Modern Periodic Table
In this page we are going to discuss about trends in the modern periodic table concept .The periodic trends of the following properties are studied here.
- Atomic size
- Metallic and Non-metallic character
- Ionization potential
- Electron affinity
- Electronegativity
The distance from the centre of the nucleus to the outermost shell of an atom is called the atomic radius of that atom.
Electrons in the same period progressively increase from left to right, as the atomic number increases, e.g., in the 3rd period, from sodium (Z = 11) to chlorine (Z = 17). It would be expected that as the number of protons, electrons and neutrons increase, the size of the atom increases. Contrary to expectations, in spite of the increased atomic number, the size of the atom gradually decreases from left to right.
The distance from the nucleus to the outermost shell depends on the electrostatic attraction (nuclear charge) that the nucleus exerts on the electrons of the outer shell. More the nuclear charge closer are the shell and electrons, hence smaller is the atomic radius of an atom.
With the increase in the atomic number (increased number of protons, electrons and neutrons) in the 3rd period, the net positive charge of the nucleus gradually increases. This increased positive charge exerts a greater attraction on the shells and attract the electrons in the shells a little closer to the nucleus. Hence, sodium has the largest atom and chlorine the smallest. This is true of other periods as well.
Conversely, elements in the same group increase in size downward. For example, in group 1 starting from lithium to sodium, potassium, rubidium and caesium, the atomic size increases because there is a gradual increase in the number of shells.
| Group I A elements | Electronic configuration | No.of elememnts |
|---|---|---|
| Hydrogen | 1 | 1 |
| Lithium | 2,1 | 2 |
| Sodium | 2,8,1 | 3 |
| Potassium | 2,8,8,1 | 4 |
| Rubidium | 2,8,18,8,1 | 5 |
| Caesium | 2,8,18,18,8,1 | 6 |
| Francium | 2,8,18,32,18,8,1 | 7 |
| Group II A elements | Electronic configuration | No.of elememnts |
|---|---|---|
| Beryllium | 2,2 | 2 |
| Magnesium | 2,8,2 | 3 |
| Calcium | 2,8,8,2 | 4 |
| Strontium | 2,8,18,8,2 | 5 |
| Barium | 2,8,18,18,8,2 | 6 |
| Radium | 2,8,18,32,18,8,2 | 7 |
| Remember |
| In the case of noble (inert) gases i.e., in the Zero group there are exceptions and the atomic size of the elements may be greater than the other atoms of the period. |
The tendency of an element to lose electrons and form positive ions (cations) is called electropositive or metallic character. For example, alkali metals are the most electropositive elements.
"The tendency of an element to accept electrons to form an anion is called its non-metallic or electronegative character." For example, chlorine, oxygen and phosphorous show greater electronegative or non-metallic character.
In each period, metallic character of elements decreases as we move to the right. Elements to the left of the periodic table have a pronounced metallic character while those to the right have a non-metallic character. Conversely, non-metallic character increases from left to right.
In the third period, sodium on the extreme left is most metallic. The metallic character decreases towards magnesium and aluminium, which are to the right. Silicon is midway between metals and non-metals. From phosphorus to sulphur to chlorine, non-metallic character gradually increases, chlorine being the most non-metallic in behaviour. In the 18 or zero group, argon does not exhibit either metallic or non-metallic character.
The elements to the left of the periodic table have a tendency of losing electrons easily as compared to those to the right. As we move from left to right of the period, the electrons of the outer shell experience greater pull of the nucleus. This greater force of attraction is because the nuclear charge increases and the size of the atom decreases from left to right. Thus, electrons of the elements to the right of the table do not lose electrons easily so are non-metallic in nature.
Metals usually have 1, 2 or 3 electrons in the outermost shell and ionize by giving out these electrons. Thus they gain positive charges equal to the number of electrons lost. Germanium, tin and lead with four electrons each in the valence shell are also included among the metals.
Non-metals usually have 5, 6 or 7 electrons in the outermost shell and ionize by accepting electrons. Thus they gain a negative charge equal to the number of electrons gained. Although carbon and silicon have four electrons each in the valence shell, they are included in the non-metals. Boron is an exception; it has three electrons in the outermost shell but is still included among non-metals.
As we move down the group the number of shells increases. This causes the effective nuclear charge to decrease due to the outer shells being further away: in effect the atomic size increases. The electrons of the outermost shell experience less nuclear attraction and so can lose electrons easily thus showing increased metallic character.
Ionization potential (or ionization energy) is the amount of energy required to remove one or more electrons from the outermost shell of an isolated atom in the gaseous state.
Atom(g) + IE 
Positive ion(g) + electron(g)
Ionization energy is also called as ionization potential because it is measured as the minimum potential required in removing the most loosely held electron from the rest of the atom. It is measured in the units of electron volts (eV) per atom or kilo joules per mole of atoms (kJ mol-1)
Thus, the ionization energy gives the ease with which the electron can be removed from an atom. The smaller the value of the ionization energy, the easier it is to remove the electron from the atom.
