1.0 Atomic structure Lister p 4 - 20. AQA AS Specification LessonsTopics 1-2 Fundamental particles be able to describe the properties of protons, neutrons.

1.0 Atomic structure

Lister p 4 - 20

AQA AS Specification

Lessons Topics

1-2 Fundamental particles• be able to describe the properties of protons, neutrons and electrons in terms of relative charge and relative mass

3 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

The atom consists of two parts:

1. The nucleus which contains:

2. Orbiting electrons.

protonsneutrons

Draw a model of an atom and label the main parts

Structure of an atom

• 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

atomic diameter ~ 10 – 10 m

If a helium atom was the size of a full stop, then the average gerbil would be the size of the

Earth.

If a helium atom was the size of a full stop, then the average gerbil would be the size of the

Earth.

Atoms: How 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?

Atoms: very small

The helium atom is not in the right proportions. The three subatomic particles are wrongly enormous in comparison to

the atom’s radius.

The helium atom is not in the right proportions. The three subatomic particles are wrongly enormous in comparison to

the atom’s radius.

How big is a nucleus?

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.

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, nElectron,

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

particle

Relative charge

Relative mass

Common depictio

n

Proton +1 1

Neutron 0 1

Electron -1 1 10-5

+

-

Subatomic particles in more detail

Subatomic

particle

Relative charge

Relative mass

Common depictio

n

Proton +1 1

Neutron 0 1

Electron -1 1 10-5

+

-

Subatomic particles in more detail

1.2 Electron arrangement• How are electrons arranged in

atoms?They are arranged in shellsThey are arranged in shells

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 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 2,5The shorthand form for, eg, Nitrogen, is 2,5

Shell Number Maximum number of electrons

1 2 x 12 = 2 x 1 = 2

2 2 x 22 = 2 x 4 = 8

3 2 x 32 = 2 x 9 = 18

4 2 x 42 = 2 x 16 = 32

5 2 x 52 = 2 x 25 = 50

Task

Complete the following table:

Shell Number Maximum number of electrons

1

2

3

4

5

Now complete worksheet 1.1

1.3 Mass number, atomic number and isotopesHow 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 protonsThe number of neutronsThe number of electrons

The number of protonsThe number of neutronsThe number of electrons

It tells us what the element is and how many electrons are present in the neutral atom

It tells us what the element is and how many electrons are present in the neutral atom

The number of nucelons is called the Mass number .

What information can we get from this?

We can find out the number of neutronsWe can find out the number of neutrons

Li7

3

No. of protons + neutronsLithium

Number of protons

Number of electrons=

Atomic number does not always equal the number of neutrons.

Atomic number does not always equal the number of neutrons.

Lithium

Electrons

3

Protons 3

Neutrons

4

Mass number

(No. of protons)

Atomic number or proton number

LithiumNumber of

protonsNumber of electrons=

Lithium

Electrons

3

Protons 3

Neutrons

4

This is because the atom is neutral. The charges balance out

-3 charge

+3 charge

But atoms can gain and lose electrons (they become ions). This changes the overall charge on the atom.

But atoms can gain and lose electrons (they become ions). This changes the overall charge on the atom.

No charge

The number of protons “defines” an element – nothing else!The number of protons “defines” an element – nothing else!

Atomic number does not always equal the number of neutrons. This can change, even in atoms of the same element. These are called isotopes.

Atomic number does not always equal the number of neutrons. This can change, even in atoms of the same element. These are called isotopes.

7-Lithium (7Li)

Electrons 3

Protons 3

Neutrons 4

Some isotopes of lithium:

4Li 4-Lithium 3 protons,1 neutron

6Li 6-Lithium 3 protons,3 neutrons



Lithium: always 3 protons!

Isotopes

Complete:

Atom P n e-

Na

Rh

phosphorus

The last of the halogens

Xe

The only liquid non-metal

Li+

F-

Carbon-14 (14C)

A helium atom

He4

2

2 protons2 electrons2 neutrons

Answers

Atom P N ENa 11 12 11

Rh 45 58 45

phosphorus 15 16 15

The last of the halogens 85 125 85

Xe 54 77 54

The only liquid non-metal 35 45 35

Li+ 3 4 2

F- 9 10 10

Carbon-14 (14C) 6 8 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?

