1. Chemical Bonding
Chapter 12

2. 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)

4. Types of Bonds 2. Covalent bond
Formed from the sharing of e- pairs between two or more nonmetals resulting in a molecule.
EX: H2 + O  H2O

5. 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.

6. 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. (atomic radius)

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
|1.31 – 3.44| = difference
What type of bond?

10. That’s right!
IONIC

11. Electrons are shared and make covalent bonds when the electronegativity difference between atoms is low, 1.69 and below.
EX: Hydrogen and Oxygen
|2.20 – 3.44| = difference
What type of Bond?
Covalent!

12. Ionic Bond Character
Ionic Character refers to the amount of time that a bond exists in the ionic form. Bonds fluctuate between existing in ionic forms and in covalent forms.
The relationship between the difference in the electronegativity values of two bonded atoms and the percent of ionic character can be expressed graphically.

13. Bond Character
Calculate e-neg differences and label as ionic or covalent.
B – P |2.04 – 2.19| =
0.15 COVALENT
Mg – N |1.31 – 3.04| =
1.73 IONIC
C – Na |2.55 –.93| =
1.62 COVALENT

14. 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.

15. 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.

16. Lewis Dot Structures
Step 3: Each “side” of symbol represents an orbital. Draw dots on sides as you would fill orbitals. Put two on the bottom side for “s” orbital and then one at a time in each of the three “p” orbitals, then pair electrons. Start with the bottom orbital and work clockwise when filling.

17. Lewis Dot Structures EX: carbon step 1: C
step 2: [He]2s22p2 = 4 valence e-
step 3: C

18. Lewis Dot Structures EX: bromine step 1: Br
step 2: [Ar]4s23d104p5 = 7 valence e-
step 3: Br
Do page from red chem

19. 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.

20. Shorthand Lewis Dot Structures
Group 1 has 1 valence, Group 2 has 2, Group 13 has 3, and so on… 8 1 3 4 5 6 7 2
Do page from red chem

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

22. Drawing Simple Ionic Bonds
Now draw them together with charge and then as neutral compounds (metals always written first).
Na+1 Cl-1 and NaCl
Ionic bonds form because of the attraction of the opposite charges.
Do page from red chem

24. Continue process until metals have lost all valence electrons and all the nonmetals have 8 valence electrons.

25. Lewis Dot Diagrams of Molecules
*1. Count the total number of valence e-. (math column)
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 creating a skeleton.

26. 4. Connect all bonded atoms in the skeleton with one bond.
*5. Subtract the number of electrons already used for the single bonds; two for each bond. (math column)
*6. Distribute the remaining electrons in pairs around the atoms, trying to satisfy the octet rule. Assign them to the most electronegative atom first. (Subtract e- used from math column)

27. 7. If you run out of electrons before all
7. If you run out of electrons before all atoms have an octet of electrons, you need to form double or triple bonds. *Zero e- on math.
8. If you have extra electrons (*e- left over in math column) and all of the atoms have an octet, put the extra electrons on the central atom in pairs.
9. If the central atom has an atomic number greater than fifteen, you are allowed to have more than eight electrons around it.

28. 10. Once math column is zero and all available electrons have been distributed, check each atom to ensure that all have 8 valence electrons. If not, make multiple bonds. *Must have math column for credit*

29. Exceptions to octet:
Hydrogen can only have 2 shared electrons (one bond, never multiple bonds)
Aluminum and Boron are satisfied with 6 shared electrons (3 bonds)
Beryllium is satisfied with 4 shared electrons (2 bonds)

30. Lewis Dot Diagrams EX: AsI3 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

31. 5. Subtract # used for bonds (26 – 6) = 20 e-.
6. Place e- around the three I atoms first because they are the most electronegative.
I –As –I I
7. We did not run out of e- so we do not have any double or triple bonds.

32. 8. We have 2 extra e- (20 starting – (3x6) = 2
8. We have 2 extra e- (20 starting – (3x6) = 2. Place them around the central atom.
I –As –I I
5+(3x7) = 26 e-
26 e-
6 e-
20 e-
-18 e-
2 e-
- 2 e-

33. Lewis Dot Diagram of CH2O
1. Count total valence e-
(C=4, H=1x2, O=6)= 12 valence e-
2. Place C in the center (only one central atom and C is least e- neg).
3. Place the 2 hydrogens and one oxygen around C.
4. Draw lines connecting H and O to central C.

