Evidence for shells

History of the atom

·        The model of the atom has changed as our observations of its behaviour and properties have increased.

·        A model is used to explain observations.  The model changes to explain any new observations.

·        The stages in the development of the atom:-

George Johnstone Stoney (1891)                   Electrolysis.  The charge of an electron.

Joseph J Thompson (1897)                             The cathode ray tube and e/m deflection.  The mass / charge of an electron.

Robert Milikan (1909)                                         Oil drop experiment.  The mass / charge of an electron.

Joseph J Thompson                                          ‘Plum – Pudding’ model of an atom.

Geiger, Rutherford and Marsden (1909)        Alpha particle deflection.  The nuclear model.

Henry Moseley (1913)                                        Atomic number

Plasma displays

A plasma unit has 100,000's of tiny cells (pixels) filled with a mixture of neon and xenon gasses. A single pixel is made up of three coloured sub-pixels, red, green and blue. Each sub pixel is driven by its own electrode, stimulates the gas to release ultraviolet light photons. The photons interact with a phosphor material coated on the inside wall of the cell. Phosphors are substances that give off light when they are exposed to other light eg. Ultraviolet light. The phosphors in a pixel give off coloured light when they are charged. The cells are situated in a grid like structure. The varying intensity of the current can create millions of different combinations of red, green and blue across the entire spectrum of colour.

• How atoms are able to give off light comes down to understanding the movement of electrons in atoms and ions:

• When electrons are removed energy is required.

Ionisation energies:        Click for DEMO (click demo once)

·        To form positive ions, electrons must be completely removed i.e. ionisation. ·        To do this the electron must completely escape the attraction of the atom. i.e. reach n=¥. ·        At n=¥ the electron has sufficient energy to escape the attraction from the nucleus.

   Definition:    

·     The 1^st ionisation energy of an element is the energy required to remove 1 electrons from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

1^st ionisation energy:

                                    Ca_(g)             à        Ca⁺_(g)        +          e^-                        1st IE = +590 KjMol^-1

2^nd ionisation energy:

                                    Ca⁺_(g)           à        Ca²⁺_(g)       +          e^-                        2nd IE = +1150 KjMol^-1

Factors affecting ionisation energy - always refer to all 3 in any explainations:

1)     The distance of the electron from the nucleus

• The further an electron is from the nucleus, the lower the force of attraction:

                    F  a  1/d²

• This means that the electron will be easier to remove which means the ionisation energy will be lower.

2)     Size of the positive nuclear charge

• The more protons in the nucleus, the higher the nuclear charge, the harder it is to remove an electron, the higher the ionisation energy.

3)     The ‘shielding’ effect by full inner shells

• A full inner shell of electrons will repel electrons in outer shells. 

• These ‘shields’ affect of the attraction from the nucleus on outer electrons. 

• The more inner shells the greater the shielding.

Ionisation energies down a Group:

• As you go down a Group, the outer electron shell is further from the nucleus - attraction decreases.

• The more inner shells the greater the shielding.

• As you go down Group 2, ionisation energy decreases.

Ionisation energies across a Period:

• As you go across a Period, electrons are removed from the same electron shell, shielding is the same.

• This means that the No of electrons : protons decreases.

• This increases the attraction pulling the electron shell in slightly.

• This increases attraction.

• This increases the successive ionisation energies

Successive ionisation energies        Click for DEMO

·         The successive ionisation energies can tell us which group an element is in.

·         For potassium the easiest electron(s) to remove before a large jump tells us the group.

·         This can also be done by looking at data. 

        Click for SUMMARY DEMO

Questions: 1 - 3 p41/ 13 p73

Shells and orbitals

Flame colours and emission spectra

See some line emission spectra

Energy levels or shells -

• We have already seen what happens when an electron is removed from an atom - the atom becomes and ion.

• What if we supply enough energy to promote electrons to higher empty electron shells but not remove it?

• One way to look at the arrangement of electrons in an atom is to disturb them and see what happens when they go back to their original arrangement.

• This is done by heating compounds in a non-luminous Bunsen flame and studying the characteristic colours emitted when the electrons fall back to their original arrangement.

• The electrons gain energy from the Bunsen and lose energy (as light ) as they go back to their original arrangement:

• This happens for any electron being promoted to any higher electron shell and falling back to any lower electron shell.

• This gives us the series of lines as electrons give out only specific energy when they move to a lower electron shell.

• A spectroscope shows us the specific energies characteristic to that element.

