1. Trends in the Periodic Table

2. Learning Outcomes Atomic radii
Explanations for general trends in values: (i) down a group (ii) across a period (main group elements only).
Electronegativity

3. Learning Outcomes First ionisation energies.
Explanations for general trends in values: (i) down a group (ii) across a period (main group elements)
Explanations for exceptions to the general trends across a period.
Second and successive ionisation energies.
Evidence for energy levels provided by successive ionisation energy values.
Trend in Chemical Reactivity
Of Alkali metals and Halogens

4. The Periodic Table- Vocab
Periods
Rows in the periodic table
Numbered from 1 to 7.
Groups
Vertical columns of elements with similar chemical properties
Periods
Groups

5. 1. Atomic Radius

6. Atomic Radius
Half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond

7. 1. Atomic Radius- How is it measured?
Not possible to measure the distance from nucleus to electron (Heisenberg Uncertainty Principle).
Distance between two nuclei is measured using X-Rays. This is the bond length.
No values for Nobel Gases as they do not form bonds together.

8. 1. Atomic Radius
What happens to the atomic radius as you move down the groups?
What happened to the atomic radius as you move across the periods?
Think about:
Number of electrons
Attraction between electrons and protons
Energy levels

9. 1. Atomic Radius

10. 1. Atomic Radius Atomic Radius increases New energy level
Screening effect

11. Atomic Radius increases down the groups
New energy levels
Additional electrons are going into a new energy level which is further away from the nucleus

12. New energy levels down a Group…H Li Be Na Mg Li 3P 4N Na
11P
12N

13. Li 3P 4N Na
11P
12N
Going down a group, there are additional energy levels to add to the size of the atom, thus the atomic radius increases

14. Atomic Radius increases down the groups
New energy levels
Additional electrons are going into a new energy level which is further away from the nucleus
Screening Effect
Electrons in the inner energy level/ levels shield the outer electrons from the positive charge in the nucleus

15. Screening effect

16. Screening effect

17. 1. Atomic Radius Atomic Radius decreases
No increase in screening effect
Increase in effective nuclear charge
Atomic Radius increases
New energy level
Screening effect

18. Atomic Radius decreases across a period
No increase in screening effect
Extra electrons added when moving across a period go into the same outer energy level
 no increase in screening
Increase in effective nuclear charge
When the number of protons in the nucleus is increasing without an increase in the screening effect
 There is a greater attractive force on the outer electrons

19. Across a Period… H Li Be Na Mg Na
11P
12N Mg
12P
12N

20. Na
11P
12N Mg
12P
Adding one proton increases the charge of the nucleus to pull the electrons towards them, therefore across a period, the atomic size decreases

21. 1. Atomic Radius Atomic Radius decreases
Increase in effective nuclear charge
No increase in screening effect
Atomic Radius increases
New energy level
Screening effect

22. p.95 Q7.2
Describe and account for the trend in atomic radii of the elements
Across the second period
Down any group, of the period table
The atomic radius decreases as you go across the second period from Li to F.
2 reasons: increasing effective nuclear charge as you go across the period means the outer electrons are held closer to the nucleus and no increase in screening effect
(ii) Increases as you go down any group as electrons are going into a new shell and the inner electrons screen the outer electrons.

23. 2. Electronegativity

24. Electronegativity
The relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond
Imagine a tug-of-war for the electrons

25. 2. Electronegativity
What happens to electronegativity as you move down the groups?
What happened to the electronegativity as you move across the periods?
Think about:
Atomic radius
Screening effect

26. 2. Electronegativity

27. 2. Electronegativity Electronegativity decreases Increasing atomic
radius
2. Screening effect

28. Electronegativity decreases down the groups
Increasing atomic radius
As atomic radius increases, outer electrons are further away from the attractive force of the nucleus
 smaller attraction between the nucleus and the shared pair of electrons
Screening Effect
Electrons in the inner energy level/ levels shield the outer electrons from the positive charge in the nucleus

29. 2. Electronegativity Electronegativity increases
Increase in effective nuclear charge
Decreasing atomic radius
Electronegativity decreases
Increasing atomic
radius
2. Screening effect

30. Electronegativity increases across a period
Increase in effective nuclear charge
Number of protons in the nucleus increases across a period
 Increasing attraction between outer electrons and nucleus
Decreasing atomic radius
Atomic radius decreases across a period
Outer electrons become closer to the nucleus
Greater attraction between outer electrons and nucleus

31. Across a Period… H Li Be Na Mg Na
11P
12N Mg
12P
12N

32. 2. Electronegativity Electronegativity increases
Increase in effective nuclear charge
Decreasing atomic radius
Electronegativity decreases
Increasing atomic
radius
2. Screening effect

33. 3. Ionisation Energy

34. First Ionisation Energy
Ionisation energy of an atom is the minimum energy required to completely remove the loosely bound electron from a neutral gaseous atom in its ground state

