1. Chapter 7 Bonding

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

3. Ionic Bonding l An atom with a low ionization energy reacts with an atom with high electron affinity. l The electron is transferred. l Opposite charges hold the atoms together.

4. Coulomb's Law l If charges are opposite, energy change is negative l exothermic l Same charge, positive E, requires energy to bring them together.

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

6. Covalent Bonding l Electrons are shared by atoms. l In polar bonds the electrons are not shared evenly. l One end is slightly positive, the other negative. Indicated using small delta 

7. Electronegativity l The ability of an atom to attract shared electrons to itself. l Tends to increase left to right. l decreases as you go down a group. l Noble gases aren’t part of it. l Difference in electronegativity between atoms tells us how polar bond is.

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

9. Dipole Moments l A molecule with a center of negative charge and a center of positive charge is dipolar (two poles), l Also called a dipole moment. l Center of charge doesn’t have to be on an atom. l It line up in the presence of an electric field.

10. How It is drawn H - F ++ --

11. Which Molecules Have Them? l Any two atom molecule with a polar bond. l With three or more atoms there are two considerations: –There must be a polar bond. –Geometry can cancel it out.

12. Geometry and polarity l Three shapes will cancel them out. l Linear

13. Geometry and polarity l Three shapes will cancel them out. l Planar triangles 120º

14. Geometry and polarity l Three shapes will cancel them out. l Tetrahedral

15. Geometry and polarity l Others don’t cancel l Bent

16. Geometry and polarity l Others don’t cancel l Trigonal Pyramidal

17. Ions l Atoms tend to react to form noble gas configuration. l Metals lose electrons to form cations l Nonmetals can share electrons in covalent bonds. l Or they can gain electrons to form anions.

18. Ionic Compounds l Ions align themselves to maximize attractions between opposite charges, l and to minimize repulsion between like ions. l Can stabilize ions that would be unstable as a gas. l React to achieve noble gas configuration

19. Na(s) + ½F 2 (g)  NaF(s) First sublime NaNa(s)  Na(g)  H = 109 kJ/mol Ionize Na(g) Na(g)  Na + (g) + e -  H = 495 kJ/mol Break F-F Bond½F 2 (g)  F(g)  H = 77 kJ/mol Add electron to F F(g) + e -  F - (g)  H = -328 kJ/mol

20. Na(s) + ½F 2 (g)  NaF(s) Lattice energy Na(s) + ½F 2 (g)  NaF(s)  H = -928 kJ/mol l An ionic compound is considered any substance that conducts electricity when melted. l Also use the generic term salt.

21. The Covalent Bond l Due to the tendency of atoms to achieve the lowest energy state. l It takes 1652 kJ to dissociate a mole of CH 4 into its ions l Since each hydrogen is hooked to the carbon, we get the average energy = 413 kJ/mol

22. l CH 3 Cl has 3 C-H, and 1 C - Cl l the C-Cl bond is 339 kJ/mol l The bond is a method of explaining the energy change associated with forming molecules. l Bonds don’t exist in nature, but are useful as a model of behavior.

23. Covalent Bond Energies l We made some simplifications in describing the bond energy of CH 4 l Each C-H bond has a different energy. CH 4  CH 3 + H  H = 453 kJ/mol CH 3  CH 2 + H  H = 435 kJ/mol CH 2  CH + H  H = 425 kJ/mol CH  C + H  H = 339 kJ/mol

24. Localized Electron Model l A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. l Three Parts 1) Valence electrons using Lewis structures 2) Prediction of geometry using VSEPR 3) Description of the types of orbitals

25. Lewis Structure l Shows how the valence electrons are arranged. l One dot for each valence electron. l A stable compound has all its atoms with a noble gas configuration. l Hydrogen follows the duet rule, Be the quartet, B the sextet. l The rest follow the octet rule. l Bonding pair is the one between the symbols.

26. Rules l Sum the valence electrons. l Use a pair to form a bond between each pair of atoms. l Arrange the rest to fulfill the octet rule (except for H, Be and B). l H 2 O l A line can be used instead of a pair.

