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Electrons in Atoms
Chapter 13
Connection to Chapter 5
• From the atomists to Rutherford, we discussed
the evolution of subatomic particles.
• The discussion continues here….
In this chapter (and the next few) we take a look at
electrons in atoms!
*Chemical properties of atoms, ions, and molecules
are related to the arrangement of electrons
Review
• John Dalton: Atoms are hard, indivisible
spheres.
• JJ Thomson: discovered negatively
charged subatomic particles called electrons
– Assumed a positive charge
was present to cause the
atom to have no overall
charge
– Plum Pudding Model
(or cookie dough model)
https://reich-chemistry.wikispaces.com/Fall.2008.MMA.Cushman.Hutchinson.Timeline
Thompson’s model does not
explain…
 How many protons and electrons?
 Where are the protons and electrons located?
 What happens when electrons are gained/lost?
Ernest Rutherford
• Gold foil experiment: (a review)
– Atom is mostly empty space
– Small, positively charged nucleus with most of
the atom’s mass
– Electrons “float” around
nucleus
Neils Bohr
• Proposed that electrons travel in circular
paths around the nucleus.
• Electrons in a path have a fixed energy
which prevents them from falling into the
nucleus.
• Energy level: The region
around the nucleus where
the electron is likely to
be moving.
Think of energy levels like a
ladder:
What is the difference between the two ladders?
•Electrons cannot exist between energy levels.
Highest Energy Level
Lowest Energy Level
Lowest energy
level is closest
to the nucleus
Highest energy is farthest
away from the nucleus
The quantum
• To move from rung (or orbit) to the next
electrons must gain or lose the right amount of
energy.
– This amount is called
a quantum of energy.
Erwin Schrödinger
• Came up with a mathematical equation to
describe the location of an electron.
• This modern description is called the
Quantum Mechanical (QM) Model of the
atom.
– Does not define a path for the electron, but a
probability of where you will find an electron.
Schrödinger Equation
• Darker areas where electron is likely to be
found
The Quantum Mechanical Model
of Electron Behavior in Atoms
• According to the quantum mechanical model of
electron behavior, the precise paths of electrons
moving around the nucleus cannot be determined
accurately.
• Instead of circular orbits, the location and energy of
electrons moving around the nucleus is specified
using the three terms principle energy level,
sublevel and orbital.
Principal Energy level
• The location of electrons is indicated by assigning
a number n to the level and all electrons located
in that level
• The value of n can be 1, 2, 3, 4, etc.
• Must be a whole number
• The higher the n value, the higher is the energy of the
shell and the contained electrons.
Sublevels
• Each principal energy level is made up of one or
more sublevels that are designated by a letter
from the group s, p, d, or f.
• The number of sublevels = the energy level (n)
– EXAMPLE:
» Energy level 3 (n = 3)
» 3 sublevels: designated 3s, 3p, and 3d.
Atomic Orbitals
• These are the types of sublevels….
• Each sublevel consists of one or more atomic orbitals
• Four types of orbitals with different shapes:
– s orbitals: Spherical in shape
– p orbitals: shaped like a dumbell
Notice the three different
orientations of the p orbitals
Atomic Orbitals continued…
• d orbitals: 5 different
orientations
• f orbitals: 7 different
orientations
Atomic Orbitals Summary
All s sublevels consist of a single s orbital.
All p sublevels consist of three p orbitals.
All d sublevels consist of five d orbitals.
All f sublevels consist of seven f orbitals.
Quantum Mechanical Model Says…
• A maximum of two electrons per atomic orbital
2e-
2e-
2e-
Maximum total of 6 electons for p orbitals
How Many in the d orbitals?
That is a lot to remember…
A Summary…
• The principle energy level (n) is an integer.
Low energies – lower n values.
• # of sublevels = n
– Example: How many sublevels in n=3?
• Ans: 3
• Types of sublevels in n=3:
– 3s, 3p, 3d
• How many orbitals in n=3?
– 3s (1) + 3p(3) + 3d(5) = 9
– Or # of orbitals = n2
• Now, each orbital can hold 2 electrons.
• How many electrons in the n=3?
– Has 9 orbitals = 9x2= 18 electrons
Let’s do another example:
• Look at n=4 (the 4th principal energy level)
• # sublevels?
– # sublevel = n = 4
• Types of sublevels?
– = 4s, 4p, 4d, 4f
• Number of orbitals?
– 4s (1), 4p (3), 4d (5), 4f (7) = 16
• Number of electrons?
– 16x2= 32 electrons fit in the 4th principal energy level
Types of problems you might
encounter
• How many electrons fit in the 3d sublevel?
– There are 5 3d orbitals, each holding 2 electrons
for 10 total electrons in this sublevel.
13.2 Electron Configurations
• Electron Configuration: the arrangement of
electrons around the nucleus
• Specifics about the principal energy levels, sublevels,
orbitals and the number of electrons
How electrons fill orbitals
(writing electron configurations)
1. Aufbau Principle
• Electrons enter orbitals of lowest energy
first.
• Look at figure 13.6. The lowest energy level
is the 1s.
• Electrons go here first
How electrons fill orbitals
(writing electron configurations)
2. The Pauli Exclusion Principle
– Only 2 electrons per orbital
– Don’t squish more than 2 electrons in a box!
– Each of the electrons must have opposite spin
– One up arrow and one down arrow
How electrons fill orbitals
(writing electron configurations)
3. Hund‘s Rule
• When electrons occupy orbitals of equal
energy (ex. 2p) one electron enters each
orbital until all orbitals contain one electron
with parallel (the same) spin.
Writing Electron Configurations
for the first 10 elements.
•
•
•
•
•
•
•
•
•
•
H
He
Li
Be
B
C
N
O
F
Ne
How do I remember that figure?
Exceptions to the rule…
• Aufbau diagrams work for the most part.
• They don’t work for Cr and Cu.
• It ends up that things are more stable when
the d orbitals are half filled so Cr and Cu
move things around to make this happen.