1. BONDING

2. VALENCE ELECTRONS
All atoms want to have their outer energy level complete. They will achieve this by losing or gaining electrons.
The electrons found in the outermost energy level are called valence electrons.

3. VALENCE ELECTRONS
When valence electrons are lost, the valence number is assigned a + sign.
If valence electrons are gained, then the valence number is assigned a – sign.

4. LEWIS DOT STRUCTURES
Valence electrons can be represented using a Lewis dot structure.
A Lewis diagram shows the atomic symbol to represent the nucleus and the core electrons. The dots represent the valence electrons.

5. LEWIS DOT STRUCTURES

6. Structures Classification
Atoms are classified as metals if they tend to lose electrons, as nonmetals if they tend to gain electrons, and as noble gases if they tend neither to gain nor lose electrons and have a complete energy level.
Reactivity
The fewer electrons an atoms tends to gain, lose, or share to fill its outer energy level the more reactive it tends to be in chemical reactions.

7. Quantum Numbers
Each electron orbital of the atom may be described by a set of four quantum numbers.
For a given orbital, n refers to size, l refers to shape; ml refers to orientation; and ms refers to spin.

8. Quantum Numbers Principal quantum number (n)
Refers to the average distance of the orbital from nucleus.
1 is closest and has the least energy.
The numbers correspond to the orbits or energy levels.

9. Quantum Numbers Angular Momentum quantum number (l)
Refers to the shape of the orbital.
The number of possible shapes is limited by the principal quantum number.
The first energy level (1), has only one possible shape, the s orbital.

10. Quantum Numbers Angular Momentum quantum number (l)
The second energy level has two possible orbital shapes, s and p.
Other orbital shapes described by the letters d and f become available when the n value is large enough to have l values of 2 and 3, respectively.

11. Quantum Numbers Magnetic Quantum Number (ml)
The number of spatial orientations possible for a given orbital is related back to the number of ml values.

12. Quantum Numbers Spin Quantum Number (ms)
It describes the spin in either of two possible directions: +1/2 or /2.
Each orbital can be filled by only two electrons, each with an opposite spin.

13. Quantum Numbers Spin Quantum Number (ms)
It describes the spin in either of two possible directions: +1/2 or /2.
Each orbital can be filled by only two electrons, each with an opposite spin.

14. Quantum Numbers Spin Quantum Number (ms)
The main significance of electron spin is explain by the Pauli Exclusion Principle, which states that in a given atom no two electrons can have the same set of four quantum numbers.

15. Aufbau Principle
When filling the orbitals of multielectronic atoms, ground-state configurations are realized when the lowest energy orbital available is filled first.

16. Hund’s Rule of Maximum Multiplicity.
When there’s more than one orbital to be filled at a particular level, only one electron will fill each orbital until each has one electron, then pairing will occur with the addition of one more electron to each orbital.

17. Electron Configuration
Remember that:
The s sublevel has one orbital and fits 2 electrons.
The p sublevel has three orbitals and fits 6 electrons.
The d sublevel has five orbitals and fits 10 electrons.
The f sublevel has seven orbitals and fits 14 electrons.

18. Electronegativity
The electronegativity of an element is a number that measures the relative strength with which the atoms of the element attract valence electrons in a chemical bond.
Electronegativity decreases down a group, and increases across a period.

19. Ionization Energy
The amount of energy needed to remove the first electron is called the first ionization energy.
Once the first electron is gone, the removal of other electrons becomes more difficult because of the decrease of repulsion between electrons.
The least electronegative element is francium, the most is flourine.

20. Today’s Activity Get into groups of three.
Grab a book, turn to chapter 8 (Chemical Bonds).
You are to come up with a mini lesson for your classmates on your assigned type of bond.