Periodicity

The 'staircase line' divides metals (on its left) from non-metals (on its right). 

Elements that touch this line, such as silicon, have a combination of metallic and non-metallic properties. They are called metalloids or semi-metals. Silicon, for example, is a non-metal, but it looks quite shiny and conducts electricity, although not as well as a metal. 

Areas of the Periodic Table are labelled s-block, p-block, d-block and f-block.

• All the elements that have their highest energy electrons in s-orbitals are in the s-block.
• All the elements that have their highest energy electrons in p-orbitals are in the p-block.
• Etc. 

Strictly speaking the transition metals and the d-block elements are not exactly the same. Scandium and zinc are not transition metals because they do not form any compound in which they have partly filled d-orbitals, which is the characteristic of transition metals. 

A group is a vertical column of elements. 

The elements in the same group form a chemical 'family' - they have similar propertites.

Elements in the same group have the same number of electrons in the outer main level.

It is common to number the groups in ordinary numbers as 1-7 and 0 (or 1-8) and sometimes as 1-18 including transition metals. 

In the s-block (metals) elements get more reactive going down a group. To the right (non-metals) tend to get more reactive going up a group. 

Transition elements are a block of rather unreactive metals. This is where the most useful metals are found. 

Lanthanides are metals which are not often encountered. They all tend to form +3 ions in their compounds and have broadly similar reactivity. 

Actinides are radioactive metals. Only thorium and uranium occur natually in the Earth's crust in anything more than trace quantities. 

Horizontal rows of elements are called periods. 

The periods are numbered starting from Period 1, which contains only hydrogen and helium. Period 2 contains the elements lithium to neon, etc. 

There are trends in physical properties and chemical behaviour as you go across a period. 

The positions of hydrogen and helium vary in different versions of the Periodic Table. 

Helium is usually placed above the noble gases (Group 0) because of its properties. However, it is not a p-block element - it's electron arrangement is 1s².

Hydrogen is sometimes placed above Group 1, but it is often placed on its own. 

It usually forms singly charged +1 (H⁺) ions like the Group 1 elements, but otherwise it is not similar to them since they are all reactive metals and hydrogen is a gas.

It is sometimes placed above the halogens because it can form H^- ions and also bond covalently. 

Periodicity is explained by the electron arrangements of the elements. 

• The elements in Groups 1, 2, and 3 (Na, Mg and Al) are metals. They have giant structures. They lose their outer electrons to form ionic compounds. 
• Silicon, in Group 4, has four electrons in its outer shell with which it forms four covalent bonds. The element has some metallic properties and is classed as a semi-metal. 
• The elements in Groups 5, 6 and 7 (P, S and Cl) are non-metals. They either accept electrons to form ionic compounds, or share their outer electrons to form covalent compounds. 
• Argon, in Group 0 is a noble gas - it has a full outer shell and is unreactive. 

There is a clear divide  between elements on the left with high melting poings (with the exception of sodium) and those on the right with low melting points. These trends are due to their structures:

• Giant structures (found on the left) tend to have high melting points and boiling points.
• Molecular or atomic structures (found on the right) tend to have low melting points and boiling points. 

The melting and boiling points of the metals increase from sodium to aluminium because of the strength of metallic bonding. As you go from left to right the charge on the ion increases so more electrons join the delocalised 'sea' of electrons that holds the giant metallic lattice together. 

The melting points of the non-metals with molecular structures depend on the sizes of the van der Waals forces between the molecules. This in turn depends on the number of electrons in the molecule and how closely the molecules can pack together. As a result the melting points of these non-metals are ordered: S8 > P4 > Cl2. Silicon, with its giant structure has a much higher melting point. Boiling points follow a similar pattern. 

The atomic radii tells us the size of atoms. 

You cannot measure the radius of an isolated atom because there is no clear point at which the electron cloud density around it drops to 0. Instead, half the distance between the centres of a pair of atoms is used. 

The atomic radius of an element can differ as it is a general term. It depends on the type of bond that it is forming (covalent, ionic, metallic, van der Waals etc.). The covalent radius is most commonly used as a measure of the size of the atom. 

(Even metals can form covalent molecules, such as Na2, in the gas phase. Since nobel gases do not bond  covalently with one another, they do not have covalent radii and so are often left out of comparisons with atomic sizes.)

You can explain why it decreases by looking at the electronic structures of the elements in a period, for example, sodium to chlorine in Period 3. 

As you move from sodium to chlorine you are adding protons to the nucleus and electrons to the outer main level, the third shell. The charge on the nucleus increases from +11 to +17. This increased charge pulls the electrons in closer to the nucleus. There are no additional electron shells to provide more shielding. So the size of the atom decreases as you go across the period. 

Going down a group in the Periodic Table, the atoms of each element have one extra complete main level of electrons compared with the one before. 

For example, in Group 1 the outer electron in potassium is in main level 4, whereas in sodium it is main level 3. 

So, going down a group, the outer electron main level is further from the nucleus and the atomic radii increase. 

The first ionisation energy is the energy required to convert one mole of gaseous atoms into one mole of singly positively charged gaseous ions. 

E(g) ---> E⁺ (g) + e^-  (g)                             (where E stands for any element)

The first ionisation energy generally increases across a period - alkali metals, such as sodium and lithium, have the lowest values and the nobel gases, such as helium and neon, have the highest. 

The first ionisation energy decreases going down any group.

As you go across a period from left to right, the number of protons in the nucleus increases, but the electrons enter the same main level. 

The increased charge on the nucleus means that it gets increasingly difficult to remove an electron.

The number of filled inner levels increases going down the group. 

This results in an increase in shielding. 

Also, the electron to be removed is at an increasing distance from the nucleus and is therefore held less strongly. 

Therefore the outer electrons get easier to remove going down a group because they are further away from the nucleus.

Moving from neon, in Period 0, (far right) with the electron arrangement 2,8 to sodium, 2,8,1 (Period 1, far left) there is a sharp drop in the first ionisation energy. 

This is because at sodium a new main level starts and so there is an increase in atomic radius, the outer electron is further from the nuclues, less strongly attracted and easier to remove. 

For the first ionisation energy:

• magnesium, 1s²2s²2p⁶3s², loses a 3s electron
• aluminium, 1s² 2s² 2p⁶ 3s² 3p¹, loses loses the 3p electron.

The p-electron is already in a higher energy level than the s-electron, so it takes less energy to remove it. 

An electron in a pair will be easier to remove than one in an orbital on its own because it is already being repelled by the other electron. 

• phosphorus, 1s²2s²2p⁶3s²3p³, has no paired electrons in a p-orbital because each p-electron is in a different orbital
• sulfur, 1s²2s²2p⁶3s²3p⁴, has two of its p-electrons paired in a p-orbital so one of these will be easier to remove than an unpaired one due to the repulsion of the other electron in the same orbital. 

If you remove the electrons from an atom one at a time, each one is harder to remove than the one before. 

When removing electrons from sodium (2,8,1) there is a sharp increase in ionisation energy between the first and second electrons. This is followed by a gradual increase over the next eight electrons and then another jump before the final two electrons. Sodium, in Group 1 of the Periodic Table, has one electron in its outer main level (the easiest one to remove), eight in the next main level and two (very hard to remove) in the innermost main level. 

The number of electrons that are easy to remove tell us the group number of the element in the Periodic Table. For example, the values of 906, 1763, 14,855, and 21,013 KJ mol^-1 for the first five ionisation energies of an element, tell us that the element is in Group 2. This is because the big jump occurs after two electrons have been removed.