≫3.2 - Periodicity ✔

Question 1

How does the periodic table arrange the known elements?

Answer 1

*According to proton number.
*All elements along a period have the same number of electron shells.
*All elements down a group have the same number of outer electrons.

Question 2

Elements are classified into blocks that show electron configurations, give the elements in each of the following blocks…
≫S-Block:
≫P-Block:
≫D-Block:

Answer 2

*Groups 1 and 2 + H₂ and He
*Groups 3 to 0.
*Transition metals.
*Radioactive metals.

Question 3

What is the trend of atomic radius along a period and explain why:

Answer 3

*Atomic radius decreases.
*This is due to an increased nuclear charge for the same number of electron shells.
*The outer electrons are pulled closer in to the nucleus as the increased charge produces a greater attraction.
*As a result, the atomic radius for that element is reduced.

Question 4

What is the trend of atomic radius down a group and explain why:

Answer 4

• Atomic radius increases.
• With each increment down a group an electron shell is added each time.
• This increases the distance between the outer electrons and the nucleus, reducing the power of attraction.
• More shells also increases electron shielding where the inner shells create a ‘barrier’ that blocks the attractive forces. Therefore the nuclear attraction is reduced further and atomic radius increases.

Question 5

Describe the trend of ionisation energy along a period:

Answer 5

• Ionisation energy increases.
• The decreasing atomic radius and increasing nuclear charge means that the outer electrons are held more strongly and therefore more energy is required to remove the outer electron and ionise the atom.

Question 6

Describe the trend of ionisation energy down a group

Answer 6

• Ionisation energy decreases.
• The nuclear attraction between the nucleus and outer electrons reduces due to increasing atomic radius and increasing amounts of shielding means less energy is required to remove the outer electron.

Question 7

Melting points across period 3

Describe the melting points across from sodium to aluminium and explain why:

Answer 7

• Melting points increase from Na-Mg-Al.
• These are all metals with metallic bonding and so MP increases due to greater positively charged ions (Na⁺, Mg²⁺, Al³⁺)
• This also means more electrons are released into the ‘sea’ so the attractive electrostatic forces increase.

Question 8

Melting points across period 3

Describe the melting point of silicon and explain why:

Answer 8

• Very high melting point.
• Silicon has a giant macromolecular structure meaning it has very strong covalent bonds between each atom requiring vast amounts of energy to break.

Question 9

Melting points across period 3

Describe the melting points across from phosphorus to chlorine and explain why:

Answer 9

• Relatively low similar melting points.
• Phosphorus, sulphur and chlorine are all simple covalent molecules held together with weak Van der Waals forces.
• These IMF don’t require much energy to overcome.
• Sulphur has the highest as it exists as S₈ meaning its large and so there are more electrons making stronger VDW forces compared to P₄ and Cl₂

Question 10

Melting points across period 3

Describe the melting point of argon and explain why:

Answer 10

• Melting point is very low and it is a gas at room temperature.
• Argon is a noble gas that exists as individual atoms with a full outer shell of electrons.
• This makes the atom very stable and the Van der Waals forces between them very weak requiring little energy to overcome.

Question 11

Sketch the graph of period 3 elements against melting point:

Answer 11

Question 12

What is periodicity?

Answer 12

• The repeating patterns of physical and chemical properties or reactions.