The Periodic Table
Newlands was one of the first chemists to describe a repeating pattern in physical and chemical properties when the elements were placed in order of atomic weight – his Law of Octaves, although it did not work for all elements, was the first periodic law.
Mendeleev’s organization of the periodic table into atomic weight order, with elements having matching properties being placed in the same column, required him to leave gaps where the properties did not match. These he attributed to elements that had not yet been discovered, and he went on use his table to successfully predict the properties of these undiscovered elements based on the properties of known elements in the same Group.
The modern periodic table is arranged in order of increasing atomic number (proton number), with a new Period (row) started every time electrons occupy a new shell. The consequence of this is that elements in each Group (column) have the same number of outer shell electrons and therefore similar physical and chemical properties.
This also means that there are repeating trends and patterns in physical and chemical properties from one period to the next, which we refer to as periodicity. Using Period 3 as an example, we can start to examine some periodic properties that show the same patterns and trends across other groups.
i) Electron arrangements in Period 3
|[Ne] 3s1||[Ne] 3s2||[Ne] 3s2 3p1||[Ne] 3s2 3p2||[Ne] 3s2 3p3||[Ne] 3s2 3p4||[Ne] 3s2 3p5||[Ne] 3s2 3p6|
Across Period 3 we see an increasing number of electrons in the n=3 shell, with the 3s subshell being filled first, followed by the 3p subshell. Exactly the same pattern is seen across Period 2 with the filling of the 2s and 2p subshells
Period 4 shows the same periodic trend too, although the pattern is there interrupted by the filling of the 3d subshell after 4s is filled.
ii) Properties that depend on electron arrangement: first ionisation energies
Consider the first ionisation energies of the first 40 elements:
There are clear patterns in the ionisation energies across one period compared to other periods – compare Period 2 (red) and Period 3 (green), for example, and note the same pattern in Period 4 (purple) if the d-block elements (grey) are ignored.
There is a generally increasing trend in ΔHi1 across Period 3, although we also see a small dip after Mg and again after P, so we need to be able to explain why ionisation of Al and S requires a little less energy than might be expected.
Across a period: Across a period the electron being removed in first ionisation comes from the same shell, and the outer shell electrons experience the same shielding from filled inner shells, but the number of protons increases and hence the nuclear attraction towards the electron being removed increases. As a result it requires an increasing amount of energy to remove an outer shell electron across the group.
In Na and Mg, first ionisation involves removal of an electron from the 3s subshell. For Al through to Ar, first ionisation involves removal of an electron from the 3p subshell. The 3p subshell is higher in energy than the 3s, with the electrons on average a little further from the nucleus and hence less strongly attracted. As a result the removal of electrons from 3p takes a little less energy than from 3s, and we see a small downwards step in first ionisation energies between Mg and Al. This is experimental evidence validating the subshell model of atomic structure.
In Al, Si and P the three 3p orbitals are all singly-filled. From S to Ar the 3p orbitals are progressively filled with two electrons. Electrons occupying the same orbital experience mutual repulsion even though the opposite spins minimize this, and hence it takes a little less energy to remove an electron from a doubly-filled orbital than a singly filled orbital in the same subshell. As a result there is a small downwards step in the first ionisation energies between P and S. This is experimental evidence validating the single filling of orbitals in the subshells before electrons are paired, in the current model of atomic structure.
Comparing periods: As we finish one period and start the next, the same overall pattern of ΔHi1 is seen, but each successive period repeats at lower ΔHi1 values. Each time we finish a period and start the next, the nuclear charge continues to increase as the atomic number increases, however this increasing attraction towards the electron being removed is outweighed by the fact that when we start a new period we start a new principle quantum shell, so the ionised electron is further from the nucleus, and we have one more filled inner shell shielding the ionised electron from the nuclear charge. For these reasons the attraction between the outer shell electrons and the nucleus is decreased and less energy is required for first ionisation.
Trend in first ionisation energies down a Group
Although it is not a periodic property, we also need to be able to explain the decreasing trend in first ionisation energies down a group.
