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Atomic Structure & the periodic table
Starter:
Draw the electron arrangement for carbon, nitrogen, lithium, oxygen, helium.
Ionisation energies
Key Words:• Subshells• Orbitals• Principle
quantum number
• Ionisation energy
• Ionisation
Objectives:
Discuss Ionisation energy
Outcomes:
D:recall and understand the definition of ionisation energies of gaseous atoms
A-C:
- Understand that they are endothermic processes
Energy levels & electron shells• Electrons in an atom are arranged in a series of
shells • Shells:
• Each shell ins described by a principle quantum number:
• The larger the value of n, the further from the nucleus you are likely to find the electron:
Ionisation Energies:• Ionisation: the complete removal of an electron
from an atom• It is an endothermic process
Since energy is needed to overcome the attractive force between the electron and the nucleus.
• Ionisation energy: the amount of energy needed to remove an electron from its atom can be measures by increasing voltage applied to a gas until
it conducts electricity & emits light – which tells you an electron has been freed
• Ground state: the lowest energy state for an atom
• The energy needed to remove the one electron from Hydrogen in its ground state is normally quoted for 1 mole and is the ionisation energy of hydrogen
• First ionisation energy: energy needed to remove the first electron from an atom
• Second ionisation energy: energy needs to remove the second electron from an atom
• The first ionisation is a measure of how tightly or loosely an outer electron is attracted to the positive nucleus.
• The more easily an electron is removed, the more reactive an atom will be.
• Total energy required to remove electrons add the 1st & 2nd ionisation energies together.
• The different energies needs to remove the 1st & subsequent electrons confirm that electrons are found on different energy levels
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
Subshells:• Quantum mechanics also tells us that each shell
may contain subshells.• Subshells: regions of differing energy within a
shell, shown by letters: s, p, d, f, g.
• The following subshells are available in each shell:
• Shell 1 is closest to the nucleus so it takes the most energy to remove electrons from this shell
• Electrons in the lowest energy subshells are closest to the nucleus
s (lowest energy) < p < d
• Each type of subshell contains one or more orbitals
• Orbitals: the region where the electrons are most likely to be found. They hold a maxiumum of 2 electrons
• All orbitals in a particular subshell are at the same energy level
• As n increases, the energy gap between successive shells gets smaller– Due to this, orbitals in neighbouring shells may
overlap– The 3d orbital has an energy level above that of 4s
orbital but below 4p orbital
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LEVELS AND SUB-LEVELS
PRINCIPAL ENERGY 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.
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LEVELS AND 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.
A study of Ionisation Energies and the periodic properties of elements suggested that the main energy levels were split into sub levels.
Level 1 was split into 1 sub level
Level 2 was split into 2 sub levels
Level 3 was split into 3 sub levels
Level 4 was split into 4 sub levels
SUB LEVELS
CONTENTSCONTENTS
PRINCIPAL ENERGY LEVELS
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 SHAPE OCCURRENCE
s spherical one in every principal level
p dumb-bell three in levels from 2 upwards
d various five in levels from 3 upwards
f 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 SHAPE OCCURRENCE
s spherical one in every principal level
p dumb-bell three in levels from 2 upwards
d various five in levels from 3 upwards
f 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
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.
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1 1s
22s
2p
4s
33s3p3d
44p
4d4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS
ORDER OF FILLING ORBITALS
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.
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1 1s
22s
2p
4s
33s3p3d
44p
4d4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS
1 1s
22s
2p
3d
33s3p4s
44p
4d4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS
ORDER OF FILLING ORBITALS
Practice:
Orbitals & shells
Key Words:• Electron spin• Electronic
configuration
Objectives:
Electrons spins & filling orbitals
Outcomes:D: recall electrons populate orbitssingly before pairing up
A-C:
- Understand electron spin
- predict the electronic structure & configuration of atoms of hydrogen to krypton inclusive using 1s …notationand electron-in-boxes notation
Electron Spin:• Electron Spin: the rotation of electrons
clockwise or anticlockwise creating a magnetic field
• An electron can spin either clockwise or anticlockwise – and because it is moving, it creates a magnetic field
• This can be represented by using a small arrow ( ) or ( ) – so showing spins in opposite directions
• 2 electrons in the same orbital cannot have the same spin
• This means each orbital can have a maximum of 2 electrons, having opposite spins
Filling the orbitals:
• Electronic configuration: the arrangement of electrons in an atom in their subshells and orbitals
• E.g: Hydrogen in it ground state has one electron 1s1
• Therefore:– Helium: 1s2
– Lithium: 1s22s1
– Be: 1s22s2
How would you
draw the boxes?