Yes – chemistry is about the movement of electronsYes – chemistry is about the movement of electrons

How can you get a fraction of a nucleon?

35.535.5

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

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?

It ionizes atoms and then sends them through an em field where they become deflected on the basis of their mass and charge

It ionizes atoms and then sends them through an em field where they become deflected on the basis of their mass and charge

Why is it important that the instrument is under vacuum?

To prevent collisions of the ions with gas moleculesTo prevent collisions of the ions with gas molecules

How are samples put into the machine?

Volatile liquids and gases can be injected directly, solids must be vapourised.

Volatile liquids and gases can be injected directly, solids must be vapourised.

http://www.youtube.com/watch?v=J-wao0O0_qM&feature=related



Mass Spectrometer

A

B

C

D

E

F

G

Mass Spectrometry - summary

Sample vapourised

Sample ionised positive ions

+ve ions in beam accelerated by electric field

Vacuum pump to keep whole apparatus at v. low pressure

+ve ions subjected to variable magnetic field

+ve ions separated according to mass: charge ratio

+ve ions detected and measured mass spectrum

LOWER m:z ratio

HIGHER m:z ratio

State what happens at each of the locations A-G

Calculating RAM of atoms

Calculate the relative atomic mass of boron.

The tallest “stick” is often (but not always) set at 100

The tallest “stick” is often (but not always) set at 100

boron-10 23 boron-11 100

(100 x 11) + (23 x 10)/123 = 10.8

boron-10 23 boron-11 100

(100 x 11) + (23 x 10)/123 = 10.8

http://www.chem.uoa.gr/applets/AppletMS/Appl_Ms2.html

QuestionHow many isotopes does this element have? What element is it?

Calculate the RAM

51.5

11.2

17.1 17.4

2.8

Answer

51.5

11.2

17.1 17.4

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 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).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.91.3 is the relative atomic mass of zirconium.

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

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?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

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

Cl2+ Cl + Cl+ Cl2+ Cl + Cl+

What can molecular chlorine ions (Cl2+ ) fragment into?

What happens to the Cl atom? If it doesn’t acquire a charge in the ionization chamber then it gets “lost” in the MS

If it doesn’t acquire a charge in the ionization chamber then it gets “lost” in the MS

What are the possible combinations of chlorine-35 and chlorine-37 atoms in a Cl2+ ion?

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 = 7035 + 37 = 7237 + 37 = 74

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 = 7035 + 37 = 7237 + 37 = 74

What would the MS look like?

Chlorine MS …

Why is there no scale on the y-axis?

Because you cannot predict how the molecules will ionize and fragment

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 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

http://www.yellowtang.org/images/electrons_atoms_pos_c_la_784.jpgThe shells represent energy

levels in atoms. Electrons can move between these levels, gaining or losing energy in the process.

The shells represent energy levels in atoms. Electrons can move between these levels, gaining or losing energy in the process.

Sublevels

Each energy level is divided into one or more sublevels. These sublevels have energies that differ slightly from that of the shell energy.

Each energy level is divided into one or more sublevels. These sublevels have energies that differ slightly from that of the shell energy.

How many types of sublevel are there?

(hint – think about he number of regions in the PT)

There are 4: s.p.d.fThe “s-block” comprises Groups 1 and 2The “p- block” comprises Groups 3 - 8

There are 4: s.p.d.fThe “s-block” comprises Groups 1 and 2The “p- block” comprises Groups 3 - 8

How many electrons can an s sublevel have in it?

How many electrons can a p sublevel have in it?

Edps

3

2

1

What is the significance of the order of the subshells?

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

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.

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

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)

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.

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

Nomenclature: the number of electrons in a particular orbital is denoted by superscript. e.g. 1s2 2s2 3p2

Electrons have a property called “spin”. This determines the way in which the degenerate levels are populated.

Electrons have a property called “spin”. This determines the way in which the degenerate levels are populated.

Aufbau

Orbitals do not always fill up in the way expectedOrbitals do not always fill up in the way expected

This is due to overlap in the energies of the sublevels

This is due to overlap in the energies of the sublevels

Look at the energy level diagram for Silicon. Which orbitals have energy levels which overlap?

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.

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)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 d-block contains the 3d elements. These follow from the 4s elements, Potassium and Calcium.

