2. Chemical Bonding Chapter 12

4. Types of Bonds 1. Ionic bond Transfer of e- from a metal to a nonmetal and the resulting electrostatic force that holds them together forms an ionic compound. EX: Na + + Cl -  NaCl (neutral)

5. Types of Bonds 2. Covalent bond Formed from the sharing of e- pairs between two or more nonmetals resulting in a molecule. EX: H 2 + O  H 2 O

6. Types of Bonds 3. Metallic bond Metals bonding with other metals do not gain or lose e- or share e- unequally. These bonds are created from the delocalized e- that hold metallic atoms together.

7. Electronegativity Remember that electronegativity is the tendency of an atom to attract an e-, specifically when bonded to another atom. Electronegativities of an element are influenced by the same factors that affect ionization energies and electron affinities of the elements.

8. Review trends Decreasing electronegativity Decreasing electron affinity

9. Bond Character Electrons of atoms are exchanged when the difference between electronegativity between the atoms is high. A difference more than 1.7 will produce ionic bonds. EX: Magnesium + Oxygen |3.50 – 1.23| = 2.27 difference What type of bond?

10. That’s right! IONIC

11. Bond Character is really called “ionic character” Let’s try 3 more: B – P |2.01 – 2.06| = 0.05 COVALENT Mg – N |1.23 – 3.07| = 1.84 IONIC C – Na |2.50 – 1.01| = 1.49 COVALENT

12. Ionic character: a spectrum, not black and white Ionic character is a way to describe the BEHAVIOR of a compound based on it’s bonding: Nonpolar polar covalent ionic covalent 0.0 increasing electronegativity difference 1.7 + increasing polarity Most compounds fall into the middle; polar covalent. The closer it gets to ionic the more “ionic character” the compound demonstrates

13. What is ionic character Ionic compounds have the following characteristics: Solid at room temphigh melting point temp. Dissolves in waterforms crystal Electrolyte in solution (conducts electricity when dissolved) The greater the E.N. difference between atoms in a compound the more ionic character it demonstrates; or, the more polar the more ionic character.

14. Properties of Nonpolar covalent compounds (molecules) Gases at room temp. Not soluble in water (won’t dissolve) Nonelectrolyte Low melting point temps. Low boiling point temps. Examples: N 2, CO 2, O 2

15. Valence Electrons The electrons that are involved in bonding are the outer most electrons. These electrons are called valence electrons. According to the Octet Rule, eight valence electrons assures stability. Atoms exchange or share electrons so that they can have eight valence electrons.

16. Lewis Dot Structures To visualize valence e-, we will use Lewis Dot Structures. Step 1: The element symbol represents the nucleus and all e- except valence. Step 2: Write the e- config. From config., select e- in the outer level. These e- are the ones with the largest principal quantum numbers.

17. Lewis Dot Structures Step 3: Each “side” of symbol represents an orbital. Draw dots on sides as you would fill orbitals. One at a time, then paired. Start with the bottom orbital and work clockwise when filling.

18. Lewis Dot Structures EX: carbon step 1: C step 2: [He]2s 2 2p 2 step 3: C

19. Lewis Dot Structures EX: bromine step 1: Br step 2: [Ar]4s 2 3d 10 4p 5 step 3: Br

20. Shorthand Lewis Dot Structures Because e- config. demonstrates a periodic trend, most valence e- can be determined by placement on periodic table. Remember oxidation states (+ or – charges of atoms), valence e- are the cause of these states.

21. Shorthand Lewis Dot Structures Group 1 has 1 valence, Group 2 has 2, Group 13 has 3, and so on… 1 2 34567 8

22. Drawing Ionic Bonds Draw Lewis dots for each element. Na + Cl Draw an arrow to show the e- exchange between atoms.

23. Drawing Simple Ionic Bonds Now draw them together. Na Cl Ionic bonds form because of the attraction of the negative charges.

