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 of 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 the trend across period 2 ?

Answer 6

Across period 2, the 2s sub shell fills with two electrons, followed by the 2p sub shell with six electrons

Question 7

What is the trend across period 3 ?

Answer 7

Across period 3, the 3s sub shell fills with two electrons, followed by the 3p sub shell with six electrons

Question 8

What is the trend across period 4 ?

Answer 8

• Across period 4, although the 3d sub shell is involved, the highest shell number is n=4.
• From the n=4 shell, only the 4s and 4p sub shells are occupied

Question 9

What are blocks ?

Answer 9

The elements in the periodic table can be divided into blocks corresponding to their highest energy sub shell

Question 10

What is ionisation ?

Answer 10

The removal of one or more electrons from an atom

Question 11

Define first ionisation energy

Answer 11

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 12

How does atomic radius affect ionisation energy ?

Answer 12

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

Question 13

How does nuclear charge affect ionisation energy ?

Answer 13

More protons creates a greater attraction between the nucleus and the other electrons

Question 14

What is shielding ?

Answer 14

It is when inner shell electrons repel outer shell electrons

Question 15

How does electron shielding affect ionisation energy ?

Answer 15

This reduces the attraction between the nucleus and the outer electrons

Question 16

Why does ionisation energy decrease going down a group ?

Answer 16

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

Question 17

Why does ionisation energy increase across a period ?

Answer 17

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

Question 18

Why do successive ionisation energies increase ?

Answer 18

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

Question 19

Define second ionisation energy

Answer 19

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 20

What causes the large jumps in successive ionisation energies ?

Answer 20

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

Question 21

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

Answer 21

• The number of electrons in the outer shell
• The group of the element in the periodic table
• The identity of the element

Question 22

Explain the trend of first ionisation energy down a group

Answer 22

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

Question 23

Explain the general trend of first ionisation energy across a period

Answer 23

• Nuclear charge increases
• Same shell so similar shielding
• Nuclear attraction increases
• Atomic radius decreases
• First ionisation energy increases

Question 24

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

Answer 24

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

Question 25

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

Answer 25

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

Question 26

Explain the trend of atomic radii across a group

Answer 26

• Nuclear charge increases
• Nuclear attraction increases
• Atomic radius decreases

Question 27

What is metallic bonding ?

Answer 27

Metallic bonding is when each atom donates an outer shell electron, which then becomes delocalised. This creates cations.

Question 28

What causes metallic bonding ?

Answer 28

Strong electrostatic force of attraction that occurs between cations and the delocalised electrons

Question 29

What are common properties of metals ?

Answer 29

• Strong metallic bonds
• High electrical conductivity
• High melting and boiling points

Question 30

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

Answer 30

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

Question 31

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

Answer 31

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

Question 32

What are the properties of giant covalent lattices ?

Answer 32

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

Question 33

Why are giant covalent lattices insoluble ?

Answer 33

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

Question 34

Why do most giant covalent lattices not conduct electricity ?

Answer 34

All four outer shell electrons are involved in covalent bonding

Question 35

Why do simple molecules have low melting points ?

Answer 35

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