1. BONDING Ch 8 & 9 – Honors Chemistry General Rule of Thumb:
metal + nonmetal = ionic
polyatomic ion + metal or polyatomic ion = ionic (both)
nonmetal + nonmetal(s) = covalent

2. Isn’t it ionic that opposites attract?
Ionic Bonds
Isn’t it ionic that opposites attract?

3. Valence Electrons
Knowing electron configurations is important because the number of valence electrons determines the chemical properties of an element.
Valence Electrons: The e- in the highest occupied energy level of an element’s atoms.

4. Valence Electrons
All elements in a particular group or family have the same number of valence electrons (and this number is equal to the group number of that element)
Examples:
Group 1 elements (Na, K, Li, H): 1 valence electron.
Group 2 elements (Mg, Ca, Be): 2 valence electrons.
Group 17 elements (Cl, F, Br): 7 valence electrons.

5. Lewis Structures
Electron dot structures show the valence electrons as dots around the element’s symbol: Li B Si N O F Ne

6. Lewis Structures
Electron dot structures show the valence electrons as dots around the element’s symbol: Li B Si N O F Ne

7. Octet Rule
Noble gas atoms are very stable; they have stable electron configurations. In forming compounds, atoms make adjustments to achieve the lowest possible (or most stable) energy.
Octet rule: atoms react by changing the number of electrons so as to acquire the stable electron structure of a noble gas.

8. Octet Rule Atoms of METALS obey this rule by losing electrons. Na:
Atoms of NONMETALS obey this rule by gaining electrons.
Cl:
Cl-:
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively stable configuration with its 4d sublevel filled (pseudo noble-gas)

9. Octet Rule Atoms of METALS obey this rule by losing electrons. Na:
Atoms of NONMETALS obey this rule by gaining electrons.
Cl:
Cl-:
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively stable configuration with its 4d sublevel filled (pseudo noble-gas)

10. Octet Rule Atoms of METALS obey this rule by losing electrons. Na:
Atoms of NONMETALS obey this rule by gaining electrons.
Cl:
Cl-:
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively stable configuration with its 4d sublevel filled (pseudo noble-gas)

11. Octet Rule Atoms of METALS obey this rule by losing electrons. Na:
Atoms of NONMETALS obey this rule by gaining electrons.
Cl:
Cl-:
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively stable configuration with its 4d sublevel filled (pseudo noble-gas)

12. Ionic Bonds
Anions and cations have opposite charges; they attract one another by electrostatic forces (IONIC BONDS)

13. Ionic Bonds
Ionic compounds are electrically neutral groups of ions joined together by electrostatic forces. (also known as salts)
the positive charges of the cations must equal the negative charges of the anions.
use electron dot structures to predict the ratios in which different cations and anions will combine.

14. Examples of Ionic Bonds
Na Cl
Al Br
K O
Mg N
K P
Na+Cl- = NaCl
Al3+Br- = AlBr3
K+O2- = K2O
Mg2+N3- = Mg3N2
K+P3- = K3P

15. Criss Cross Method
Criss-cross ionic charges down as subscripts (without +/-) and reduce (to determine lowest whole # ratio of cation:anion)
ANY TIME YOU ADD A SUBSCRIPT TO A POLYATOMIC ION, YOU MUST FIRST PUT THAT ION IN PARENTHESES.

16. Covalent Bonds
The joy of sharing!

17. Covalent Bonds
Covalent bonds: occur between two or more nonmetals; electrons are shared not transferred (as in ionic bonds)
The result of sharing electrons is that atoms attain a more stable electron configuration.

18. Covalent Bonds Most covalent bonds involve:
2 electrons (single covalent bond),
4 electrons (double covalent bond, or
6 electrons (triple covalent bond).

19. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO

20. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO

21. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO

22. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO

23. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO

24. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO

25. Octet Rule
Octet Rule: The representative elements achieve noble gas configurations (8 electrons) by sharing electrons.

26. Writing Lewis Structures
Select a skeleton for the molecule (the least electronegative element is usually the central element).
Calculate N (the # of valence e- need by all atoms in the molecule of polyatomic ion.
Calculate A (the # of electrons available).
Calculate S (the # of electrons shared in the molecule) S = N – A
Place the S electrons as shared pairs in the skeleton.
Place the additional electrons as unshared pairs to fill the octet of every representative elements (except hydrogen!).

27. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

28. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

29. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

30. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

31. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

32. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

33. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

34. Lewis Structure Examples:
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

35. Electronegativity
We’ve learned how valence electrons are shared to form covalent bonds between elements. So far, we have considered the electrons to be shared equally. However, in most cases, electrons are not shared equally because of a property called electronegativity.

36. Electronegativity
The ELECTRONEGATIVITY of an element is: the tendency for an atom to attract electrons to itself when it is chemically combined with another element.
The result: a “tug-of-war” between the nuclei of the atoms.

