periodicity (inorganic)

Question 1

How are elements in the periodic table arranged?

Answer 1

Elements are arranged according to their proton number

Question 2

What is a period on the periodic table?

Answer 2

The horizontal rows

Question 3

What is a group on the periodic table?

Answer 3

The vertical columns

Question 4

What does the group number indicate on the periodic table?

Answer 4

The number of outer electrons of an electron

Question 5

What are the 4 blocks of the periodic table?

Answer 5

s-block
p-block
d-block
f-block

Question 6

What elements are in each block of the periodic table?

Answer 6

s-block = groups 1 and 2.
p-block = groups 3 to 0.
d-block = transition metals.
f-block = radioactive elements

Question 7

What is periodicity?

Answer 7

The study of trends within the periodic table. Often these trends are linked to elements’ electronic configurations

Question 8

What is the trend in atomic radius along a period?

Answer 8

Along a period, atomic radius decreases

Question 9

Why does the atomic radius decrease along a period?

Answer 9

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

Question 10

What is the trend in atomic radius going down a group?

Answer 10

Going down a group, atomic radius increases

Question 11

Why does atomic radius increase down a group?

Answer 11

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

Question 12

What is the trend in ionisation energy along a period?

Answer 12

Along a period, ionisation energy increases

Question 13

Why does ionisation energy increase along a period?

Answer 13

It increases because atomic radius decreases, hence nuclear change increases. This means that the outer electrons are held more strongly so more energy is required to remove the outer electron and ionise the atom

Question 14

What is the trend in ionisation energy going down a group?

Answer 14

Going down a group, ionisation energy decreases

Question 15

Why does ionisation energy decrease going down a group?

Answer 15

The nuclear attraction between the nucleus and outer electrons reduces and shielding also increases. Both of these factors mean less energy is required to remove the outer electron

Question 16

What does the melting point of period 3 elements depend on?

Answer 16

The structure of the element.
The bond strength

Question 17

What happens to melting points across period 3 between sodium and aluminium?

Answer 17

Sodium, magnesium and aluminium are all metals Mg=+2, Al=+3). This means more electrons are with metallic bonding. Their melting points increase due to greater positive charge of their ions (Na=+1, released in the form of delocalised electrons. This increases the attractive electrostatic forces from Na to Al, therefore more energy is needed to break them

Question 18

Why does the melting point increase dramatically for silicone in period 3?

Answer 18

Silicon has a very strong covalent structure so more energy is required to break the strong covalent bonds - giving it a very high melting point

Question 19

Why does the melting point decrease in period 3 between phosphorus and chlorine?

Answer 19

Phosphorus, sulfur and chlorine are all simple covalent molecules held with weak van der Waals forces. Less energy is needed to overcome these weak intermolecular forces, so these molecules have relatively low melting points.

Question 20

Why does argon have an even lower melting point than chlorine?

Answer 20

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. As a result, less energy is needed to overcome these weak van der waals forces so argon exists as a gas at room temp.