• Note: the empty p orbitals are shown. • It doesn’t matter which 2p orbital is filled first
as they all have the same energy
Hund’s Rule
Writing electronic configurations:
• For an ion: you simply add or subtract the right number of electrons from the outer shell
• Remember Hund’s Rule when removing electrons: one electron comes out of each completely filled orbital in the outer shell before any unpaired electrons and removed.
• E.g: O2- is: 1s22s22p6
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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This states that…
“ELECTRONS ENTER THE LOWEST AVAILABLE
ENERGY LEVEL”
THE ‘AUFBAU’ PRINCIPAL
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
d orbitals
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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HYDROGEN
1s1
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
Hydrogen atoms have one electron. This goes into a vacant orbital in the lowest available energy level.
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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HELIUM
1s2
THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
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
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
LITHIUM
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.
1s2 2s1
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
BERYLLIUM
Beryllium atoms have four electrons so the fourth electron pairs up in the 2s orbital. The 2s sub level is now full.
1s2 2s2
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
BORON
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.
1s2 2s2 2p1
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
CARBON
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.
1s2 2s2 2p2
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
NITROGEN
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.
1s2 2s2 2p3
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
OXYGEN
With all three orbitals half-filled, the eighth electron in an oxygen atom must now pair up with one of the electrons already there.
1s2 2s2 2p4
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
FLUORINE
The electrons continue to pair up with those in the half-filled orbitals.
1s2 2s2 2p5
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
NEON
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.
1s2 2s2 2p6
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
SODIUM - ARGON
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.
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
SODIUM - ARGON
Na 1s2 2s2 2p6 3s1
Mg 1s2 2s2 2p6 3s2
Al 1s2 2s2 2p6 3s2 3p1
Si 1s2 2s2 2p6 3s2 3p2
P 1s2 2s2 2p6 3s2 3p3
S 1s2 2s2 2p6 3s2 3p4
Cl 1s2 2s2 2p6 3s2 3p5
Ar 1s2 2s2 2p6 3s2 3p6
Remember that the 3p configurations follow Hund’s Rule with the electrons remaining unpaired to give more stability.
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
POTASSIUM
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.
1s2 2s2 2p6 3s2 3p6 4s1
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
CALCIUM
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.
1s2 2s2 2p6 3s2 3p6 4s2
‘Aufbau’
Principle
‘Aufbau’
Principle
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
SCANDIUM
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.
1s2 2s2 2p6 3s2 3p6 4s2 3d1
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
TITANIUM
1s2 2s2 2p6 3s2 3p6 4s2 3d2
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
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.
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
VANADIUM
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.
1s2 2s2 2p6 3s2 3p6 4s2 3d3
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
CHROMIUM
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.
1s2 2s2 2p6 3s2 3p6 4s1 3d5
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
MANGANESE
The new electron goes into the 4s to restore its filled state.
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1s2 2s2 2p6 3s2 3p6 4s2 3d5
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
IRON
Orbitals are filled according to Hund’s Rule. They continue to pair up.
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1s2 2s2 2p6 3s2 3p6 4s2 3d6
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
COBALT
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1s2 2s2 2p6 3s2 3p6 4s2 3d7
Orbitals are filled according to Hund’s Rule. They continue to pair up.
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
NICKEL
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1s2 2s2 2p6 3s2 3p6 4s2 3d8
Orbitals are filled according to Hund’s Rule. They continue to pair up.
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
COPPER
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.
1s2 2s2 2p6 3s2 3p6 4s1 3d10
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
ZINC
The electron goes into the 4s to restore its filled state and complete the 3d and 4s orbital filling.
1s2 2s2 2p6 3s2 3p6 4s2 3d10
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
GALLIUM - KRYPTON
The 4p orbitals are filled in exactly the same way as those in the 2p and 3p sub levels.
HUND’S RULEOF
MAXIMUM MULTIPLICITY
HUND’S RULEOF
MAXIMUM MULTIPLICITY
1 1s
22s
2p
4s
3
3s
3p
3d
44p
4d
4f
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THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS
GALLIUM - KRYPTON
Ga - 4p1
Ge - 4p2
As - 4p3
Se - 4p4
Br - 4p5
Kr - 4p6
Remember that the 4p configurations follow Hund’s Rule with the electrons remaining unpaired to give more stability.
Prefix with…
1s2 2s2 2p6 3s2 3p6 4s2 3d10
Practice:1. The electronic structure of an atom of an element in Group 6 of the Periodic Table could be:A 1s2 2s2 2p2
B 1s2 2s2 2p4
C 1s2 2s2 2p6 3s2 3p6 3d6 4s2
D 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
2.
3.
