1. Ionic, Metallic and Covalent Bonding
Chemical Bonding
Ionic, Metallic and Covalent Bonding

2. Chemical Bonding
Atoms in compounds are held together by chemical bonds.
Chemical bonds result from the sharing or transfer of electrons between pairs of atoms.

3. Valence electrons
All elements in the same group behave similarly because they have the same number of valence electrons.
Valence electrons are the electrons in the highest occupied energy level of an element’s atoms.

4. Remember….
The number of valence electrons in a representative element is the same as the group number.
Examples:
Na has 1 valence electron,
Mg has 2,
Al has 3,
Si has 4,
P has 5, etc.

5. III. Covalent Bonding
In covalent bonding, the bonded atoms share electrons.
Molecules form as atoms share electrons in covalent bonds.

6. Describing electron pairs in electron dot diagrams
In Lewis dot structures, “dot-dot” pairs are called unshared pairs or lone pairs.
“Dashes” are called shared pairs or bonding pairs.

7. Lewis dot structures
Only the valence electrons are shown in electron dot structures.

8. S [ S -2] K [K+1] Practice Problems:
3) a) Draw the electron dot notation for a potassium atom.
b) Draw the electron dot notation for a potassium ion.
(4) a) Draw the electron dot notation for a sulfur atom.
b) Draw the electron dot notation for a sulfur ion. K
[K+1] S
[ S -2]

9. The Octet Rule
In forming compounds, atoms react to gain the electron configuration of a noble gas.
An octet is 8. Remember, most noble gases have 8 valence electrons (except helium).

10. Electron configurations for cations
cation (positive ion) forms when an atom loses electrons.
Na+ and achieves the same electron structure as neon.

11. “Isoelectronic” Na+ ion is isoelectronic with a neon atom.
Na (1s2 2s2 2p6 3s1) 
Na+ (1s2 2s2 2p6) + 1e-
Ne (1s2 2s2 2p6)

12. Electron configurations for anions
An anion (negative ion) forms when an atom gains electrons.
Cl- has the same electron structure as argon.
The Cl- ion is isoelectronic with the argon atom.

13. I. Ionic Bonding
Ionic bonds form from the transfer of electrons between atoms.
Metals give electrons to nonmetals. The metal forms a cation (+) and the nonmetal forms an anion (-).

14. Ionic Bonding
Ionic bonds are the forces of attraction between positive and negative ions in ionic compounds.
Remember: opposite charges attract.
The total negative charges equal the total positive charges, so ionic compounds are neutral.

15. How to Represent an Ionic Bond
Electron Dot Notations:
Na Cl  [ ] [ ]
Practice Problems:
(1) Draw the electron dot notation for the formation of an ionic compound between aluminum and bromine.
(2) Draw the electron configuration notation for the formation of an ionic compound between magnesium and nitrogen.
Na +1
Cl -1

17. Properties of Ionic Compounds
Ionic compounds form between metals and nonmetals.
They have high melting points.
They are generally soluble in water.

18. Properties of Ionic Compounds
Ionic compounds are good conductors of electricity when melted or dissolved in water.
For a compound to conduct electricity, it must have charged particles (ions) that are free to move (in liquid state or solution).

19. Properties of Ionic Compounds
5. Ionic compounds are crystalline solids at room temperature.
The ions in this beautiful CuSO4 crystal are arranged in a repeating, three-dimensional pattern.

20. Ionic crystals
The coordination number of an ion is the number of opposite charged ions that surround the ion.
In NaCl, each Na+ ion is surrounded by 6 Cl- ions.

21. Crystalline Patterns

22. II. Bonding in Metals
Metals are made of closely packed cations, instead of neutral atoms.
Because of a metal’s low ionization energy, the valence electrons become mobile, and drift freely.

23. Bonding in metals
Metallic bonding is the attraction of the free-floating valence electrons for the positively charged metal ions. This force of attraction holds metals together.

24. The mobile valence electrons of metals explain why…
Metals are malleable (can be pounded into sheets) and ductile (can be pulled into wires).
The metal cations can slip past each other, being separated by the sea of free floating electrons.

