Periodicity Flashcards by Alan Kader

How is the periodic table arranged?

Arranged in increasing atomic number.

How well did you know this?

Not at all

Perfectly

What are metalloids/semimetals? (2)

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

How well did you know this?

Not at all

Perfectly

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

⇒ 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)

How well did you know this?

Not at all

Perfectly

What is the reactivity trend down groups? (3)

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

How well did you know this?

Not at all

Perfectly

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

⇒ Rare metals (1)

⇒ Tend to from 3+ ions (1)

How well did you know this?

Not at all

Perfectly

What are Actinides? (1)

⇒ Radioactive metals

How well did you know this?

Not at all

Perfectly

What are periods?

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

How well did you know this?

Not at all

Perfectly

What are groups?

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

How well did you know this?

Not at all

Perfectly

Explain the position of H and He. (4)

⇒ [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)

How well did you know this?

Not at all

Perfectly

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

⇒ They are metals, have giant structures (1)

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

How well did you know this?

Not at all

Perfectly

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

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

How well did you know this?

Not at all

Perfectly

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

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

How well did you know this?

Not at all

Perfectly

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

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

How well did you know this?

Not at all

Perfectly

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

⇒ 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)

How well did you know this?

Not at all

Perfectly

Explain melting and boiling point of Silicon (3)

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

How well did you know this?

Not at all

Perfectly

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

Molecular structures (1)

- Weak VdW ℱ between molecules (1)

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

⇒ 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)

Explain melting and boiling point of Ar? (2)

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

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

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

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

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

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

⇒ 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)

What happens to atomic radius when starting the next period?

It jumps and increases

Why does atomic radius increase down a group?

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

What is first ionisation energy? (2)

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

Explain trend of first IE across a period (4)

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

Explain trend of first IE down a group (3)

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

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

⇒ 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)

What is successive IE, explain the trend?

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

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

⇒ 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)

Explain the drop between Group 2 and Group 3

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

Explain the drop between Group 5 and Group 6

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

What can successive IE tell us? (1)

⇒ 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.

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

⇒ 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)

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

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

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

⇒ Beryllium / Be (1)

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)

⇒ 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)

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

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