Task: work out the electronic configuration of atoms from hydrogen to argon
B, C, D
1s1
1s2
1s2 2s1
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
HHeLiBeBCNOFNeNaMgAlSiPSClArKCaScTiVCrMnFeCoNiCuZn
ELECTRONIC CONFIGURATIONS OF ELEMENTS 1-30
ELECTRONIC CONFIGURATION OF IONS
• Positive ions (cations) are formed by removing electrons from atoms• Negative ions (anions) are formed by adding electrons to atoms• Electrons are removed first from the highest occupied orbitals (EXC. transition metals)
SODIUM Na 1s2 2s2 2p6 3s1 1 electron removed from the 3s orbital
Na+ 1s2 2s2 2p6
CHLORINE Cl 1s2 2s2 2p6 3s2 3p5 1 electron added to the 3p orbital
Cl¯ 1s2 2s2 2p6 3s2 3p6
ELECTRONIC CONFIGURATION OF IONS
• Positive ions (cations) are formed by removing electrons from atoms• Negative ions (anions) are formed by adding electrons to atoms• Electrons are removed first from the highest occupied orbitals (EXC. transition metals)
SODIUM Na 1s2 2s2 2p6 3s1 1 electron removed from the 3s orbital
Na+ 1s2 2s2 2p6
CHLORINE Cl 1s2 2s2 2p6 3s2 3p5 1 electron added to the 3p orbital
Cl¯ 1s2 2s2 2p6 3s2 3p6
FIRST ROW TRANSITION METALS
Despite being of lower energy and being filled first, electrons in the 4s orbital are removed before any electrons in the 3d orbitals.
TITANIUM Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2
Ti+ 1s2 2s2 2p6 3s2 3p6 4s1 3d2
Ti2+ 1s2 2s2 2p6 3s2 3p6 3d2
Ti3+ 1s2 2s2 2p6 3s2 3p6 3d1
Ti4+ 1s2 2s2 2p6 3s2 3p6
Electrons & orbitals
Key Words:• Electron
density• Electron cloud• Electron
density map
Objectives:
Electrons density maps and the shape & orientation of orbitals
Outcomes:
A-C:
- Recall what electron density maps are
- Know & recall the shapes and orientations of orbitals: p and d
Electron configuration & chemical properties
Key Words:• Periodic law• Groups• Periods• Transition
metals• Metalloids• Lanthanides• Actinides
Objectives:
- electronic structure determines the chemical properties of an element
- periodic table is divided into blocks
Outcomes:D: recall that chemical properties are related to electronic structure- Know the blocks of the periodic tableA-C:- Know the chemical properties of:
- s-block elements- d-block elements- p-block elements
Periods & groups:• The reactivity of an element, and how it
combines with other elements, is determined by its arrangement of electrons in its outer shell
• The periodic table arranges elements in order of their atomic number
• Groups: the vertical columns in the periodic table
• Periods: the horizontal rows in a periodic table
• All the elements in a period have the same number of electron shells.
• So, the elements in each group and period show particular characteristics and trends in their chemical and physical properties
Periodic Law: the properties of the elements are a function of their atomic numbers
Blocks
Block Groups Subshell:s 1 + 2 Outer electrons in s subshells
p 3+4+5+6+7+0 Outer electrons in p subshells
d Transitional metals
Outer electrons in d subshells
f Lanthanides + actinides
Outer electrons in f subshells
s-block elements
• Reactive metals• Lower melting temperature• Lower boiling temperature• Lower density• Conduct electricity• Include hydrogen and helium – but usually
treated as a separate group.
Than other metals
d-block elements• Called Transitional metals• Less reactive that Group 1+ 2 metals – this is
because the inner d orbital is being filled while the outer s orbital is full
• All conduct electricity and heat• Are shiny, and hard• Ductile – pulled into shape• Malleable – hammered into shape• Mercury is the only exception – low melting
temperature liquid at room temperature
f-block elements• Lanthanides – are all similar• Actinides – all radioactive
– Only the actinides up to uranium are naturally occurring
– The others have all been synthesises by scientists and have extremely short half-lives
p-block elements• All the non-metals and metalloids• Include Tin and Lead
– Form positive ions– Form ionic bonds with non-metals
• Many metals in p block do not have strong metallic characteristics– All conduct heat and electricity– Called post transitional metals generally
unreactive
• Metalloids occur in a diagonal block• Mostly like non-metals• Conduct electricity – but poorly• Silicon and germanium are responsible for
microchips
• Non-metals all form covalent bonds with other non-metals & ionic bonds with metals
• Majority do not conduct electricity• Some elements form giant covalent structures
Practice:
Trends in the Periodic Table
Key Words:• Atomic radius• Ionisation
energy• Melting
temperature
Objectives:
- Understand trends in the periodic table
Outcomes:
D: understand and describe the trends in the periodic table
A-C: Explain the trends in the periodic table - - -- ionization energy based on given data or recall of the shape of the plots of ionization energy versus atomic number using ideas of electronic structure and the way that electron energy levels vary across the period.
