I AM A MOLE AND I LIVE IN A HOLE!

PERIODICITY 

I) FINDING PERIODIC PATTERNS

Elements can be ordered in many ways. They could be classified by their states at a particular temperature or as metals and non-metals; there may be patterns in their reactions with oxygen or water or other chemicals.

They could be sorted due to their atomic masses.

The first three alkali metals, lithium, sodium and potassium all react similarly with water; however the reaction becomes more vigorous down the group.

• Johann Dobereiner’s Law of Triads stated that the middle element in a group has the mean relative atomic mass of the other two elements. This law works for the first three alkali metals and the halogens.
• John Newlands presented his idea for the Law of Octaves, where elements were arranged in order of increasing atomic mass displayed a repetition in properties every eighth element.
• Dmitri Mendeleev also arranged the elements in increasing atomic mass, however unlike Newlands; he left gaps in the table so that similar elements could always appear in the same group. Also he said that the spaces would be filled by elements not then discovered. Furthermore he predicted the properties of these elements, based on the properties of known elements in the same group.

• In 1913, Henry Mosely was able to show that the real sequence in the Periodic Table is not the order of relative atomic masses, but instead due to atomic number. This answered questions about variations in chemical and physical properties, however many properties depend upon the distribution of electrons in atoms. 

II) VERSIONS OF THE PERIODIC TABLE

The vertical groups of elements are labelled 1, 2, and 3 up to 7 with the noble gases in group 0

The horizontal periods are labelled 1, 2, 3, etc.

The Periodic Table is split into blocks by the type of orbital most affecting the properties:

• Groups 1 and 2 are in the s-block
• Groups 3 to 7 and group 0 are in the p-block
• The transition elements are included in the d-block
• The lanthanide and actinide elements are included in the f-block.

III) PERIODIC PATTERNS OF PHYSICAL PROPERTIES OF ELEMENTS

All group 2 (s-block) elements have an outer shell configuration of s².

In the s-block, the outermost electrons are in an s orbital; in the p-block, the outermost electrons are in p orbitals

• Elements in the same group have the same number of electrons in their outer shell

• For the elements in Groups 1 to 7, the number of outer-shell electrons is the same as the group number; for example, chlorine in group 7 has seven outer shell electrons: 1s² 2s² 2p⁶ 3s² 3p⁵, a total of 7 electrons in the third shell

• Group 0 elements, the noble gases, have a full outer shell of eight electrons.

Atomic radii

The size of an atom cannot be measured precisely as electrons do not have a clear outer limit.

A measure of an atom’s size is its atomic radius. It can be either the covalent radius or metallic radius. Covalent radius is half the distance between the nuclei of neighbouring atoms in molecules. Metallic radius is half the distance between the nuclei of neighbouring atoms in metallic crystals.

• Across the periods, atomic radii decrease as the size of the nuclear charge increases, causing outer shell electrons to be attracted closer to the nucleus.

• Down groups, atomic radii increase as the addition of extra shells increases the distance from the nucleus to the outer shell. Also the inner shell electrons shield the outer shells from the nuclear attraction, causing them to be repelled further away from the nucleus.

Boiling points

The lowest boiling points in the periodic table are accompanied by elements that exist in diatomic molecules (e.g. H₂, N₂, O₂, F₂, Cl₂, Br₂) or single atoms (He, Ne, Ar, Kr). The forces of attraction between the particles are very weak as they are intermolecular forces.

The highest boiling points are in the metals and group 4 elements. Metals have metallic bonding and carbon, silicon and germanium exist in giant covalent compounds. A large amount of energy is required to dissociate these bonds. Across group 1 and 2, boiling points increase as there are more electrons in the outer shell to take part in bonding.

IV) PERIODIC PATTERNS OF FIRST IONISATION ENERGIES

• The elements with the highest first ionisation energies are the noble gases

• The atoms with the lowest first ionisation energies are the alkali metals
• There is a general increase in first ionisation energies across a period, although the trend is uneven.
• The first ionisation energies of the d-block elements show little variation.

The patterns in first ionisation energies are due to:

• The increase in positive nuclear charge across a period.

• The forces of attraction between the positive nuclear charge and electrons decrease as the quantum number of the shells increase. The further the shell is from the nucleus the lower the first ionisation energy is

• Filled inner shells of electrons shield outer electrons. The outer electron shells are repelled by the electrons in the filled inner shells, so the first ionisation energy falls.

Across a period

There is a general trend of increasing ionisation energies across a period. This is because the nuclear attraction increases across a group.

It is expected that boron would have a higher ionisation than beryllium. However as it is easier to remove electrons from a p-orbital because they are of a higher energy level, boron’s first ionisation energy is less than beryllium’s.

The first ionisation energy of oxygen is unexpectedly lower than nitrogen. The first ionisation energy is lower for oxygen because it involves the removal of an electron from a p-orbital with two electrons in it. As electrons are constantly trying to repel each other, less energy is required to remove the electron. In the nitrogen atom, there is less electron repulsion, so more energy is required to remove the electron.

In groups

First ionisation energies generally decrease down the groups.

With increasing proton number:

• The positive nuclear charge increases

• The atomic radius increase so the distance of the outer electrons from the nucleus also increases with each new shell

• The shielding effect of the filled inner electron shells increases as the number of inner shells grows.

• The combined distance and shielding effects overcome the effect of increasing nuclear charges. Therefore the first ionisation energies decrease down a group.

In atoms with low first ionisation energies, reactions tend to be more vigorous, as less energy is required to form the ion, the reaction can happen easier.

V) SUCCESSIVE IONISATION ENERGIES AND THE PERIODIC TABLE

Successive ionisation energy trends can be used to identify an element’s position in the Periodic Table:

• They provide evidence for the general pattern of electron shells

• Predict the simple electronic configuration of an element

• Confirm the position of an element in the Periodic Table.

A large ionisation energy will increase will occur between an element’s n^th   and n^th + 1 ionisation energy, where n is the group number.

So for carbon, in group 4, the greatest increase will be from the fourth ionisation energy to the fifth ionisation energy.

GLOSSARY

Atomic radius: half the distance between the nuclei of two covalently bonded atoms.