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The modern picture of the electronic configuration of atoms takes the form of the ‘atomic shell model’ whereby the configurations of all atoms are considered as modified versions of the configuration of the hydrogen atom, the simplest atom in the periodic table. The atomic shell model states that the electrons orbit the nucleus of the atom in shells of well-defined energy that increase in radius from the nucleus outward, just like the shelled layers of an onion. Then, on closer examination, each of these larger shells contain ’sub-shells’ which are basically, smaller intervals of energy difference in the range defined by the particular shell being considered.

Each shell is labelled by both a letter and a number. The number is the principle quantum number ‘n’, and has values from 0 to infinity in integer steps. The letters range from J (corresponding to n = 0) ‘upward’ in relation to the values of n; K represents n = 1, L represents n = 2 and so on.

A letter and a number also define each sub-shell. The number here is the orbital angular momentum quantum number ‘l’ and has values that range from 0 to n-1 again in integer steps. The letter that corresponds to the number in this case is slightly different. For the numbers l = 0 to l = 3 the corresponding letters are s, p, d, f, and for the numbers where l = 4+, the letters just follow the pattern of the alphabet (continuing on from f) so l = 4 has the letter g, l = 5 the letter h, and so on.

The periodic table is an ordered chart of the known elements based upon this model of the atoms electrons, and also more importantly, their chemical and physical properties. In the table, only the elements ground state (lowest energy) electron configurations are represented, partly because this is the usual state of most elements atoms in their pure form under standard conditions, but mostly because a table that included any excited states wouldn’t really have much significance, and would be impractical.

The table can basically be considered as an ordered list of increasing atomic number, which is read from left to right and has a new line (or period) for every shell that is being filled with electrons. The table is also arranged in columns. In the most general sense each column is of varying width and causes the table to be divided into 4 blocks, each of which corresponds to a particular sub-shell being filled. The blocks are labelled as s, p, d, and f, as there are no stable atoms that have electrons in the g and beyond sub-shells in their ground states.

Each sub-shell can hold a different number of electrons (which is why the very general column widths vary), which is given by the number of orbitals it has. Each orbital can hold 2 electrons whose spins are anti-parallel, and the number of orbitals in each sub-shell is given by the magnetic quantum number, m_l. The values that the magnetic quantum number can take (and hence the number of orbitals that can be filled) is given by the range 0 to ± l, so for example the d sub-shell has l = 2, so m_l ranges from –2 to +2 in integer steps, and therefore there are five orbitals that can be filled.

Generally the table takes the following form, note that the f block is removed from the main part of the table, this is mainly for convenience, because the f block elements are generally rare, radioactive, not really in common usage and as they are so different, they disturb the regularity of the rest of the table.

Also you can see that H and He are included in the s and p blocks respectively, this isn’t a strictly true representation, but for alternative reasons. H is included at the head of the first column of the s-block as it is filling the s sub-shell, however it has very few similar properties when compared to the other s-block elements. He is included at the head of the last column of the p-block, this is because its properties are similar to the rest of the elements in that column. Although He isn’t actually filling its p sub-shell, it does have a full s sub-shell and consequently a full outer shell of electrons. Overall these elements should be detached from the bulk of the table, as they only fit some of the trends that the others do.

Each of these four blocks has its separate columns and each column corresponds to the position of the last electron to orbit the nucleus. For example the very first column in the period table is the first column of the s-block, this means that the last electron to orbit the nucleus is an s-electron and we can also deduce that there is one possible orbital for it to be in and that it’s the first electron in that orbital. The same applies for the second column in the s-block, but here we can further say that the spin of this electron is opposite to whatever that of the first electron in this orbital was due to the Pauli Principle.

When looking at the columns across the table we generally only consider the columns that contain what are known as the main group elements. These are the elements contained within the s and p-blocks only and as a result we label each column as groups 1 – 8 from left to right across the table. In some cases we also consider the ’sub-groups’ formed from the columns of the d-block elements, this adds another 10 groups to the table, making a total of 18.

So, each group relates atoms with a similar final electron configuration. As the final configuration also designates what chemical and physical properties the various elements have, the groups also relate elements with similar characteristics, and this results in a number of readily observable trends within the periodic table, with emphasis on the main group elements.