37. Electronegativity
Electronegativities are given numerical values (the most electronegative element has the highest value; the least electronegative element has the lowest value)
**See Figure 6-18 p. 169 (Honors)
Most electronegative element: Fluorine (3.98)
Least electronegative elements:
Fr (0.70), Cs (0.79)

38. Electronegativity Notice the periodic trend:
As we move from left to right across a row, electronegativity increases (metals have low values nonmetals have high values – excluding noble gases)
As we move down a column, electronegativity decreases.
The higher the electronegativity value, the greater the ability to attract electrons to itself.

39. Nonpolar Bonds
When the atoms in a molecule are the same, the bonding electrons are shared equally.
Result: a nonpolar covalent bond
Examples: O2, F2, H2, N2, Cl2

40. Polar Bonds
When 2 different atoms are joined by a covalent bond, and the bonding electrons are shared unequally, the bond is a polar covalent bond, or POLAR BOND.
The atom with the stronger electron attraction (the more electronegative element) acquires a slightly negative charge.
The less electronegative atom acquires a slightly positive charge.

41. Polar Bonds H Cl Example: HCl Electronegativities: - + H = 2.20

42. Polar Bonds
Example: H2O
Electronegativities:
H = 2.20
O = 3.44

43. Polar Bonds
Example: H2O
Electronegativities:
H = 2.20
O = 3.44

44. Polar Bonds
Example: H2O
Electronegativities:
H = 2.20
O = 3.44

45. Electronegativity Difference
Predicting Bond Types
Electronegativities help us predict the type of bond:
Electronegativity Difference
Type of Bond
Example
0.00 – 0.40
H-H
0.41 – 1.00
H-Cl
1.01 – 2.00
H-F
2.01 or higher
Na+Cl-
covalent
(nonpolar)
covalent
(slightly polar)
covalent
(very polar)
ionic

46. Polar Molecules
A polar bond in a molecule can make the entire molecule polar
A molecule that has 2 poles (charged regions), like H-Cl, is called a dipolar molecule, or dipole.

47. Polar Molecules
The effect of polar bonds on the polarity of a molecule depends on the shape of the molecule.
Example: CO2
O = C = O shape: linear
*The bond polarities cancel because they are in opposite directions; CO2 is a nonpolar molecule.

48. Polar Molecules
The effect of polar bonds on the polarity of a molecule depends on the shape of the molecule.
Water, H2O, also has 2 polar bonds:
But, the molecule is bent, so the bonds do not cancel.
H2O is a polar molecule.

49. Resonance
A molecule or polyatomic ion for which 2 or more dot formulas with the same arrangement of atoms can be drawn is said to exhibit RESONANCE.

50. Resonance Example CO32- 3 resonance structures can be drawn for CO32-
the relationship among them is indicated by the double arrow.
the true structure is an average of the 3.

51. Resonance Example CO32- 3 resonance structures can be drawn for CO32-
the relationship among them is indicated by the double arrow.
the true structure is an average of the 3.

52. Resonance Example CO32- 3 resonance structures can be drawn for CO32-
the relationship among them is indicated by the double arrow.
the true structure is an average of the 3

53. Resonance Structures
Another way to represent this is by delocalization of bonding electrons:
(the dashed lines indicate the 4 pairs of bonding electrons are equally distributed among 3 C-O bonds; unshared electron pairs are not shown)
See p. 256

54. valence shell electron pair repulsion
VSEPR
valence shell electron pair repulsion

55. Molecular Shape
Lewis structures (electron dot structures) show the structure of molecules…but only in 2 dimensions (flat).
BUT, molecules are 3 dimensional!
for example, CH4 is:

56. Molecular Shape
Lewis structures (electron dot structures) show the structure of molecules…but only in 2 dimensions (flat).
BUT, molecules are 3 dimensional!
but in 3D it is:
a tetrahedron!
= coming out of page
= going into page
= flat on page

57. Why do molecules take on 3D shapes instead of being flat?
Valence Shell Electron Pair Repulsion theory
“because electron pairs repel one another, molecules adjust their shapes so that the valence electron pairs are as far apart from another as possible.”

58. Why do molecules take on 3D shapes instead of being flat?
Valence Shell Electron Pair Repulsion theory
Remember: both shared and unshared electron pairs will repel one another.
Non-Bonding
Pairs
H—N — H —
Bonding
Pairs H

59. 5 Basic Molecule Shapes
Linear
Example: CO2

60. 5 Basic Molecule Shapes Bent or angular Example: H2O
Notice electron pair repulsion

61. 5 Basic Molecule Shapes
tetrahedral
example: CH4

62. 5 Basic Molecule Shapes Pyramidal Example: NH3
(note: unshared pair of electron repels, but is not considered part of overall shape; no atom there to contribute to the shape)

63. 5 Basic Molecule Shapes Trigonal planar or planar triangular
Example: BF3

64. Geometry and polarity Three shapes will cancel out polarity.
Shape One: Linear

65. Geometry and polarity 120º Three shapes will cancel out polarity.
Planar triangles
120º

66. Geometry and polarity Three shapes will cancel out polarity.
Tetrahedral

67. Geometry and polarity
Others don’t cancel
Bent

68. Geometry and polarity
Others don’t cancel
Trigonal Pyramidal