Periodic+Trends

= = There are three major trends in the periodic table that you will need to explain. They are


 * 1) atomic and ionic radii,
 * 2) Ionisation energy and
 * 3) Electronegativity.

When attempting to explain these trends you will need to consider the relative size of the electrostatic attraction between the protons in the nucleus and the valence electrons. This electrostatic attraction depends on
 * the number of protons in the nucleus (electrostatic attraction between positive nucleus and valence electrons increases with the size of the nuclear charge)
 * the distance of the valence electrons from the nucleus (electrostatic attraction decreases as the distance between the positive and negative charges increases)
 * the amount of electron-electron repulsion. (When the number of electron shells increases this effect is often described as increased shielding from the inner shells of electrons. This results in decreased electrostatic attraction between the protons and valence electrons).

As a general rule when **going across a period** the effect of the **increased nuclear charge** is the most important factor whereas when **going down a group** the increased number of electron shells (i.e. **increased radius**) is the more important factor.

The atomic radius of a metal is half the distance between the nuclei in two adjacent atoms. For elements that exist as diatomic molecules the atomic radius is half the distance between the nuclei of the two atoms in a molecule.
 * 1. Atomic Radii **

The diagram below shows the relative size of some atoms.

Two trends in atomic size are obvious: This is because as you go down a group the valence electrons are occupying orbitals which are at higher energy levels and are further from the nucleus. This factor is greater than the effect of the increased nuclear charge, particularly as the inner shells of electrons provide additional shielding. Going across a row there is an increase in the number of protons in the nucleus and therefore an increase in the nuclear charge. The additional electrons are all added to the same energy level, and therefore provide no additional shielding. As a result, the increased electrostatic attraction results in
 * a) Down a group the atomic size increases **
 * b) Across a row (or period) the atomic radius decreases **

the electrons being attracted more closely to the nucleus and a decrease in atomic size.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The relative radii of positive and negative ions is shown on the right. The radius of a positive ion is always smaller than the radius of the atom from which it was formed. Removing one or more electrons reduces electron-electron repulsion so thesame attractive force from the nuclear charge results in the electrons being closer to the nucleus. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Also the formation of cations often results in the loss of all the electrons in the outer valence shell. For example, Na, 1//s//2 2//s//22//p//6 3//s//1 forms a Na+ ion by losing the electron from its 3//s// orbital. Since the valence electrons in the sodium ion are only in the 2nd energy level (2//s// and 2//p// orbitals) the radius of the ion is considerably smaller than its atom. <span style="font-family: Arial,sans-serif; font-size: 11pt;">When an anion forms the increased electron-electron repulsion resulting from the additional electron increases the size of the electron cloud. The increased radius results in decresased electrostatic attraction between nucleus and valence electrons. media type="youtube" key="8dDVxqI_0I4" width="420" height="315"
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">2. Ionic Radii **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The first ionisation energy is the energy required to remove the least tightly held electron from each <span style="font-family: Arial,sans-serif; font-size: 11pt;">atom in one mole of gaseous atoms in their ground state. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Na//(g)// ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Na+//(g)// + 1e -
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">3. Ionisation Energy **



<span style="font-family: Arial,sans-serif; font-size: 11pt;">(The second and third ionisation energies are the energies needed to remove the second and third electrons respectively.) Ionisation energies are all positive values as the breaking of the electrostatic attraction between the valence electron and the nucleus must be an endothermic process.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The graph of ionisation energy against atomic number on the right shows the periodic trend both down and across the periodic table.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The explanations for the trends are essentially the same as the explanations for atomic size. <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. although both the nuclear charge **and** the number of filled energy levels increases down a group, the shielding from the inner shells of electrons more than compensates for the increased nuclear charge, so it requires less energy to remove electrons.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">This is because going across a row there is an increase in nuclear charge, but the electrons are added into the same energy level with no additional shielding. The increase in electrostatic attraction between the nucleus and the valence electron results in an increase in ionisation energy.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">a) Across the table there is an increase in ionisation energy. **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Down a group there is an increase in nuclear charge, but the electrons are being added to energy levels at higher energies further from the nucleus with additional shielding from inner shells of electrons. The electrostatic attraction between nucleus and valence electron therefore decreases and this results in a decrease in the ionisation energy.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">b) Down a group the ionisation energy decreases. **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">An obvious consequence of the above trends is that the low ionisation energies of metal atoms means they readily form positive ions, whereas non-metal atoms do not form positive ions. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Also the decrease in ionisation energy going down a group (e.g. Li to Fr) can be related to the increase in reactivity of metals as it requires less energy to form the cation.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The removal of an electron from an inert gas is very difficult, reflecting the very stable arrangement resulting fom a full energy level.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTE: **

[|For more on Ionisation energies including an explanation of anomalies, clink here]

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<span style="font-family: Arial,sans-serif; font-size: 11pt;">The electronegativity of an element is a measure of the ability of an atom to attract towards itself the <span style="font-family: Arial,sans-serif; font-size: 11pt;">A chemist named Linus Pauling devised a scale of relative electronegativities. Some values are shown below. You are not expected to recall these values in the exam, but you should be able to explain the trend in electronegativity both down a group and across a row.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">4. Electronegativity **
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">electrons shared in a chemical bond. **



<span style="font-family: Arial,sans-serif; font-size: 11pt;">Electronegativity values show the same trends as 1 <span style="font-family: Arial,sans-serif;">st <span style="font-family: Arial,sans-serif; font-size: 11pt;"> ionisation energies, but their interpreted values are much more convenient to use than the precise measured values of ionisation energies. <span style="font-family: Arial,sans-serif; font-size: 11pt;">This means that an atom such as fluorine that has a high ionisation energy (does not lose electrons <span style="font-family: Arial,sans-serif; font-size: 11pt;">easily) also has a high electronegativity. In fact **fluorine is the most electronegative element.**
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTES: **
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">In contrast to metals, the reactivity of non-metals **increases up** each group of the periodic table as the greater electronegativity values reflect the greater attraction for electrons required to gain a complete energy level.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">The elements of Group 18 were not given electronegativity values as they do not tend to form bonds with other elements.


 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTE: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Atoms of elements with widely different electronegativities tend to form ionic bonds with each other since the atom of the less electronegative element gives up its electron(s) to the atom of the more electronegative element. Atoms of elements with more similar electronegativities tend to form covalent bonds. If the atoms have the same, or almost the same electronegativities the covalent bond is non-polar and the electrons of the covalent bond are shared equally between the two atoms. The greater the difference in electronegativity the more polar the covalent bond, and the atom which has the greater electronegativity will form the slightly negative end of the polar bond. The chlorides across any row of the periodic table show a continuum from ionic bonding between a metal and a non-metal with a big electronegativity difference (e.g. NaCl) through to polar covalent bonding between non-metal atoms of slightly different electronegativities (e.g. SiCl4) to non-polar covalent bonding between atoms of identical electronegativity (e.g. Cl2).