Transcript ORBITALS

1.0 Atomic structure
Lister p 4 - 20
AQA AS Specification
Lessons
1-2
3
Topics
Fundamental particles
• be able to describe the properties of protons, neutrons and electrons in terms of
relative charge and relative mass
Electron arrangement
• know that early models of atomic structure predicted that atoms and ions with
noble gas electron-arrangements should be stable
4-7
Mass number and isotopes
• be able to recall the meaning of mass number (A) and atomic(proton) number (Z)
• be able to explain the existence of isotopes
• understand the principles of a simple mass spec,limited to ionisation, acceleration,
deflection and detection, and its use for identifying elements and RMM
8-11
Electron arrangement
• know the electron configurations of atoms and ions up to Z = 36 in terms of levels
and sub-levels (orbitals) s, p and d
• know the meaning of the term ionisation energy.
• understand how ionisation energies in Period 3 (Na – Ar) and inGroup 2 (Be – Ba)
give evidence for electron arrangement in sub-levels and in levels
The Atom
Draw a model of an atom and label the
main parts
The atom consists of two parts:
1. The nucleus which contains:
protons
neutrons
2. Orbiting electrons.
Structure of an atom
atomic diameter ~ 10 – 10 m
• An atom consists of a
central positively charged
nucleus containing protons
and neutrons (nucleons)
• Diameter approx. 10-15 m
(1 femtometre)
• Electrons surround the
nucleus
• Atomic diameter approx.
10-10 m roughly 100 000 x
nucleus diameter
nucleus diameter ~ 10 – 15 m
Atoms: How small?
If a helium atom was the size of a full stop, then
the average gerbil would be the size of the
Earth.
Atoms: very small
Now let’s pretend that the helium atom on the right
is the size of the Earth.
What’s wrong with this simple picture?
How big is a nucleus?
The helium atom is not in the right proportions. The three
subatomic particles are wrongly enormous in comparison to
the atom’s radius.
Most of the atom is empty space!
If you imagine an atom being the size of Wembley
stadium, the nucleus would be about the size of a
football on the centre spot.
The electrons would be two peas flying around the whole
stadium. The rest of it: emptiness.
Properties of subatomic particles
Property
Proton, p
Neutron, n
Electron, e-
Mass/ kg
1.673 x 10-27
1.675 x 10-27
0.911 x 10-31
Charge/C
+1.602 x10-19
0
-1.602 x 10-19
Position
In the
nucleus
In the
nucleus
Around the
nucleus
Subatomic particles in more detail
Subatomic
particle
Relative
charge
Relative
mass
Common
depiction
Proton
+1
1
+
Neutron
0
1
Electron
-1
1  10-5
-
Subatomic particles in more detail
Subatomic
particle
Relative
charge
Relative
mass
Common
depiction
Proton
+1
1
+
Neutron
0
1
Electron
-1
1  10-5
-
1.2 Electron arrangement
• How are electrons arranged in
atoms?
How do we know how many
electrons are in each shell?
The shells are numbered outward from
the nucleus.
The maximum number of electrons
found in each shell can be calculated
from 2n2 where n is the shell number.
The shorthand form for, eg, Nitrogen, is
2,5
They are arranged in
shells
Task
Complete the following table:
Shell Number
Maximum number of
electrons
1
2 x 12
=2x1 =2
2
2 x 22
=2x4
=8
3
2 x 32
=2x9
= 18
4
2 x 42
= 2 x 16 = 32
5
2 x 52
= 2 x 25 = 50
Now complete worksheet 1.1
1.3 Mass number, atomic number and isotopes
How can we describe an atom in terms
of it’s subatomic structure?
What information do we need to know?
The number of protons is called the
Atomic number .
What is significant about the number of
protons in the nucleus?
The number of nucelons is called the
Mass number .
What information can we get from this?
The number of protons
The number of neutrons
The number of electrons
It tells us what the
element is and how
many electrons are
present in the neutral
atom
We can find out the
number of neutrons
Lithium
No. of protons + neutrons
Mass number
7
Li
3
Atomic number or
proton number
(No. of protons)
Number of
protons
=
Number of
electrons
Lithium
Electrons
3
Protons
3
Neutrons
4
Atomic number does not always equal
the number of neutrons.