25. What happens when you pound an ionic crystal vs
What happens when you pound an ionic crystal vs. a metal crystal with a hammer?

26. The mobile valence electrons of metals explain why…
Metals are good conductors of electricity.
To be a conductor of electricity, charged particles must be free to move….
the electrons can flow freely within the metal.

27. Metal Alloys
Two metals can be mixed together to form alloys.

28. Remember: The octet rule applies to covalent bonding, too!

29. Draw an F2 molecule.

30. Draw an F2 molecule.
A single covalent bond forms when 2 atoms share a pair of electrons.

31. Draw an O2 molecule.

32. Draw an O2 molecule.
A double covalent bond forms when two atoms share two pairs of electrons

33. Draw an N2 molecule.

34. Draw an N2 molecule.
A triple covalent bond forms when two atoms share three pairs of electrons.

35. Sigma and pi bonds
Sigma bonds are bonds that lie directly on the bond axis (from one atom’s center to the other atom’s center).
Pi bonds do not lie on the bond axis.
A single covalent bond is 1 sigma bond.
A double covalent bond is 1 sigma + 1 pi bond.
A triple covalent bond is 1 sigma + 2 pi bonds.

36. Central vs. terminal atoms
The central atom in a molecule is the atom with the most metallic character (least electronegative).
Hydrogen cannot be a central atom because it has only one electron to use to form bonds. It must be a terminal atom.

37. Draw: Methane, CH4 Ammonia, NH3 Water, H20 Carbon dioxide, CO2
Ethylene, C2H4
Acetylene, C2H2

38. Draw carbon monoxide.

39. Draw carbon monoxide.
A coordinate covalent bond forms when one atom contributes both bonding electrons in a covalent bond.

40. Draw ozone.

41. Draw ozone.
Resonance occurs when 2 or more equally valid Lewis dot structures can be drawn for a molecule.

42. Bond Properties Bond order
Bond order = 1 when there is a single bond (1 sigma bond)
Bond order = 2 when there is a double bond (1 sigma + 1 pi)
Bond order = 3 when there is a triple bond (1 sigma + 2 pi)

43. Bond Properties Single bonds are long and weak.
2. Bond length
Single bonds are long and weak.
Double bonds are shorter and stronger.
Triple bonds are shortest and strongest.

44. Bond Properties Bond energy increases as bond order increases.
Atoms are held more tightly when there are multiple bonds.

45. Bond Properties
3. Bond polarity
When the two atoms involved in a bond have the same value of electronegativity, the electron pair is shared equally, and the bond is described as a nonpolar bond.

46. Bond Properties
3. Bond polarity (continued)
When the two atoms in a bond have different electronegativity values, the pair is not shared equally, and it is described as a polar covalent bond.
The electron pair is attracted toward the more electronegative atom, giving it a slightly negative charge.

50. Exceptions to the octet rule:
Draw BF3.
Exception: Fewer than 8 valence electrons
Draw NO2.
Exception: Odd number of electrons
Draw PF5 and SF6.
Exception: Expanded valence

51. VSEPR Theory
Valence Shell Electron Pair Repulsion Theory allows us to predict the molecular structures of compounds.
Because pairs repel each other, they move as far apart as possible, yet they are “tied” to the central atom. The pairs orient themselves to make the angles between them as large as possible.

52. VSEPR Theory
Structural Pairs (electron domains) consist of sigma bonds and lone pair electrons.
Pi bonds lie beside sigma bonds and do not take up any space.

53. VSEPR Theory 2 structural pairs = linear geometry
3 structural pairs = trigonal planar geometry
4 structural pairs = tetrahedral geometry
5 and 6 structural pairs lead to expanded valence (trigonal bipyramidal for 5 and octahedral for 6)

56. Molecular polarity
A molecule is described as polar if it meets two criteria:
1) it is made of polar bonds, and
2) it is unsymmetrical.
A molecule can be made of polar bonds yet still be a nonpolar molecule because of symmetry (CO2, CH4, etc.)