[3.2.1] Periodicity

Question 1

How are elements classified on the period table?

Mark on the periodic table where these blocks can be found.

Answer 1

• Classified as s, p or d block elements according to which orbitals the highest energy electrons are in.
  
  • S block elements have their outer electron filling an s-sub shell..
  • P block elements have their outer electron filling a p-sub shell.
  • D block elements have their outer electron filling a d-sub shell.

Question 2

Describe and explain the trend in atomic radius from Na-Ar i.e. period 3.

(Is the trend the same or different for period 3 elements?)

Answer 2

• Atomic radii decrease from left to right across a period because the increased number of protons create more positive charge attraction electrons which are in the same shell with similar shielding.
• (Exactly the same trend in period 2).

Question 3

Describe and explain the trend in the first ionisation energy from Na-Ar i.e. period 3.

Answer 3

• There is a general trend across to increase.
• This is due to increasing number of protons as the electrons are being added to the same shell.

Question 4

Using the graph below, explain the exceptions in the trend in ionisation energy from Na-Ar i.e. period 3.

(Are you able to apply these same explanations to exceptions in ionisation energy in period 2?)

Answer 4

SMALL DROP BETWEEN MG AND AL

• Mg has its outer electrons in the 3s subshell, whereas Al is starting to fill the 3p subshell.
• Al’s electron is slightly easier to remove because the 3p electrons are higher in energy.
• Thus Al has a lower ionisation energy than Mg.

SMALL DROP BETWEEN P and S

-S’s outer electron is being paired up with another electron in the same 3p orbital.
- When the second electron is added to an orbital, there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove.
- Thus S has a lower ionisation energy than P.

(Exactly the same trend in period 2 with drops between Be & B and N to ) for same reasons - make sure to change 3s and 3p to 2s and 2p in explanation)

Question 5

Describe and explain the trend in the melting & boiling points from Na-Ar i.e. period 3.

Refer to the structure and bonding of the elements.

(Is the trend similar in period 2?)

Answer 5

Na, Mg & Al

• Metallic bonding.
• Gets stronger the more electrons there are in the outer shell that are released to the sea of electrons.
• A smaller-sized ion with a greater positive charge also makes the bonding stronger.
• Higher energy is needed to break bonds and thus Na, Mg & Al have high melting and boiling points.
• In order of increasing melting and boiling points: Na > Mg > Al.

Si

• Macromolecular.
• Many strong covalent bonds between atoms.
• High energy is needed to break bonds and thus Si has very high melting and boiling points, higher than that of Na, Mg & Al.

Cl₂ (g), S₈ (s) & P₄ (s)

• Simple molecular.
• **Weak van der Waals **between molecules.
• Little energy needed to overcome intermolecular forces and thus Cl₂ (g), S₈ (s) & P₄ (s) all have low melting and boiling points (lower than Na, Mg, Al & Si).
  
  • S₈ has a higher melting point than P₄ because it has more electrons so has stronger van der Waals between molecules.

Ar

• Monoatomic.
• Weak van der Waals between atoms.
• Lowest melting and boiling point out of all period 3 elements.#

(Similar trend in period 2:

• Li & Be - metallic bonding - high mp & bp.
• B & C - macromolecular - very high mp.
• N₂ & O₂ - molecular - low mp & bp due to small vdw.
• Ne - monoatomic gas - low mp.)