1. Electrons
Chemistry: the study of matter and the changes that it undergoes.
If all atoms are composed of the same fundamental building blocks, then what explains differences in chemical behavior among elements?

2. Structure of the atom 3 subatomic particles. Atomic number (Z)
Number of electrons?
Isotopes
Atomic mass

3. What does an atom look like?

4. The best model – Electron Orbitals
Orbital: 3-D area in space where you are 90 % likely of finding an electron.
These areas are defined by a wavefunction equation.
2 electrons maximum can fit in an orbital.

5. These probable locations are called orbitals.
Electron energies and locations are best described with math formulas called wavefunctions.
When graphed, the solutions to these formulas describe “probable locations” of the electrons.
These probable locations are called orbitals.
sometimes called “electron clouds”

6. Orbital Shapes s sublevels have spherical orbitals.
p sublevels have dumb-bell orbitals.

7. More Shapes
d sublevels have clover-leaf orbitals.

8. Energy Levels, Sublevels and Orbitals
Each orbital is assigned to a level and a sublevel.
Orbitals in higher energy levels are larger in volume. 2s

9. Energy Levels, Sublevels and Orbitals
Energy order of first 12 sublevels:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s
The energy level is represented by the number. The larger the number, the further the orbital is from the nucleus. Electrons in high-level orbitals are not tightly held to the atom, and are easily lost to other atoms.
The letter represents the sublevel of the orbital, as well as the orbital shape.
s sublevels have 1 orbital, p sublevels have 3, and d have 5.

10. What does an electron configuration “look like?”
Sodium atoms have 11 electrons:
1s2 2s2 2p6 3s1
What would Aluminum look like? e- e- e- e- e- e- e- e- e- e- e-

11. Orbitals “stack together.”

12. Electron Configurations
No need to memorize the order of levels and sublevels! Use the PT!

13. Noble Gas Shortcut Ex: Magnesium: [Ne] 3s2 Bromine: [Ar] 4s2 3d10 4p5
Are all of an atom’s electrons involved in chemical bonding?
Electron configurations can be “shortened” to emphasize the valence electrons with the Noble Gas Shortcut.
Ex: Magnesium: [Ne] 3s2
Bromine: [Ar] 4s2 3d10 4p5

14. Orbital Diagrams 3 rules: Aufbau, Hund, Pauli
Each box represents an orbital.
Arrows represent electrons.
3 rules: Aufbau, Hund, Pauli

15. This idea leads us to our next topic, Ionic Bonding.
Predicting Charge
Octet rule: Atoms gain or lose electrons to achieve a filled outer
energy level of electrons.
This is most often 8 electrons in the outer energy level.
General rules: Atoms with 3 or fewer valence e- lose those electrons, forming a
positive ion, or cation.
Atoms with 5 or more valence e- gain electrons, forming a negative ion, or anion.
This idea leads us to our next topic, Ionic Bonding.

16. Cations
Positively charged ions
Result from loss of valence electrons.
Ex: Magnesium: [Ne] 3s2 will lose its 2 valence electrons, to result in Mg2+ ion.
Named _____________ ion.
insert element name

17. Anions
Negatively charged ions
Result from gain of valence electrons from another atom.
Ex: Nitrogen: 1s22s22p3 will gain 3 valence electrons from another atom, to result in N3- ion.
Named ____________ -ide
insert element name, dropping last syllable

18. Polyatomic ions
Primarily negatively charged ions, but there are some positive ones.
They are molecules that carry a charge.
The back of your PT lists the ones used in this course.
Their names end in –ate or –ite for the most part.

19. Variable-charge cations
Some metals can form more than one ion.
Most of these metals are in the d block, but some are low in the p-block.
Iron is a common example. It forms either a 2+ or a 3+ charge.
Roman numerals are used to distinguish between the ions.

20. Bohr Model Only specific energy level values are allowed.
Quantization of energy
Energy is gained and lost when transitions occur.
See it: Bohr/applet_files/Bohr.html