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Introduction to Atomic Orbitals

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					                       Introduction to Atomic Orbitals

Or, where do the electrons “live”?

The locations where the electrons are can be thought of as an address: the
street, the dormitory, the room, the bed.

Principal Quantum Number (n): - basically the numbers down the left side of
the Periodic Table: 1, 2, 3, 4 etc. – indicates the main energy level

Angular Momentum Quantum Number (l): while scientists say these are
0,1,2,3, etc, it’s probably more useful to think of them as the letters
corresponding to each of these numbers, which are s, p, d, f. (There could be
more, but the known elements never need them) Think of it as the shape
(coming from “angle”) – starting with a single s orbital, the number of each
type increases by 2: so there are 3 p’s, 5 d’s, and 7 f’s

Magnetic Quantum Number (ml): These tell us something about the
orientation of the orbital in space. (see the pictures on pp 102, 103). As
we’re considering the filling of orbitals, this is not too critical; when we talk
about bonding, we’ll come back to it.

Spin Quantum Number (ms): Think of an electron as a spinning top: it can
either point up (+1/2) or down (-1/2) – no other choices.
As we build up electron configurations when atomic number increases, we’ll
see how these four numbers make up an “address” for an electron.

The rules for filling orbitals include the Aufbau principle, which simply says
that electrons go into the lowest energy orbital available:



               1s

               2s    2p

               3s    3p     3d

               4s    4p     4d     4f

               5s    5p     5d     5f

               6s    6p     6d     6f

               7s    7p     7d     7f




SOOOOO, fill the 1s, then follow the diagonal arrows and the dotted
“return” lines to fill the orbitals using the following additional rules:

The Pauli exclusion principle, which says that no two electrons can have the
same four principal quantum numbers: or, so to speak, even if electrons pair
up in a “room” (orbital), they will have different beds (spin).


Hund’s Rule, which says orbitals of the same energy (for example the three
p’s at level 2 – called degenerate orbitals) will each accept one electron
before they begin to pair up.
EXCEPTIONS: (there’s always at least one, right?) – the two sublevels
circled, 5d and 6d, each get one electron, and then the 4f and 5f levels fill,
and there are several other anomalies: look at chromium and molybdenum for
example – this problem occurs in the transition elements.

**************************

Writing (or picturing) electron configurations and orbital notation – ways of
reading the “address”:

Electron configuration notation just collects the s’s, p’s, d’s and f’s by
principal quantum number, thus:

Na = 1s2 2s2 2p6 3s1         eleven nice electrons, all comfortable at home

Note that we passed through the electron configuration of Neon, which is:

Ne = 1s2 2s2 2p6    situations where the outer s’s and p’s are filled are called
                   an octet

We can therefore use a shorthand called the noble gas configuration, thus:

Na = [Ne] 3s1

We can also illustrate this in orbital notation, where we use a horizontal line
for each orbital, which can hold a pair of electrons. Electrons are illustrated
by up and down arrows – some chemists use half arrowheads, some use full:


   Na ____ ____ ____ ____ ____ ____
       1s 2s     2px 2py  2pz 3s




Practice: P 116 Section review Q 4 & 5, P 118 22 - 30

				
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