1. 1.4 Learning Outcomes down a group
 explain the general trends in values of atomic radii (covalent radii only)
down a group
across a period (main group elements only)
define and explain first ionisation energy
explain the general trends in first ionisation energy values:
across a period (main group elements) and
explain the exceptions to the general trends across a period
define and explain second and successive ionisation energies  

2. 1.4 b Learning Outcomes
describe how second and successive ionisation energies provide evidence for energy levels
recognise the relationship and trends in successive ionisation energies of an individual element
explain how chemical properties of elements depend on their electronic structure
explain how atomic radius, screening effect and nuclear charge account for general trends in properties of elements in groups I and VII  

3. Trends in the Periodic Table of Elements
Atomic radius
Ionisation energy
Electronegativity
Two reasons required for each trend
- down a group
- across a period

4. Atomic Radius Hard to determine using one atom
Why? can’t define the size of the atom exactly because orbitals are areas of high probability of finding an electron
So position of electron hard to measure
Nucleus easier to measure or pinpoint

5. Atomic Radius [covalent radius]
Half the distance / between the nuclei of two atoms / of the same element / joined by a single covalent bond.
Distance between two hydrogen nuclei
H-H = nm [1 nm = 10-9 m ]
Atomic radius = nm

6. Across a Period [left to right] Atomic Radius gets smaller Reasons
Increasing nuclear charge so electrons pulled in more
No increase in screening effect from electrons already there

7. Trends in the Periodic Table Terms
The nuclear charge [i.e. the charge in the nucleus as a result of the number of protons in the nucleus] – check cartoon
The screening effect of inner electrons - The screening effect is observed in atoms with an atomic number of greater than 2. In these atoms, electrons in inner energy levels shield outer electrons from the attraction of protons in the nucleus

8. Electronic Screening Effect
Electrons on inner sublevels shield the attraction of protons in the nucleus for outer electrons

9. Down a Group Radius gets bigger Reasons
New main energy level [shell] filling
Increased nuclear charge cancelled out to some extent by increased screening effect of an extra layer of e-s in sublevels nearer the nucleus
e- is also further away to begin with.

11. Atomic (Neutral) versus Ionic Size
Metals
Nonmetals
Group 1 Al
143 50 e-
Group 13
Group 17 e- e-
152
186
227 Li Na K 60
Li+ F-
136 F Cl Br 64 99
114 e- e- 95
Na+
Cl-
181
Al3+ e- e-
133 K+
Br-
195
Cations are smaller than parent atoms
Anions are larger than parent atoms

12. First Ionisation Energy
The energy required to completely remove the ‘most loosely bound’ electron from a [mole of] ‘neutral gaseous atom’ [s], in its [their] ‘ground state’
Na(g) = Na+(g) + e kJ mol-1
States must be given (e.g. (g))

13. Across a Period Ionisation energy Increases
(i) Increasing nuclear charge
Number of protons increases so attraction increases
Electrons in same main energy level
(ii) Decreasing atomic radius
Atomic radius decreases so electrons are closer to positive nucleus so held tighter.

14. Down a Group Decreases [gets less]
(i) Increasing atomic radius means electron further from attractive force of the nucleus
(ii) Screening effect of inner electrons Positive charge of nucleus is increasing but inner e- shells shield the outer electron from this increased charge.

15. Exceptions to general trend
Look at the general trend.
Argon has the highest value.
Why?
e- from ‘full’ 3p sublevel
First ionisation Energies
Are any elements higher than general trend?
Yes
Mg and P
Why?
Mg: e- from ‘full’ 3s sublevel
P: e- from ‘half-filled’ 3p sub-level
Must state 3 in all cases

16. Second Ionisation Energy
The removal of the second electron from a monopositive positive ion (formed when the first electron is removed.)
Na+(g) = Na2+(g) + e-
(States must be given)

17. Further evidence for the existence of Energy Levels
Take a Mg atom and remove the electrons one by one using energy – measure the amount used in each case.
Plot a graph of the successive ionisation energies.
Use log10 because values so big.

