Focus on the electrons:
Bohr model (Bohr recieved the 1922 Nobel prize)
The idea of quantization and orbitals
Why did Bohr believe in quantization - it explains "line spectra"?
The energy of light is determined by its color (more blue more energy, more red less energy)
White light - continuous spectrum - all energies
Demo: elements, specific colors.
Specific colors are due to transitions between orbitals of specific energies.
Different elements are used to give fireworks their color!
To learn more about the science of fireworks - see: http://www.pbs.org/wgbh/nova/kaboom/
Modern models - idea of orbitals and energy quantization still their but the orbitals are not the planetary type orbitals of Bohr
1. What do the orbitals look like and how do we label them?
Rules of atomic orbitals - developing the blueprint:
1. Electrons occupy orbitals of discrete energy. Orbitals belong to a primary energy level (coarse measure of energy) and an energy sub-level (finer measure of energy).
2. Principle energy level - n - determines the size of the orbital (in the Bohr model, larger n larger radius of orbit).
3. Each primary energy level can be occupied by up to 2n2
4. Each primary energy levels contains a distinct of set orbitals that belong to various sub-levels (types of orbitals).
5. The types or sublevels are labeled as follows - show pictures. Each sublevel contains a certain number of orbitals.
s - spherical in shape (1 orbital)
6. Each primary energy level has n types of orbitals (s,p,d or f) and a total of n2 orbitals.
Some movies of orbitals from: http://wunmr.wustl.edu/EduDev/Orbitals/movie.html
The drawings of s,p,d and f orbitals shown in your book and in class are probability surfaces. Basically, 95% of the time an electron is found within the region defined by these surfaces. We cannot actually know the exact position of an electron, we can only say what the probability is that it will be in a particular location.
7. Energy ranking - mainly determined by primary energy level, but there is some overlap between sublevels of different primary energy levels. (Show Fig. 3.5).
Shorthand for remembering energy ranking.
Note: there are exceptions to this order of filling, but we don't have time to go into them.
We are now ready to write the "electron configuration" of an atom (element). First we need some rules:
1. Aufbau principle - electrons enter orbitals of lower energy first.
2. Pauli exclusion principle - two electrons per orbital with opposite "spins". One electron is spinning in one direction and the other the opposite direction. Represented by up or down arrows.
3. Hund's rule: fill orbitals of a given energy one at a time before pairing electrons up.
Examples: Let's use these to write the "electronic configuration" of an atom (also fill out fig 3.5)
Hydrogen - 1 electron - goes into the lowest orbital first
Shorthand (show on transparency of Fig 3.5 what this means)
1s1 - the superscipt one means one electron in the 1s orbital.
Lithium - 3 protons, 3 electrons
1s22s1 (show on transparency what this means)
Carbon - p=6, n=6, e=6
1s22s22px12py1 (not 2px2 by Hund's rule). Generally, in this shorthand we do not worry about different p-orbitals (or d or f), but remember it when filling out a diagram such as Fig 3.5
Potassium - 19 electrons
1s22s22p63s23p64s1 (note: 4s fills before 3d because it is slightly lower in energy)
Why do we care???
The electron configuration largely determines the chemistry of an element. In particular, the outer most (highest energy) electrons.