Lithium
Number of
protons
=
Number of
electrons
Lithium
This is because the
atom is neutral. The
charges balance out
Electrons
3
-3 charge
Protons
3
+3 charge
Neutrons
4
No charge
But atoms can gain and lose electrons (they become
ions). This changes the overall charge on the atom.
Isotopes
Some isotopes of lithium:
4Li
6Li
4-Lithium
6-Lithium
The number of protons “defines” an
element – nothing else!
3 protons,
1 neutron
3 protons,
3 neutrons
10Li
3 protons,
10-Lithium
7 neutrons
11Li
3 protons,
11-Lithium
8 neutrons
7-Lithium (7Li)
Electrons
3
Protons
3
Neutrons
4
Lithium:
always 3
protons!
Atomic number does not always equal the number of
neutrons. This can change, even in atoms of the same
element. These are called isotopes.
Complete:
Atom
Na
Rh
phosphorus
The last of the halogens
Xe
The only liquid non-metal
Li+
FCarbon-14 (14C)
4
2
He
A helium atom
P
n
e-
2 protons
2 electrons
2 neutrons
Answers
Atom
Na
Rh
phosphorus
The last of the halogens
Xe
The only liquid non-metal
Li+
FCarbon-14 (14C)
P
11
45
15
85
54
35
3
9
6
N
12
58
16
125
77
45
4
10
8
E
11
45
15
85
54
35
2
10
6
Chemical properties of isotopes
Would you expect the
isotopes of lithium to
have the same
chemical properties?
What is the Mass
number, Z, of Chlorine?
How can you get a
fraction of a nucleon?
Yes – chemistry is about the movement
of electrons
35.5
The relative abundance of two chlorine
isotopes is similar, hence the mass
number on the PT is an average
number determined by the abundances
of the isotopes
1.4 Mass spectrometry
What does a mass
spectrometer do?
Why is it important
that the instrument
is under vacuum?
How are samples
put into the
machine?
It ionizes atoms and then sends them
through an em field where they
become deflected on the basis of
their mass and charge
To prevent collisions of the ions with
gas molecules
Volatile liquids and gases can be
injected directly, solids must be
vapourised.
http://www.youtube.com/watch?v=Jwao0O0_qM&feature=related
Mass Spectrometer
Mass Spectrometry - summary
State what happens
at each of the
locations A-G
HIGHER m:z ratio
LOWER m:z ratio
A
Sample vapourised
B
Sample ionised  positive ions
C
+ve ions in beam accelerated by electric field
D
Vacuum pump to keep whole apparatus at v. low pressure
E
+ve ions subjected to variable magnetic field
F
+ve ions separated according to mass: charge ratio
G
+ve ions detected and measured  mass spectrum
Calculating RAM of atoms
Calculate the relative
atomic mass of boron.
boron-10
boron-11
23
100
(100 x 11) + (23 x 10)/123 = 10.8
The tallest “stick” is often
(but not always) set at
100
http://www.chem.uoa.gr/applets/AppletMS/
Appl_Ms2.html
Question
How many isotopes does this element have?
What element is it?
51.5
17.1
11.2
Calculate the RAM
17.4
2.8
Answer
51.5
17.1
17.4
11.2
2.8
Step 1: Find the total mass of these 100 typical atoms:
(51.5 x 90) + (11.2 x 91) + (17.1 x 92) + (17.4 x 94) + (2.8 x 96)
= 9131.8
Step 2: find the average mass of these 100 atoms :
9131.8 / 100 = 91.3 (to 3 sig fig).
91.3 is the relative atomic mass of zirconium.
Question
The mass spectrum of uranium has 3 peaks: at 234 m/z,
235 m/z and 238 m/z. The abundance of the isotopes is
0.006%, 0.72% and 99.2% respectively. What is the
average relative atomic mass of uranium?
240
237.0
237.8
238
Question
Chlorine has two isotopes, 35Cl
and 37Cl, in the approximate ratio
of 3 atoms of 35Cl to 1 atom of
37Cl. Draw the stick diagram for
Chlorine
Question
Chlorine has two isotopes, 35Cl
and 37Cl, in the approximate ratio
of 3 atoms of 35Cl to 1 atom of
37Cl. Draw the stick diagram for
Chlorine
Wrong!
