Chapter 7 - Periodicity

Question 1

Across each period what name is given to the repeating trend in properties?

Answer 1

Periodicity

Question 2

What is ionisation energy?

Answer 2

Ionisation energy measures how easily an atom loses electrons to form positive ions

Question 3

What is the definition of the first ionisation energy?

Answer 3

The first ionisation energy is 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 4

What factors affect the ionisation energy?

Answer 4

a. Atomic radius – The greater the distance between the nucleus and the outer electrons the less the nuclear attraction. The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect.
b. Nuclear charge – The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons.
c. Electron shielding – Electrons are negatively charged, and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons.

Question 5

How many ionisation energies does an element have?

Answer 5

An element has as many ionisation energies as there are electrons. For example, helium has two electrons and two ionisation energies.

Question 6

What is the definition of the second ionisation energy?

Answer 6

The second ionisation energy is 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 7

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

Answer 7

a. Down the group, the first ionisation energy decreases.
i. Atomic radius increases

ii. More inner shells so shielding increases
iii. Nuclear attraction on outer electrons decreases
iv. First ionisation energy decreases

Question 8

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

Answer 8

a. Across a period, the first ionisation energy increases.
i. Nuclear charge increases

ii. Same shell: similar shielding
iii. Nuclear attraction increases
iv. Atomic radius decreases
v. First ionisation energy increases

Question 9

Why is the first ionisation energy of boron less than the first ionisation energy of beryllium?

Answer 9

a. The fall in first ionisation energy from beryllium to boron marks the start of filling the 2p sub-shell.
i. The 2p sub-shell in boron has a higher energy that the 2s sub-shell in beryllium. Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium. As a result, the first ionisation energy of boron is less than the first ionisation energy of beryllium.

Question 10

Why is the first ionisation energy of oxygen less than the first ionisation energy of nitrogen?

Answer 10

a. The fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbitals of the 2p sub-shell.
i. In nitrogen and oxygen, the highest energy electrons are in a 2p sub-shell.

ii. In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom.
iii. Therefore, the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen.

Question 11

What are the elements near to the metal/non-metal divide called?

Answer 11

They are called semi-metals or metalloids.

Question 12

At room temperature what state are metals in and what is the one constant properties of all metals?

Answer 12

a. At room temperature, all metals except mercury are solids.
b. Secondly, one constant property of all metals is their ability to conduct electricity. This is a remarkable property for a solid, as charge must be able to move within a rigid structure for conduction to take place.

Question 13

What is the type of bonding in metals and how does it work?

Answer 13

a. Metals have a type of bonding called Metallic bonding and it is the strong electrostatic attraction between cations and delocalised electrons.
i. The cations are fixed in position, maintaining the structure and shape of the metal.

ii. The delocalised electrons are mobile and are able to move throughout the structure. Only the electrons move.
b. When drawing a diagram of metallic bonding make sure the delocalised electrons cancel the net positive charge of the cations.

Question 14

What type of structure is held within metals?

Answer 14

Billions of metal atoms are held together by metallic bonding in a “giant metallic lattice”.

Question 15

List what properties most metals have

Answer 15

a. Strong metallic bonds – attraction between positive ions and delocalised electrons
b. High electrical conductivity
c. High melting and boiling points

Question 16

Why can metals conduct electricity and which states can they conduct electricity in?

Answer 16

Metals conduct electricity in solid and liquid states. When a voltage is applied across a metal, the delocalised electrons can move through the structure, carrying charge.

Question 17

What melting and boiling points are experienced within metals and what does the melting point depend on?

Answer 17

a. For most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons. This strong attraction results in most metals having high melting and boiling points.
b. The melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice.

Question 18

How soluble are metals?

Answer 18

Metals do not dissolve. It might be expected that there would be some interaction between polar solvents and the charges in a metallic lattice as with ionic compounds, but any interactions would lead to a reaction rather than dissolving, as with sodium and water.

Question 19

What is a simple molecular lattice?

Answer 19

Many non-metallic elements exist as simple covalently bonded molecules. In solid state, these molecules form a simple molecular lattice structure held together by weak intermolecular forces. These structures therefore have low melting and boiling points.

Question 20

What is a giant covalent lattice?

Answer 20

The non-metals boron, carbon and silicon have a giant covalent lattice. Many billions of atoms are held together by a network of strong covalent bonds to form a giant covalent lattice.

Question 21

How many covalent bonds surround a carbon and silicon atom in the giant covalent lattice structure?

Answer 21

a. Carbon in its diamond allotrope and silicon uses all their four outer shell electrons to form covalent bonds to other carbon and silicon atoms.
b. The result is a tetrahedral structure, with bond angles of 109.5 by electron-pair repulsion.

Question 22

What properties do giant covalent lattice structures have?

Answer 22

a. Giant covalent lattices have high melting and boiling points as they have strong covalent bonds. High temperatures are necessary to provide the large quantity of energy needed to break the strong covalent bonds.
b. Giant covalent lattices are insoluble in almost all solvents. The covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.

c. Giant covalent lattices are also non-conductors of electricity, the exceptions being the other allotropes of carbon (graphene and graphite).
i. In carbon (diamond) and silicon, all four outer shell electrons are involved in covalent bonding, so none are available for conducting electricity.

ii. Carbon is special in forming several structures in which one of the electrons is available for conductivity. Graphene and graphite are able to conduct electricity.

Question 23

Describe graphene.

Answer 23

a. Graphene is a giant covalent structure with bond angles of 120 by electron-pair repulsion. Only three of the four outer shell electrons are used in covalent bonding. The remaining electrons are released into a pool of delocalised electrons shared by all atoms in the structure.
b. Graphene is a single layer of graphite, composed of hexagonally arranged carbon atoms linked by strong covalent bonds.
c. Graphene has the same electrical conductivity as copper and is the thinnest and strongest material ever made.

Question 24

Describe graphite.

Answer 24

a. Graphite is a giant covalent structure with bond angles of 120 by electron-pair repulsion. Only three of the four outer shell electrons are used in covalent bonding. The remaining electrons are released into a pool of delocalised electrons shared by all atoms in the structure. Therefore, graphite can conduct electricity.
b. Graphite is composed of parallel layers of hexagonal arranged carbon atoms, like a stack of graphene layers. The layers are bonded by weak London forces.

Question 25

25) What are the periodic trends in melting points of period 2 and 3 and explain why?

Answer 25

a. Across period 2 and period 3, the melting point increases from group 1 to group 4.
b. There is a sharp decrease in melting point between group 4 and group 5.
c. The melting points are comparatively low from group 5 to group 0.

d. The sharp decrease in melting point marks a change from giant to simple molecular structures.
i. On melting, giant structures have strong forces to overcome so have high melting points. Simple molecular structures have weak forces to overcome, so have much lower melting points.