• Looking at the line emission spectra 2 thing become apparent:

                1)  That the lines are specific colours representing specific energies / electron shells.

                2)  That as the move up in energy (towards the violet end) they get closer together (converge)

• This tells us that the electron shells are not nicely spread out but get closer together as they get further away from the nucleus:

• Your teacher will draw on these diagrams to show how the spectra is formed:

• That the electron shells can be converted into an energy level diagram showing the Principle Quantum Number, n

• The number of electrons in each Principle Quantum Number, n can be calculated:

Atomic orbitals

• At GCSE you will have been given a planetary model of the atom.  This model assumes electrons are solid particles.

• Electrons are not solid particles, (more like a wave packet of energy or electron cloud):

• The very nature of the electron means that electrons cannot orbit around around the nucleus.

• Heisenberg's uncertainty principle states that  we cannot determine both position of the electron is and its momentum.

• This makes it impossible to determine the orbit

• And so the planetary model has been replaced with the orbital model.

• These are regions around the nucleus in which the electron is likely to be found.

• Each electron shell is made up of orbitals.

• Each orbital can hold a maximum of 2 electrons.

• Imagine each electron as an electron cloud with the shape of an orbital.

• Each electron in an orbital would be an electron cloud of the same shape.

• 2 electrons in an orbital would be the same shape but twice as dense.

s - Orbitals:

• These are spherical in shape and represent the likelihood of finding the electron in that region

• All electron shells contain an 's' orbital.

• As each orbital can hold a maximum of 2e the s - orbital has a total of 2e in every electron shell.

p - Orbitals:

• Each p orbital is the shape of a dumb - bell.

• Each p orbital can hold a maximum of 2 electrons.

à

• As there are 3 p orbitals, this means that the p orbitals in a shell can hold a maximum of 6 electrons.

• All electron shells from n = 2 upwards contain 3 p orbitals.

d and f orbitals:

• These orbitals are more complex.  You do not need to know the shapes of these.

• Each shell from n = 3 upwards contains 5 d orbitals (5 x 2 = 10 electrons).

• Each shell from n = 4 upwards contains 7 f  orbitals (7 x 2 = 14 electrons).

A summary of orbitals:

• Each orbital can hold a maximum of 2e

Complete the table below: 

Representing electrons in orbitals

• With different types of orbitals and having different shapes, we represent an orbital with a box.

• As a box is an orbital, each box can hold 2e.

• If electrons are negative and repel each other how can they occupy the same orbital?

• This can only be explained by looking in more detail at the electrons.

• Electrons spin on their axis.

• When anything with a charge spins a magnetic field is produced.

• One of the electrons in an orbital will spin one way and the other spins in the opposite direction.

• This gives opposing magnetic fields which we represent with an arrow.

• This is the only way electrons can share an orbital as there is now some attraction between them.

Questions 1 - 2 p43

Sub - shells and energy levels

Sub - shells

• An electron shell is made up from orbitals.

• Orbitals of the same kind are called sub - shells.

• This means that each sub - shell contains the same type of orbitals, each of which holds a maximum of 2e:

n = 1 shell: maximum 2 electrons
Sub - shell	1s
Orbital
Electrons	2e
n = 2 shell: maximum 8 electrons
Sub - shell	2s	2p
Orbital
Electrons	2e	2e	2e	2e
n = 3 shell: maximum 18 electrons
Sub - shell	3s	3p			3d
Orbital
Electrons	2e	2e	2e	2e	2e	2e	2e	2e	2e

Electrons and the Periodic Table

Need for a more complex model:

• Across a period the general trend is for the 1^st ionisation energies to increase. 

• This is due to increasing nuclear charge on electrons in the same principal quantum number (electron shell).

• Each group of 8 is made up of smaller breaks of 2-3 and 5-6. 

• This tells us that the 8 electrons are not exactly the same (in terms of energies).  If they were we would expect a smooth increase.

a) 2 - 3 split:  s and p orbitals

• The first clear distinction is that the 3rd electron across a Period is easier to remove than the second.

• The 3rd electron is the first electron in the p sub-shell.

• This suggests that this electron is slightly higher in energy (or slightly further from the nucleus) than the 2nd electron in the s subshell.

• This gives us evidence that the sub-shells are similar but not identical in energies:

• This also happens for the n = 3 shell:

b) 5 - 6 split:  filling the p orbital

• The next clear distinction is that the 6th electron across a Period is easier to remove than the second.

