1. Section 5.1—Types of Bonds

2. Why atoms bond
Atoms are most stable when the outer shell of electrons is full.
Atoms bond to fill this outer shell
The outer shell electrons of an atom are called the Valence Electrons. (outermost s or s and p)
Most atoms want 8 electrons in their valence.
We call this the Octet Rule
Common exceptions are Hydrogen and Helium which can only hold 2 electrons.

3. One way valence shells become full- - - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - -
Sodium has 1 electron in it’s valence shell
Chlorine has 7 electrons in it’s valence shell
Some atoms give electrons away to reveal a full level underneath.
Some atoms gain electrons to fill their current valence shell.

4. One way valence shells become full- - - + - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - -
The sodium now is a cation (positive charge) and the chlorine is now an anion (negative charge).
These opposite charges are now attracted, THIS FORMS AN IONIC BOND!

5. Ionic Bonding—Metal + Non-metal
Metals have low valence counts. For Example: Group 1 Metals have ONE valence electron.
Metals also have low ionization energies. This means it doesn’t take much energy to remove their valence electrons.
Therefore, metals tend to lose their electrons and non-metals gain electrons
Metals become cations (positively charged)
Non-metals become anions (negatively charged)
The cation & anion are attracted because of their charges—forming an ionic bond – ionic means formed from ions!!!

6. The Ionic Crystal – take note of how the ion charges force the atoms into a repetitive structure called a Crystal!

7. Covalent Bonding - between Non-metals
When two non-metals bond, they share electrons
Non-metals that share electrons evenly form non-polar covalent bonds
Non-metals that share electrons un-evenly form polar covalent bonds

10. Metallic Bonding – Metal to Metal Example: The Gold Atoms in a bar of Gold
The metals do not (as previously mentioned) hold onto their electrons well.
The valence electrons in metallic bonding are free to move throughout the structure - like a sea of electrons. This makes the atom appear to have a positive center.

12. Bond type affects properties
The type of bonding affects the properties of the substance.

13. Melting/Boiling Points
Ionic bonds tend to have very high melting/boiling points as it’s hard to pull apart those electrostatic attractions
They’re found as solids under normal conditions

14. Ionic vs. Covalent

15. Melting/Boiling Points
Polar covalent bonds have the next highest melting/boiling points
Most are solids or liquids under normal conditions

16. Melting/Boiling Points
Non-polar covalent bonds have lower melting/boiling points
Most are found as liquids or gases

18. Solubility in Water
Ionic & polar covalent compounds tend to be soluble in water

20. Solubility in Water
Non-polar & metallic compounds tend to be insoluble

22. How does soap work?

26. Conductivity of Electricity
In order to conduct electricity, charge must be able to move or flow
Metallic bonds have free-moving electrons—they can conduct electricity in solid and liquid state
Ionic bonds have free-floating ions when dissolved in water or when they are molten (liquid) form that allow them conduct electricity
Covalent bonds are NOT formed from charges and therefore cannot conduct electricity in any situation

27. Electrical Conductivity of solutions
No Light Dim Light Bright Light Bright Light
Question: Do all of these dissolve in water? If so, do all water soluble compounds conduct a current? Why or Why Not?

30. Distribution of electron density in H2.
Figure 9.13
Distribution of electron density in H2.
At some distance (bond length), attractions balance repulsions.
Electron density is high around and between the nuclei.

32. Figure 9.12
Covalent bond formation in H2.

33. Bonding Pairs and Lone Pairs
Atoms share electrons to achieve a full outer level of electrons. The shared electrons are called a shared pair or bonding pair.
The shared pair is represented as a pair of dots or a line: H •• or
H–H
An outer-level electron pair that is not involved in bonding is called a lone pair, or unshared pair. •• F
F–F or

34. Properties of a Covalent Bond
The bond order is the number of electron pairs being shared by a given pair of atoms.
A single bond consists of one bonding pair and has a bond order of 1.
The bond energy (BE) is the energy needed to overcome the attraction between the nuclei and the shared electrons. The stronger the bond the higher the bond energy.
The bond length is the distance between the nuclei of the bonded atoms.

35. Trends in bond order, energy, and length
For a given pair of atoms, a higher bond order results in a shorter bond length and higher bond energy.
For a given pair of atoms, a shorter bond is a stronger bond.
Bond length increases down a group in the periodic table and decreases across the period.
Bond energy shows the opposite trend.

36. Table 9.2 Average Bond Energies (kJ/mol) and Bond Lengths (pm)

37. Table 9.3 The Relation of Bond Order, Bond Length, and Bond Energy

38. Bond length and covalent radius.
Figure 9.14
Bond length and covalent radius.
Internuclear distance
(bond length)
Covalent radius
72 pm
Internuclear distance
(bond length)
Covalent radius
114 pm
Internuclear distance
(bond length)
Covalent radius
100 pm
Internuclear distance
(bond length)
Covalent radius
133 pm

40. Sample Problem 9.2
Comparing Bond Length and Bond Strength
PROBLEM:
Using the periodic table, but not Tables 9.2 or 9.3, rank the bonds in each set in order of decreasing bond length and decreasing bond strength:
(a) S–F, S–Br, S–Cl
(b) C=O, C–O, CΞO
PLAN:
S is singly bonded to three different halogen atoms, so the bond order is the same. Bond length increases and bond strength decreases as the atomic radius of the halogen increases.
The same two atoms are bonded in each case, but the bond orders differ. Bond strength increases and bond length decreases as bond order increases.

41. Sample Problem 9.2
SOLUTION:
(a) Atomic size increases going down a group, so F < Cl < Br.
Bond length: S–Br > S–Cl > S–F
Bond strength: S–F > S–Cl > S–Br
(b) By ranking the bond orders, we get
Bond length: C–O > C=O > CΞO
Bond strength: CΞO > C=O > C–O