The Atom
Objectives
Previous Unit
To know the three elementary particles which
compose atoms.
To understand the development of the atomic
model.
Let’s take this one step further…
The Atom Today
Over time, the model of the atom evolved.
Two early models we saw were:
Thomson’s Plum Pudding Model
Rutherford’s Nuclear Atom
The Atom Today
The model proposed by Niels Bohr was the Planetary
Model
The central nucleus (like the sun) surrounded by orbiting
electrons (like the planets)
Explained that electrons don’t fall into nucleus because they
have fixed energy
The Bohr model was an improvement, but was replaced by the…
Quantum Mechanical Model
Modern model of the atom
No longer are electrons treated like they travel like
“classical” particles (bowling balls)
Focus is on electrons
Electrons are found not in orbits but in “clouds”
Based on probabilities
Area where there is a 90% chance of finding an electron
Electron clouds
Are 3-dimensional
Come in several shapes
Are tied to specific energies
Energy levels fill in special order
Organizing Electrons
Electrons in an atom are organized into
different energy levels
1. Principal energy levels have sublevels
2. Sublevels take the form of atomic orbitals
3. Orbitals “contain” electrons
Principal Energy Levels (n)
Principal energy levels (n) are numbered
Maximum number of electrons in a level is 2n2
n = 1
n = 2
n = 3
2 x 12 = 2 electrons
2 x 22 = 8 electrons
Each principal level has a number of sublevels
equal to the level number
n = 1
n = 2
1 sublevel
2 sublevels
Sublevels
Each sublevel contains a certain number
of atomic orbitals
Orbitals are regions where it is likely an electron
will be found
Each orbital can hold a maximum of 2 electrons
Letters are used to denote orbitals
Orbitals have characteristic shapes
Atomic Orbitals
Orbital
How many per Shape
sublevel?
s
1
spherical
p
3
peanut
d
5
4 are double-peanut
1 is a donut-ringed peanut
f
7
flower
Orbitals are regions in
Space
The energy level of
an electron is the
region around the
nucleus where the
electron is likely to
be moving.
S orbital is a sphere
P orbitals
Each P orbital can hold two electrons, but they
need to have opposite spins
The P sublevel holds 6 electrons
D level orbitals
Each The D sublevel can hold 10 electrons.
Each orbital holds 2 electrons with opposite
spins
The F sublevel has 7
orbitals
Each orbital can hold
2 electrons with
opposite spin
The F sublevel holds
14 electrons
The F Sublevel
Atomic Orbitals
s
p
d
f
Electrons can change orbitals
Electrons can change orbital, by
absorbing energy. When an
electron absorbs a quantum of
energy, it moves up to a higher
orbital.
When the electron falls from a high
orbital to a lower orbital, energy is
released, and we see light.
Wintergreen mint is an example
We will also see this in our
spectroscopy and flame test
labs!
Atomic Orbitals
Electron configurations
A series of numbers and letters to show which orbitals
contain electrons for a given element
Before we proceed…
Aufbau Principle
Electrons enter orbitals of the lowest energy first
Pauli Exclusion Principle
No two electrons can have the same “state”
• State = electron’s orbital and its spin
• Electrons in the same orbital will have opposite spin
Hund’s Rule
When electrons occupy orbitals of equal energy, one
electron enters each orbital until all orbitals contain one
electron with spin in the same direction
Aufbau Diagram
Orbital Diagrams
A box stands for one orbital
An arrow stands for one electron
Each box (orbital) can only hold 2 arrows
(electrons)
or to show direction of spin
*Note: There is some overlap of energy levels, so
sublevels do not fill in exact numerical order.
Orbital Diagrams
Orbital Diagram for A Nitrogen
Atom
N
1s
2s
2p
3s
Orbital Diagram for A Fluorine
Atom
F
1s
2s
2p
3s
Orbital Diagram for A
Magnesium Atom
Mg
1s
2s
2p
3s
Learning Check O1
Write the orbital diagram for the electrons
in an oxygen atom.
Solution O1
Write the orbital diagram for the electrons
in an oxygen atom.
1s
2s
2p
3s
Learning Check O2
Write the orbital diagram for the electrons
in an iron atom.
Solution O2
Write the orbital diagram for the electrons
in an iron atom.
1s
2s
2p
3d
3s
3p
Electron Configuration
Once electrons have been placed in boxes, the electron
configuration is written.
Write the energy level and letter for every sublevel
that holds electrons
Use a superscript to indicate the number of electrons
in each sublevel
Example: for Mg – 1s2 2s2 2p6 3s2
Mg
1s
2s
2p
3s
Exceptional Electron
Configurations
Filled energy levels are the most stable
Half filled energy levels are the second
most stable
Copper, Chromium and Silver are very
important examples of how “exceptions” to
filling the Aufbaum diagram leads to
understanding the charge of transition
metals
Examples
Write electron configurations for: P, Cr, Sn
Now write the electron configuration without the
boxes: Mg, Ni
What element has the following electronic
configuration:
1s2 2s2 2p6 3s2 3p1
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
For fun at home:
http://www.lon-capa.org/~mmp/period/electron.htm
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