• If we remove the s electrons you can see that electrons 1 - 3 in the p sub-shell follow the pattern expected.

• The 4th electron in the p sub-shell is easier to remove than the 3rd

• This gives us a clue to how the p sub-shell is filled:

• The first 3 electrons in the p sub shell spin in one direction and occupy p_x, p_y and p_z.

• After this the electrons have to pair up with those spinning the opposite way.

• In a sub-shell, electrons will remain unpaired in the orbitals until they have to pair up. 

• This is the same in the d orbital.

Electron shells overlap

• From previous knowledge, we know that:

1. The larger the value of n, the higher the energy level and the further from the nucleus.

2. That the electron shells get closer together as you move away from the nucleus.

3. As we move from n = 1 to n = 4, a new type of sub shell is added.

4. Within a shell the energies of the sub - shells increases in the order s, p, d and f.

• After 3p it gets a bit complicated:

The 4s energy level is below the 3d energy level. This means that the 4s orbital fills before the 3d orbitals. (according to the Aufbau Principle)

• Remember the orbitals fill up in energy level order from the bottom up:-

Electron energy levels

• Within an electron shell, the sub - shells have slightly different energies as explained above.

• The sub - shell increases in energy in the order: s, p, d and f.

Filling shells and sub - shells:

• The electron configuration is the arrangement of electrons within an atom.

• This can be worked out by following a set of rules called The Aufbau Principle:

The Aufbau Principle:

1. Electrons are added one at a time to 'build up' the atom.

2. The lowest available energy level fills first.

3. Each energy level must be full before the next, higher energy level can be filled.

4. Each orbital in a sub - shell is filled by single electrons before pairing up.

5. Each orbital can hold 2e of opposite spin

Filling the orbitals - Using the Aufbau Principle:

• The first 3 electrons in the p sub - shell spin in one direction and occupy p_x, p_y and p_z.

• After this the electrons have to pair up with those spinning the opposite way.

• In a sub - shell, electrons will remain unpaired in the orbitals until they have to pair up.  This is the same in the d orbital

Electron configuration

• We use a shorthand to show how the electrons are arranged in an atom.

• Hydrogen is the simplest atom so we will use this to look at the shorthand:

   

  

   

Questions 1 - 2 p45

Sub shells and the Periodic Table:

• Fill in the following table:   

Things to notice:

• The Period the element is in on the Periodic Table is equivalent to its highest Principle Quantum Number, n (and the electron shell of the outer electrons)

• All s orbital elements are in Groups 1 and 2 (except H and He)

• All p orbital elements are in groups 3,4,5,6,7 and 0

Extending further:

• The Periodic Table is arranged in blocks of 2, 6, 10 and 14

• s block elements are all in Gp1 and 2 = 2

• p block elements are all in Gp3,4,5,6,7 and 0 = 6

• d block elements are all in the Transition Metals = 10

• f block elements are all in the Lanthanides and Actinides = 14

Using the Periodic Table for electron configurations

• So, the electron configuration can be worked out from the Periodic Table, filling from left to right then top to bottom:

Other examples:-

1)     Na   1s² 2s² 2p⁶ 3s¹

2)     Sc  1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s²  (remember 4s fills before 3d)

Shortening an electron configuration

• For atoms with many electrons, the electron configuration can be long:

                    Sc          1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s²

• We can write a shorthand version by using the closest noble gas configuration for the inner shells:

                    Sc          [1s² 2s² 2p⁶ 3s² 3p⁶]3d¹ 4s²

                                           Inner shells

                                    [Ar]  has the same electron configuration as the inner shells

                                    [Ar] 3d¹ 4s²    where [Ar] represents the electronic configuration of argon.

Electronic configurations in ions

·        Follow the same principle as for atoms but add / remove electrons:-

Eg Sodium ion Na⁺

                                                         Na atom =   1s² 2s² 2p⁶ 3s¹

To make a sodium ion 1 electron is removed

                                                         Na ion =   1s² 2s² 2p⁶

Eg Chlorine ion Cl^-

                                                         Cl atom =   1s² 2s² 2p⁶ 3s²2p⁵

To make a chlorine ion 1 electron is added

                                                         Cl^- ion =   1s² 2s² 2p⁶ 3s²2p⁶

Things to note:

• All s and p block ions have a noble gas configuration.

• Electrons are taken from 4s before 3d (as it fills 3d before 4s)

Questions  1 - 4 p47 / 1,2 and 13 p73 / 1 p74