1. Electronic Structure of Atoms
atomic radii,
ionisation energies
and trends within groups

2. The size of an atom
It is difficult to measure the size of an atom directly.
The exact position of the furthest electron from the nucleus cannot be found!

3. However.. Scientists can measure the length of a covalent bond between two atoms using various methods
X-ray scattering techniques are a family of non-destructive analytical techniques which reveal information about the crystallographic structure, chemical composition, and physical properties of materials and thin films. X ray diffraction - These techniques are based on observing the scattered intensity of an x-ray beam hitting a sample as a function of incident and scattered angle, polarization, and wavelength or energy.
Elecron diffraction - Electron diffraction is a technique used to study matter by firing electrons at a sample and observing the resulting interference pattern. This phenomenon occurs due to the wave-particle duality, which states that a particle of matter (in this case the incident electron) can be described as a wave. For this reason, an electron can be regarded as a wave much like sound or water waves. This technique is similar to the X-ray diffraction and neutron diffraction.
Electron diffraction is most frequently used in solid state physics and chemistry to study the crystal structure of solids. These experiments are usually performed in a transmission electron microscope (TEM), or a scanning electron microscope (SEM) as electron backscatter diffraction. In these instruments, the electrons are accelerated by an electrostatic potential in order to gain the desired energy and determine their wavelength before they interact with the sample to be studied.
The distance is .074 nm

4. Atomic radii
The atomic radius is half the distance between two atoms of the same element joined together by a single covalent bond.
The ionic radius of an element is its share of the distance between adjacent ions in an ionic solid.
  For example, the distance between the magnesium cation and the oxide anion in magnesium oxide is 211 pm, the radius of the Mg2+ ion is calculated from 205pm pm = 65 pm.
The size of an ion also varies somewhat with its environment so average values for ionic radii are often quoted.

5. It only works if it is the same atoms.. And a single covalent bond

6. The atomic radii of the elements

7. A pattern in atomic radii
Descrease going across a period
Ìncrease going down a group

8. As you go down a group… The atomic radii increases..3
1. the atoms have extra electrons which have to go into a new shell further away from the nucleus. 11
2. the atoms have a bigger nuclear charge. You might expect that this would pull the electrons closer to the nucleus but the electrons in the extra shells are being shielded from the extra nuclear charge by the inner shells of electrons so this doesn’t happen!
“The shielding effect” 19
The atomic radii increases..
Potassium

9. As you go across a period..4 5 3 6
Aluminium
Carbon
Beryllium
Lithium
As you go across a period..
1. the nuclear charge in the atom increases. The increasing nuclear charge pulls the electrons closer to the nucleus
2. the atoms have extra electrons but they do not go into new shells! There is therefore no extra shielding effect as you go across a period!
Due to more protons.
The atomic radii decreases..

10. Ionisation energy
Higher level only
Ionisation energy is the minimum energy needed to remove the most loosely bound electron from a neutral gaseous atom.
In its ground state and isolated,
Not e the position of hydrogen is kind of incorrect

11. Check the ionisation energy
What is the pattern
1) as you go down a group
2) as you go across a period

12. Decreases as you go across any period
Higher level only
Ionisation energy – the pattern
Decreases as you go across any period
Increases as you go down any group
Notice the same trend as before!!!!

13. Ionisation energy – explaining the trend
Higher level only
Ionisation energy – explaining the trend
Increases as you go down any group
Notice the same trend as before!!!!

14. As you go down any group…3
1. The atomic radii increase so the furthermost electron gets further away from the nucleus
2. The atoms have a bigger nuclear charge. But the furthermost electron is shielded from the extra nuclear charge by the inner shells of electrons
“The shielding effect” 11
The ionisation energy decreases.. It gets easier to pull the electron away from the nucleus 19

15. Decreases as you go across any period
Ionisation energy – explaining the trend
Higher level only
Decreases as you go across any period
Notice the same trend as before!!!!

16. As you go across a period..
Lithium
Beryllium
Aluminium
Carbon 3 4 5 6
As you go across a period..
1. the nuclear charge in the atom increases. The increasing nuclear charge pulls the furthermost electron tighter to the nucleus
Due to more protons.
2. The atomic radii decrease so the furthermost electron is closer to the nucleus
The ionisation energy increases.. It gets harder to pull the electron away from the nucleus

17. Plot of atomic number against ionisation energy
The second period
Maybe an activity where they spot the trends and the exceptions and revise what they have just done!
Higher level only

18. Ionisation energies across the second period
Higher level only
Ionisation energies across the second period
Generally they go up - but there are exceptions…
Beryllium and Nitrogen have higher ionisation energies than you would expect.

19. The electronic configuration of:
Higher level only
The electronic configuration of:
Lithium
It has three electrons.
The electronic configuration is 1s2 2s1

20. The electronic configuration of:
Higher level only
The electronic configuration of:
Beryllium
It has 4 electrons.
The electronic configuration:
1s2 2s2
It has a full 2s sublevel and this gives it extra stability!
It would be extra hard to pull the electron away so the ionisation energy is especially high.
More excercises, all 36 elements!