18. Being held by average of 1 proton 2 layers of shielding
Remove first electron
Being held by average of 1 proton
2 layers of shielding e-
Energy
Requird 1 2 3 4 5 6 7 8 9 10 11 12
738kJ
12+

19. Mg+ Remove first electron Being held by average of 1 proton
2 layers of shielding e-
Energy
Requird 1 2 3 4 5 6 7 8 9 10 11 12
738kJ
12+
Notice that the radius of the ion has decreased
Mg+

20. Mg2+ Remove second electron
Energy
Requird 1 2 3 4 5 6 7 8 9 10 11 12
Remove second electron
Being held by average of 1.1 protons (Increased effective nuclear charge)
Still 2 layers of shielding
Atom slightly smaller
Harder to remove [+712]
738kJ
1450 kJ
12+
Notice that the radius of the ion is much smaller
Mg2+

21. Remove third electron
Being held by average of 1.2 protons (Increased effective nuclear charge)
Now only 1 layer of shielding
Radius now much smaller
Removing e- from a full shell
Much harder to remove [+6282] e-
Energy
Requird 1 2 3 4 5 6 7 8 9 10 11 12
738kJ
1450 kJ
7732 kJ
12+

22. Remove fourth electron
Being held by average of 1.33 protons (Increased effective nuclear charge)
Now only 1 layer of shielding
Atom a bit smaller
Harder to remove [+ 2808] e-
Energy
Requird 1
738 kJ 2
1 450 kJ 3
7 732 kJ 4 5 6 7 8 9 10 11 12 kJ
12+

23. Huge jump e- removed from full n = 1 main level
Energy [log]
Huge jump e- removed from full n = 1 main level
135000
6000
Big jump e- removed from full n = 2 main level
700
Electrons removed

24. Successive Ionization Energies
Potassium - 1s2,2s2,2p6,3s2, 3p6, 4s1
1s2
Very Important: Proof of Existence of E levels
2s1
2p6
3s1
3p6
4s1

25. Trends in Electronegativity Values
Electro negativity: Is the relative attraction of an atom for shared pairs of electrons in a covalent bond.
The most commonly used scale is the Pauling Scale which runs from zero [ 4 is high electronegativity] to four

26. Going across a period, Electronegativity increases

27. WHY ???
Increasing nuclear charge
Decreasing atomic radius

28. Going down a group, Electronegativity decreases

29. WHY ???
Increasing atomic radius
Increasing Screening effect

30. Electronegativity Values

31. Summary of Trends

32. Exam Q’s 2002

33. Exam Q’s 2002

34. Exam Q’s 2004

35. Exam Q’s 2004

36. Exam Q’s 2009

37. Exam Q’s 2009

38. Exam Q’s 2009

39. Exam Q’s 2013

40. Exam Q’s 2013

41. Exam Q’s 2013

42. Exam Q’s 2014
Bohr
E level: The definite energy of an electron in an atom / shell / orbit
Orbital: region where the probability of finding an electron is 95%
1s2, 2s2, 2p6, 3s2, 3p2

43. Exam Q’s 2014
Write the electron configuration (s, p) of an
atom of silicon showing the distribution of
electrons in atomic orbitals in the ground state. (6)
Hence, state how many (i) main energy levels,
(ii) atomic orbitals, are occupied in the silicon
atom in its ground state.
1s2, 2s2, 2p6, 3s2, 3p2
(i) main energy levels = 3
(ii) atomic orbitals = 8

44. Exam Q’s 2014
(b) Define first ionisation energy. (6)
Explain why the first ionisation energy value of silicon is:
(i) greater than that of aluminium,
(ii) less than that of carbon. (9)
The minimum energy required to remove the most loosely bound electron // from a gaseous atom in its ground state
i)Greater nuclear charge // smaller atomic radius
ii)Greater atomic radius so most loosely bound electron is further from nucleus // greater screening as more shells

45. Exam Q’s 2014 Line emission spectra
Sharp increase from 4th to 5th shows that this is the first electron to be removed from the second shell
Sharp increase from 12th to 13th shows that this is the first electron to be removed from the first shell
Line emission spectra