1. Explaining periodicity
When asked to compare atoms or ions of different elements start by comparing the fundamental factors and determine how they affect the other properties.
What are the fundamental factors?

2. Fundamental Factors Nuclear charge of atom or ion (number of protons)
Effective nuclear charge (= atomic number – number of electrons on filled inner principal energy levels)
energy level being filled/ number of filled energy main energy levels/ shielding effect
number of electrons: the more electrons , the greater the repulsion

3. Atomic Radius
Describe the trends in the atomic radius across period 3 and also down group 1 You need to refer to all three fundamental factors

4. Across period 3 from left to right
Atomic size/radius decreases as you move across because:
the nuclear charge increases;
electrons are added to same main energy level /similar shielding effect as the number of complete inner energy levels is the same.
As a result of the above two factors the effective nuclear charge increases and the outer electrons are attracted more strongly.
The increased repulsion between the electrons which would cause the radius to increase is cancelled out by the increased nuclear charge.

5. Down the group Increases as you go down because:
the number of main energy levels increases (more electrons) which cancels out the increase in nuclear charge; outermost electron is placed in a higher energy level which is further away from the nucleus.
the shielding effect increases as you go down as there are more filled energy levels.

6. Ionic radius The atomic radius changes when atoms form ions
Positive ions always have a smaller ionic radius that the original atom.
Because: the loss of electron(s) means that the remaining electrons each have a greater share of the positive charge of the nucleus so are more tightly bound
And when an ion in formed, a whole ion shell is usually lost

7. Ionic Radius
Negative ion has a larger atomic radius than that of the original atom
even though the extra electrons are in the same electron shell, the addition of the negative charge means that the electrons are less tightly bound to the nucleus
So the atomic radius is larger

8. Questions Arrange the following from shortest to longest atomic radium
Sodium, Lithium, Potassium
Magnesium, sulfur, Germanium
Argon, Potassium or calcium
Cl-, S2-, P3-, Ar
Na+, Ne or F-

9. Answers Li  Na  K S  Mg  Ge Ar  Ca  K Ar Cl-  S2-  P3-
Na+  Ne  F-

10. Electron Affinity
Electron affinity of an element is the enthalpy change that occurs when one electron is gained by each atom in a mole of gaseous atoms of the element to give one mole of gaseous ions, each with a single negative charge (at standard temperature and pressure).
Electron affinity can be negative (usually e.g. halogens) which means it is an exothermic process as energy is released or it can be positive (like some group 2 elements or the noble gases in group 18) indicating an endothermic process.
The equation defines first electron affinity: X (g) e  X- (g)
The more negative the electron affinity value the greater the ability of the atoms of the element to accept electrons.

11. Horizontal trend
Overall electron affinities become more negative as you go across a period because when moving to the across the periods:
the positive nuclear charge increases;
the atomic radius decreases so the nucleus can attract other electrons better;
the difference in shielding effect is minimal as all have the same number of filled inner shells.

12. Vertical trend
Trends differ in different groups but generally electron affinity values become less negative as you go down a group because ...
the number of main energy levels increases, increasing the shielding effect on free electrons (=electrons part of another atom),
increased nuclear charge but this is cancelled out by the increased shielding – similar effective nuclear charge;
increased atomic radius;

13. Electronegativity
Electronegativity is the ability for an atom to attract a bonding pair of electrons (a shared pair in a diatomic covalent bond). The more commonly used scale is the Pauling scale in which all values are measured relative to fluorine that has the maximum electronegativity of 4.0.

14. Horizontal trend
Increases when moving to the across the periods because:
the positive nuclear charge increases;
the atomic radius decreases so the nucleus can attract other electrons better;
the difference in shielding effect is minimal as all have the same number of filled inner shells.

15. Vertical trend Decreases as you go down groups:
the number of main energy levels increases, increasing the shielding effect on free electrons (=electrons part of another atom),
increased the atomic radius;
the two above cancel out the increased nuclear charge.

16. Ionization energy
Ionization energy refers to the minimum amount of energy required to remove an electron from one mole of gaseous atoms (or ions); it is measured in kilojoules per 1 mole and is defined by the following equation:
atom in ground state (g) + IE atom + (g) e -
Ionization energy has a positive value as it is an endothermic process – energy is needed.

17. Horizontal trend
Overall trend = ionisation energy increases when moving across because more energy is needed because:
the increased nuclear charge;
smaller atomic radius (outermost electron closer to nucleus);
electrons go in the same energy level (similar shielding effect).
The result is a stronger attraction that pulls the valence electrons closer to the nucleus/stronger attraction.

18. Vertical trend
Ionisation energy decreases as you go down groups because:
Atomic radius increases/outer electron to be removed is further away from nucleus, reducing the attraction between the valence electrons and the nucleus (this offsets the increased nuclear charge).
Increased shielding as there are more energy levels so there is less effective nuclear charge; increased shielding also offsets increased nuclear charge.

19. Summary physical property horizontal trend vertical trend
atomic radius
decreases
increases
ionization energy
electronegativity
cation radius
anion radius
electron affinity

20. Summary questions

22. On mini white boards

23. Answers A B C

24. Questions

25. Answers D