Why?
The problem is that chlorine consists of molecules, not individual atoms.
When chlorine is passed into the ionisation chamber, an electron is
knocked off the molecule to give a molecular ion, Cl2+. Doubly charges
ions can also form.
These ions aren’t very stable, and some will fall apart to give a chlorine
atom and a Cl+ ion. The term for this is fragmentation
Chlorine MS
What can molecular chlorine ions (Cl2+ ) fragment into?
Cl2+  Cl + Cl+
What happens to the Cl atom?
What are the possible
combinations of chlorine-35 and
chlorine-37 atoms in a Cl2+ ion?
What would the MS look like?
If it doesn’t acquire a charge in
the ionization chamber then it
gets “lost” in the MS
Both atoms could be 35Cl, both atoms
could be 37Cl, or you could have one
of each sort.
Masses of the Cl2+ ion:
35 + 35 = 70
35 + 37 = 72
37 + 37 = 74
Chlorine MS …
Why is there no
scale on the y-axis?
Because you cannot
predict how the
molecules will ionize
and fragment
1.5 Electron configurations
Why is the periodic table broken up into sections?
What links each of these sections?
The distribution of electrons within the shells is, in most cases, more
complicated than simple spheres. The regions within the PT closely
follow the patterns of these distributions – or probabilities of electron
density
The shells represent energy
levels in atoms. Electrons can
move between these levels,
gaining or losing energy in the
process.
http://www.yellowtang.org/images/electrons
_atoms_pos_c_la_784.jpg
Sublevels
Each energy level is divided
into one or more sublevels.
These sublevels have
energies that differ slightly
from that of the shell energy.
What is the significance of the order of
the subshells?
E
3
How many types of sublevel
are there?
2
(hint – think about he
number of regions in the PT)
1
There are 4: s.p.d.f
The “s-block” comprises Groups 1 and 2
The “p- block” comprises Groups 3 - 8
d
p
s
How many electrons can
an s sublevel have in it?
How many electrons can
a p sublevel have in it?
Atomic orbitals
Sublevels aren’t all the same.
The s-sublevel has the lowest
energy and so is filled first. It
can hold a maximum of two
electrons.
The s orbital is spherical and
represents the probability of
finding the electrons within its
boundary
The p-, d- and f- sublevels are
degenerate, ie further broken down
into more sublevels of almost
equivalent energy.
If a p-orbital can hold 6
electrons in total, how many
degenerate orbitals are there?
Orbital shapes
Spin
Electrons have a property
called “spin”. This determines
the way in which the
degenerate levels are
populated.
Electron s are either spin up, or spin
down – ie clockwise or anticlockwise.
(corresponding to a spin quantum
number of +1/2 and -1/2)
Degenerate orbitals of the same
energy fill up first. Parallel spins go in
first followed by antiparallel spin.






1s
2s
2px 2py 2pz




Nomenclature: the number of electrons in a particular
orbital is denoted by superscript. e.g. 1s2 2s2 3p2
Aufbau
Silicon (Si)
Energy
Orbitals do not always fill
up in the way expected
4f
4d
4p
This is due to overlap in
the energies of the
sublevels
3d
4s
3p
3s
2p
2s
Look at the energy level diagram
for Silicon. Which orbitals have
energy levels which overlap?
1s
Distance from nucleus
The 4s orbital has an energy
between that of the 3p and 3d
orbitals. This means that the 4s
orbital fills before the 3d orbital.
Filling orbitals
Electrons enter the lowest energy orbital
available (Aufbau principle)
In the periodic table, the
transition elements make
up the “d-block”.
The first row in the dblock contains the 3d
elements. These follow
from the 4s elements,
Potassium and Calcium.
Complete worksheet 1.5
Electrons prefer to occupy orbitals on their
own, and only pair up when no empty
orbitals of the same energy are available
(Hund's Rule)
THE BOHR ATOM
Ideas about the structure of the atom have
changed over the years. The Bohr theory
thought of it as a small nucleus of protons and
neutrons surrounded by circulating electrons.
Each shell or energy level could hold a maximum
number of electrons.
