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 periodic trend in electron configuration?

Answer 3

The sub shells of n energy level fill up.

Question 4

What is ionisation?

Answer 4

The removal of one or more electrons from an atom.

Question 5

Define first ionisation energy

Answer 5

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 6

How does atomic radius affect ionisation energy?

Answer 6

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

Question 7

How does nuclear charge affect ionisation energy?

Answer 7

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

Question 8

How does electron shielding affect ionisation energy?

Answer 8

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

Question 9

Explain why successive ionisation energies always increase.

Answer 9

As each electron is removed, the outer shell is drawn closer to the nucleus. Nuclear attraction is greater and more energy is needed to remove the next electron

Question 10

Define second ionisation energy

Answer 10

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 11

What causes the large jumps in successive ionisation energies?

Answer 11

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

Question 12

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

Answer 12

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 13

Explain the trend of first ionisation energy down a group

Answer 13

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

Question 14

Explain the general trend of first ionisation energy across a period

Answer 14

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

Question 15

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

Answer 15

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

Question 16

In period 2, explain the fall from nitrogen to oxygen

Answer 16

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 17

Explain the trend of atomic radii decreasing across a group

Answer 17

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

Question 18

Define metallic bonding

Answer 18

The strong electrostatic attraction between cations and delocalised electrons

Question 19

What are common properties of metals?

Answer 19

Strong metallic bonds
High electrical conductivity
High melting and boiling points

Question 20

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

Answer 20

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

Question 21

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

Answer 21

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

Question 22

What are the properties of giant covalent lattices?

Answer 22

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

Question 23

Why are giant covalent lattices insoluble?

Answer 23

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

Question 24

Why do most giant covalent lattices not conduct electricity?

Answer 24

All four outer shell electrons are involved in covalent bonding.

Question 25

Why do simple molecules have low melting points?

Answer 25

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

Question 26

Define periodicity

Answer 26

A repeating trend in the properties of the elements across each period

Question 27

Why do metals have melting points?

Answer 27

Hight temperatures are needed to overcome the strong electrostatic attraction between the cations and the electrons

Question 28

Explain how the bonding in a simple molecular lattice differs from that in a giant covalent lattice.

Answer 28

A simple molecular lattice has London forces between molecules. A giant covalent lattice has covalent bonds between them