24. Exchange of Electrons

25. Lewis Dot Diagrams of Molecules 1. total number of valence e-. 2. Determine the central atom. The following are guides: Often the unique atom (only one of it) is the central atom. Or put the least electronegative element in the middle. 3. Arrange the other atoms around the central atom.

26. 4. Connect all atoms with one line (bond). 5. Subtract 2 electrons for each bond from total valence electrons. 6. Distribute the remaining electrons in pairs around the atoms to satisfy the octet rule. Assign them to the most electronegative atom first.

27. 4 7. If you run out of electrons before all atoms have an octet of electrons, you need to form double or triple bonds. 1. 8. If you have extra electrons and all of the atoms have an octet, put the extra electrons on the central atom in pairs. 2. 9. If the central atom has an atomic number greater than fifteen, you are allowed to have more than eight electrons around it.

28. Lewis Dot Diagrams EX: AsI 3 1. Count valence e- [5 + (3 x 7)]= 26 2. Place As (least in number) in center. 3. Place the three Iodines around As. 4. Draw lines (bonds connecting them) I –As –I I

29. 5. Subtract # used for bonds (26 – 6) = 20 e-. 6. Place e- around the three I atoms first because they are the most electronegative 7. all I have 8 valence e-. We used 18 of the 20 meaning there 2 leftover e-s. These are placed on As to complete octet :IAs I: :I:

30. 8. We have 2 extra e- (20 starting – (3x6) = 2. Place them around the central atom. I –As –I I

31. Lewis Dot Diagram of CH 2 O 1. Count total valence e- (C=4, H=1x2, O=6)= 12 valence e- 2. Place C in the center. 3. Place the 2 hydrogens and one oxygen around C. 4. Draw lines connecting H and O to C. 5. Subtract number of bonded e- from total. (12-6) = 6

32. 6. Place e- around the oxygen atom first because it is the most electronegative. 7. We ran out of e- before carbon satisfied its octet. This means that we will have a double bond between the carbon and oxygen. (Hydrogen cannot form double bonds). 8. There will be no unshared e-.

33. Lewis Dot Diagram of CH 2 O HC O H

34. Lewis Dot Diagram of Polyatomic Ions Polyatomic ions are groups of covalently bonded atoms that carry an overall charge. Constructing Dot Diagrams for the polyatomic ions is the same, except the difference in charge (+ or -) must be accounted for.

35. Lewis Dot Diagram of ClO 3 1- 1. Count valence e-, including charge 7 + (3x6) + 1 = 26 total 2. Place Cl in the center. 3. Arrange the three O around it. 4. Draw bonds from O to Cl. 5. Subtract e- used in bonds (26 – 6) = 20 e-. 6. Place remaining e- around the three oxygens to satisfy octet.

36. 7. There are two e- left over so we will not have multiple bonds. 8. Place the last two e- on the central atom. *For polyatomic ions: place structure in brackets with the charge indicated on outside as demonstrated

37. 1- OCl O O

38. Resonance Structures Some molecules and ions cannot be represented by a single Lewis Structure. EX: Ozone O 3 O==O—O or O—O==O Draw a 2-way arrow to show resonance

39. Resonance of NO 2

40. Shapes of Molecules Molecules assume their shapes due to electron repulsion and the interplay between atomic orbitals. The first step in identifying the shape of a molecule is to draw its Lewis Dot Diagram.

41. EX: H 2 O H –O –H 2 pair of e- are involved in the bonding, these are shared pairs. 2 pair are not bonded, these are called unshared pairs. What type of shape does this molecule assume? Can it be predicted?

42. Electron Pair Repulsion Each bond and each unshared pair of e- form a charge cloud that repels other charge clouds. Electron pairs spread as far apart as possible to minimize repulsive forces.

43. VSEPR Theory Valence shell, electron-pair repulsion (VSEPR) Theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. Using VSEPR, we can predict the geometries of molecules.

