2. Periodicity Physical Properties

3. Ionisation energies Li Rb Kr K Ar Na Ne He

4. The elements in group 1 have the lowest value in each period. As we descend group 1 from Li  Cs the values decrease as the outer electron is further from the nucleus and already in a higher energy level so less energy is needed to remove it. Across a period as each energy level is being filled, an extra proton is also being added to the nucleus. As each electron is added it is more strongly attracted to the increased charge on the nucleus so becomes lower in energy and is harder to remove. The exceptions in this trend are due to the existence of sub- levels within the main energy level. The noble gases have the highest first ionisation energy in each period.

5. Electronegativity Do not learn these values!

6. Group 0 do not have values as they do not form compounds. The value of the electronegativity is related to the size of the atom. Generally, as the atoms become smaller, the nucleus will tend to attract an electron pair more strongly.

7. Atomic radius It is actually impossible to measure atomic radius. Why? What is used instead of an actual the radius of an individual atom?

8. Going down a group the atomic radius increases. This is because the outer electron is in an energy level that is progressively further from the nucleus. e.g Li: 2, 1Na: 2, 8, 1K: 2, 8, 8, 1 etc. The atomic radius decreases across a period. As the number of outer electrons increases so does the number of protons in the nucleus. This increase in charge on the nucleus increases the attraction to the outer energy level. This results in the outer energy level becoming closer to the nucleus.

9. Ionic Radius When positive ions are formed, the radius becomes smaller. There is one fewer electrons than protons so the nucleus attracts the remaining electrons more strongly and there is one fewer energy level. When negative ions are formed there are more electrons in the outer shell so more electron – electron repulsion. As the number of protons remains the same, each electron is less strongly attracted.

10. For both negative and positive ions the size of the ion increases down the as the outer energy level is progressively further from the nucleus. Across a period: For both positive and negative ions we need to compare isoelectronic ions (same number of electrons) All have 10 electrons Number of protons increases from Na + (10 p) to Al 3+ (13 p). So as we move across the period the ions become smaller as there is a stronger attraction for outer energy level by the increasing number of protons.

11. All of these ions have 18 electrons but the number of protons increases from P 3- (15 p) to Cl - (17 p). So the ions decrease in size as there is an increase in the attraction between the outer energy level and the increasing nuclear charge.

12. Melting Points Down a group: ElementMelting point / o C Li181 Na98 K64 Rb39 Cs29 ElementMelting point / o C F2F2 -220 Cl 2 -101 Br 2 -72 I2I2 114 There is a decrease in melting point because as the atoms get larger, the forces of attraction between them decrease. Metallic bonding increase with valence electrons. There is an increase in melting point as the molecules get bigger down the group due to increased van der Waals forces (see later).

13. Across a period:

14. Linked to the bond strength and structure of the elements. Na  Al These elements are metals. The melting and boiling points increase from sodium to aluminium. This is because the atoms are smaller and have an increasing nuclear charge so there is a stronger attraction for the delocalised electrons. Therefore, the strength of the metal-metal bonds increase. Alternatively: Increase in charge of metal ion  increase in attraction for delocalised electrons. Silicon has the highest m.p. and b.p. as it is a macromolecule with a diamond-like structure. Strong covalent bonds link all of the atoms in 3 dimensions. A large amount of energy is required to break these bonds.

15. Phosphorus, sulphur and chlorine are all molecular substances. The m.p.of each is determined by the strengths of van der Waals forces (which in turn depends upon size of the molecule). Each has a low m.p. as the van der Waals forces are weak. n.b P 4 but S 8 and then Cl 2 ! Argon has the lowest m.p and b.p as it exists as single atoms. It has few electrons so and van der Waals forces are very weak.