21. The electronic configuration of:
Higher level only
The electronic configuration of:
Boron
It has 5 electrons.
The electronic configuration:
1s2 2s2 2p1
More excercises, all 36 elements!

22. The electronic configuration of:
Higher level only
The electronic configuration of:
Carbon
It has 6 electrons.
The electronic configuration:
1s2 2s2 2p2
More excercises, all 36 elements!

23. The electronic configuration of:
Higher level only
The electronic configuration of:
Nitrogen
It has 7 electrons.
The electronic configuration:
1s2 2s2 2p3
It has a half full 2p sublevel and this gives it extra stability!
It would be extra hard to pull the electron away so the ionisation energy is especially high.
More excercises, all 36 elements!

24. The electronic configuration of:
Higher level only
The electronic configuration of:
Oxygen
It has 8 electrons.
The electronic configuration:
1s2 2s2 2p4
The trend continues for fluorine and neon!

25. Plot of atomic number against ionisation energy
The third period
Maybe an activity where they spot the trends and the exceptions and revise what they have just done!
Higher level only

26. Ionisation energies across the third period
Higher level only
Ionisation energies across the third period
Generally they go up - but there are exceptions…
Magnesium and Potassium have higher ionisation energies than you would expect.

27. The electronic configuration of:
Higher level only
The electronic configuration of:
Magnesium
It has 12 electrons.
The electronic configuration:
1s2 2s2 2p6 3s2
It has a full 3s sublevel and this gives it extra stability!
It would be extra hard to pull the electron away so the ionisation energy is especially high.
The trend continues for fluorine and neon!

28. The electronic configuration of:
Higher level only
The electronic configuration of:
Potassium
It has 15 electrons.
The electronic configuration:
1s2 2s2 2p6 3s2 3p3
It has a half full 3p sublevel and this gives it extra stability!
It would be extra hard to pull the electron away so the ionisation energy is especially high.
The trend continues for fluorine and neon!

29. Check your learning.. What is an energy level? What is an orbital?
What is an energy sublevel?
Define atomic radius
Define ionisation energy

30. Today’s objectives
Second and successive ionisation energies and evidence for energy levels
The chemical properties of the elements based on atomic radii and ionisation energies

31. Second ionisation energies
Higher level only
Second ionisation energies
Def the energy needed to remove the second electron from a singly charged positive ion.
First ionisation energy
needed for:
Be Be+ + e
After the first electron has already been moved!
Second ionisation energy
needed for:
Be Be2+ +e

32. Beryllium: Successive ionisation energies
Higher level only
Beryllium: Successive ionisation energies + 4 4
This cntinues for the third ionisation energies and so on…
Beryllium+ (positive ion)
Beryllium (neutral atom)
The second ionisation energy is always higher than the first because in an ion the atomic radius decreases and the electrons are closer to the nucleus – harder to take away!

33. Beryllium: Successive ionisation energies
Higher level only
Beryllium: Successive ionisation energies
Beryllium
There is a dramatic increase in ionisation energy needed to remove the third electron!

34. Beryllium: Successive ionisation energies
Higher level only
Beryllium: Successive ionisation energies
There is a large ionisation energy needed because the electron is removed from a new energy level!
The electronic configuration:
1s2 2s2

35. Ionisation values provide evidence for the existence of energy levels:
Higher level only
Ionisation values provide evidence for the existence of energy levels:
The ionisation energy required to remove an electron from a new shell jumps dramatically! This is because it is much harder to remove an electron from an inner shell as it is closer to the nucleus and has less shielding from the nuclear charge than before.
Inner shell is also more stable!

36. Chemical Properties of the elements
The chemical properties of each element depends on its electronic structure.
Elements in the same group in the Periodic Table have similar electronic structures – they all have the same number of electrons in their outermost shell - and so have similar chemical properties.
(how they tend to behave in chemical reactions)

37. Periodic table

38. Group 1- The Alkali metals
Why do they always react to lose an electron?
Can you explain why reactivity increases going down the group?
They all have one electron in their outermost shell, which they have a tendency to lose.
Potassium

39. Group 1 – the Alkali metals3
Going down the group the atomic radius increases
Going down the group the nuclear charge also increases, but screening also increases so this effect is cancelled out.
The result is that as you go down the group the outermost electron becomes easier to remove, and the element is therefore more reactive! 11 19
Potassium

40. Alkali metals reacting with water
Lithium
Sodium
Potassium

41. Periodic table

42. Group 17- The Halogens 9 17 24
Why do these elements always gain an electron?
Can you explain why reactivity decreases as you go down this group?
They all have 7 electrons in their outermost shell and have a tendency to gain one electron in chemical reactions

43. Group 17- The Halogens 9
A high nuclear charge and small atomic radius attracts electrons to the halogens
Going down the group the atomic radius increases
Going down the group the nuclear charge also increases, but screening also increases so this effect is cancelled out.
The result is that as you go down the group the ability to gain an electron decreases… and the element is therefore less reactive! 17
Flourine >chlorine>bromine 24