The energy of levels became greater as they got
further from the nucleus and electrons filled
energy levels in order.
The theory couldn’t explain certain aspects of
chemistry.
Maximum electrons
per shell
1st shell
2
2nd shell
8
3rd shell
18
4th shell
32
5th shell
50
LEVELS AND SUB-LEVELS
INCREASING ENERGY / DISTANCE FROM NUCLEUS
PRINCIPAL
ENERGY
LEVELS
4
3
2
1
During studies of the spectrum of
hydrogen it was shown that the energy
levels were not equally spaced. The
energy gap between successive levels
got increasingly smaller as the levels
got further from the nucleus. The
importance of this is discussed later.
LEVELS AND SUB-LEVELS
INCREASING ENERGY / DISTANCE FROM NUCLEUS
PRINCIPAL
ENERGY
LEVELS
4
SUB LEVELS
During studies of the spectrum of
hydrogen it was shown that the energy
levels were not equally spaced. The
energy gap between successive levels
got increasingly smaller as the levels
got further from the nucleus. The
importance of this is discussed later.
3
A study of Ionisation Energies and the
periodic properties of elements suggested
that the main energy levels were split
into sub levels.
2
Level 1 was split into 1 sub level
Level 2 was split into 2 sub levels
Level 3 was split into 3 sub levels
1
Level 4 was split into 4 sub levels
CONTENTS
RULES AND PRINCIPLES
HEISENBERG’S UNCERTAINTY PRINCIPLE
“You cannot determine the position and momentum of an electron at the same time.”
This means that you cannot say exactly where an electron is. It put paid to the idea of
electrons orbiting the nucleus in rings and introduced the idea of orbitals.
THE AUFBAU PRINCIPLE
“Electrons enter the lowest available energy level.”
PAULI’S EXCLUSION PRINCIPLE
“No two electrons can have the same four quantum numbers.”
Two electrons can go in each orbital, providing they are of opposite spin.
HUND’S RULE OF MAXIMUM MULTIPLICITY
“When in orbitals of equal energy, electrons will try to remain unpaired.”
Placing two electrons in one orbital means that, as they are both negatively charged,
there will be some electrostatic repulsion between them. Placing each electron in a
separate orbital reduces the repulsion and the system is more stable. It can be
described as the “SITTING ON A BUS RULE”!
ORBITALS
An orbital is... a region in space where one is likely to find an electron.
Orbitals can hold up to two electrons as long as they have opposite spin; this
is known as PAULI’S EXCLUSION PRINCIPAL.
Orbitals have different shapes...
ORBITALS
An orbital is... a region in space where one is likely to find an electron.
Orbitals can hold up to two electrons as long as they have opposite spin; this is
known as PAULI’S EXCLUSION PRINCIPAL.
Orbitals have different shapes...
ORBITAL
s
p
d
f
SHAPE
spherical
OCCURRENCE
one in every principal level
dumb-bell
three in levels from 2 upwards
various
five in levels from 3 upwards
various
seven in levels from 4 upwards
ORBITALS
An orbital is... a region in space where one is likely to find an electron.
Orbitals can hold up to two electrons as long as they have opposite spin; this is
known as PAULI’S EXCLUSION PRINCIPAL.
Orbitals have different shapes...
ORBITAL
s
p
d
f
SHAPE
spherical
OCCURRENCE
one in every principal level
dumb-bell
three in levels from 2 upwards
various
five in levels from 3 upwards
various
seven in levels from 4 upwards
An orbital is a 3-dimensional statistical shape showing where one is most likely to
find an electron. Because, according to Heisenberg, you cannot say exactly where
an electron is you are only able to say where it might be found.
DO NOT CONFUSE AN ORBITAL WITH AN ORBIT
SHAPES OF ORBITALS
s orbitals
• spherical
• one occurs in every principal energy level
SHAPES OF ORBITALS
p orbitals
• dumb-bell shaped
• three occur in energy levels except the first
SHAPES OF ORBITALS
d orbitals
• various shapes
• five occur in energy levels except the first and second
ORDER OF FILLING ORBITALS
SUB LEVELS
4
4f
4d
4p
4s
3
3d
3p
3s
2
1
2p
2s
1s
INCREASING ENERGY / DISTANCE FROM NUCLEUS
PRINCIPAL
ENERGY
LEVELS
Orbitals are not filled in numerical order because the principal energy levels get
closer together as you get further from the nucleus. This results in overlap of sub
levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.
