1. Electrons and Bonding

2. Valence Electrons The electrons that are located in the outer energy shell of each atom These electrons are available to be shared, lost, or gained in the formation of a chemical compound.

3. Electron Shells Remember as you move down a period on the periodic table you add an energy level (shell) to that atom. Each shell has a maximum number of electrons that it can hold before becoming full (and stable) All Noble Gases are stable, this means their outer energy levels are full. They do not want to share, gain or lose any of their electrons.

4. Heisenberg Uncertainty Principle This is the theory that states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. In other words, this principle is stating that it is impossible to know where the electron is exactly, how fast it is going nor the direction it is moving. Velocity is the speed and direction of the electron

5. Quantum Theory Schrodinger derived an equation that described energy & position of electrons in atom Schrodinger along with other scientists laid the foundation for the modern quantum theory, which describes mathematically the wave properties of electrons and other very small particles.

6. Orbitals and Quantum Numbers Electrons exist around the nucleus in certain regions called atomic orbitals - 3D regions around the nucleus that indicates the probable location of an electron. Quantum numbers are used to describe atomic orbitals. The quantum numbers are like an address for the electrons.

7. Orbitals There are 4 main types of orbitals s p d f Each one has: a certain number of electrons that they can hold A certain shape

8. s-orbital 1 s orbital for every energy level Spherical shaped can hold 2 electrons Called the 1s, 2s, 3s, etc.. orbitals.

9. P orbitals Start at second energy level 3 different directions 3 different shapes Each can hold 2 electrons

10. P orbitals

11. D orbitals Start at third energy level 5 different shapes Each can hold 2 electrons

12. F orbitals Start at fourth energy level 7 different shapes 2 electrons per shape

13. Summary of Orbitals Orbital S P D F # of Shapes Max electrons Starts at energy level 1 3 5 7 2 6 10 14 1 2 3 4

14. By energy level Orbitals do not fill up in a neat order. Energy levels overlap Lowest energy fill first.

15. Electron Configuration Way electrons are arranged in atoms. Aufbau principle- electrons occupy the lowest-energy orbital that can receive it. This means electrons enter the lowest energy first.

16. This causes difficulties because of overlap of orbitals of different energies. Pauli Exclusion Principle- states that no 2 electrons in the same atom can have the same set of four quantum numbers. Basically, at most there are 2 electrons per orbital with different spins.

17. Hund’s Rule- Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. This means when electrons occupy orbitals of equal energy they don’t pair up until they have to. (BUS RULE)

19. Start at the beginning of the periodic table and write down the energy level, subshell and number of electrons for each layer that you pass until you reach the element that you want. Example: Manganese would be 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Remember each s orbital can hold 2 electrons, p can hold 6, etc… Add up all of the subscripts to find the number of electrons. 2+2+6+2+6+2+5=25 2s22s2 Energy Level or shell Shape or subshell # of electrons

20. Manganese would be 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5

21. What is the electron configuration for bromine? 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4p 5 4s 2

22. Ionic Bonding Transfer of electrons Involves the attraction of oppositely charged particles or ions (cations &anions) Referred to as “salts” or ionic compounds Occur between a metal and a nonmetal Electronegativity difference of 1.6 or higher Individual piece called a formula unit

23. Characteristics/Properties of an Ionic Bond Called Salts Forms a crystalline solids at room temperature High melting and boiling points Generally soluble in water Brittle Conduct electricity when molten or dissolved in water Breaks into ions when dissolved Electrolytes

24. Covalent Bonds Sharing of electrons Occurs between two or more nonmetals Referred to as molecular compounds Individual piece called a molecule Two types Polar Unequal sharing Electronegativity difference of 0.2 to 1.6 Nonpolar Electronegativity difference of 0 to 0.2 Equal sharing

25. Characteristics/Properties of Covalent Bonds Found as amorphous solids, liquids & gases Low melting and boiling points Generally only slightly to insoluble in water Generally do NOT conduct electricity Do not break into pieces when dissolved Remain as molecules

26. Metallic Bonding Made of closely packed cations Electron Sea Model – electrons are free to move between the atoms – this explains the properties. Electrons can become delocalized Characteristics/Properties Good conductors of electricity & heat Malleable Ductile