In the periodic table, the transition elements make up the “d-block”.

The first row in the d-block contains the 3d elements. These follow from the 4s elements, Potassium and Calcium.

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)

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)

Complete worksheet 1.5

THE BOHR ATOMTHE 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

1

2

3

4

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SLEVELS AND SUB-LEVELSLEVELS 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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2

3

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SLEVELS AND SUB-LEVELSLEVELS 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



SUB LEVELS

CONTENTSCONTENTS


RULES AND PRINCIPLESRULES 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”!

ORBITALSORBITALS

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...

ORBITALSORBITALS

An orbital is... a region in space where one is likely to find an electron. a region in space where one is likely to find an electron.



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

ORBITALSORBITALS

An orbital is... a region in space where one is likely to find an electron. a region in space where one is likely to find an electron.



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 ORBITALSSHAPES OF ORBITALS

s orbitals

• spherical

• one occurs in every principal energy level


p orbitals

• dumb-bell shaped

• three occur in energy levels except the first


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


SUB LEVELS

ORDER OF FILLING ORBITALSORDER OF FILLING ORBITALS


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1 1s

22s

2p

4s

33s3p3d

44p

4d4f


SUB LEVELS

1 1s

22s

2p

3d

33s3p4s

44p

4d4f


SUB LEVELS



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1 1s

22s

2p

4s

33s3p3d

44p

4d4f


SUB LEVELS

1 1s

22s

2p

3d

33s3p4s

44p

4d4f


SUB LEVELS


THE FILLING ORDER

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s 7p

HOW TO HOW TO REMEMBER ..REMEMBER ....

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’ PRINCIPALTHE ‘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 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 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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STHE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 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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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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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


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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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HUND’S RULEOF


HUND’S RULEOF


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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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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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

INC

RE

AS

ING

EN

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/ D

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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

INC

RE

AS

ING

EN

ER

GY

/ D

IST

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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

INC

RE

AS

ING

EN

ER

GY

/ D

IST

AN

CE

FR

OM

NU

CL

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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

INC

RE

AS

ING

EN

ER

GY

/ D

IST

AN

CE

FR

OM

NU

CL

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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

INC

RE

AS

ING

EN

ER

GY

/ D

IST

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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

INC

RE

AS

ING

EN

ER

GY

/ D

IST

AN

CE

FR

OM

NU

CL

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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


HUND’S RULEOF


1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

ER

GY

/ D

IST

AN

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NU

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TITANIUM

1s2 2s2 2p6 3s2 3p6 4s2 3d2

HUND’S RULEOF


HUND’S RULEOF


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

INC

RE

AS

ING

EN

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/ D

IST

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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


HUND’S RULEOF


1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

ER

GY

/ D

IST

AN

CE

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OM

NU

CL

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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


HUND’S RULEOF


1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

ER

GY

/ D

IST

AN

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CL

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MANGANESE

The new electron goes into the 4s to restore its filled

state.

HUND’S RULEOF


HUND’S RULEOF


1s2 2s2 2p6 3s2 3p6 4s2 3d5

1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

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GY

/ D

IST

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IRON

Orbitals are filled according to Hund’s Rule. They continue to pair up.

HUND’S RULEOF


HUND’S RULEOF


1s2 2s2 2p6 3s2 3p6 4s2 3d6

1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

ER

GY

/ D

IST

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COBALT

HUND’S RULEOF


HUND’S RULEOF


1s2 2s2 2p6 3s2 3p6 4s2 3d7


1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

ER

GY

/ D

IST

AN

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NU

CL

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NICKEL

HUND’S RULEOF


HUND’S RULEOF


1s2 2s2 2p6 3s2 3p6 4s2 3d8


1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

ER

GY

/ D

IST

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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

INC

RE

AS

ING

EN

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/ D

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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

INC

RE

AS

ING

EN

ER

GY

/ D

IST

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GALLIUM - KRYPTON

The 4p orbitals are filled in exactly the same way as

those in the 2p and 3p sub levels.