ORDER OF FILLING ORBITALS
SUB LEVELS
4
4f
4d
4p
4s
3
3d
3p
3s
2
1
2p
2s
1s
INCREASING ENERGY / DISTANCE FROM NUCLEUS
PRINCIPAL
ENERGY
LEVELS
PRINCIPAL
ENERGY
LEVELS
4
3
2
1
SUB LEVELS
4f
4d
4p
3d
4s
3p
3s
2p
2s
1s
Orbitals are not filled in numerical order because the principal energy levels get
closer together as you get further from the nucleus. This results in overlap of sub
levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.
ORDER OF FILLING ORBITALS
4
3
2
1
SUB LEVELS
4f
4d
4p
4s
3d
3p
3s
2p
2s
1s
INCREASING ENERGY / DISTANCE FROM NUCLEUS
PRINCIPAL
ENERGY
LEVELS
PRINCIPAL
ENERGY
LEVELS
4
3
SUB LEVELS
4f
4d
4p
3d
4s
3p
3s
HOW TO
REMEMBER ...
THE FILLING ORDER
1s
2s
3s
2
2p
2s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
7s
1
2p
6d
7p
1s
Orbitals are not filled in numerical order because the principal energy levels get
closer together as you get further from the nucleus. This results in overlap of sub
levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.
INCREASING ENERGY / DISTANCE FROM NUCLEUS
THE ‘AUFBAU’ PRINCIPAL
4
4f
This states that…
4d
“ELECTRONS ENTER THE
LOWEST AVAILABLE
ENERGY LEVEL”
4p
3d
4s
3
3p
3s
2p
2
2s
The following sequence will
show the ‘building up’ of the
electronic structures of the
first 36 elements in the
periodic table.
Electrons are shown as half
headed arrows and can spin
in one of two directions
or
s orbitals
p orbitals
1
1s
d orbitals
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
HYDROGEN
4
4d
1s1
4p
3d
4s
3
3p
Hydrogen atoms have one
electron. This goes into a
vacant orbital in the lowest
available energy level.
3s
‘Aufbau’
2p
2
2s
1
1s
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
HELIUM
4
4d
1s2
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
Every orbital can contain 2
electrons, provided the
electrons are spinning in
opposite directions. This is
based on...
PAULI’S EXCLUSION
PRINCIPLE
The two electrons in a
helium atom can both go in
the 1s orbital.
‘Aufbau’
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
LITHIUM
4
4d
1s2 2s1
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
1s orbitals can hold a
maximum of two electrons
so the third electron in a
lithium atom must go into
the next available orbital of
higher energy. This will be
further from the nucleus in
the second principal
energy level.
The second principal level
has two types of orbital (s
and p). An s orbital is
lower in energy than a p.
‘Aufbau’
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
BERYLLIUM
4
4d
1s2 2s2
4p
3d
4s
3
3p
Beryllium atoms have four
electrons so the fourth
electron pairs up in the 2s
orbital. The 2s sub level is
now full.
3s
2p
2
2s
1
1s
‘Aufbau’
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
BORON
4
4d
1s2 2s2 2p1
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
As the 2s sub level is now
full, the fifth electron goes
into one of the three p
orbitals in the 2p sub level.
The 2p orbitals are slightly
higher in energy than the
2s orbital.
‘Aufbau’
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
CARBON
4
4d
1s2 2s2 2p2
4p
3d
4s
3
3p
3s
2p
2
2s
The next electron in
doesn’t pair up with the
one already there. This
would give rise to
repulsion between the
similarly charged species.
Instead, it goes into
another p orbital which
means less repulsion,
lower energy and more
stability.
HUND’S RULE
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
NITROGEN
4
4d
4p
3d
4s
3
3p
3s
1s2 2s2 2p3
Following Hund’s Rule,
the next electron will not
pair up so goes into a
vacant p orbital. All three
electrons are now
unpaired. This gives less
repulsion, lower energy
and therefore more
stability.
