1. Chapter 12
Chemical Bonding

2. Objectives
To learn about ionic and covalent bonds and explain how they are formed
To learn about the polar covalent bond
To understand the nature of bonds and their relationship to electronegativity
To understand bond polarity and how it is related to molecular polarity

3. A. Types of Chemical Bonds
Bond – force that holds groups of atoms together and makes them function as a unit
Bond energy – energy required to break a chemical bond

4. A. Types of Chemical Bonds
Ionic Bonding
Ionic compound results when a metal reacts with a nonmetal.

5. A. Types of Chemical Bonds
Covalent Bonding
A covalent bond results when electrons are shared by nuclei.

6. A. Types of Chemical Bonds
Covalent Bonding
A polar covalent bond results when electrons are shared unequally by nuclei.
One atom attracts the electrons more than the other atom.

7. B. Electronegativity
Electronegativity – the relative ability of an atom in a molecule to attract shared electrons to itself
Increases from left to right across a period
Decreases down a group of representative elements

8. Concept Check
If lithium and fluorine react, which has more attraction for an electron? Why?
In a bond between fluorine and iodine, which has more attraction for an electron? Why?
The fluorine has more attraction for an electron than does lithium. Both have valence electrons in the same principal energy level (the 2nd), but fluorine has a greater number of protons in the nucleus. Electrons are more attracted to a larger nucleus (if the principal energy level is the same).
The fluorine also has more attraction for an electron than does iodine. In this case the nuclear charge of iodine is greater, but the valence electrons are at a much higher principal energy level (and the inner electrons shield the outer electrons).

9. Concept Check
What is the general trend for electronegativity across rows and down columns on the periodic table?
Explain the trend.
Generally speaking, the electronegativity increases in going from left to right across a period because the number of protons increases which increases the effective nuclear charge. The electronegativity decreases going down a group because the size of the atoms increases as you go down so when other electrons approach the larger atoms, the effective nuclear charge is not as great due to shielding from the current electrons present.

10. B. Electronegativity
The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond.

11. B. Electronegativity

12. a) N–F O–F C–F b) C–F N–O Si–F c) Cl–Cl B–Cl S–Cl a) C–F, N–F, O–F
Exercise
Arrange the following bonds from most to least polar: 
a) N–F O–F C–F
b) C–F N–O Si–F
c) Cl–Cl B–Cl S–Cl
a) C–F, N–F, O–F
b) Si–F, C–F, N–O
c) B–Cl, S–Cl, Cl–Cl
The greater the electronegativity difference between the atoms, the more polar the bond.
C-F, N-F, O-F
Si-F, C-F, N-O
B-Cl, S-Cl, Cl-Cl

13. Concept Check
Which of the following bonds would be the least polar yet still be considered polar covalent?
Mg–O C–O O–O Si–O N–O
The correct answer is N-O. To be considered polar covalent, unequal sharing of electrons must still occur. Choose the bond with the least difference in electronegativity yet there is still some unequal sharing of electrons.

14. Concept Check
Which of the following bonds would be the most polar without being considered ionic?
Mg–O C–O O–O Si–O N–O
The correct answer is Si-O. To not be considered ionic, generally the bond needs to be between two nonmetals. The most polar bond between the nonmetals occurs with the bond that has the greatest difference in electronegativity.

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

16. C. Bond Polarity and Dipole Moments
Water molecule dipole moment

17. C. Bond Polarity and Dipole Moments
The polarity of water affects its properties.
Permits ionic compounds to dissolve in it
Causes water to remain liquid at higher temperature

18. Objectives
To learn about stable electron configurations
To learn to predict the formulas of ionic compounds
To learn about the structures of ionic compounds
To understand factors governing ionic size

19. A. Stable Electron Configurations and Charges on Ions

20. A. Stable Electron Configurations and Charges on Ions

21. A. Stable Electron Configurations and Charges on Ions
Atoms in stable compounds usually have a noble gas electron configuration.
Metals lose electrons to reach noble gas configuration.
Nonmetals gain electrons to reach noble gas configuration.

