Chapter 7 - Periodicity

Question 1

Who created the modern periodic table?

Answer 1

Dmitri Mendeleev

Question 2

How were the elements ordered by Mendeleev?

Answer 2

Increasing atomic mass

Question 3

What is the name for the vertical columns of the periodic table?

Answer 3

Groups

Question 4

What is the name for the horizontal rows of the periodic table?

Answer 4

Periods

Question 5

What is the periodic trend in electron configuration?

Answer 5

The sub shells of n energy level fill up.

Question 6

What is ionisation?

Answer 6

The removal of one or more electrons from an atom.

Question 7

Define first ionisation energy

Answer 7

The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

Question 8

How does atomic radius affect ionisation energy?

Answer 8

Greater distance between the nucleus and outer electrons, the attraction is lower.

Question 9

How does nuclear charge affect ionisation energy?

Answer 9

More protons creates a greater attraction between the nucleus and the outer electrons. therefore higher ionisation energy

Question 10

How does electron shielding affect ionisation energy?

Answer 10

Inner shell electrons repel outer shell electrons, called shielding. This reduces the attraction between the nucleus and the outer electrons. therefore lower ionisation energy

Question 11

Why does ionisation energy decrease going down a group?

Answer 11

More electrons shells so the outer electrons are further away and there is a greater shielding effect.

Question 12

Why does ionisation energy increase across a period?

Answer 12

The number of protons in the nucleus increases so nuclear charge increases causing atomic radius to decrease, whilst shielding stays the same.

Question 13

Why do successive ionisation energies increase?

Answer 13

There are less electrons so the nuclear attraction on the remaining electrons will be greater.

Question 14

Define second ionisation energy

Answer 14

The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.

Question 15

What causes the large jumps in successive ionisation energies?

Answer 15

Moving down to a closer shell, as these electrons are closer so experience a greater nuclear attraction.

Question 16

What predictions can be made from a graph of successive ionisation energies?

Answer 16

The number of electrons in the outer shell, the group of the element in the periodic table and thus the identity of the element.

Question 17

Explain the trend of first ionisation energy down a group

Answer 17

Atomic radius increases,
More inner shells so shielding increases,
Nuclear attraction on outer electrons decreases,
First ionisation energy decreases.

Question 18

Explain the general trend of first ionisation energy across a period

Answer 18

Nuclear charge increases,
Same shell: similar shielding,
Nuclear attraction increases,
Atomic radius decreases,
First ionisation energy increases.

Question 19

In period 2, explain the fall from beryllium to boron of first ionisation energies

Answer 19

The new electron enters the 2p sub shell, which is slightly further away from the nucleus than the 2s sub shell.

Question 20

In period 2, explain the fall of ionisation energy from nitrogen to oxygen

Answer 20

Nitrogen’s electrons in the 2p sub shell are unpaired so oxygen’s 8th electron is paired, causing repulsion and a lower ionisation energy.

Question 21

Explain the trend of atomic radii decreasing across a group

Answer 21

Nuclear charge increases,
Nuclear attraction increases,
Atomic radius decreases.

Question 22

What is metallic bonding?

Answer 22

Each atom donates an outer shell electron, which becomes delocalised. This creates cations.

Question 23

What causes metallic bonding?

Answer 23

Strong electrostatic attraction between the fixed cations and the delocalised electrons.

Question 24

What are common properties of metals?

Answer 24

Strong metallic bonds
High electrical conductivity
High melting and boiling points

Question 25

Why does the melting point and boiling point increase across the metals of a period?

Answer 25

Number of delocalised electrons per atom and charge on cation increase, so stronger electrostatic attraction.

Question 26

In period 3, why does silicon have the highest melting point?

Answer 26

Forms a giant covalent lattice, where each atom is covalently bonded to four others.

Question 27

What are the properties of giant covalent lattices?

Answer 27

High melting and boiling points
Insoluble in almost all solvents
Do not conduct (except for graphite and graphene)
formed from carbon, silicon anf boron

Question 28

Why are giant covalent lattices insoluble?

Answer 28

Covalent bonds holding it together are too strong to be broken by interaction with solvents.

Question 29

Why do most giant covalent lattices not conduct electricity?

Answer 29

All four outer shell electrons are involved in covalent bonding.

Question 30

Why do simple molecules have low melting points?

Answer 30

Weak induced dipole-dipole forces between molecules are easy to break.