1. BONDING 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: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
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: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
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: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
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. Octet Rule Atoms of METALS obey this rule by losing electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
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)

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

14. 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.

15. Ex. of Ionic Bonds: Determine charge & CRISS CROSS
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
DETERMINE THE CHARGE OF THE ION, CRISS CROSS CHARGES so the compound is NEUTRAL!!! (compounds DO NOT have charges)
Cation + always goes 1st, Anion – always goes last!!

16. Making Ionic Compounds
+ CATION always goes first in a compound!!!
- ANION always goes last in a compound!!!

17. Another Example: reducing to the simplest formula
What is the ionic compound formula for the calcium ion with the oxide ion?
Ca O

18. Ionic Compounds w/ polyatomic ions
DO NOT CHANGE THE FORMULA OF A POLYATOMIC ION!!!
Determine the formula for:
Li+ and PO43-
NH4+ and N3-

19. More practice
Determine the charge on an Aluminum ion, then pair it with the sulfite ion, SO32-

20. More practice…
Determine the charge on a Calcium ion, then pair it with the sulfate ion, SO42-

21. Covalent Bonds
The joy of sharing!

22. 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.

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

24. MOLECULES
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)

25. Octet Rule
Octet Rule: The representative elements achieve noble gas configurations (8 electrons) by sharing electrons.
THERE ARE A FEW EXCEPTIONS!!!
Hydrogen can only have 1 bond (2 electrons around it)

26. 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

27. 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

28. 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

29. 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

30. 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

31. 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

32. Writing Lewis Formulas (for molecules)
How to:
Add up all valence electrons for EACH atom in the molecule
Attach atoms with a single bond (skeleton drawing) (C = ALWAYS CENTRAL, H = ALWAYS ON OUTSIDE OF STRUCTURE)
Subtract out 2 electrons for each single bond you drew (EACH BOND = 2 electrons)
Distribute remaining electrons (in pairs) around atoms to obtain octet rule (except H)
If there’s not enough electrons to satisfy the octet rule, make MULTIPLE BONDS (double, triple)

33. Now let’s get more complex…

34. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

35. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

36. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

37. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

38. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

39. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

40. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

41. Lewis Structure Examples (remember your 5 steps):
CBr3-
N22-
CO32-
NH4+
CO2-
OH-
NO3-
SO42-

42. 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.

43. 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.

44. Electronegativity
Electronegativities are given numerical values (the most electronegative element has the highest value; the least electronegative element has the lowest value)
*** see Periodic Table, in lower part of element box
Most electronegative element:
Fluorine (4.0)
Least electronegative elements:
Fr (0.70), Cs (0.79)

45. 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.

46. 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

47. 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.

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

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

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

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

52. 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

53. 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.

54. 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.

55. 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.

56. RESONANCE STRUCTURES

57. 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.

58. 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.

59. 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.

60. 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

61. 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)

62. valence shell electron pair repulsion
VSEPR
valence shell electron pair repulsion

63. 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:

64. 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

65. 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.”

66. 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 (unshared electron pairs repel MORE than shared electron pairs in bonds)
Non-Bonding
Pairs
H—N — H —
Bonding
Pairs H

67. 5 Basic Molecule Shapes
1. tetrahedral
example: CH4

68. 5 Basic Molecule Shapes 2. 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)

69. 5 Basic Molecule Shapes 3. Bent or angular Example: H2O
Notice electron pair repulsion

70. 5 Basic Molecule Shapes
4. Linear
Example: CO2

71. 5 Basic Molecule Shapes 5. Trigonal planar or planar triangular
Example: BF3

72. CHEAT SHEET… Let’s make this easy to remember shapes
Linear= only 2 regions of SHARED electron pairs (no unshared electrons) around the CENTRAL atom
Bent = 2 SHARED electron pairs, 2 UNSHARED electron pairs around the CENTRAL atom
Tetrahedral = 4 SHARED electron pairs (meaning 4 bonds) around the CENTRAL atom
Trigonal Planar = 3 SHARED electron pairs (meaning 3 bonds) around the CENTRAL atom
Pyramidal = 3 SHARED electron pairs, 1 UNSHARED pair of electrons around the CENTRAL atom

73. Geometry and polarity Three shapes will cancel out polarity.
Linear = CO2 = NONPOLAR

74. Geometry and polarity 120º Three shapes will cancel out polarity.
Trigonal Planar = BF3 = NONPOLAR
120º

75. Geometry and polarity Three shapes will cancel out polarity.
Tetrahedral = CH4
= NONPOLAR

76. Geometry and polarity
Others don’t cancel
Bent = H2O = POLAR

77. Geometry and polarity
Others don’t cancel
Pyramidal = NH3 = POLAR