1. Chapter 6.2 (Partial), 6.4 and 7.1 (Partial)
Covalent Compounds
Naming Covalent Compounds
Metallic Bonds

2. Covalent Bonds
Molecules are formed as elements make covalent compounds.
Covalent bonding decreases potential energy because each atom achieves an electron configuration like a noble gas.
Individual atomic orbitals of each element overlap where the covalent bond occurs.
This results in a concept called molecular orbitals.

3. Covalent Bonding

4. Molecules
Carbon Dioxide
CO2
Adenine

6. Covalent Bonds Potential energy determines the length of the bond.
At minimum potential energy, the distance between two bonded atoms is called bond length.
The bonded atoms vibrate about the bond, therefore, bond length is an average length.

7. Covalent Bonds
Polar molecules have partially positive and partially negative ends.
Such molecules are called “dipoles”.
The positive end is designated as δ+ and the negative end as δ-
example: δ+H--Fδ-

8. Naming Covalent Compounds
First name: name of first element in formula
usually least electronegative
requires a prefix if more than one of them
Second name: ends in –ide
the prefix “mono” is sometimes omitted completely, but we will use it on the second element in the compound.

9. Prefixes 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa
9 nona
10 deca

10. Naming Covalent Compounds

11. Metallic Bonding
Section 6.4

12. Properties of Substances with Different Types of Bonds

13. Introduction
Delocalized: refers to valence electrons that are free to move throughout the structure
In metals, valence electrons are delocalized and can move throughout the metal (shared by all atoms).
Metals are therefore composed of cations surrounded by delocalized electrons.

14. Continued Sometimes referred to as the electron-sea model
Metallic bonding is the electrostatic attraction between the metal ions and the delocalized electrons
Delocalized electrons = stabilized structure

15. Visual of the Electron Sea

16. Physical Properties of Metals
Ductile: can be drawn into a wire
Malleable: can be hammered into thin sheets
Both of these properties are attributed to the metal's ability to have layers pushed so they slide over one another without disrupting the metallic bonding
Some metals are better at this than others

17. Thermal and Electrical Conductivity
When a voltage is applied, the delocalized electrons are free to move
This also aids in conducting heat as the movement of the electrons carries kinetic energy

18. Comparison of Properties
Chapter 6

19. Introduction
The physical properties of a substance depend on the forces between the particles of the chemical species of which it is composed.
The stronger the bonding and the intermolecular forces, the harder the substances, and the higher the boiling point

20. Continued
The presence of impurities always lowers the melting point as the regular lattice is disrupted.
Volatility: a qualitative measure of how readily a liquid or solid is vaporized upon heating or evaporation

21. Metallic Structure Variable hardness, malleable rather than brittle
Melting points and boiling points vary, depending on the number of valence electrons, but is generally high.
Good electrical and thermal conductivity as solids and liquids
Generally insoluble, except in other metals to form alloys (mixture of metals).
Electron sea model: mass of positive nuclei in a sea of electrons

22. Ionic Structure Hard, but brittle High melting and boiling points
DO NOT conduct heat and electricity as solids, but DO when molten or in solution.
Often quite soluble in water compared to other solvents.
(+) + (-) [From PT, Metal + Nonmetal]

23. Covalent Structure Usually soft and malleable unless H-bonded
Low melting and boiling points. Liquids and gases at room temperature are usually molecular covalent.
Do not conduct electricity or heat in any state
More soluble in non-aqueous solvents, unless they can H-bond to water or react with it
Examples: CO2, CH3CH2OH, I2
Shared e-; [From PT: nonmetal + nonmetal]