1. Chapter 11
Chemical Bonding

2. 11.1 Characteristics of Chemical Bonds
Chemical bond – the force that holds two or more atoms together and makes them function as a unit
Bond energy – energy required to break a given chemical bond

3. Types of Chemical Bonds
Ionic Bond – attraction between oppositely charged ions
Electrons lost/gained

4. Covalent Bond – type of bond in which atoms share electrons

5. Polar covalent bond – a covalent bond in which the electrons are not shared equally
More electronegative atom pulls electrons closer

6. Electronegativity
Electronegativity - The tendency of an atom in a molecule to attract shared electrons to itself
Increases across a period
Decreases down a group

8. Rank the following in order of increasing electronegativity (without looking at the values): Al, Cl, S, F, Na

9. Polarity of a bond depends on the difference between the electronegativity values of the atoms

10. Rank the following in order of increasing polarity: H-H, O-H, Cl-H, S-H, F-H

11. Bond Polarity and Dipole Moments
A dipole moment results when a polar molecule has a center of positive charge separate from a center of negative charge

12. Water has a dipole moment which is the reason you are sitting here RIGHT NOW!!!

13. Liquid at room temperature
“universal solvent”
Like dissolves like
Liquid at room temperature

14. 11.2 Characteristics of Ions and Ionic Compounds

15. Ions form to achieve a noble gas electron configuration
Octet rule – atoms will lose, gain, or share electrons to acquire a full set of 8 valence electrons
Metals
nonmetals

16. Neutral atom Na
1s22s22p63s1 B
1s22s22p1 P
1s22s22p63s23p3 F
1s22s22p5

17. Ion
Na1+
1s22s22p6 (Ne)
B3+
1s2 (He)
P3-
1s22s22p63s23p6 (Ar)
F1-

18. 1 = s1 (1 ve, easily lost, ion formed = )
13 = s2p1
14 = s2p2
15 = s2p3
16 = s2p4
17 = s2p5
18 = s2p6

19. Ionic Bonding and Structures of Ionic Compounds
LiF = formula unit
Actual compound contains a lot of ions packed together to maximize attraction
Every + ion surrounded by – ions
Every – ion surrounded by + ions

21. Ionic compounds containing polyatomic ions
Behave the same way as monatomic ions
Covalent bonds hold polyatomic ion together so it behaves as a unit

22. 11.3 Lewis Structures
Lewis structure – representation of a molecule or polyatomic ion showing how electrons are arranged among the atoms
Only include valence electrons

23. Bonding pair – a pair of electrons shared between 2 atoms
Lone pair – electron pairs in a Lewis structure that are not involved in bonding

24. If atoms need 1 electron, it will usually form 1 covalent bond.
H and Halogens typically only form one bond
If atoms need 2 electrons, it will usually form 2 covalent bonds.
If atoms needs 3 electrons, it will usually form 3 covalent bonds.

25. Steps for writing Lewis Structures:
Obtain the sum of the valence electrons from all of the atoms
Use one pair of electrons to form a bond between each pair of bound atoms
Arrange the remaining electrons to satisfy the octet rule
Exception:

26. Single Covalent Bonds
One pair of electrons is shared
Cl2 HF

27. Double Covalent Bond Two pairs of electrons are shared O2
Typically C, N, O, S O2

28. Triple covalent bond Three pairs of electrons are shared N2
Typically C or N N2

29. Write the Lewis structures for the following:
NH3
CO2

30. C2H4

31. Lewis Structures of Polyatomic Ions
Number of dots in the molecule must take the ions charge into account
NH41+

32. PO43-

33. Resonance Structures
A molecules shows resonance when more than one Lewis structure can be drawn for the molecule

34. O3
CO32-

35. Exceptions to the octet rule
Suboctets – stable configurations with fewer than 8 electrons around an atom
Boron typically forms a suboctet
BH3

36. Expanded octets – central atoms contain more than eight valence electrons
Phosphorous and sulfur can form expanded octets
PCl5

37. 12.4 Structures of Molecules
VSEPR model
V alence
S hell
E lectron
P air
R epulsion
The shape of a molecule is determined by minimizing repulsion between lone pairs

38. Bond angle
VSEPR forces cause atoms in a molecule to be positioned at fixed angles relative to one another

39. Linear – 180 degrees
Binary molecules or molecules with 2 atoms bonded to the central atom and no lone pairs on the central atom
CO2

40. Bent – degrees
Two atoms bonded to central atom, two lone pairs on central atom
H2O

41. Trigonal Planar – 120 degrees
Three atoms bonded to central atom, no lone pairs on central atom
BH3

42. Tetrahedral – 109.5 degrees Pyramid shaped
4 atoms bonded to central atom, no lone pair on central atom
CH4

43. Trigonal Pyramidal – 107.3 degrees
A triangle that is not flat
Three atoms bonded to central atom, one lone pair on central atom
NH3

44. Polar and Nonpolar Molecules
Polar bond = electrons shared unequally (dipole)
More electronegative atom pulls electrons closer to it and therefore has a slightly negative charge
Less electronegative atom has a slightly positive charge

45. Polar molecules always have:
At least one polar covalent bond
Asymmetric geometry
Lone pairs on the center atom
Different atoms bonded to the center atom

46. HCl
H2O

47. Not every molecule with polar bonds is polar!!!
CH4
CO2

48. Polar Bond vs. Polar Molecule
Polar bond = electrons shared unequally between molecules
Polar molecule = entire molecule has different partial charges on opposite sides of the molecule
Depends on shape

49. CCl4
NH3
H20
Linear vs. bent