1. Valence Electrons, Lewis Dot Structures, and Electronegativity
Ionic Bonding
Valence Electrons,
Lewis Dot Structures,
and
Electronegativity

4. Valence Electrons
valence electrons – the outermost electrons.

5. 6 protons = C = carbon
Outer electrons = 4 e–available for bonding
Inner electrons = 2 e– not available for bonding

6. 4 electrons in valence shell

7. Lewis Dot Structures
Lewis dot structures are a convenient way to show how many valence electrons an atom has.
Example: Draw the Lewis dot structure for hydrogen. H

8. Ne He Lewis Dot Structures dots = number of valence electrons.
The maximum number of dots is 8.
Look for the number at the top of the column (e.g. 5A).
Exception: Helium only has 2 dots. Ne He

9. More Lewis Structure Practice
Draw the Lewis structure for oxygen.
Draw the Lewis structure for magnesium.
Draw the Lewis structure for chlorine. O Mg Cl

10. Even More Lewis Structure Practice
Draw the Lewis structure for carbon.
Draw the Lewis structure for potassium.
Draw the Lewis structure for phosphorus. C K P

11. Electronegativity
electronegativity – how much an atom wants to keep hold of its electrons.
ionization energy – the energy required to remove an electron from an atom.

13. Lower
electronegativity
Greater
electronegativity

14. Metals

15. Nonmetals

16. Role Models: The Noble Gases
An atom’s electrons are at their most stable when they reorganize their electrons to more closely resemble the electron configuration of a noble gas.
All atoms want to have stable electron configurations.

17. Role Models: The Noble Gases
All atoms wish their electrons were like the noble gases’ electrons.

18. Example: Beryllium and Oxygen

19. The 2 valence e–
Now beryllium’s matches that of helium
4 protons =
Be = beryllium

20. Now oxygen’s matches that of neon
The 6 valence e–
8 protons =
O = oxygen

21. Same Thing, but in Lewis Dot StructureBe

23. Semi-metals
or metalloids
Nonmetals
Metals

24. Three General Bonding Types
Metal with Nonmetal
- form ionic compounds
Metal with Metal
- form metallic compounds
Nonmetal with Nonmetal
- form covalent compounds

25. Metal with Nonmetal Bonding
Ionic compounds – the metal gives all of its valence e– to the nonmetal.
Known as – Salts, ions

27. Ions
ion – an atom that gained or lost electrons to become more like a noble gas.
+ Metals lose electrons to become positively charged ions. We call them cations (cat-ions) the “t” looks like a “+”. [e.g. 2A  +2]
– Nonmetals gain electrons to become negatively charged ions. We call them anions (an-ions) “n” for negative “–”.
[e.g. (8 – 6A) × -1 -2]

28. Writing the Charges Na+ Cl– Mg2+ O2–
Write out the ion that sodium forms.
Na+
Write out the ion that chlorine forms.
Cl–
Write out the ion that magnesium forms.
Mg2+
Write out the ion that oxygen forms.
O2–

29. So where do the electrons go?
Usually atoms that become cations give their electrons to anions.
Now both the cations and anions resemble noble gases, however now both have net charges. 2– 2+ O Be

30. Basic Electrical Charge Laws
+ and – : Attract
(pull together)

31. Naming (aka nomenclature)
Metals keep their names unchanged. (e.g. sodium, aluminum, calcium)
Transition metals have their charge shown as roman numerals in parenthesis after the name.
Fe2+  iron (II) Cu1+  copper (I) Fe3+  iron (III) Cu2+  copper (II)
Nonmetals have the last one or two syllables of their names altered with an –ide ending.
fluorine  fluoride nitrogen  nitride chlorine  chloride oxygen  oxide

32. Naming (aka nomenclature)
Metals keep their names unchanged. (e.g. sodium, aluminum, calcium)
If there are more than one possible charge for a metal (the transition metals), the charge will be specified in roman numerals after the name Fe2+  iron (II) Cu1+  copper (I) Fe3+  iron (III) Cu2+  copper (II)

33. Naming Continued
Nonmetals have the last one or two syllables of their names altered with an –ide ending.
Examples: carbon  carbide fluorine  fluoride nitrogen  nitride chlorine  chloride oxygen  oxide bromine  bromide sulfur  sulfide iodine  iodide phosphorus  phosphide