An electron is held in an atom by the electrostatic force of the positively charged protons in the nucleus and the negative charge of the electrons. By supplying enough energy, it is possible to remove an electron from an atom. The element is first brought into the vapour state. Then the electron is removed by supplying energy equivalent to the ionization potential.
In referring to the Periodic table of ionization potentials (in electron volts) shown below, the following conclusions may be arrived at:
- Metals usually have low ionization potential whereas non-metals have high ionization potential. Metalloids have intermediate ionization potential.
- The inert gases have very high ionization potential, due to the stability of the outer shell. Helium has the highest ionization potential.
- Within a group, the ionization potential generally decreases with increasing atomic number. Increasing atomic number results in increasing atomic radii. Thus, the electrons of the outer shell are further away than those of the previous element. The effective nuclear charge decreases as atomic size increases. Thus it is easier to pull one electron from the outermost shell of the atom.
- Ionization potential does not necessarily vary uniformly from one element to another. But it is a periodic property. It increases from group 1 to group 18. But the increase is not very regular. Ionization potential increases across the period because of increase in nuclear charge due to which the atomic size decreases. Thus, more energy is required to pull away the electron from the outermost shell of the atom of smaller size.
Electron affinity is the amount of energy released when an electron is added to an isolated gaseous atom.
Atom (g) + electron (g) Anion (g) + energy
Electron affinity is the ability of an atom to hold an additional electron. If the atom has more tendency to accept an electron then the energy released will be large and consequently the electron affinity will be high. Electron affinities can be positive or negative. It is taken as positive when an electron is added to an atom. It is expressed as electron volts per atom (eV per atom) or kilo joules per mole.
Electron affinity depends on:
- Extent of nuclear charge
- Size of the atom
- Electronic configuration
- As a result of the gain in electrons, the atom gains one negative charge. In the case of halogens, all the elements have a high electron affinity, as they need one electron to complete the octect of their outermost shell.
From chlorine to iodine, which ionize by accepting one electron there is a decrease in the electron affinity or the energy released. The lower electron affinity of fluorine when compared to chlorine is not fully understood.
| Element | Electron Affinity |
|---|---|
| Fluorine | 3.62 eV |
| Chlorine | 3.79 eV |
| Bromine | 356 eV |
| Iodine | 3.28 eV |
f the electron affinity is low, the electron is weakly bound; if the electron affinity is high, the electron is strongly bonded, e.g., electron affinity of chlorine is 3.79 which is higher than that of iodine i.e., 3.28. Hence, chlorine accepts the electrons more easily than iodine.
- Electron affinity increases from left to right across the period because of increase in nuclear charge and decrease in atomic size. This causes the incoming electron to experience a greater pull of the nucleus thus giving a higher electron affinity.
- Electron affinity decreases down the group because the number of shells increases i.e., the atomic size increases and the effective nuclear charge decreases. This causes the incoming electron not to experience much attraction of the nucleus thus giving a lower electron affinity.
- The electron affinity of completely filled atoms is almost zero. An atom does not accept an electron in its outermost shell if it already has a stable configuration i.e. a duplet or octet, as in the case of inert gases.
Electronegativity is the tendency of an atom to attract electrons towards itself in a molecule of a compound. The value of electronegativity of an element describes the ability of its atom to compete for electrons with the other atom to which it is bonded. Electronegativity is however not the property of an isolated atom.
Electronegativity increases from left to right in each period ending at group 17.
In the 3rd period, electronegativity increases from sodium to chlorine i.e., chlorine can accept electrons most easily in that period followed backwards by sulphur, phosphorus, silicon, aluminium, magnesium and sodium. All the atoms of the above mentioned elements have three shells but chlorine has the smallest atomic radii. Hence chlorine experiences more positive charge from the nucleus than all other atoms in that period. So, if one electron is available, chlorine can attract it most easily.
When the molecule is formed by transfer of electrons (ionic bonding) the transfer takes place from electropositive atom to electronegative atom. In the example below, Na is electropositive and Cl is electronegative.
If the molecule is formed by sharing of electrons (covalent bond) the bonded pair of electrons shifts towards more electronegative atom resulting in the formation of polar molecule. In the example below, chlorine atom is more electronegative as compared to hydrogen atom, resulting in a covalent bond where the shared pair of electron shifts towards the more electronegative atom. This results in polar molecules.
The electron pair is closer to the chlorine atom and so the molecule gets polarized i.e., the chlorine atom gets a negative charge while the hydrogen atom gets a positive charge.
A summary of periodic properties and their variation in groups and periods is given below:
| Remember |
| Fluorine is the most electronegative element. |
