chemistry module 3 periodicity

what is periodicity

the properties of the elements show trends that repeat themselves in each period of the periodic table

These trends are known as periodic trends, and the study of these trends is known as periodicity

what is the first ionisation energy

the energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

what is the first ionisation energy of sodium

Na (g) → Na+ (g) + e-

three factors that influence the ionisation energy

atomic radius, nuclear charge, electron shielding

how does atomic radius influence ionisation energy

electrons in shells that are further away from the nucleus are less attracted to the nucleus so the further the outer electron shell is from the nucleus, the lower the ionisation energy

how does nuclear charge affect ionisation energy

the nuclear charge increases with increasing atomic number, which means that there are greater attractive forces between the nucleus and outer electrons, so more energy is required to overcome these attractive forces when removing an electron

how does electron shielding affect ionisation energy

the shielding effect is when the electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge so the greater the shielding of outer electrons by inner electron shells, the lower the ionisation energy

what is the trend in ionisation anergy as you move down a group

As you move down a group, the nuclear charge increases but the ionisation energy decreases

This is due to the following factors:

The atomic radius increases

The shielding (by inner shell electrons) increases

Therefore, the attraction between the nucleus and the outer electrons decreases

what is the Trend in first ionisation energy across a period

increases due to the following factors:

• Across a period, the nuclear charge increases

• The distance between the nucleus and outer electron remains reasonably constant (no significant change in atomic radius)

  • The shielding by inner shell electrons remains the same

why is there a rapid decrease in ionisation energy between the last element in one period and the first element in the next period

The increased distance between the nucleus and the outer electrons

The increased shielding by inner electrons

These two factors outweigh the increased nuclear charge

why is there a slight decrease in first ionisation energy between beryllium and boron

the fifth electron in boron is in the 2p subshell, which is further away from the nucleus than the 2s subshell of beryllium

Beryllium has a first ionisation energy of 900 kJ mol-1 as its electron configuration is 1s2 2s2

Boron has a first ionisation energy of 801 kJ mol-1 as its electron configuration is 1s2 2s2 2p1

why is there a slight decrease in first ionisation energy between nitrogen and oxygen

the paired electrons in the 2p subshell of oxygen repel each other, making it easier to remove an electron in oxygen than nitrogen.

Nitrogen has a first ionisation energy of 1402 kJ mol-1 as its electron configuration is 1s2 2s2 2p3

Oxygen has a first ionisation energy of 1314 kJ mol-1 as its electron configuration is 1s2 2s2 2p4

metallic bonding

Metal atoms are tightly packed together in lattice structures

When the metal atoms are in lattice structures, the electrons in their outer shells are free to move throughout the structure

The free-moving electrons are called delocalised electrons and they are not bound to their atom

When the electrons are delocalised, the metal atoms become positively charged ions

The positive charges repel each other and keep the neatly arranged lattice in place

There are very strong forces between the positive metal centres and the ‘sea’ of delocalised electrons

covalent lattices

Covalent bonds are bonds between nonmetals where there is a shared pair of electrons between the atoms

In some cases, it is not possible to satisfy the bonding capacity of a substance in the form of a molecule

The bonds between atoms continue indefinitely, and a large lattice is formed

There are no individual molecules and covalent bonding exists between all adjacent atoms

Such substances are called giant covalent substances

The most important examples are the carbon allotropes graphite, diamond and graphene as well as silicon(IV) oxide

diamond

Diamond is a giant covalent lattice (or macromolecule) of carbon atoms

Each carbon is covalently bonded to four others in a tetrahedral arrangement with a bond angle of 109.5o

The result is a giant lattice structure with strong bonds in all directions

Diamond is the hardest substance known

For this reason, it is used in drills and glass-cutting tools

graphite

In graphite, each carbon atom is bonded to three others in a layered structure

The layers are made of hexagons with a bond angle of 120o

The spare electrons are delocalised and occupy the space between the layers

All atoms in the same layer are held together by strong covalent bonds

However, the layers are held together by weak intermolecular forces

These weak intermolecular forces allow the layers to slide over each other

graphene

Some substances contain an infinite lattice of covalently bonded atoms in two dimensions only to form layers.

Graphene is an example

Graphene is made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons

Graphene is one million times thinner than paper; so thin that it is actually considered two dimensional

Silicon(IV) oxide

Silicon(IV) oxide is also known as silicon dioxide, but you will be more familiar with it as the white stuff on beaches!

Silicon(IV) oxide adopts the same structure as diamond - a giant covalent lattice / macromolecular structure made of tetrahedral units all bonded by strong covalent bonds

Each silicon is shared by four oxygens and each oxygen is shared by two silicons

This gives an empirical formula of SiO2

Properties of metallic substances

Due to the delocalised ‘sea’ of electrons, metallic structures have some characteristic properties:

High melting and boiling point: as a lot of energy is required to overcome the strong electrostatic forces of attraction between positive ions and the 'sea' of delocalised electrons

Solubility: metals do not dissolve. There is some interaction between polar solvents and charges in the metallic lattice but these lead to reactions, rather than dissolving e.g. sodium and water

Electrical conductivity: conduct electricity in both solid and liquid states. This is due to the delocalised electrons which are free to move / carry charge around the structure

Properties of giant covalent substances

Giant covalent lattices have very high melting and boiling points

These compounds have a large number of covalent bonds linking the whole structure

A lot of energy is required to break the lattice

The compounds can be hard or soft

Graphite is soft as the intermolecular forces between the carbon layers are weak

Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds

Most compounds are insoluble with water

Most covalent substances do not conduct electricity

For example, diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom is involved in a covalent bond , so there are no free electrons available

Graphite has delocalised electrons between the carbon layers, which can move along the layers when a voltage is applied

Graphene is an excellent conductor of electricity due to the delocalised electrons