2p
2
2s
HUND’S RULE
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
OXYGEN
4
4d
1s2 2s2 2p4
4p
3d
4s
3
3p
With all three orbitals halffilled, the eighth electron in
an oxygen atom must now
pair up with one of the
electrons already there.
3s
‘Aufbau’
2p
2
2s
1
1s
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
FLUORINE
4
4d
1s2 2s2 2p5
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
The electrons continue to
pair up with those in the
half-filled orbitals.
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
NEON
4
4d
1s2 2s2 2p6
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
The electrons continue to
pair up with those in the
half-filled orbitals. The 2p
orbitals are now
completely filled and so is
the second principal
energy level.
In the older system of
describing electronic
configurations, this would
have been written as 2,8.
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
SODIUM - ARGON
4
4d
4p
3d
4s
3
3p
3s
2p
With the second principal
energy level full, the next
electrons must go into the
next highest level. The
third principal energy level
contains three types of
orbital; s, p and d.
The 3s and 3p orbitals are
filled in exactly the same
way as those in the 2s and
2p sub levels.
2
2s
‘Aufbau’
Principle
1
1s
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
SODIUM - ARGON
4
4d
Na
1s2 2s2 2p6 3s1
Mg
1s2 2s2 2p6 3s2
4p
3d
3
4s
Al
1s2 2s2 2p6 3s2 3p1
3p
Si
1s2 2s2 2p6 3s2 3p2
3s
P
1s2 2s2 2p6 3s2 3p3
S
1s2 2s2 2p6 3s2 3p4
Cl
1s2 2s2 2p6 3s2 3p5
Ar
1s2 2s2 2p6 3s2 3p6
2p
2
2s
1
1s
Remember that the 3p
configurations follow Hund’s
Rule with the electrons
remaining unpaired to give
more stability.
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
POTASSIUM
4
4d
1s2 2s2 2p6 3s2 3p6 4s1
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
In numerical terms one
would expect the 3d
orbitals to be filled next.
However, because the
principal energy levels get
closer together as you go
further from the nucleus
coupled with the splitting
into sub energy levels, the
4s orbital is of a LOWER
ENERGY than the 3d
orbitals so gets filled first.
‘Aufbau’
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
CALCIUM
4
4d
1s2 2s2 2p6 3s2 3p6 4s2
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
As expected, the next
electron pairs up to
complete a filled 4s orbital.
This explanation, using
sub levels fits in with the
position of potassium and
calcium in the Periodic
Table. All elements with an
-s1 electronic configuration
are in Group I and all with
an -s2 configuration are in
Group II.
‘Aufbau’
Principle
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
SCANDIUM
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d1
4p
3d
4s
3
3p
3s
With the lower energy 4s
orbital filled, the next
electrons can now fill the
3d orbitals. There are five d
orbitals. They are filled
according to Hund’s Rule BUT WATCH OUT FOR
TWO SPECIAL CASES.
2p
2
2s
HUND’S RULE
OF
MAXIMUM MULTIPLICITY
1
1s
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
TITANIUM
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d2
4p
3d
4s
3
3p
3s
The 3d orbitals are filled
according to Hund’s rule
so the next electron
doesn’t pair up but goes
into an empty orbital in the
same sub level.
HUND’S RULE
2p
2
2s
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
VANADIUM
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d3
4p
3d
4s
3
3p
3s
The 3d orbitals are filled
according to Hund’s rule
so the next electron
doesn’t pair up but goes
into an empty orbital in the
same sub level.
HUND’S RULE
2p
2
2s
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
CHROMIUM
4
4d
1s2 2s2 2p6 3s2 3p6 4s1 3d5
4p
3d
4s
3
3p
3s
2p
2
One would expect the
configuration of chromium
atoms to end in 4s2 3d4.
To achieve a more stable
arrangement of lower
energy, one of the 4s
electrons is promoted into
the 3d to give six unpaired
electrons with lower
repulsion.
2s
HUND’S RULE
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
MANGANESE
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d5
4p
3d
4s
3
The new electron goes into
the 4s to restore its filled
state.
3p
3s
HUND’S RULE
2p
2
2s
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
IRON
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d6
4p
3d
4s
3
Orbitals are filled
according to Hund’s Rule.
