1. Trends in the periodic table

2. Atomic radius

3. Atomic radii trends and explanations Atomic radius decreases across a period because each successive element has one more proton in its nucleus and one more electron is added to the same valence shell. Therefore, each electron experiences a greater effective nuclear charge so is attracted more strongly to the nucleus resulting in a smaller atomic radius. Atomic radius increases down a group because as we progress down a group the valence electrons are found in another shell much further from the nucleus.

4. Ionic radius

5. Atomic vs ionic radius Cations: When cations form, all the valence electrons are removed from the outer shell, so the ions have one less shell than the atom. This results in a smaller radius than the atom. Anions: When anions form, electrons are added to the existing valence shell. Greater repulsion between valence electrons results in a larger radius than the atom.

6. Ionic radii trends and explanations Ionic radius increases down a group for the same reason atomic radius increases. Ionic radius decreases across a period for the cations for the same reason atomic radius increases. However, here is a big jump in ionic radius between cations and anions because the anions have one more shell than the cations. The trend resumes for the anions.

7. 1 st ionization energy

8. 1 st ionization energy definition 1 st ionization energy = amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms (unit is kJ/mol). Endothermic process. Example equation: Na (g) Na + (g) + e -

9. Trends and anomalies 1 st I.E increases across period because as effective nuclear charge increases, valence electrons are held more tightly, so more energy is needed to remove an electron. Anomalies exist in trend due to more stable electron configurations which require more energy to remove electrons from half full or completely full subshells. 1 st I.E increases down a group because valence electrons are found in a shell much further from the nucleus and are also shielded from the nucleus so are not held as tightly and require less energy to remove.

10. Successive ionization energies e.g sodium

11. Explanation of sodium successive I.E Use the graph on the previous slide to explain the successive ionization energies for sodium. In your answer you should: Describe what 1 st, 2 nd, 3 rd etc ionisation energy means. Describe the overall trend. Explain the jumps in I.E by referring to which shells and sub-shells electrons are being removed from.

12. Electronegativity A measure of the attraction an atom has for electrons in a bond. The four most electronegative elements are F, O, N and Cl. These elements have a high attraction for electrons in a bond because their atomic radius is relatively small and they have a high effective nuclear charge. Essentially F, O, N and Cl nuclei attract a pair of bonding electrons more strongly than the nuclei of other elements they are bonded to. Their nuclei are closer to the bonding electrons and they have a higher effective nuclear charge to attract those electrons.

13. Electronegativity and polarity of bonds The difference in electronegativity helps to determine whether a bond will be polar or not.