Chapter 7 Periodicity

Question 1

What is periodicity?

Answer 1

The repeating pattern of properties across a period

Question 2

What are the vertical columns on the periodic table known as?

Answer 2

Groups

Question 3

What are the horizontal rows on the periodic table known as?

Answer 3

Periods

Question 4

How are elements arranged in order of within the periodic table?

Answer 4

Atomic number

Question 5

What do the groups on the periodic table tell us?

Answer 5

The group number shows the number of electrons in the outer shell, therefore this means that elements in the same group have similar chemical properties.

Question 6

What do the periods on a periodic table tell us?

Answer 6

Tells us the number of the highest energy electron shell in an elements atom

Question 7

What is the trend across a period in the periodic table?

Answer 7

Each period starts with an electron in a new highest energy shell.
For example, across period 2, the 2s sub shell fills with two electrons, followed by the 2p sub-shell with six electrons

Question 8

What is first ionisation energy?

Answer 8

The energy that is required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous +1 ions.
Eg Na (g) —> Na+ (g) + e-

Question 9

What are the factors that affect ionisation energies?

Answer 9

Atomic radius
Nuclear charge
Electron shielding

Question 10

How does atomic radius impact ionisation energy?

Answer 10

The greater the distance between the nucleus and outer electrons, the larger the atomic radius. Therefore, the electron being removed is further away, so there is less attraction between the nucleus and the electron being removed.

Question 11

How does nuclear charge impact ionisation energy?

Answer 11

The more protons there are in the nucleus, the greater the nuclear charge. Therefore, the stronger the attraction between the nucleus and the electron being removed is.

Question 12

How does electron shielding impact ionisation energy?

Answer 12

The more inner shells of electrons there are between the nucleus and the outer electrons, the greater repulsion there is between the inner electrons and the electron being removed. This reduces the attraction between the nucleus and electron being removed so first ionisation energy decreases.

Question 13

How many ionisation energies does an element have?

Answer 13

An element has as many ionisation energies as there are electrons.

Question 14

Explain the ionisation energy of Helium.

Answer 14

Helium has two ionisation energies, however the second ionisation energy is greater than the first ionisation energy. This is because there are two protons attracting the two electrons in the 1s sub-shell. After the first electron is lost, the single electron is pulled closer to the nucleus due as there is a stronger nuclear charge. Therefore, more ionisation energy is needed to remove the second electron.

Question 15

What is the second ionisation energy?

Answer 15

The energy that is required to remove one mole of electrons from one mole of +1 gaseous ions to produce one mole of gaseous +2 ions.

Question 16

What do successive ionisation energies allow us to predict?

Answer 16

The number of electrons in the outer shell based on the number of electrons before the large jump signalling the move to the next shell
The group of the element on the periodic table
The identity of the element

Question 17

What is the trend in first ionisation energies down a group?

Answer 17

First ionisation energies decrease down a group.
This is because the atomic radius increases and there are more inner shells so shielding also increases. Therefore, the nuclear attraction to outer electrons decreases.

Question 18

What is the trend in first ionisation energies across a period?

Answer 18

First ionisation energies increase across a period. This is because the nuclear charge increases as an extra proton is added each time you move across the period, there is the same number of shells so similar electron shielding occurs , due to this the nuclear attraction between the nucleus and outer electrons is increased so the atomic radius decreases.

Question 19

Why is there a drop in ionisation energy from group 2 to group 3?

Answer 19

This marks the start of filling the p sub-shell in the Group 3 element which is of higher energy than that of the s sub-shell in the group 2 element. Therefore, the p sub shell electron of the group 3 element is easier to remove than one of the s sub shell electrons in the group 2 element.
The first ionisation energy of group 3 is less than the first ionisation energy of group 2.

Question 20

Why is there a drop in ionisation energy from group 5 to group 6?

Answer 20

Moving from group 5 into group 6, the electrons begin to pair up in a p sub shell, this increases repulsion between the electrons therefore making it easier to remove an electron from a group 6 element than a group 5 element.
The first ionisation energy of a group 6 element is less than the first ionisation energy of a group 5 element

Question 21

What is metallic bonding?

Answer 21

The electrostatic forces of attraction between the positive metal cations which are fixed in position and the negative sea of delocalised electrons that are mobile.
This forms a giant metallic lattice

Question 22

What are the properties of metals?

Answer 22

High electrical conductivity- electrons are mobile and free to carry the charge when solid or liquid
High melting and boiling points- due to the electrostatic attraction between the cation and delocalised electrons which require a lot of energy to overcome
Insoluble- as the metallic bonds are too strong to be overcome by the interactions of a solvent
They don’t dissolve in polar solvents as the interactions between the solvent and metal would lead to a reaction

Question 23

What non metals form giant covalent lattices?