22. A. Stable Electron Configurations and Charges on Ions
Predicting Formulas of Ionic compounds
Chemical compounds are always electrically neutral.

23. B. Ionic Bonding and Structures of Ionic Compounds
Ions are packed together to maximize the attractions between ions.

24. B. Ionic bonding and Structures of Ionic Compounds
Cations are always smaller than the parent atom.
Anions are always larger than the parent atom.

25. B. Ionic bonding and Structures of Ionic Compounds
Ionic Compounds Containing Polyatomic Ions
Polyatomic ions work in the same way as simple ions.
The covalent bonds hold the polyatomic ion together so it behaves as a unit.

26. Objectives
To learn to write Lewis structures
To learn to write Lewis structures for molecules with multiple bonds

27. A. Writing Lewis Structures
In writing Lewis structures, we include only the valence electrons.
Most important requirement
Atoms achieve noble gas electron configuration (octet rule, duet rule).

28. A. Writing Lewis Structures
Bonding pairs are shared between 2 atoms.
Unshared pairs (lone pairs) are not shared and not involved in bonding.

29. A. Writing Lewis Structures

30. B. Lewis Structures of Molecules with Multiple Bonds
Single bond – covalent bond in which 1 pair of electrons is shared by 2 atoms
Double bond – covalent bond in which 2 pairs of electrons are shared by 2 atoms
Triple bond – covalent bond in which 3 pairs of electrons are shared by 2 atoms

31. B. Lewis Structures of Molecules with Multiple Bonds
A molecule shows resonance when more than one Lewis structure can be drawn for the molecule.

32. B. Lewis Structures of Molecules with Multiple Bonds
Some Exceptions to the Octet Rule
Boron – incomplete octet
Molecules containing odd numbers of electrons – NO and NO2

33. Draw a Lewis structure for each of the following molecules:
Concept Check
Draw a Lewis structure for each of the following molecules: H2 F2 HF

34. Draw a Lewis structure for each of the following molecules:
Concept Check
Draw a Lewis structure for each of the following molecules:
NH3
CO2
CCl4

35. Draw a Lewis structure for each of the following molecules:
Concept Check
Draw a Lewis structure for each of the following molecules:  CO
CH3OH
OCN–

36. Objectives
To understand molecular structure and bond angles
To learn to predict molecular geometry from the number of electron pairs
To learn to apply the VSEPR model to molecules with double bonds

37. A. Molecular Structure
Three dimensional arrangement of the atoms in a molecule
Water - bent

38. A. Molecular Structure
Linear structure – atoms in a line
Carbon dioxide

39. A. Molecular Structure
Trigonal planar – atoms in a triangle
BF3

40. A. Molecular Structure
Tetrahedral structure
methane

41. B. The VSEPR Model
Valence shell electron pair repulsion (VSEPR) model
Molecular structure is determined by minimizing repulsions between electron pairs.

42. B. The VSEPR Model
Two Pairs of Electrons
BeCl2
180o - linear

43. B. The VSEPR Model
Three Pairs of Electrons
BF3
120o – trigonal planar

44. B. The VSEPR Model
Four Pairs of Electrons
CH4
109.5o – tetrahedral

45. B. The VSEPR Model

46. B. The VSEPR Model

47. C. Molecules with Double Bonds
When using VSEPR model to predict molecular geometry of a molecule
a double bond is counted the same as a single electron pair

48. Concept Check
Determine the molecular structure for each of the following molecules, and include bond angles:
HCN
PH3
CCl4
HCN – linear, 180o
PH3 – trigonal pyramid, 109.5o (107o)
CCl4 – tetrahedral, 109.5o
HCN – linear, 180o
PH3 – trigonal pyramid, 107o
CCl4 – tetrahedral, 109.5o