Down a group, the electron being removed during ionisation comes from a shell progressively further out from the nucleus. The electron is also shielded from the nuclear charge by progressively more filled inner shells, and both of these factors cause a decreasing attraction between the outer shell electrons and the nucleus, outweighing the effect of the increased nuclear charge as the number of protons increases down the group. With the decreasing attraction, less energy is needed to remove an outer shell electron, and hence the first ionisation energies decrease down a Group.
iii) Properties that depend on electron arrangement: atomic radius
The electron cloud surrounding an atom does not have a specific ‘end’, so it is not possible to measure to the edge of the electron cloud to measure the atomic radius. One solution is measure the distance between bonded atoms, taking half the bond length as the atomic radius. This is why we can’t measure it for noble gases.
|Atomic radius (nm)||0.190||0.160||0.130||0.118||0.110||0.102||0.099||N/A|
Note: 1nm = 1 x 10-9m
The atomic radius decreases across each period. Across the period there are the same number of filled inner shells and hence the same shielding, but an increasing number of protons in the nucleus and hence greater nuclear charge. This leads to increasing attraction between the nucleus and the outer shell electrons across the period, and as a result the outer shell electrons are pulled closer to the nucleus, decreasing the atomic radius.
Bonding and Structure
*bonding becomes increasingly covalent in character across the period
Bonding in the Period 3 elements and their compounds becomes increasingly covalent in character across the period. Aluminium shows ionic bonding in some compounds and polar covalent bonding in others. This is because Aluminium is more electronegative than Group 1 or 2 metals, so there is less of a difference in electronegativity between aluminium and the non-metals it bonds with. The same trend towards increasingly covalent bonding character is seen in the other Periods.
The pattern in types of structure across Period 3 is also seen in the other groups, with the first elements in the period being giant metallic, then giant covalent, and the remaining elements having simple molecular structure.
v) Properties that depend on bonding/structure: melting and boiling points
The patterns in bonding and structure form the basis for the patterns in melting and boiling points across Period 3, which are also repeated in other Periods.
Melting points relate to the type of structure. The more energy needed to break down the structure so that the particles can move around, the higher the melting point will be.
Elements with a giant structure (giant metallic lattice Na, Mg and Al, and the giant covalent lattice of Si) all require a lot of energy to break the strong bonds holding the atoms/ions together in the lattice, so the melting points are high.
From sodium to aluminium the melting and boiling points increase because:
- the ions have greater charge and so are more strongly attracted to the delocalised electrons [Na+, Mg2+, Al3+]
- there are an increasing number of delocalised electrons for the metal ions to be attracted to [Na → Na+ + e–, Mg → Mg2+ + 2e–, Al → Al3+ + 3e–]
- so more energy is need to overcome these strong electrostatic attractions in order to break the metallic bonds
Elements with simple molecular structures have only weak intermolecular forces holding the molecules together. To overcome these so that the molecules can move around requires little energy, so the melting points are low. We don’t have to break the strong covalent bonds holding atoms together inside each molecule. Trends in boiling points follow the same pattern as the melting points, at a higher temperature.
For P4, S8 and Cl2 and Ar the size of the intermolecular forces (and therefore how high the melting and boiling points are) depends on the number of electrons in the molecule. The more electrons there are, the stronger the intermolecular forces (London forces – we’ll meet these in more detail later).
The trend in melting points is therefore S8 > P4 > Cl2 > Ar since the numbers of electrons in the molecules are 128, 60, 34, 18 respectively.
vi) Properties that depend on bonding/structure: electrical conductivity
Electrical conductivity is also related to the type of bonding. Na, Mg and Al have metallic bonding. They are good conductors because they possess mobile delocalised electrons which can carry the charge. Al is a better conductor (of both heat and electricity) than Na because it donates 3 outer shell electrons rather than 1 into the sea of delocalised electrons.
The remaining elements in the period have covalent bonding, and do not contain any mobile delocalised electrons, so they are poor conductors. They get increasingly poor because an electron has to be removed from the outer shell in order for conduction to take place, and the first ionisation energies increases across the period.