Periodicity

Question 1

How is the periodic table arranged?

Answer 1

Arranged in increasing atomic number.

Question 2

What are metalloids/semimetals? (2)

Answer 2

⇒ Elements that have a combination of metallic and non metallic properties (1)
⇒ These elements lay on the “staircase line” (1)
eg. Silicon

Question 3

How are elements classified in the periodic table? (2)

Answer 3

⇒ s(sharp), p(principal), d(diffuse) block or f(fine)-block (1)
⇒ The block they are in depends on, which orbital their highest energy electrons are in (valence electron) (1)

Question 4

What is the reactivity trend down groups? (3)

Answer 4

⇒ Metals get more reactive going down the group(1)
⇒ Non metals get less reactive going down the group(1)
⇒ Transition elements are unreactive(1)

Question 5

What are Lanthanides, what ions do they form? (2)

Answer 5

⇒ Rare metals (1)

⇒ Tend to from 3+ ions (1)

Question 6

What are Actinides? (1)

Answer 6

⇒ Radioactive metals

Question 7

What are periods?

Answer 7

Horizontal rows of elements, elements of that group will have the same number of electron shells/similar shielding

Question 8

What are groups?

Answer 8

Vertical columns where elements of that group will have the same number of electrons in their outermost shell.

Question 9

Explain the position of H and He. (4)

Answer 9

⇒ [He]: above noble gases due to its property (full outer shell,unreactive) (1)
⇒ - However it is not a p-block element, the highest energy electron is in s-block (1s2) (1)
⇒ [H]: placed on its own (1)
⇒ Can form H+ ion or H- ion but it is a gas, not a metal (1)

Question 10

Explain properties of Group 1, 2 and 3 elements. (2)

Answer 10

⇒ They are metals, have giant structures (1)

⇒ Lose outer e- to form ionic compounds (1)

Question 11

Explain properties of Group 4 element (e.g silicon) (3)

Answer 11

Silicon:
⇒ Macromolecular structure
⇒ 4 e- in outer shell that form 4 covalent bonds
⇒ Classified as metalloid/semi metal as it has metallic and non metallic properties

Question 12

Explain properties of Group 5, 6 and 7 elements (3)

Answer 12

• P, S, Cl (P4, S8, Cl2)
• Non metals
• Can accept e- to form ionic compounds
• Can share e- to form covalent compounds

Question 13

Explain properties of Group 0 element (e.g: Ar) (3)

Answer 13

Ar
⇒ Noble gas, full outer shell
⇒ Inert (unreactive)
⇒ Monoatomic

Question 14

Explain melting and boiling point of the metals going across a period (Na, Mg, Al) (4)

Answer 14

⇒ strength of metallic bonding ↑’s, thus increasing the MP,BP (1)
⇒ Ionic charge increases, more electrons donated per atom (1), so stronger electrostatic attraction between the sea of delocalised e- and positive ions(1)
⇒ More energy is required to break the bonds (1)

Question 15

Explain melting and boiling point of Silicon (3)

Answer 15

⇒ Giant macromolecular structure
⇒ Has many strong covalent bonds between atoms
⇒ Requires lots of energy to break the bonds

Question 16

Explain melting and boiling point of the non metals (P, S, Cl) (2)

Answer 16

• Molecular structures (1)

- Weak VdW ℱ between molecules (1)

Question 17

Explain why the BP from P increases as you go to S then Cl (3)

Answer 17

⇒ M.p and b.p depends on size of van der waals forces (1)
⇒ S has a higher m.p and b.p than P due to having more electrons so it has stronger van der waals forces (1)
⇒ S8 > P4 > Cl2 (1)

Question 18

Explain melting and boiling point of Ar? (2)

Answer 18

It is a monatomic molecule that has a full outer shell so very weak VdW forces so low BP

Question 19

State and explain the trend in MP of Group II elements Ca-Ba? (3)

Answer 19

⇒ Trend: Decreases (1)
⇒ Increase in size of atom/ion (1)
⇒ Weaker attraction for delocalised/free/sea of electrons / weaker metallic bonding (1)

Question 20

State and explain the trend in MP of Group II elements Ca-Ba? (3)

Answer 20

⇒ Decreases (1)
⇒ Increase in size of ion/atom (1)
⇒ Weaker attraction for delocalised/free/sea of electrons / weaker metallic bonding (1)

Question 21

What are the trends in atomic radius across the period? (4)

Answer 21

⇒ decreases across a period (1)
⇒ Increased (atomic charge) number of protons (1)
⇒ Stronger electrostatic attraction between electrons and nucleus (1)
⇒ Electrons are in the same shell with similar shielding (1)

Question 22

What happens to atomic radius when starting the next period?

