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 ionisation?

Answer 5

The removal of one or more electrons from an atom.

Question 6

Define first ionisation energy

Answer 6

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 7

How does atomic radius affect ionisation energy?

Answer 7

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

Question 8

How does nuclear charge affect ionisation energy?

Answer 8

More protons creates a greater attraction between the nucleus and the outer electrons.

Question 9

How does electron shielding affect ionisation energy?

Answer 9

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

Question 10

Why does ionisation energy decrease going down a group?

Answer 10

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

Question 11

Why does ionisation energy increase across a period?

Answer 11

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

Question 12

Why do successive ionisation energies increase?

Answer 12

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

Question 13

Define second ionisation energy

Answer 13

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 14

What causes the large jumps in successive ionisation energies?

Answer 14

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

Question 15

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

Answer 15

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 16

Explain the trend of first ionisation energy down a group

Answer 16

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

Question 17

Explain the general trend of first ionisation energy across a period

Answer 17

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

Question 18

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

Answer 18

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

Question 19

In period 2, explain the fall from nitrogen to oxygen

Answer 19

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 20

What happens in metallic bonding?

Answer 20

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

Question 21

Define metallic bonding?

Answer 21

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

Question 22

What are common properties of metals?

Answer 22

Strong metallic bonds
High electrical conductivity
High melting and boiling points

Question 23

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

Answer 23

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

Question 24

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

Answer 24

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

Question 25

What are the properties of giant covalent lattices?

Answer 25

High melting and boiling points
Insoluble in almost all solvents
Do not conduct (except for graphite and graphene)

Question 26

Why are giant covalent lattices insoluble?

Answer 26

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

Question 27

Why do most giant covalent lattices not conduct electricity?

Answer 27

All four outer shell electrons are involved in covalent bonding.

Question 28

Why do simple molecules have low melting points?

Answer 28

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

Question 29

State three examples of Giant covalent structures

Answer 29

Diamond
Graphite
silicon dioxide

Question 30

what is an allotrope

Answer 30

so you have two compounds with the same element, but their arrangement of atom is different, for example, diamond and graphite are carbon compounds

Question 31

state the properties of diamond

Answer 31

1) high strength
2) HMP
3) insulator
4) insoluble

Question 32

state the properties of graphite

Answer 32

1) low strength
2) HMP
3) good conductor
4) insoluble

Question 33

state the properties of silicon dioxide

Answer 33

1) high strength
2) HMP
3) insulator
4) insoluble

Question 34

what is a graphite composed of?

Answer 34

composed of parallel layers of hexagonally arranged carbon atoms and between the layers are London forces, so holds the layers together.

Question 35

whats between the layers of graphite

Answer 35

we have delocalised electrons that are donated by the carbon atoms and also have London forces between the layers

Question 36

What is graphene

Answer 36

a single layer of graphite

Question 37

Do giant covalent structures conduct electricity

Answer 37

No, except graphene and graphite.

Question 38

Are giant covalent structures soluble

Answer 38

no they are insoluble in almost all solvents