The Periodic Table
The periodic table is a list of all the elements in order of increasing atomic number. We can predict the properties of an element from its position in the table.
Between the boundary of metals and non-metals on the periodic table are elements which have properties of metals and non-metals. They are called metalloids or semi-metals. Silicon is an example of this. It is a non-metal but looks shiny and conducts electricity.
Areas of the table are labelled S block, p block, d block and f block.
All elements that have their highest energy electrons in s-orbitals are in the s-block, e.g sodium (1s2, 2s2, 2p6, 3s1)
All elements that have their highest energy electrons in p-orbitals are in the p-block, e.g carbon (1s2, 2s2, 2p2)
All elements that have their highest energy electrons in d-orbitals are in the d-block, e.g iron (1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d6)
Trends in the Properties of Elements in Period 3
Periodicity is explained by the electron arrangements of the elements.
Elements in groups 1,2 and 3, sodium, magnesium and aluminium are metals. They have giant structures. They lose their outer electrons to form ionic compounds.
Silicon in Group 4 has 4 electrons in its outer shell with which it forms 4 covalent bonds. It has some metallic properties and is classed as a semi-metal.
Elements in groups 5, 6 and 7, phosphorus, sulfur and chlorine, 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, has a full outer shell and therefore is unreactive.
There is a clear break in the middle of the table between elements with high melting points (on the left with Sodium as an exception) and those with low melting points, on the right. These are due to their structures.
Trends in the Properties of Elements in Period 3
Groups 1, 2 and 3 are metals and have a giant metallic structure. Apart from sodium, they have high melting and boiling points. These are due to the strength of metallic bonding. Charge of the ion increases so more electrons join the delocalised "sea" that holds the lattice together.
Group 4 contains Silicon which is a semi-metal and macromolecular and has a higher boiling/melting point than Group 3.
Group 5, 6, 7 and 0 are molecular (atomic for Argon) and have low melting/boiling points. These depend on the van der Waals forces between the molecules. This depends on the number of electrons in the molecule and how closely the molecules can pack together.
Melting points of non metals: S8 > P4 > Cl2
Silicon has a giant covalent structure and has a much higher melting point.
More Trends in the Properties of the Elements of P
Atomic Radii - We cannot measure the radius of an isolated atom because there is no clear point at which the electron density will be 0. Instead, we use half the distance between centres of a pair of atoms.
Atomic radius can differ due to the type of bond. The covalent radius is the most commonly used measure of the size of the atom.
Metals can also form covalent molecules in gas phase. Noble gases do not bond covalently with each other so they are left out in these comparisons.
Atomic radii decreases across each period and there is a jump when we start the next period.
Atoms get larger as we go down a group.
As we move from sodium to chlorine, we are adding protons to the nucleus and electrons to the outer main level (third shell). The charge on the nucleus increases from +11 to +17. The increased charge pulls electrons closer to the nucleus. No additional shells to provide more shielding, therefore atom decreases in size as we go across a period.
More Trends in the Properties of the Elements in P
Radii of atoms increase down a group because as we go down a group, the atoms of each element have one extra complete main level of electrons compared with the one before. Outer main level is further from the nucleus and the atomic radii increases.
First ionisation energies have periodic patterns. They generally increase across a period. Alkali metals have the lowest value whereas noble gases have the highest values. This is because the number of protons in the nucleus increases but the electrons enter the same main level. Increased charge means it gets increasingly difficult to remove an electron.
First ionisation energies decreases going down a group. This is because the number of filled inner levels increase which results in an increase in shielding. Also the electron to be removed is at an increasing distance from the nucleus and is held less strongly. Outer electrons get easier to remove as we go down a group as they are further away from the nucleus.
There is a sharp drop in I.E from one period to the next as we start a new main level so there is an increase in radius, outer electron is further away from nucleus
A Closer Look At Ionisation Energies
There is a drop in first ionisation energy between Group 2 and Group 3 due to the sub-level from which the first electron is removed.
Magnesium: 1s2, 2s2, 2p6, 3s2 loses a 3s electron.
Aluminium: 1s2, 2s2, 2p6, 3s, 3p1 loses a 3p electron.
3p is a higher sub-level than 3s so it takes less energy to remove it.
There is a drop in first ionisation energy between Group 5 and Group 6 due to repulsion between electron in an orbital.
Phosphorus - 1s2, 2s2, 2p6, 3s2, 3p3 - no paired electrons as each p electron is in a different orbital.
Sulfur - 1s2, 2s2, 2p6, 3s2, 3p4 - it has two of its p electrons paired so one of these will be easier to remove than an unpaired one as they repel each other.
A Closer Look at Ionisation Energies
If we remove the electrons from an atom one at a time, each one is harder to remove than the one before.
For sodium, there is a sharp increase in ionisation energy between the first and second electrons followed by a gradual increase over the next 8 electrons and then another jump before the final two electrons.
Sodium is in Group 1 and has one electron in its outer shell, eight in the next main level and two (very hard to remove) in the inner most shell.
The number of electrons that are relatively easy to remove tells us the group number in the periodic table.
906, 1763, 14855, 21013 kJ mol^-1 are the first five ionisation energies of an element. The element belongs in group 2 since the sharp jump is after the second ionisation energy (after two electrons were removed).