35. 3. First Ionisation Energy
Energy needed to remove the outermost electron

36. 3. First Ionisation Energy
X(g) → X+(g) + e-
The units of ionisation energy are kilojoules per mole.
The second ionisation energy is the energy required to remove a second electron from an atom.
X+(g) → X2+(g) + e-
NB: need to include (g) symbol as definition refers to ‘gaseous state’

37. 3. First Ionisation Energy
What happens to the ionisation energy as you move down the groups?
What happened to the ionisation energy as you move across the periods?
Think about:
Where are the ‘most loosely bound electrons’ located?
Atomic radius
Screening effect
Effective nuclear charge

38. 3. First Ionisation Energy

39. 3. First Ionisation Energy
First Ionisation Energy decreases
Increasing atomic
radius
2. Screening effect

40. First Ionisation Energy decreases down the groups
Increasing atomic radius
As atomic radius increases, outer electrons are further away from the attractive force of the nucleus
 easier to remove an electron from outer energy level
Screening Effect
Electrons in the inner energy level/ levels shield the outer electrons from the positive charge in the nucleus

41. 3. First Ionisation Energy
First Ionisation Energy increases
Increase in effective nuclear charge
Decreasing atomic radius

42. First Ionisation Energy increases across a period
Increase in effective nuclear charge
Number of protons in the nucleus increases across a period
Increasing attraction between outer electrons and nucleus
Harder to remove electron from outer energy level
Decreasing atomic radius
Atomic radius decreases across a period
Outer electrons become closer to the nucleus
Greater attraction between outer electrons and nucleus

43. 3. First Ionisation Energy
First Ionisation Energy increases
Increase in effective nuclear charge
Decreasing atomic radius
First Ionisation Energy decreases
Increasing atomic
radius
2. Screening effect

44. First Ionisation Energy Exceptions to the Trend
Why?? N P Be Mg

45. First Ionisation Energy Exceptions to the Trend
Be has a high value as the electron is coming from a full sub-level.
Be = 1s2, 2s2.
Mg has a high value as the electron is also coming from a full sub-level.
Mg = 1s2, 2s2, 2p6, 3s2.

46. First Ionisation Energy Exceptions to the Trend
N has a high value as the electron is coming from a half-filled sub-level.
N = 1s2, 2s2, 2p3
P has a high value as the electron is also coming from a half-filled sub-level.
P = 1s2, 2s2, 2p6, 3s2, 3p3

47. Second Ionisation Energy
Potassium has a low first ionisation energy but a very high second ionisation energy. Why?
The first electron is being removed from the 4th shell where it is easily removed so K can have a full outer shell.
The second is very high as it is coming from a full shell and is hard to overcome the nuclear charge.

48. LCH Exam Q
The table shows the first and second ionisation energies of nitrogen, oxygen, neon and sodium
Account for the decrease in ionisation energy between nitrogen and oxygen
Explain why the second ionisation energy of sodium is significantly (about nine times) higher than the first, while the increase in the second ionisation energy of neon compared to its first is relatively small (less than twice the first)
Element
First Ionisation Energy
(kJ mol -1)
Second Ionisation Energy
Nitrogen
1400
2860
Oxygen
1310
3390
Neon
2080
3950
Sodium
494
4560

49. LCH Exam Q
The table shows the first and second ionisation energies of nitrogen, oxygen, neon and sodium

50. Trends within Groups
The chemical properties of a group are determined by how many electrons they have in their outside shell.
Group 1 = 1
Group 2 = 2 etc…….

51. Group 1- Alkali Metals Properties: Very reactive – stored under oil
Extracted from compounds or ores
Low ionisiation energy and electronegativity value.
Form ionic bonds as tend to loose electrons
Reactivity increase down the group Li Na K Rb Cs
Increasing Reactivity

52. Group 1- Alkali Metals 2K + ½ O2 → K2O
Group 1 react with oxygen to form oxides.
2Li + ½ O2 → Li2O
2Na + ½ O2 → Na2O
2K + ½ O2 → K2O

53. Group 1- Alkali Metals
Group 1 react with water to form hydroxides and hydrogen gas.
Li H2O → LiOH ½H2
Na + H2O → NaOH + ½H2
K H2O → KOH ½H2

54. Group 1- Alkali Metals
Group 1 react with HCl to form chlorides and hydrogen gas.
Li + HCl → LiCl ½H2
Na + HCl → NaCl + ½H2
K + HCl → KCl ½H2

55. Group 7- Halogens
Highly electronegative and most reactive on Periodic Table
Do not exist free in nature.
F is most reactive than Cl, Br, I, etc.
The reactivity decreases down the group. F Cl Br I At
Increasing Reactivity

56. Group 7- Halogens
Boiling point increase down the group as Van der Waals forces exist between atoms and bigger atoms need more energy to break them apart.

57. Noble Gases Inert gases – do not form compounds
Boiling point increase as you go down the group due to Van der Waal forces exist between atoms and they are stronger between bigger atoms.