Periodicity

Question 1

What decides which element is in which block?

Answer 1

Depending on which orbital the highest energy electrons are in, (e.g. s-orbital, p-orbital, d-orbital).

Question 2

What is a group?

Answer 2

The vertical column of elements.

Question 3

What happens to the reactivity of s-block metals as you go do down the group?

Answer 3

They get more reactive as you go down.

Question 4

What happens to the reactivity of p-block non metals as you go down the group?

Answer 4

They get less reactive as you go down.

Question 5

What is a period?

Answer 5

Horizontal rows of elements are called periods.

Question 6

What is periodicity explained by?

Answer 6

The electron arrangement of an element.

Question 7

What type of elements are those that are in group 1, 2, 3

Answer 7

Metals

Question 8

What structures do the elements in group 1, 2 and 3 have.

Answer 8

Giant structures

Question 9

What properties does silicon have?

Answer 9

They have some metallic properties and is therefore classed as a semi-metal.

Question 10

Why do the melting and boiling points increase from sodium to aluminium?

Answer 10

The metallic bonding strength increases due increased ion charges and the amount of electrons increase.

Question 11

What do the melting points of non-metals depend on?

Answer 11

Intermolecular forces

Question 12

What does the atomic radii tell us?

Answer 12

The size of atoms.

Question 13

What is the atomic radius affected by?

Answer 13

The type of bond.

Question 14

Why do nobles gases not get compared in atomic radii?

Answer 14

They cannot form covalent bonds.

Question 15

Why can you not measure the radius of a lone atom?

Answer 15

There is no clear point where electron cloud density drops to 0.

Question 16

When can metals form covalent bonds?

Answer 16

In a gas phase.

Question 17

What happens to atomic radii across a period?

Answer 17

It decreases.

Question 18

Why does the atomic radii change across a period?

Answer 18

The higher amount of proton charge has a stronger pull on electrons to the nucleus.

Question 19

What happens to atomic radii down a group?

Answer 19

It increases.

Question 20

Why does the atomic radii increase down a group?

Answer 20

The outer electron is further away.

Question 21

What is the first ionization energy?

Answer 21

The energy needed to convert one mole of gaseous atoms into one mole of +1 positive charged gaseous ions.

Question 22

What is the ionization equation?

Answer 22

M(g) -> M+(g) + e-(g)
or
M(g) + e- -> M+(g) + 2e-(g)

Question 23

What happens to the first ionization energy down a group?

Answer 23

It decreases

Question 24

What happens to the first ionization energy across a group?

Answer 24

Generally increases.

Question 25

Why does the first ionization energy increase a group?

Answer 25

The amount of protons increases (stronger positive charge) but the electrons stays in the same level.

Question 26

Why does the first ionization energy decrease down a group?

Answer 26

More inner levels means increased shielding so the outer electrons are easier to remove.

Question 27

Why does oxygen have a lower ionization energy than nitrogen?

Answer 27

Nitrogen has half shell stability whereas oxygen doesn't.

Question 28

Why does boron have a lower ionization energy than beryllium?

Answer 28

Removing an electron from boron means it loses the p orbital.

Question 29

Why does the successive ionization energy increase?

Answer 29

As you remove electrons the positive charge on the ion increases and the energy needed to remove an orbital is higher.