Periodicity

Structure of the periodic table

• Elements in the periodic table are arranged according to their atomic number (the number of protons in the nucleus).

• The rows in the periodic table are referred to as periods and the columns are called groups.

• All of the elements within a period have the same number of electron shells — sodium and magnesium are both in period 3 and have three electron shells.

• There are trends in the physical and chemical properties of elements as you go across the period, such as a decreasing atomic radius. These trends are referred to as periodicity.

• Elements that are in the same group of the periodic table have the same number of outer electrons — carbon and silicon are both in group 4 and have four outer electrons.

• Since the number of outer electrons dictates how an element will react with other molecules, elements in the same group will take part in similar chemical reactions (i.e. they have similar chemical properties).

Blocks

The periodic table is organised into blocks.

• Groups 1 and 2 make up the s block — these elements have their outermost electron in an s sub-shell.

• The d block consists of the transition elements, which have their outermost electron in a d sub-shell.

• Groups 3-8 make up the p block since their highest energy (outermost) electron is in a p sub-shell.

• There’s also an f block consisting of the lanthanides which are found in that landing strip right at the bottom. Luckily for us, at A-level we don’t need to concern ourselves with these f block elements.

Ionisation energy

Ionisation energy is the energy required to remove an electron from an atom, molecule or ion. If an electron is being removed from a neutral atom to form an ion with a +1 charge, it’s called first ionisation energy. If the electron is being removed from a +1 ion to form a +2 ion, it’s called second ionisation energy. Ionisation energy is always measured in gaseous states.

• First ionisation energy: the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

• Second ionisation energy: the energy required to remove one mole of electrons from one mole of gaseous +1 ions to form one mole of gaseous 2+ ions.

The higher the ionisation energy, the harder it is to remove the electron. Ionisation energy depends on a few different factors:

• Nuclear charge – the more protons there are in the nucleus, the greater the nuclear charge and the greater the attraction between the nucleus and the outer electron. This makes it harder to remove the outer electron.

• Atomic radius – the greater the distance between the outer electron and the nucleus, the weaker the attraction between the nucleus and the electron and the easier it is to remove.

• Electron shielding – the more electrons located between the nucleus and the outer electron, the more they will ‘shield’ the charge from the nucleus. This makes the attraction between the nucleus and the electron weaker, making it easier to remove.

As you go across the period, ionisation energy increases. This is because atomic number increases along the period (each element has one more proton in its nucleus than the last.) This means they have a greater nuclear charge and a greater hold of their outer electrons. In addition, the atomic radius decreases along the period since the greater positive charge is able to pull the outer shell electrons more tightly towards the nucleus. The decreasing atomic radius further increases the nuclear attraction to its outer electrons. Since all of the elements in the same period have the same number of electron shells electron shells, they will all experience similar shielding.

As you go down a group, ionisation energy decreases. Although the elements have more protons in their nucleus, their atomic radii are getting bigger. This reduces the nuclear attraction to the electrons and makes it easier to remove an electron from their outer shell. It also means that electron shielding is increased since each additional electron shell puts more electrons between the outer electrons and the nucleus, deflecting the nuclear charge from the outer shell. This means that as you go down the group, each element becomes more easily ionised.

So we’ve seen that as we go across a period, ionisation energy increases. This is only a general rule and there are a few exceptions to this trend:

• There’s a small drop in ionisation energy between groups 2 and 3 — the outer electron is occupying a p orbital rather than an s orbital. Because p orbitals have a slightly higher energy than the s orbitals, they are slightly further from the nucleus than s orbitals. They experience a weaker nuclear charge due to the increased atomic radius and the electron shielding from the s electrons. These two factors override the effect of increased nuclear charge from having one extra proton in its nucleus.

• There is also a dip in ionisation energy between groups 5 and 6. For elements in group 5, they have three electrons in a p sub-shell which are occupying orbitals singly. Once we get to group 6, the electrons will have to start pairing up. Being negatively charged, pairs of electrons repel each other. This repulsion between the two electrons in an orbital means that electrons are easier to remove from shared orbitals.

Periodic trends in structure and boiling point

Metallic bonding is the type of attraction that occurs within a metal, such as a lump of copper or sodium. It is the electrostatic attraction between positive metals ions and delocalised electrons. Electrons within a metal are delocalised when they detach themselves from the atom they normally belong to and are free to move through the structure. Since the metal atoms have lost electrons, they become positive metal ions. The structure of a metal is a giant lattice. Its positive metal ions are arranged in a regular layered arrangement. This type of structure gives metals their properties:

• High melting and boiling points – due to the strong metallic bonding

• Good conductors of heat and electricity – due to the delocalised electrons which are able to move and carry a current

• Malleability (can be bent or hammered into shape without breaking) – due to the layers of metal ions which can slide over each other

Giant covalent molecules, such as diamond, graphite and graphene are another type of giant structures. Diamond is a giant covalent structure made up of carbon atoms bonded to four other carbon atoms. These atoms are held together by covalent bonds, making diamond very hard and giving it a high melting and boiling point. It is not able to conduct electricity because it does not have any charged particles which are free to move.

Graphite is another giant covalent structure, but unlike diamond, each carbon atom is bonded to only three others in a planar hexagonal arrangement. This means that one of carbon’s outer electrons is not being used and it instead becomes delocalised, which means that it can move through the structure. Because these electrons are free to move and carry a charge, this means that graphite is able to conduct electricity. Since the carbon atoms are held together by strong covalent bonds, this means that graphite has a high melting and boiling point.

Graphite exists in layers of hexagons which are held together by weak intermolecular forces. These forces break easily, separating the layers from each other as they slide over each other. It is this property that makes graphite a useful substance in pencils – when you apply pressure and move the pencil, layers of graphite slide away, leaving a mark on the paper.

Graphene is just a single layer of graphite and so has similar properties – it is able to conduct electricity and has a high melting and boiling point. These properties mean that graphene is a useful material in electronics and composite materials.

Trends in melting and boiling point

As you go across a period, the type of bond formed between atoms of an element changes. Since the type of bonding affects how much energy is needed to break the bonds, this means there are trends in melting and boiling points as you go across the period:

• Metals (left-hand side of periodic table) have a high melting and boiling point. This increases across the period as the number of delocalised electrons increase (and metallic bonds get stronger).

• For giant covalent lattice structures, like carbon and silicon, the melting and boiling point is also high due to the high number of strong covalent bonds which need to be broken.

• Simple covalent structures (on the right-hand side of the periodic table) have a low melting and boiling point. This is because the strong covalent bonds within the molecule stay intact – it’s the weak intermolecular forces that are breaking when these substances change state.

• The noble gases have the lowest melting and boiling points because they are held together by the weakest intermolecular forces. They exist as single atoms (monoatomic) rather than molecules, so they have fewer electrons than the diatomic halogens. Fewer electrons mean weaker intermolecular forces.