1. Chapter 8 Bonding and Molecular Structure

2. Overview Chemical Bond Formation Covalent Bonding and Lewis Structures Atom Formal Charges in Covalent Molecules and Ions Resonance Exceptions to the Octet Rule Molecular Shapes Bond Polarity and Electronegativity Bond and Molecular Polarity Bond Properties: Order, Length, Energy

3. Chapter Goals Understand the difference between ionic and covalent bonds. Draw Lewis electron dot structures for small molecules and ions. Use VSEPR to predict the shapes of simple molecules and ions and to understand the structures of more complex molecules.

4. Chapter Goals Use electronegativity and formal charge to predict the charge distribution in molecules and ions, to define the polarity of bonds, and to predict the polarity of molecules. Understand the properties of covalent bonds and their influence on molecular structure.

5. Valence Electrons There are two types of electrons in an atom: valence and core electrons. Valence electrons are in the outermost shell and core electrons are the rest. For main group elements, the group number tells you the number of valence electrons the element has. Lewis electron dot symbols show the number of valence electrons in the atom.

6. Check and See Give the number of valence electrons for Ca and Se. Draw the Lewis dot symbol for each element. Give the number of valence electrons for Ba, As and Br. Draw the LDS for each element.

7. Chemical Bond Formation When a chemical reaction occurs between two atoms, their valence electrons are reorganized so that a net attractive force— a chemical bond—occurs between atoms. Bonds are either ionic or covalent and we can depict them through LDS.

8. Chemical Bond Formation An ionic bond forms when one or more valence electrons is transferred from one atom to another creating positive and negative ions. Sodium loses an electron to chlorine to form Na + and Cl -. The bond is the attractive force between the cation and anion.

9. Chemical Bond Formation Covalent bonding involves the sharing of valence electrons between atoms. Two chlorine atoms share a pair of electrons to form a covalent bond.

10. Covalent Bonding & Lewis Structures A covalent bond results when one or more electron pairs are shared between two atoms. The electron pair bond between the two atoms of a H 2 molecule is represented by a pair of dots or a line. H : H or H—H These representations are called Lewis structures.

11. Covalent Bonding & Lewis Structures Most atoms in molecules have bond pairs and lone pairs of electrons. Some atoms also share more than one pair of electrons. They are referred to as double and triple bonds. Examples are CO 2 and N 2.

12. Octet Rule Each atom is surrounded by an octet of eight electrons. Hydrogen is an example of an atom that will not have an octet of eight electrons around it.

13. Formal Charge Formal Charge = Group # - (LPE + ½ BPE)

14. Check and See Draw Lewis structures for ammonia, the hypochlorite ion and the nitrite ion. Draw Lewis structures for the ammonium ion, carbon monoxide and the sulfate ion.

15. Mini Quiz Name the shape of each molecular geometry.

16. Drawing Lewis Structures It is often helpful to understand how atoms “generally” bond. Carbon—4 bonds, no lone pairs Nitrogen—3 bonds, 1 lone pair Oxygen—2 bonds, 2 lone pairs Fluorine—1 bond, 3 lone pairs

17. Check and See Draw Lewis electron dot structures for CCl 4 and NF 3. Predict Lewis structures for methanol, CH 3 OH, and hydroxylamine, NH 2 OH. The formulas of these compounds are written to guide you in choosing the correct arrangement of atoms. Draw the Lewis structure for the anion, H 2 PO 4 -, derived from phosphoric acid.

18. Isoelectronic Species Draw NO +, N 2, CO and CN -. What do you notice about them all?

19. Check and See Is the acetylide ions, C 2 2-, isoelectronic with N 2 ?

20. Resonance Draw O 3. Notice that you can place the double bond in one of two places. To take care of this, Linus Pauling came up with the idea of resonance and resonance structures.

21. Resonance Resonance structures represent bonding in a molecule or ion when a single Lewis structure fails to describe accurately the actual electronic structure. The different structures have the same energy and are said to exist simultaneously as a hybrid.

22. Check and See Draw resonance structures for the carbonate ion. The  seems to imply that a reaction is taking place or that the structure is alternating between forms. These ideas, however, are not true. Draw resonance structures for the nitrite ion. Are the N—O bonds single, double or intermediate?

23. Exceptions to the Octet Rule Some central atoms can have fewer or more than eight electrons around them. Boron (B) typically does this for the less than rule. Boron only has 3 valence electrons and often it will share these 3 with other elements to give it a total of 6 around it. Like in BF 3 and B(OH) 3.

24. Exceptions to the Octet Rule However, because of this electron deficiency it will be strongly attracted to a lone pair of electrons from some other molecules, say NH 3. When they are held tightly, this is said to be a coordinate covalent or dative bond. On the other hand, a lot of central atoms can have more than 8 electrons around them and still be okay. Any non-metal in the 3 rd row or beyond can do this.

25. Exceptions to the Octet Rule The reason for this is elements in the 2 nd row have only four orbitals in which to hold electrons—1s and 3p orbitals. However, elements in the 3 rd row and beyond have d-orbitals also that can accommodate the extra electrons.

26. Check and See 1.Sketch the Lewis structure of the [ClF 4 ] - ion. 2.Sketch the Lewis structures for [ClF 2 ] + and [ClF 2 ] - ions. How many lone pairs and bond pairs surround the Cl atom in each ion?

27. Molecules with an Odd # of e- When counting up the number of possible electrons to place for a molecule or ion— most of the time you will get an even number. This is because electrons like to be in pairs. However, sometimes you get those naughty ions and molecules that have an odd number of electrons.