For example, group eight contains the nobel gasses. They all have full shells and are therefore in their most stable (lowest energy) state. The result of this is that they are all monatomic gasses (as either ionic or covalent ‘attraction’ between the atoms has to occur for them to form liquids or solids, and although there may be some very momentary ‘van der Waals’ attraction between the atoms at room temperature, this is not long lived enough to cause liquefaction or solidification!) that barely react with anything (as for a ‘normal’ reaction to take place there has to be a net loss of energy between the two reacting atoms, via the transfer or ’sharing’ of an electron(s) between them, as the shells are full there is very little opportunity for these atoms to ‘release’ or ‘accept’ electrons).

All other atoms (aside from those in group eight) ‘want’ to be in their lowest possible energy configurations and as a result they will try to ‘accept’, ‘loose’ or ’share’ each other’s electrons to achieve this. For example the reaction between Cs and F gives the most stable ‘ionic’ compound known:

Group one is the alkali metals, they all have one ‘outer’ electron and as a result they would ‘prefer’ to loose this to form a singly charged positive ion, which has a full outer shell of electrons, just like the noble gasses. As you descend group 1 the proton number Z increases, as does the corresponding number of electrons (as single atoms are electrically neutral). This means that the distance of the final electron from the nucleus at the bottom of the group is as large as possible and, as the electrostatic influence of the nucleus decreases exponentially with distance, the last electron on the atom at the bottom of group 1 is very ‘loosely bound’ (as shielding effects tend to cause each electron to see only one proton, the only force acting on that electron is the electrostatic attraction between the one proton and the said electron. This is not the full picture as the electron also ‘momentarily’ feels the pull of the whole nucleus, but overall one electron is bound by one proton). Cs is the last element at the bottom of group one (that isn’t radioactive), and as a result can loose its outer electron very easily.

Conversely the atoms in group seven all have a single ‘gap’ for an electron to go into to form a full outer shell, and therefore a negative ion. The effect of size increase down the group implies that the effect of the nucleus on the electrons is greatest at the top (where the atoms are smaller), therefore the atoms with the strongest pull for an electron (high electronegativity) are at the top of group seven. Fluorine is at the top of group seven and therefore has the strongest pull for an electron.

It should be noted that any ion is actually at a higher energy than the neutral atom even if it does have a full shell of electrons, but this energy is recouped when a positive ion(s) meets a negative ion(s) and forms an ionic lattice. The formation of the lattice releases a vast amount of energy that compensates for the ionisations of the component atoms and actually leads to a net decrease in the energy of the system. The quantity of energy released also depends upon how well the ions can pack together in the lattice, which is also why a very large Cs⁺ ion along with a very small F^- ion makes for a very favourable combination overall.

All s-block atoms can react in the same way with p-block atoms in groups six and seven (and partially group 5) to form ionic compounds with different ratios between the numbers of individual ions involved (depending on where the atoms come from). For example Mg has two electrons to loose to form a full shell, and again F can only accept one, but two F atoms can react with one Mg atom to form MgF₂.

Reactions with the atoms of groups three and four and some of five are fairly different as they form the borderline cases whereby they don’t ‘want’ to loose electrons, but they don’t want to gain them either. When these atoms react they have to share electrons between themselves to lead to a net energy decrease. This is covalency and an example would be Mg₃Al₂ for s-block atoms reacting with p-block, but as these atoms are ambiguous they can also react with any of the other atoms in the p-block, CO₂ being another example.

The majority of elements are metallic, and a number of physical properties depend upon the strength of the metallic bonding between the atoms, which also depends upon the electronic structure. As you move from left to right across the d-block valences increase. This means that the number of electrons available to form a (delocalised) metallic bond over the whole structure increases, and therefore the ’strength’ of the bond increases. The greater the metallic bond, the higher the electrical and thermal conductivity, and the melting and boiling points. A similar effect happens across the main block elements, until they change from being metals to non-metals. So overall these properties increase as you move from left to right.

As you move across a period (up until you reach non-metals) the density of the elements increases, this is due to the fact that no new shells are started across a period, but atomic number (and therefore atomic mass) increases with every element. So there is a vast mass increase with only a tiny increase in the atomic volume, leading to a greater density.

And lastly, the d-block elements are also the most coloured of all the elements, and patterns in the coloration which correspond to their electronic configurations can be noted as you move across the block.

Conclusion:

There are many more trends in chemical and physical properties throughout the periodic table than have been mention here, and the majority of them are due to the electronic structure of the elements. The ordering of the table, based upon the ground state electronic structure of the atoms, helps us to see these trends easily and makes the periodic table one of the most useful tools available when studying the properties of elements, whether it be from a chemical or a physical point of view.