1. Quantum Numbers
Section 3.5

2. First - The Principal Number, n
The symbol, n, is used to label the orbits (a.k.a. energy level)
Bohr used n to denote the orbitals/main energy levels

3. Second - The Secondary Quantum Number, l (lower case L)
Value of l
Sublevel symbol
Number of orbitals s 1 p 3 2 d 5 f 7
Describes additional electron energy sublevels or subshells that form part of the main energy level
And the shape of the orbital
Bohr’s theory explained the line spectra of atoms, but did not explain why bright line spectra were actually composed of more than one line, a theory was introduced in 1915 that the orbital shapes were actually elliptical and each main energy level was composed of several small steps. Thus s orbitals are circular and all others are elliptical

5. The Aufbau principle tells us that when building energy diagrams we always start at the lowest energy level and then build up

6. Each orbital has its own energy level
“s” orbitals are lower in energy than “p” orbitals and so on

7. Third - The Magnetic Quantum Number, ml
Tells us the direction of the electron orbit
It ranges from +l to –l (tells use the # of orientations of orbits possible)
Line spectra can be split into even more lines when a gas tube is placed next to a magnet, the ml explains this phenomenon
The orbits differ only in their orientation in space (different angles)

8. Fourth - The Spin Quantum Number ms
Tells us in what direction the electron is spinning, they always go in opposite directions – this is known as the Pauli exclusion principle.
The value of ms can be only – ½ or + ½
More line splitting was noticed near magnets and other kinds of weak magnetism
In 1925 Wolfgang Pauli suggested that certain atoms could act as weak magnets (due to unpaired electrons) b/c each electron is spinning on its axis while moving around the nucleus

10. Seatwork
Pg. 182 #3
Pg. 184 #3, 4, 5, 7
Pg. 197 # 7, 8, 14
Assign the quantum numbers for the last electron in the following elements:
Copper
Fluorine
Uranium
Calcium