Answer 22

It jumps and increases

Question 23

Why does atomic radius increase down a group?

Answer 23

The atoms of each element will have one extra complete main energy level of e- compared to the one before

Question 24

What is first ionisation energy? (2)

Answer 24

The ethalpy required to remove one mole of electrons from one mole of atoms (1) in a gaseous state (1)

Question 25

Explain trend of first IE across a period (4)

Answer 25

⇒ IE increases (1) ⇒ increase in atomic charge (Stronger attraction to the electron) (1) ⇒ decrease in atomic radius (Stronger attraction) (1) ⇒ same shielding (1)

Question 26

Explain trend of first IE down a group (3)

Answer 26

⇒ IE decreases: ⇒ Increase in shielding ⇒ Bigger atomic radius → less attraction

Question 27

Why is there a drop in IE from one period to the next? (3)

Answer 27

⇒ New main energy level starts → increase in atomic radius (1) ⇒ Outer e- is further from nucleus (1) ⇒ Weaker electrostatic attraction therefore it is easier to remove (1)

Question 28

What is successive IE, explain the trend?

Answer 28

Removing further electrons from a charged ion Each electron will be harder to remove than the one before as they are being removed from a charged ion, thus the strength of the electrostatic force is greater and requires more energy to overcome

Question 29

What are the deviations from the general trend of IE? (4)

Answer 29

⇒ First IE drops between Group 2 and Group 3 (1) →Al has lower IE than Mg (1) ⇒ IE drops between Group 5 and Group 6 (1) →S has lower IE than P (1)

Question 30

Explain the drop between Group 2 and Group 3

Answer 30

Al has 1 e- in the 3p orbital therefore it is easier to remove as it is in a higher energy level and being shielded by 3s orbital thus requires less energy to remove

Question 31

Explain the drop between Group 5 and Group 6

Answer 31

S has one paired e- in the 3p orbital, this will be easier to remove due to the natural repulsion from the other electrons in the same orbital

Question 32

What can successive IE tell us? (1)

Answer 32

⇒ The group that the element is in (1) → eg. values of first five IE = 906, 1763, 14855, 21013 KJmol-1 → This is in group 2 as there is a big jump after the first 2 e- have been removed.

Question 33

Explain why the second ionisation energy of sodium is greater than the second ionisation energy of magnesium. (3)

Answer 33

⇒ Na(2+) requires loss of e– from a 2(p) orbital, Mg(2+) requires loss of e– from a 3(s) orbital (1) ⇒ Less shielding (in Na); (1) ⇒ e(–) closer to nucleus/ more attraction (of electron to nucleus) (in Na) (1)

Question 34

Explain why the second ionisation energy of calcium is lower than the second ionisation energy of potassium. (2)

Answer 34

⇒ In Ca⁺ valence electron further from the nucleus (1) | ⇒ More shielding in Ca⁺ (1)

Question 35

Identify the S-block metal that has the highest first ionisation energy. (1)

Answer 35

⇒ Beryllium / Be (1)

Question 36

Use your knowledge of structure and bonding to explain why the melting point of iodine is low (113.5 °C) and why that of hydrogen iodide is very low (–50.8 °C). (6)

Answer 36

⇒ I₂ is molecular. (1) ⇒ HI is molecular (1) ⇒ IMF holds the molecules together. (1) ⇒ Weak IMF forces hence the melting point is low in both substances. (1) ⇒ I₂ bigger molecule than HI so I₂ has more electrons. (1) ⇒Therefore stronger van der Waals between molecules in I₂ that need more energy to break causing the melting point to be higher. (1)

Question 36

State which of the elements magnesium and aluminium has the lower first ionisation energy? (3)

Answer 36

⇒ Al (1) ⇒ Al's outer electron in the 3p orbital (1) ⇒ Further away from nucleus (1)