34. 5. Subtract number of bonded e- from total. (12-6) = 6
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-.

35. Lewis Dot Diagram of CH2O
H C O H
4 + (2x1) + 6 = 12 e-
12 e-
6 e-
Do red chem

36. Lewis Dot Diagram of Polyatomic Ions
Polyatomic ions are groups of covalently bonded atoms (two or more nonmetals) 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.

37. Lewis Dot Diagram of ClO31-
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.

38. 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

39. 1-
O Cl O O
7 + (3x6) + 1 = 26 e-
26 e-
6 e-
20 e-
-18 e-
2 e-

40. Resonance Structures
Some molecules and ions cannot be represented by a single Lewis Structure.
EX: Ozone O3
O==O—O or O—O==O
Draw a 2-way arrow to show resonance

41. Resonance of NO2
All possible resonance structures with double arrow are required for full credit when any molecule demonstrates resonance.
Not many molecules demonstrate resonance.

42. 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.

43. EX: H2O
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?

44. 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.

45. 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.

46. 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 any atoms bonded to it by the letter B and any unshared pair on central atom by a single E.
Step 4: Now refer to Table 6-5 VSEPR and Molecular Geometries

47. Lewis Dot Diagram of BCl3
EX: BCl3
1. Lewis Dot: B has 3 valence and Cl has (7 x 3) for a total of 24 e-.
2. Place B in center.
3. Place Cl around Al.
4. Draw bonds connecting B and Cl.
5. Subtract bonds from total e-
(24 – 6) = 18 e-

48. 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 B only forms three bonds.

49. Cl
Cl B Cl

50. Predicting Molecular Geometry of BCl3
1. Refer to structure.
2. B will be represented by letter A.
3. Cl will be represented by letter B.
4. We have AB3. Refer to Table 6-5 and look for AB3. No E because B does not have any unshared electron pairs.
5. According to the Table, the molecular geometry for BCl3 is Trigonal-planar.

51. Trigonal-planar Geometry

52. Next Example: PCl3 1. Lewis Dot Diagram: Cl Cl P Cl
2. P is represented by A. The unshared pair on central P is represented by E.
3. Cl is represented by B.
4. We have AB3E. Refer to Table 6-5.

53. Trigonal-pyramidal Geometry

54. Molecular Geometry of H2O
Remember, H2O 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 on central oxygen represented by E2.

55. AB2E2 is bent.
Vsepr red chem

57. Covalent Bonds Covalent bonds are created from the sharing of e-.
Covalent bonds have low % ionic bond character.
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 atoms are partially negative and the others are partially positive.

58. Polarity of water
Take for example H2O. When we draw the molecular geometry it looks like: O
H H
The electronegativities of H and O are 2.20 and O is the more electronegative and will pull the shared e- more of the time. We illustrate this by this symbol

59. The arrow is pointing from the least electronegative atoms to the most electronegative atoms.
The arrows show movement of electrons. O
H H

60. When electrons are being pulled so that the forces of the pulls are not even, the molecules have dipole moments.
These dipoles give a partial positive and negative charge to atoms in a polar molecule.

61. Polar
EX: O
H H
The oxygen has a partial negative charge and the hydrogens have a partial positive charge.
So……H2O is a polar molecule.

62. EX: H
H C H H
Carbon is more electronegative than Hydrogen, but this molecule is nonpolar because the charges are pulled in opposite directions equally and that causes the forces of electron pull to cancel.
(draw in arrows with pointer pen)

63. General guidelines to determine polarity:
If the central atom has 1 or more lone pair of electrons, the molecule is polar.
If there are different types of elements bonded to a central atom, the molecule is polar. EX: CH3Cl.
If two different atoms are bonded together, the molecule is polar. EX: HF

64. 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).

65. 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-.

66. 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.

67. 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.

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

69. A C6H12O6 production