Download Topic 12.1 Electron Configuration

12.1 Electron Configuration HL
Starter
 Element Bingo
References
 Moodle Powerpoints
 Workbook
Assessment Objective
 12.1.3 State the Relative energies of the s,p,d and f
orbitals in a single energy level.
 12.1.4 state the maximum number of orbitals in a given
energy level.
 12.1.5 Draw the shape of an s – orbital and the shapes
of the px. Py and pz orbitals.
 12.1.6 Apply the Aufbau principle to electron
configurations, Hind Rule and the Pauli exclusion
Principle to write the electron configuration for atoms
up to Z =54
Why care about where
electrons are located?
 Allows us to understand how atoms bond together
to form chemicals and compounds.
 These lessons will be very helpful when we look at
bonding…………………
Introduction
 Electrons are found in orbitals
 Orbital: A region in an atom where there is a high
probability of finding electrons. (Electrons clouds)
 Orbital’s are filled from low to high energy or
starting closest to the nucleus
 ORBITS ARE NOT LIKE A PLANETS!!!!!
Introduction
 The quantum mechanical
model is the current
model of the atom that
describes the
probability of finding an
electron in a given
region of space around
the nucleus.
 It is called the electron
cloud model.
 Quantum numbers are the 4 numbers that
specify the properties of atomic orbitals and
the electrons that reside in them.
1.
2.
3.
4.
The principle quantum number (shell): electrons occupy
the specific energy levels.
The angular momentum quantum number (orbital
shape): specifies the shape of the orbital.
The magnetic quantum number (orbital orientation):
specifies how this shape is arranged in three
dimensions around the nucleus.
The spin quantum numbers: specifies in which direction
the electrons are spinning.
1. The Principle Quantum Number
 Electrons can occupy only specific energy levels.
 These levels are numbered 1 – 7.
 The higher energy levels indicate a higher energy
state for that electron and a location further away
from the nucleus.
 In other
words, electrons
near the nucleus
are low energy electrons.
Increasing
Energy
1. The Principle Quantum Number
 The energy level indicates the size of the
electron cloud. The formula 2n2 indicates the
total possible electrons in an energy level.
1. The Principle Quantum Number
 Example: Calculate the total possible
electrons in the 1st through 4th energy levels.
Remember: 2n2
Energy Level
1
2
3
4
# of Electrons
2. Orbital Shape
 An energy level is actually made of many energy
states called orbitals (subshells or sublevels).
s orbital
 An orbital is a three
dimensional region
around the nucleus
that indicates the
probable location of
one pair of electrons.
These orbitals are s,
p, d and f . The
second quantum
number indicates the
shape of the orbital.
p orbital
d orbitals
Assessment Objective
 12.1.3 State the Relative energies of the
s,p,d and f orbitals in a single energy
level.
 As move up
each level
and sub level
the relative
energy
increases
Increasing energy
7 Different Main Energy Levels
Increasing energy
Sub levels in each level e.g level 2 has s and p sub levels
Assessment Objective
 12.1.5 Draw the shape of an s – orbital
and the shapes of the px. Py and pz
orbitals.
 The s-orbital is
 sphere shaped.
Need to Draw these!
The p-orbital is dumb
bell shaped
Need to Draw these!
d-orbitals
look like leaves
of clover
 f-orbitals look like
flowers
3 Orientation
 The magnetic number indicates the orientation
of an orbital around the nucleus. Since an ‘s’
orbital is spherical, it can have only 1
orientation.
3 Possible Orientations for p-Orbitals
d-Orbitals
have 5 Possible
Orientations
 f-Orbitals have 7 Possible Orientations
Orbital Orientation
Orbital
Type
MAX #
Electrons
per Orbital
Possible
Max # of
Orientations Electrons if
for Each
Each Orbital
Type of
Orientation
Orbital
is Filled
s
2
1
2
p
2
3
6
d
2
5
10
f
2
7
14
4 Spin
 The spin quantum
number indicates that
the 2 electrons
occupying a single
orbital must have
opposite spin.
Assessment Objective
 12.1.4 state the maximum number of
orbitals in a given energy level.