34. Naming Continued
Now put the metal and nonmetal ion names together and you get the name for the ionic compound.
Examples: LiF  lithium fluoride NaCl  sodium chloride KBr  potassium bromide MgS  magnesium sulfide CuI  copper (I) iodide CuO  copper (II) oxide FeN  iron (III) nitride

35. + 6 – 6 Sodium Chloride – NaCl Na+ Cl– Na+ Cl– Cl– Na+ Cl– Na+ Na+ Cl–
Na+
Cl–
Cl–
Na+

36. + 12 – 12 – 6 + 6 Magnesium Chloride Mg2+ Cl– Mg2+ Cl– Cl– Cl– Cl– Cl–
– 12
– 6
Cl–
Mg2+
Mg2+
Cl–
+ 6
Cl–
Cl–
Cl–
Cl–
Mg2+
Mg2+
Cl–

37. Magnesium Chloride
Mg2+
6 x
Cl–
12 x
Mg6Cl12
6 1
12 2 =
MgCl2

38. MgCl2 Magnesium Chloride Empirical Formula
Formula Unit (f.u.) – the smallest amount of an ionic compound that still has the same ratio of ions as in the formula.

39. Fe3+ O2– 2 x (+3) = +6 3 x (–2) = –6 criss-cross Fe2O3
Iron (III) Oxide
Fe3+ O2–
2 x (+3) = +6
3 x (–2) = –6
criss-cross
Fe2O3

40. Sodium Cloride
Na+ Cl–
criss-cross
Na1Cl1
NaCl

41. Magnesium Sulfide
Mg2+ S2–
criss-cross
2 1 =
Mg2S2
MgS

42. Sodium Oxide
Na+ O2–
criss-cross
Na2O

43. Fe3+ O2– Fe2O3 REMEMBER Top right corner: Charge Bottom right corner:
How many atoms/ions

44. Polyatomic Ions Poly – many Atomic – having to do with atoms
Polyatomic ions – ions made from multiple atoms
List on p.257

45. Polyatomic Ions (List on p.257)
Ion name Formula Acetate CH3COO– Ammonium NH4+ Carbonate CO32– Chromate CrO42– Cyanide CN– Dichromate Cr2O72– Hydroxide OH–
Ion name Formula
Nitrate NO3–
Nitrite NO2–
Permanganate MnO4–
Peroxide O22–
Phosphate PO43–
Sulfate SO42–
Sulfite SO32–
Thiosulfate S2O32–

46. Magnesium Nitrate
Mg2+ NO3–
criss-cross
Mg (NO3)2
1 x Mg
2 x N 6 x O

47. Mg2+ NO3– Mg (NO3)2 REMEMBER Top right corner: Charge
Bottom right corner:
How many atoms/ions

48. Sodium Chloride – NaCl
Na+
Cl–
Na+
Cl–
Cl–
Na+
Cl–
Na+
Na+
Cl–
Cl–
Na+

49. Sodium Chloride – NaCl

50. Sodium Chloride – NaCl

51. Sodium Chloride – NaCl

52. Sodium Chloride – NaCl

53. Sodium Chloride – NaCl

54. Magnesium Chloride (MgCl2)

55. Magnesium Chloride (MgCl2)

56. Magnesium Chloride (MgCl2)

57. Calcium Fluoride (CaF2)

58. Calcium Fluoride (CaF2)

59. Calcium Fluoride (CaF2)

60. Calcium Fluoride (CaF2)

61. Calcium Fluoride (CaF2)

62. Calcium Fluoride (CaF2)

64. He H B C N O Ne F Li Be P Al Si S Cl Ar Na Mg Br Kr K Ca I Xe

65. 4 e– in valence shell

66. Ions ion – an atom that gained or lost electrons.
metal ions lose electrons to become more positively charged. (e.g. 2A  +2)
Nonmetal ions gain electrons to become more negatively charged. (e.g. 8 – 6A –2)

67. Ions ion – an atom that gained or lost electrons.
metal ions lose electrons to become more positively charged. (e.g. 2A  +2)
Nonmetal ions gain electrons to become more negatively charged. (e.g. 8 – 6A –2)

68. + Metal ions = cations (cat-ions) the “t” looks like a “+”.
Ions - again
ion – an atom that gained or lost electrons to become more like a noble gas.
+ Metal ions = cations (cat-ions) the “t” looks like a “+”.
– Nonmetal ions = anions (an-ions). “–”