Periodicity Flashcards
How is the periodic table arranged?
Arranged in increasing atomic number.
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 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)
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)
What are Lanthanides, what ions do they form? (2)
⇒ Rare metals (1)
⇒ Tend to from 3+ ions (1)
What are Actinides? (1)
⇒ Radioactive metals
What are periods?
Horizontal rows of elements, elements of that group will have the same number of electron shells/similar shielding
What are groups?
Vertical columns where elements of that group will have the same number of electrons in their outermost shell.
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)
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)
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
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
Explain properties of Group 0 element (e.g: Ar) (3)
Ar
⇒ Noble gas, full outer shell
⇒ Inert (unreactive)
⇒ Monoatomic
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)
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
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)