They continue to pair up.
3p
3s
HUND’S RULE
2p
2
2s
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
COBALT
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d7
4p
3d
4s
3
Orbitals are filled
according to Hund’s Rule.
They continue to pair up.
3p
3s
HUND’S RULE
2p
2
2s
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
NICKEL
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d8
4p
3d
4s
3
Orbitals are filled
according to Hund’s Rule.
They continue to pair up.
3p
3s
HUND’S RULE
2p
2
2s
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
COPPER
4
4d
1s2 2s2 2p6 3s2 3p6 4s1 3d10
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
One would expect the
configuration of chromium
atoms to end in 4s2 3d9.
To achieve a more stable
arrangement of lower
energy, one of the 4s
electrons is promoted into
the 3d.
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
ZINC
4
4d
1s2 2s2 2p6 3s2 3p6 4s2 3d10
4p
3d
4s
3
3p
3s
2p
2
2s
1
1s
The electron goes into the
4s to restore its filled state
and complete the 3d and
4s orbital filling.
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
GALLIUM - KRYPTON
4
4d
4p
3d
4s
3
The 4p orbitals are filled in
exactly the same way as
those in the 2p and 3p sub
levels.
3p
3s
HUND’S RULE
2p
2
2s
1
1s
OF
MAXIMUM MULTIPLICITY
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
4f
INCREASING ENERGY / DISTANCE FROM NUCLEUS
GALLIUM - KRYPTON
4
4d
Prefix with…
4p
1s2 2s2 2p6 3s2 3p6 4s2 3d10
3d
4s
3
3p
3s
Ga
- 4p1
Ge
- 4p2
As
- 4p3
Se
- 4p4
Br
- 4p5
Kr
- 4p6
2p
2
2s
1
1s
Remember that the 4p
configurations follow Hund’s
Rule with the electrons
remaining unpaired to give
more stability.
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
1s1
ELECTRONIC
1s2
CONFIGURATIONS
1s2 2s1
OF ELEMENTS 1-30
1s2 2s2
1s2 2s2 2p1
1s2 2s2 2p2
1s2 2s2 2p3
1s2 2s2 2p4
1s2 2s2 2p5
1s2 2s2 2p6
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s2 3p1
1s2 2s2 2p6 3s2 3p2
1s2 2s2 2p6 3s2 3p3
1s2 2s2 2p6 3s2 3p4
1s2 2s2 2p6 3s2 3p5
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s1
1s2 2s2 2p6 3s2 3p6 4s2
1s2 2s2 2p6 3s2 3p6 4s2 3d1
1s2 2s2 2p6 3s2 3p6 4s2 3d2
1s2 2s2 2p6 3s2 3p6 4s2 3d3
1s2 2s2 2p6 3s2 3p6 4s1 3d5
1s2 2s2 2p6 3s2 3p6 4s2 3d5
1s2 2s2 2p6 3s2 3p6 4s2 3d6
1s2 2s2 2p6 3s2 3p6 4s2 3d7
1s2 2s2 2p6 3s2 3p6 4s2 3d8
1s2 2s2 2p6 3s2 3p6 4s1 3d10
1s2 2s2 2p6 3s2 3p6 4s2 3d10
ELECTRONIC CONFIGURATION
Electrons
are arranged
Levels
Shells
the arrangement
of in
theEnergy
electron
in theoratom.
around the nucleus of an atom.
nl
x
no.of
electrons
Sub
Main
energy energy
level level
f
1 2 p3 d 4 d
p
p
s
s s
nucleus
s
2e- 8e18e-
f=7
x 2 = 14
d = 5 x 2 = 10
32e-
p = 3x 2 = 6
s=1 x2=2
Atomic orbital
-
1 Atomic orbital = 2 e
Questions
1. Which orbital would the electrons fill first? The 2s or 2p orbital?
2. Can you have an electron in between two orbitals?
3. How many d orbitals are there in the d subshell?
4. How many electrons can the p orbital hold?
5. Why can two electrons occupy the same orbital?
1. The 2s orbital would be filled before the 2p orbital because orbitals that
are lower in energy are filled first and the 2s orbital is lower in energy
than the 2p orbital.