27. Practice l 1 st atom is commonly the central atom l POCl 3 l SO 4 -2 l SO 3 -2 l PO 4 -3 l SCl 2

30. Exceptions to the octet l SF 6 l Third row and larger elements can exceed the octet l Because they use 3d orbitals l I 3 -

31. Exceptions to the octet l When we must exceed the octet, extra electrons tend to go on central atom. l ClF 3 l XeO 3 l ICl 4 - l BeCl 2

34. Resonance l Sometimes there is more than one valid structure for an molecule or ion. l NO 3 - l Use double arrows to indicate it is the “average” of the structures. l It doesn’t switch between them. l NO 2 - l Localized electron model is based on pairs of electrons, doesn’t deal with odd numbers.

35. Formal Charge l For molecules and polyatomic ions that exceed the octet there are several different structures. l Use charges on atoms to help decide which. l Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

36. Formal Charge l The difference between the number of valence electrons on the free atom and that assigned in the molecule. l We count half the electrons in each bond as “belonging” to the atom. l SO 4 -2 l Molecules try to achieve as low a formal charge as possible. l Negative formal charges should be on electronegative elements.

37. More Examples l XeO 3 l NO 4 -3 l SO 2 Cl 2

40. VSEPR l Lewis structures tell us how the atoms are connected to each other. l They don’t tell us anything about shape. l The shape of a molecule can greatly affect its properties. l Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

41. VSEPR l Molecules take a shape that puts electron pairs as far away from each other as possible. l Have to draw the Lewis structure to determine electron pairs. l Bonding or nonbonding lone pairs l Lone pair take more space.

42. VSEPR l The number of pairs determines –bond angles –underlying structure l The number of atoms determines –actual shape

43. VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar 4109.5° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octahedron

44. Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent

45. Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 550trigonal bipyrimidal 541See-saw 532T-shaped 523linear

46. Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 660Octahedral 651Square Pyramidal 642Square Planar 633T-shaped 621linear

47. If no central atom… l Can predict the geometry of each angle. l build it piece by piece.

48. How well does it work? l Predictions are almost always accurate. l But, like all simple models, it has exceptions.

49. Atomic Orbitals Don’t Work l to explain molecular geometry. l In methane, CH 4, the shape is tetrahedral. l The valence electrons of carbon should be two in s, and two in p. l the p orbitals would have to be at right angles. l The atomic orbitals change when making a molecule

50. l When two things come off. l One s and one p hybridize. l linear

51. sp hybridization l End up with two lobes 180º apart. l p orbitals are at right angles Makes room for two  bonds and two sigma bonds. l A triple bond or two double bonds.

52. sp 2 hybridization l When three things come off atom. l trigonal planar l 120º on  one  bond

53. sp 2 hybridization l C 2 H 4 l Double bond acts as one pair. l trigonal planar l Have to end up with three blended orbitals. l Use one s and two p orbitals to make sp 2 orbitals. l Leaves one p orbital perpendicular.

54. sp 3 Hybridization l We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry. l We combine one s orbital and 3 p orbitals. l sp 3 hybridization has tetrahedral geometry.

55. How we get to hybridization l We know the geometry from experimentation l We know the orbitals of the atom l hybridizing atomic orbitals can explain the geometry. l So if the geometry requires a tetrahedral shape, it is sp 3 hybridized l This includes bent and trigonal pyramidal molecules because one of the sp 3 lobes holds the lone pair.

56. sp 3 d hybridization l The model predicts that we must use the d orbitals. l sp 3 d hybridization l sp 3 d hybridization, ex: PCl 5 l Trigonal bipyrimidal can only  bond. can’t  bond. l basic shape for five bonds.

57. sp 3 d 2 l gets us to six bonds around l octahedral

58. Two types of Bonds l Sigma bonds from overlap of orbitals. l Between the atoms. Pi bond (  bond) above and below atoms l Between adjacent p orbitals. l The two bonds of a double bond.

59. CO 2 C can make two  and two  O can make one  and one  COO