44. VSEPR Theory to Predict Molecular Geometries Step 1: Write the Lewis Dot Structure for the molecule. Step 2: Represent the central atom in molecule by the letter A. Step 3:Represent the atoms bonded to it by the letter B. Step 4: Now refer to Table 6-5 VSEPR and Molecular Geometries

45. Lewis Dot Diagram of AlCl 3 EX: AlCl 3 1. Lewis Dot: Al has 3 valence and Cl has 7 x 3 for a total of 24 e-. 2. Place Al in center. 3. Place Cl around Al. 4. Draw bonds connecting Al and Cl. 5. Subtract bonds from total e- (24 – 6) = 18 e-

46. 6. Now put electrons around the most electronegative atoms (Cl) first, try to satisfy octet rule. 7. There are no remaining e-. Note: this molecule is an exception to the octet rule because in this case Al only forms three bonds.

47. Cl ClAl Cl

48. Predicting Molecular Geometry of AlCl 3 1. Refer to structure. 2. Al will be represented by letter A. 3. Cl will be represented by letter B. 4. We have AB 3. Refer to Table 6-5 and look for AB 3. 5. According to the Table, the molecular geometry for AlCl 3 is Trigonal-planar.

49. Trigonal-planar Geometry

50. Next Example: PCl 3 1. Lewis Dot Diagram: Cl Cl P Cl 2. P is represented by A. The unshared pair is represented by E. 3. Cl is represented by B. 4. We have AB 3 E. Refer to Table 6-5.

51. Trigonal-pyramidal Geometry

52. Molecular Geometry of H 2 O Remember, H 2 O has a Lewis Structure of H—O—H The O is represented by A and the two H are represented by B. There are two unshared pairs represented by E 2.

53. AB 2 E 2 is bent.

54. Covalent Bonds Covalent bonds are created from the sharing of e-. Since some atoms are more electronegative than others, this sharing is often unequal. This results in polarity for a molecule where there are moments when one atom is more negative than the others.

55. Polarity of water Take for example H 2 O. When we draw the structure it looks like: O H H The electronegativities of H and O are 2.20 and 3.50. O is the more electronegative and will pull the shared e- more of the time. We illustrate this by this symbol

56. The arrow is pointing to the most electronegative atom. Arrows pointing in opposite directions cancel each other out. Arrows pointing in the same direction indicate polarity.

57. Polar EX: O H H H 2 O is a polar molecule.

58. EX: H HC H H Carbon is more electronegative than Hydrogen, but this molecule is nonpolar because the charges are pulled in opposite directions and cancel. ( draw in arrows with pointer pen)

59. In Summary The tendency of a bonded atom to attract shared electrons to itself when bonded to another atom is called electronegativity (e-neg) Large differences in e-neg lead to the formation of ions, atoms that have gained or lost e-, and then ionic compounds, compounds bonded ionically (attraction of opposite charges).

60. In Summary Little or no difference in e-neg leads to covalent bonds (sharing of e-) Molecules are held together by covalent bonds. The structure of a molecule or a polyatomic ion can be represented by a Lewis Dot Diagram, which shows the pattern of shared and unshared pairs of e-.

61. In Summary The shape of a molecule can be predicted by taking into account the repulsion of e- pairs. The shape of a molecule containing three or more atoms is determined by the number and type of e- clouds in outer levels of the atoms.

62. In Summary Two atoms sometimes share more than one pair of e-, forming double or triple bonds. VSEPR is a method of using Lewis Dot Diagrams to predict molecular shapes.

63. In Summary Polarity of a molecule is based on e-neg of the bonded atoms. Polarity leads to momentary charge of a molecule. Arrows pointing to the most e-neg atom are used as shorthand. If these arrows are pointing in opposite directions, the molecule is nonpolar. If these arrows are pointing in the same direction, the molecule is polar.

64. A C6H12O6 production