- melting temperature of the elements based on given data
Key Words• Ionization energy: the amount of energy it takes to
strip away the first electron• Electronegativity: a measure of how tightly an atom
holds onto its outer shell electrons• Nuclear charge: the attractive force between the
positive protons in the nucleus and the negative electrons in the energy levels. The more protons, the greater the nuclear charge.
• Shielding: inner electrons tend to shield the outer electrons from the attractive force of the nucleus. The more energy levels between the outer electrons and the nucleus, the more shielding.
• Atomic radius: measure of the size of atoms, usually measured from the nucleus to the outer shell
• Ionic radius: the size of ions
Key points:• Across periodic table: elements gain electrons
• Down a group : elements gain electron shells.
This changes the diameter of atoms which affects their physical and chemical properties
• The atomic radius generally decreases across a period:
The nuclear charge becomes increasingly positive as the number of protons in the nucleus increases.
The number of electrons also increases BUT they are all in the same shellThis means that they are attracted more strongly to
the nucleus – so reducing the atomic radius across a period
• The atomic radius generally increases down a group:
The outer electrons enter new energy levels down a group
So, even though the nucleus has more protons, the electrons are further away and they are screened by more electron shells.So, they are not held so tightly and the atomic radius
increases
Atoms to ion:
• The atomic radius changes when atoms form ions
• Positive ions always have a smaller ionic radius that the original atom.– Because: the loss of electron(s) means that the
remaining electrons each have a greater share of the positive charge of the nucleus so are more tightly bound
– And when an ion in formed, a whole ion shell is usually lost
• Negative ion has a larger atomic radius than that of the original atom
even though the extra electrons are in the same electron shell, the addition of the negative charge means that the electrons are less tightly bound to the nucleusSo the atomic radius is larger
Periodic Trends in Ionisation Energy:• Then more tightly held the outer electrons, the higher the ionisation
energy
3 main factors affecting ionisation energy of an atom:
The attraction between the nucleus & the outermost electron – decreases as the distance between them increases reducing the ionisation energy
The size of the positive nuclear charge - a more positive nucleus has a greater attraction for the outer electron so higher ionisation energy
Inner shells of electrons repel the outer electron, screening or shielding it from the nucleus - the more electron shells there are between the outer electrons and the nucleus, the less firmly held the outer electron is lower ionisation energy
Ionisation energy & periods:• Ionisation energy increases across a period• It becomes harder to remove an electron
• This is because:Increasing positive nuclear charge across the period
o Without the addition of extra electron shells to screen the outer electrons
The atomic radius gets smaller & electrons are held more firmly – so it requires more energy to make ionisation happen
The end of each period is marked by the high ionisation energy of a noble gas – this is a result of a stable electronic structure & indicates their unreactive natures
• (b) shows that First ionisation energies do not increase smoothly across a period
• This is because of subshells within each shell• E.g: the first ionisation energy of Be is larger
than B, Mg has a larger first ionisation energy than Al – why?– For Be or Mg, an electron must be removed from a
full s-shell. – Full subshells are particularly stable – so it requires
more energy than removing a single p electron from B or Al
• Nitrogen & phosphorous have unexpectedly high first ionisation energies:– They both have a half-full outer p subshell.– Half full subshells seem to have greater stability – So requires more energy
Ionisation energy decreases down a group – it becomes easier to lose an electron
Patterns in physical properties• The physical properties are closely linked to the
structure and bonding of atoms
• Melting temperature: the temperature at which the pure solid is in equilibrium with the pure liquid, at atmospheric pressure.– this is affected by the packing & binding of atoms
within a substance– It changes as you go across a period
• The relatively high melting temperatures of the metals (e.g. Li, Mg, Al) are due to their metallic structure.– The atoms are held tightly together is a ‘sea of
electrons’– It takes a lot of energy to separate them
• Giant molecular structures (metalloids-silicon, carbon-in form of diamond):– Strong covalent bonds between atoms which hold
them tightly in a crystal structure– Very difficult to remove individual atoms– So very high melting temperature
• Simple molecular structures:– Most non-metals found on right of periodic table– Small, individual molecules– Strong covalent bonds within molecules– But, molecules are held together by weak
intermolecular forces– Can be separated easily– Low melting temperature
Practice