Answer 23

Boron
Silicon
Carbon

Question 24

What are the properties of Giant covalent lattices?

Answer 24

High melting and boiling point- due to the strong covalent bonds which require a lot of energy to overcome
Insoluble in almost all solvents- the covalent bonds are too strong to be overcome by the interaction of solvents
Poor conductivity (with the exception of Graphite and Graphene) - as there are no free mobile charge carriers.

Question 25

Explain the drop in ionisation energy between Nitrogen and Oxygen

Answer 25

In nitrogen and oxygen the highest energy electrons are in a 2p sub-shell. In nitrogen, all three p orbitals are singularly occupied however in oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen than a nitrogen atom.
Therefore, the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen

Question 26

Explain the drop in first ionisation energy between Beryllium and Boron

Answer 26

This marks the start of filling the 2p sub shell in Boron, which is of higher energy that that of the 2s sub shell in Beryllium. Therefore, itv is easier to remove an electron from the 2p orbital of Boron than it is the 2s orbital of Beryllium.
This means Boron has a lower first ionisation energy than Beryllium

Question 27

What is the structure of diamond?

Answer 27

Giant covalent lattice
Four carbon atoms are covalently bonded to four other carbon atoms forming a tetrahedral shape, bond angle 109.5.

Question 28

What elements have a simple molecular structure across period 2 and 3?

Answer 28

Nitrogen Phosphorus
Oxygen Sulfur
Fluorine Chlorine
Neon Argon

Question 29

What elements have a giant metallic structure across period 2 and 3?

Answer 29

Lithium
Sodium
Beryllium
Magnesium
Aluminium

Question 30

How do you know when a substance is simple molecular or giant covalent?

Answer 30

If it is a gas, liquid or a low melting point solid then it is a simple molecular substance. However, if it is a high melting point solid then it is giant covalent.

Question 31

Do Giants covalent structures form intermolecular forces?

Answer 31

No, they only have covalent bonds.

Question 32

Explain why the melting point of Sulfur is higher than that of Chlorine.

Answer 32

Sulfur exists as S8 molecules and Chlorine as Cl2 molecules. This means that the London Forces are stronger between S8 molecules as there are more electrons than Cl2 molecules.

Question 33

Why is the Group 4 element Germanium a good conductor of electricity?

Answer 33

Has a giant metallic lattice rather than a giant covalent lattice like Carbon and Silicon, therefore is has negative delocalised electrons which are free to move and carry the charge throughout the structure.

Question 34

Why does Arsenic have a higher melting point than nitrogen and phosphorus?

Answer 34

Has a giant covalent lattice whereas in period 2 and 3, group 5 had a simple molecular lattice.

Question 35

Explain, in terms of its structure, whether you would also expect diamond to be a good conductor of electricity.

Answer 35

No because it has a giant covalent lattice structure so all outer electrons are used up in covalent bonds.

Question 36

Explain whether you would expect gallium to be soluble or insoluble in water.

Answer 36

Insoluble because the metallic bonds are too strong to be overcome by the interactions of a solvent.

Question 37

Explain in terms of bonding why Magnesium has a higher melting point than sodium

Answer 37

Magnesium has more delocalised electrons within its metallic lattice attracting the positive 2+ cations so the electrostatic attraction requires a lot more energy to overcome in the Magnesium atom than the Sodium atom.

Question 38

Describe the structure and bonding of aluminium. Include the names of the particles involved in the bonding within your answer.

Answer 38

Aluminium has a giant metallic lattice structure. This is due to its positive 3+ cations which are fixed in position attracting the negative delocalised electrons which are mobile and free to move around the structure. This forms the metallic structure due to the electrostatic attraction between the positive cations and negative delocalised electrons.

Question 39

Explain why the first ionisation energy of nitrogen is higher than the first ionisation energy of carbon.

Answer 39

Because across a period there is a general increase in first ionisation energy, the nuclear charge increases as nitrogen has more protons in its nucleus. The number of shells remains the same so inner electron shielding stays similar, however the increased nuclear attraction reduces the atomic radius, pulling the outer electrons closer slightly. Therefore, there is a stronger attraction between the nucleus and the electron being removed of Nitrogen than Carbon.

Question 40

The variation in first ionisation energies across a period of the Periodic Table provided evidence for what structures within an atom?

Answer 40

Sub shells

Question 41

From the following elements: Lithium, Beryllium and Fluorine.
Predict which one will have the largest second ionisation energy. Explain your answer.

Answer 41

Lithium
The atomic radius decreases as the only electron is the outer shell is lost
-therefore, there is only 2 electrons remaining in the final shell
-this decreases the amount of inner electron shielding so it has little effect
-there is a greater nuclear attraction between the nucleus and the electron being removed so second ionisation energy increases