Principle
Quantum #
Main
Energy
Level (n)
Types of
Orbitals in
Main
Energy
Level (n)
# of
Orientations
per Orbital
Type
# of
Orbitals per
Main
Energy
Level
(n2)
Max #
Electrons
In Filled
Orbitals
Max # of
Electrons
per Main
Energy
Level (2n2)
1
s
1
1
2
2
2
s
p
1
3
4
2
6
8
3
s
p
d
1
3
5
9
2
6
10
18
4
s
p
d
f
1
3
5
7
16
2
6
10
14
32
Principle
Quantum
Number:
Main Energy
Level (n)
5
6
7
Types of
Orbitals in
Main Energy
Level (n)
# Of
Orientations
per Orbital
Type
# of Orbitals
Max # of
Max # of
Electron per electrons per
Per Main
Energy Level Filled Orbital Main Energy
Level
s
p
d
f
1
3
5
7
(n - 1)2
s
p
d
1
3
5
(n - 3)2
s
p
1
3
(n –
16
9
5)2
4
2
6
10
14
2(n -1)2
2
6
10
2(n – 3)2
18
2
6
2(n – 5)2
8
32
The row numbers on the periodic table
are the same as the principle quantum
numbers or energy levels (n).
Writing Electron Configurations
Assessment Objective
 12.1.6 Apply the Aufbau Principle to electron
configurations, Hind Rule and the Pauli
Exclusion Principle to write the electron
configuration for atoms up to Z =54
The Aufbau Principle
 Energy levels
overlap so the
diagram shows the
order of the
sublevels. The
Aufbau principle
states that an
electron occupies
the lowest energy
orbital that can
receive it.
Writing Electron Configurations
 Using the atomic number as the total number of
electrons you can write the electron configuration for
all of the elements.
 The number of electrons in each sublevel is written as
a exponent.
Remember…………….
 Orbitals in each level
 Level 1 s only,
 level 2 s and p,
 level 3 s. p and d
 Level 4 and 5 s, p, d and f
 Level 6 s, p and d
 Level 7 s and p
 Number of electrons in each type
of orbital.
 s =2 , p = 6, d = 10, f = 14
 Electrons fill from lowest energy
levels first (Aufbau Principle)
(but energy levels overlap)
Example
 Boron (in level 2 so s and
p orbitals)
 Number of protons = 5
 Therefore number of
electrons = 5
 Answer B 1s22s22p1
 Phosphorous ( in level 3
so s , p and d orbitals)
 Number of protons =
number of electrons = 15
 Answer
P 1s22s22p63s23p5
 Lead ( level 6 so has s, p, d
and f )
 Atomic number = 82 = 82
protons = 82 electrons
 Pb
1s22s22p63s23p64s23d104p6
5s24d105p66s24f145d106p2
Using the periodic table
 Write the electron configuration for
 magnesium #12
 gallium #31,
 element #35.
Periodic Table Orbital Filling Method
Noble – Gas Notation
 The Group VIII elements, helium, argon, krypton,
xenon, and radon are called the noble gases. The
configurations of the noble gases are often used
as a shorthand method for writing longer electron
configurations.
 For Example: Sodium – Na has 11 electrons
 Electron Configuration
1s22s22p63s1
 Noble – Gas Configuration
[Ne]3s1
Example 2:
 Arsenic, As, 33 electrons
 Electron Configuration
1s22s22p63s23p64s23d104p3
 Noble – Gas Configuration
[Ar]4s23d104p3
 Example 3: Barium, Ba, 56 electrons
 Example 4: Rubidium, Rb
 Use the periodic table to write the electron
configurations for the following atoms.
 Example 1: Nitrogen, 7 electrons
 Example 2: Phosphorus, 15 electrons
 Example 3: Cerium, 58 electrons
Assessment Objective
 12.1.6 Apply the Aufbau principle to electron
configurations, Hind Rule and the Pauli exclusion
Principle to write the electron configuration for
atoms up to Z =54
Pauli and Hind
 Although the following do not effect the writing
of electronic configurations it is important to
note how these electrons are arranged within the
orbitals
Pauli Exclusion Principle
 The maximum of two electrons can occupy a single
atomic orbital
 Electrons move in 2 direction (spin)
 (no two electrons in a given atom can
have the same 4 quantum numbers. i.e have same
energy level, shape, orientation and spin)
Hind Rule
 Single electrons with same spin must
occupy equal energy orbital before
additional electrons can be added in the
opposite spin.
 In practice let us consider
to element oxygen
 N 1s22s22p3
 There are 3 orientations for
the p orbital (x, y and z)
 According to Pauli, only two
electrons in each orientation
(so max of six electrons for
p sub-orbital)
 According to Hind these electron MUST put one
electron in each orientation in the same spin
direction first.
 If we were to expand out the electron
configuration for nitrogen
N 1s22s2 2px1py1pz1
(all p orbitals in same spin direction)
 For oxygen, we start filling up the sub orbital with
electrons in the other spin direction
O 1s22s22p4
 If we expand this out we get:
O 1s22s2 2px2py1pz1
With the 2px orbital containing electrons with
opposite spins.