HUND’S RULEOF


HUND’S RULEOF


1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

RE

AS

ING

EN

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/ D

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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

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

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

ELECTRONIC CONFIGURATIONS OF ELEMENTS 1-30

2

ELECTRONIC ELECTRONIC CONFIGURATIONCONFIGURATIONthe arrangement of the electron in the atom.Electrons are arranged in Energy Levels or Shells

around the nucleus of an atom.

nucleus

1 3 4s

df

sp

sp d

ps

Atomic orbital

f = 7

d = 5

p = 3

s = 1

1 Atomic orbital = 2 e-

x 2 = 2

x 2 = 6

x 2 = 10

x 2 = 14

2e- 8e- 32e-

Main energy level

Subenergy level

no.of electrons

18e-

Questions1. 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. 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 2p41s2 2s2 2p4

If oxygen was ionized, which electron would be removed first?

The antiparallel spin electron has a slightly higher energy. Due to the repulsion from the other electron in the 2px orbital.

The antiparallel spin electron has a slightly higher energy. Due to the repulsion from the other electron in the 2px orbital.

Why?

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)

The energy needed to remove this electron is known as the First Ionisation Energy (IE)

Successive ionisations require more and more energySuccessive ionisations require more and more energy

A logarithmic plot is needed for successive ionisation energies due to the scale. log 1 = 10log 5 = 100,000

Successive Ionisations

2.0

2.5

3.0

3.5

4.0

4.5

5.0

5.5

6.0

0 2 4 6 8 10 12 14 16 18 20

electron removed

log10 of ionisation

energy

Notice the “jump” in energy needed to remove the 2nd electronNotice the “jump” in energy needed to remove the 2nd electron

Successive ionisation of potassium

Successive ionisation of potassium

2.0

2.5

3.0

3.5

4.0

4.5

5.0

5.5

6.0

0 2 4 6 8 10 12 14 16 18 20

electron removed

log10 of ionisation

energy

Successive ionisation energies for potassium

The different “jumps” are evidence for the arrangement of electrons in energy levels and sub-levels

level 1

level 2

level 3

level 4

0200400600800

1000120014001600

Na Mg Al Si P S Cl Ar

1st i

on

isat

ion

en

erg

y (k

J/m

ol)

Periodicity of ionisation energy

What trend would you expect ionisation energy to have as you move across a period? B

A

C

What does region “A” represent?

2 x s electrons

2 x s electrons

What does region “B” represent?

3 x p electrons

3 x p electrons

Which three p electrons are these?

px1 py

1 and pz1px

1 py1 and pz

1

What else do you notice about the

graph?

The slopes of A, B and C are almost the same

The slopes of A, B and C are almost the same

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

0

200

400

600

800

1000

Be Mg Ca Sr Ba

1st

ion

isat

ion

en

erg

y

Group 2

Explain why the first ionisation energy decreases as you move down a group

Describe the graph

The initial decrease is steep, but then the graph flattens out

The initial decrease is steep, but then the graph flattens out

ShieldingAs you move down a group, the distance of the outer electrons from the nucleus increases

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

As you move down a group, the distance of the outer electrons from the nucleus increases

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

Group 4 – the jump is to remove the 5th electron


0

5000

10000

15000

20000

25000

30000

35000

40000

45000

50000

0 1 2 3 4 5 6 7

electron removed

kJ/mol

0

2000

4000

6000

8000

10000

12000

14000

16000

18000

20000

0 1 2 3 4 5 6 7

electron removed

kJ/mol

Group 2 – the jump is to remove the 3rd electron

Group 2 – the jump is to remove the 3rd electron

Question

Identify the groups that these atoms belong to





0

2000

4000

6000

8000

10000

12000

14000

16000

18000

20000

0 1 2 3 4 5 6 7

electron removed

kJ/mol

0

2000

4000

6000

8000

10000

12000

14000

0 1 2 3 4 5 6 7

electron removed

kJ/mol

Question

Identify the group that this atom belongs to

Group 1 – the jump is to remove the 2nd electron

Group 1 – the jump is to remove the 2nd electron

The number of the electron whose removal causes a jump is one more than the group number that the element belongs to.

The number of the electron whose removal causes a jump is one more than the group number that the element belongs to.

0

2000

4000

6000

8000

10000

12000

0 1 2 3 4 5 6 7

electron removed

kJ/mol

Write a general rule for identifying groups from the pattern in ionisation energy

**

1.0 Atomic structure Lister p 4 - 20. AQA AS Specification LessonsTopics 1-2 Fundamental particles be able to describe the properties of protons, neutrons.

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