2. You cannot have an electron in between two orbitals. The electron will
either be in one orbital or the next.
3. There are 5 d orbitals in the d subshell.
4. A p orbital can hold 6 electrons.
5. Two electrons can occupy the same orbital because they each have a
different spin. There cannot be two electrons that have the same exact
orbital configuration and spin.
1.6 Ionization energy
Draw the electron configuration
of oxygen
1s2 2s2 2p4






If oxygen was ionized, which electron
would be removed first?
Why?
The antiparallel spin electron has
a slightly higher energy. Due to
the repulsion from the other
electron in the 2px orbital.
Will the next electron be
easier to remove?
Which one will it be?
The energy needed to
remove this electron is
known as the First Ionisation
Energy (IE)
Successive ionisations require
more and more energy
Successive Ionisations
A logarithmic plot is
needed for successive
ionisation energies
due to the scale.
log 1 = 10
log 5 = 100,000
6.0
5.5
5.0
4.5
log10 of
ionisation 4.0
energy
Successive
ionisation of
potassium
3.5
3.0
2.5
2.0
0
2
4
6
8
10
12
electron removed
Notice the “jump” in energy needed to
remove the 2nd electron
14
16
18
20
Successive ionisation energies for potassium
6.0
5.5
level 1
5.0
4.5
level 2
log10 of
ionisation 4.0
energy
3.5
level 3
3.0
2.5
level 4
2.0
0
2
4
6
8
10
12
14
16
18
20
electron removed
The different “jumps” are evidence for the arrangement of electrons in
energy levels and sub-levels
What trend would you expect
ionisation energy to have as
you move across a period?
What does region
“A” represent?
2xs
electrons
What does region
“B” represent?
3xp
electrons
Which three p
electrons are these?
What else do you
notice about the
graph?
1st ionisation energy
(kJ/mol)
Periodicity of ionisation energy
1600
1400
1200
1000
800
600
400
200
0
C
B
A
Na
Mg
Al
px1 py1 and pz1
The slopes of A, B and C are
almost the same
Si
P
S
Cl
Ar
Across the periodic table
Describe the graph
What causes the change in
the pattern at A = 21
Predict the shape of a
graph showing the trend
of first ionization energy
down a group
Trends of first ionization energy in groups
Group 2
1st ionisation energy
1000
800
Describe the graph
600
400
200
0
Be
Mg
Ca
Sr
Explain why the first ionisation
energy decreases as you move
down a group
Ba
The initial decrease is steep,
but then the graph flattens
out
Shielding
_
e
As you move down a group, the
distance of the outer electrons
from the nucleus increases
1) distance from nucleus
+
2) nuclear charge
3) shielding (repulsion) by electrons
in inner shells between nucleus
and outer electron
The inner electrons also shield
the outer electrons from the full
effect of the positive nuclear
charge and repel each other.
They are less tightly bound to
the nucleus and so are more
easily removed
Question
Identify the groups that
these atoms belong to
50000
20000
45000
18000
40000
16000
35000
14000
30000
12000
kJ/mol 25000
kJ/mol 10000
20000
8000
15000
6000
10000
4000
5000
2000
0
0
0
1
2
3
4
electron removed
Group 4 – the jump is to
remove the 5th electron
5
6
7
0
1
2
3
4
electron removed
Group 2 – the jump is to
remove the 3rd electron
5
6
7
Question
Identify the groups that
these atoms belong to
20000
14000
18000
12000
16000
10000
14000
12000
8000
kJ/mol
kJ/mol 10000
6000
8000
6000
4000
4000
2000
2000
0
0
0
1
2
3
4
electron removed
Group 3 – the jump is to
remove the 4th electron
5
6
7
0
1
2
3
4
5
electron removed
Group 5 – the jump is to
remove the 6th electron
6
7
Question
Write a general rule for identifying
groups from the pattern in
ionisation energy
Identify the group that this
atom belongs to
The number of the electron
whose removal causes a
jump is one more than the
group number that the
element belongs to.
12000
10000
8000
kJ/mol 6000
4000
2000
0
0
1
2
3
4
5
6
7
electron removed
Group 1 – the jump is to
remove the 2nd electron
**