10-12-2012, 06:46 PM
Quantum Mechanics
quantum-II.pptx (Size: 162.42 KB / Downloads: 40)
Quantum Numbers
Read on pg. 200 from “The theory of quantum…” (about third paragraph) to “The Magnetic Quantum Number, ml” on pg. 201. Do PE 3
The subshells of n = 3 are l = 0(s), 1(p), 2(d)
for n = 4: l = 0(s), 1(p), 2(d), 3(f)
the magnetic quantum number
Recall: we are looking at the first three of four quantum numbers: n, l, ml, ms
The magnetic quantum number is ml, it further divides subshells into “orbitals”
Recall that even though you can visualize these divisions as spherical regions around the nucleus, they really refer to different waveforms
ml ranges from - l to + l, in intervals of one
when l = 1, the values of ml are -1, 0, 1
More practice with quantum #s
Complete the chart on the study sheet
Look at the last two columns of the chart.
A maximum of two electrons can fit in each orbital.
For n = 3, a maximum of 18 electrons can fit in this shell (2 + 6 + 10)
This is equivalent to 2n2 : 2(3)2 = 18.
From now on, you can determine the # of electrons in a shell by using this “2n2” rule.
Summary
Read pg. 202
Figure 6.19 indicates the energies of subshells and the number of orbitals in each. We will see that each of these orbitals can hold exactly 2 electrons
Note that some shells overlap with respect to energy.
If we extend a Bohr-like model to represent this we would see shells being split into subshells causing some shells to overlap…