28. Molecules with an Odd # of e- They are called free radicals and are very reactive One example is NO 2 that converts to poisonous NO from car exhaust that cause environmental pollution.

29. Check and See 1.Draw NO 2. (It has two resonance structures) 2.The reaction that is adverse to the environment is NO 2 reacting with O 2 and sunlight to produce NO gas and ozone (O 3 )—write the balanced equation.

30. Bond Polarity & Electronegativity When two dissimilar atoms form a covalent bond, the electron pair will be unequally shared. The result is a polar covalent bond. Bonds are polar because they all atoms don’t hold onto their electrons equally. Electronegativity (Χ) is the measure of how atoms attract electrons to themselves. Page 376 has a PT that shows the Χ values from which we can determine how unequally atoms are sharing electrons.

31. Bond Polarity & Electronegativity The bigger the difference in Χ, the more polar the bond is—meaning the more unequal the share of electrons between atoms. If there is a great enough “unsharing” the stronger atom will take the weaker atom’s electron and the bond will be ionic.

32. Check and See Decide which bond in each pair below is more polar 1.B—F or B—Cl 2.Si—O or P—P 3.C=O or C=S

33. Electroneutrality Principle This principle declares that the electrons in a molecule are distributed in such a way that the charges on the atoms are as close to zero as possible. For example, how do we draw OCN -, which one is in the center and where are the multiple bonds?

34. Mini Quiz 1.Draw PCl 5, SF 4, ClF 3 and I 3 -. 2.How many total electron pairs are on each central atom? 3.What is the electron geometry of each atom? 4.What is the molecular geometry of each atom?

35. Bond Properties The bond properties we will look at are bond order, bond length and bond energy. The order of a bond is the number of bonding electron pairs shared by two atoms in molecule.

36. Bond Properties Single covalent bonds have an order of 1. Double bonds have an order of 2. Triple bonds have an order of 3. When an atom shares a pair of electrons with one atom and two pairs of electrons with another, the order is 1.5. We calculate bond order by… B.O. = # of shared pairs / number of links Draw ozone… Here we see… B.O. = 3 / 2 = 1.5

37. Bond Properties Bond length changes with the number of shared pairs. Single bonds are longer than double bonds which are longer than triple bonds. The bigger the bond order, the shorter the bond length for similar atoms.

38. Bond Properties Look at the CO 3 2- ion. There is one double bond and 2 single bonds for an overall bond order of… B.O. = # of shared pairs / number of links 4/3 or 1.33

39. Bond Properties Look at page 387…bond lengths and find lengths for C—O and C=O… We would expect that the bond length would be somewhere between a C—O bond and a C=O bond (143pm and 122pm). The measured bond length is 129pm Bond length is measured by a process called X-ray diffraction…read up if you are really interested in this.

40. Check and See Give the bond order of each of the following bonds and arrange them in order of decreasing bond distance: C=N, C-TB-N, C—N BO = #pairs/links 2, 3 and 1 Draw the resonance structures for NO 2 -. What is the NO bond order in this ion? Consult Table 9.8 for N—O and N=O bond lengths. Compare these with the NO bond length in NO 2 - (124pm).

41. Mini Quiz Hi…my name is bromine pentafluoride…please draw a Lewis Structure of me and name my electron geometry and my molecular geometry. Thank you…my friend here is named xenon tetrafluoride. Could you please draw her and name her electron and molecular geometries? Nice…we had a kid and name him sulfur hexafluoride…cool name, eh? Anyway, please draw him with his electron and molecular geometries.

42. Bond Energy ΔH rxn =Σ ΔH bonds brkn – Σ ΔH bonds frmd

43. Check and See 1.Acetone can be converted to isopropanol by hydrogenation. Calculate the enthalpy change for the following reaction using bond energies… H 3 CCOCH 3 + H 2  H 3 CCHOHCH 3 2.Estimate the heat of combustion of methane.

44. Molecular Shapes Study Table 8.10 on page 394. Draw each molecule and see how they match up with their shapes in the chapter.

45. Mini Quiz 1.CF 4 2.BrF 5 3.BrF 3 4.NH 3 5.CO 2 6.XeCl 3 - 7.SO 3 8.PF 5 9.SF 6 10.ICl 2 - 11.ICl 4 - 12.SF 4

46. Molecular Polarity Because most bonds are polar, the molecules that are made up of these polar bonds are polar. If you put polar molecules in an electric field and turn it on, the negative ends of the molecules will line up with the positive end of the field. (see page 380— Fig 8.12)

47. Molecular Polarity The extent to which molecules line up with the electric field is called their dipole moment. The unit of the dipole moment is the Coulomb meter (C·m) or the debye (D) after Peter Debye. Basically, the more polar the molecule the larger the dipole moment.

48. Hybridization This is a big messed up subject that is covered just a bit on the AP. After we talk about it here and you see it on the AP exam, it is quite possible you will never see it again. Anyway, here is the gist… Hybridization is a mixing of orbitals during bond formation. There are s, p and d orbitals involved in bonding.

49. Hybridization When there are 2 bonds (links) around a central atom, the s orbital and one of the p orbitals mix (hybridize) and make 2-sp hybridized orbitals—instead of an s orbital bonding and a p orbital bonding. When there are 3 pairs, an s and 2 p’s hybridize making sp 2 hybridization. Four pairs? sp 3 hybridization Five pairs? sp 3 d hybridization Six pairs? sp 3 d 2 hybridization

50. Check and See Are nitrogen trifluoride, dichloromethane, and sulfur tetrafluoride polar or non- polar? If polar, indicate the negative and positive sides of the molecule.