0 a + b = 3 + 1 a + b = 2 + 2 = 4 = 4 = 4

52. Tetrahedral Geometry Molecular Geometries produced by variable non-bonding electron pairs

53. Trigonal Bipyramidal Five electron pairs (trigonal bipyramidal arrangement) This structure results in both 90o and 120o bond angles

54. Trigonal Bipyramidal Other molecular geometries are possible when one or more of the electron pairs is a lone pair

55. Octahedral Geometry Six electron pairs (Octahedral arrangement) This octahedral arrangement results in: 90o bond angles

56. Other Geometries Six electron pairs (octahedral arrangement)

57. Practice Problem In the ICl4� ion, the electron pairs are arranged around the central iodine atom in the shape of a. a tetrahedron b. a trigonal bipyramid c. a square plane d. an octahedron e. a trigonal pyramid Ans: a

58. Dipole Moment andMolecular Geometry The dipole moment is a measure of the degree of charge separation in a molecule The polarity of individual bonds within a molecule can be viewed as vector quantities Thus, molecules that are perfectly symmetric have a zero dipole moment. These molecules are considered nonpolar

59. Dipole Moment andMolecular Geometry However, molecules that exhibit any asymmetry in the arrangement of electron pairs would have a nonzero dipole moment. These molecules are considered polar

60. Dipole Moment andMolecular Geometry

61. Practice Problem The Nitrogen atom would be expected to have the positive end of the dipole in the species a. NH4+ b. Ca3N2 c. HCN d. AlN e. NO+ Ans: e

62. Practice Problem Which of the following molecules is polar? a. BF3 b. CBr4 c. CO2 d. NO2 e. SF6 Ans: d

63. Practice Problem Which of the following compounds is nonpolar? a. XeF2 b. HCl c. SO2 d. H2S e. N20 Ans: a HCL is ionic and very polar SO2 has AX2E1 Trigonal Planar Bent geometry with a dipole moment (polar) H2S has AX2E2 Tetrahedral Bent geometry and with a dipole moment (polar) N2O has AX2E0 linear with asymmetric geometry.Since oxygen is more EN than N, the molecule is polar XeF2 has AX2E3 Trigonal Bypyramidal Geometry, but linear molecular geometry (nonpolar)

64. Equation Summary