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Periodic Table
Periodic Trends
The specific patterns in the properties of chemical elements present in the periodic table are known as periodic trends. The important trends are,
- Ionization energy
- Metallic character
- Atomic Radii
- Electronegativity
- Ionic radius
- Electron affinity
- Chemical reactivity
- Shielding effect
These trends arise due to changes in the structure of atoms of the elements within their groups and periods. A few exceptions exist, for example, the ionization energy of groups 3 and 6.
Periodic Law
Periodic law forms the basis for periodic trends. According to periodic law, “the chemical elements are listed in an order of increasing atomic number, and main properties thus undergo cyclic changes. Elements having similar chemical properties re-occur in regular intervals”
This principle was given by Dmitri Mendeleev. He also stated that the periodic table was not just based on the atomic weights, but also based on various physical and chemical properties of elements.
Later it was also found that the recurrence of properties was due to the recurrence of similar electronic configurations in the outer shells of atoms.
1. Ionization Energy
The ionization potential can be defined as,
“Minimum energy required by an isolated atom to remove one electron in its neutral or gaseous state”
As one goes across the period, the ionization energy increases. The reason behind this is that the nuclear charge across the period increases and thus the electrons are strongly held by the nucleus.
But as one goes down the group, the ionization energy decreases down the group. The reason behind this is, down the group the valence electrons go farther away from the nucleus, thus the nuclear charge decreases.
Factors affecting ionization energy
Various Factors that Affect the Ionization Energy Levels
Nuclear Charge
Lower the nuclear charge lower the force of attraction between the nucleus and valence electrons, thus low ionization energy.
Shielding Effect
The shielding effect increases as the nuclear charge increases, thus with an increase in shielding effect the ionization energy also increases.
Atomic Radius
As the atomic radius increases the force of attraction between the nucleus and valence electrons also decreases. Thus, with an increase in atomic radius, the ionization decreases.
Half-Filled Valence Shells
Pseudo-filled or half-filled valence shells have high ionization energy.
A simple principle that can be used is that, if the principal quantum number is low, then the ionization number will be high for the electron present in that shell.
Exceptions
All the elements in the oxygen and boron family are an exception to the above-stated periodic trend. They require a little less energy than the usual trend.
2. Metallic Property
The metallic property of an element can be defined as its ability to conduct electricity. The metallic properties increase down the group as the nuclear charge decreases down the group. Since the valence electron is loosely bounded by the nuclei, they are able to conduct electricity well.
But across a period, the metallic character decreases as the nuclear charge increases. This causes the force of attraction between the valence electrons and the nuclei to increase, thereby inhibiting them from conducting electricity or heat.
3. Atomic Radii
The atomic radius is the distance between the atomic nucleus and the outermost stable electron orbital of an atom which is at equilibrium. Across a period the atomic radius decreases, as the nuclear charge increases. The reason for the decrease is as nuclear charge increases, the force of attraction between the nucleus and the valence electrons also increases, and the nucleus holds the electron tightly, thereby decreasing the atomic radii.
In a group, the atomic radius increases down the group. The reason is, new shells are being added and thus the nuclear charge decreases. But the atomic radii also increase diagonally causing some exceptions.
Example:
Along the Period – Li> Be > B > C > N > O > F
Down the Grp – Li < Na < K < Rb < Cs
4. Electronegativity
Electronegativity can be defined as the ability of an atom or a molecule to attract a pair of electrons. The bond formed due to this is mainly determined by the difference between the electronegativity of the atoms.
Across the period, the electronegativity increases as the nuclear charge increases. Moving down a group, the electronegativity decreases as the nuclear charge decreases. The reason being the distance between the nucleus of the atom and the valence electrons is long and thus the electrons are easily lost.
Example:
Along the Period- Li < Be < B < C < N < O < F
Down the Grp – Li > Na > K > Rb > Cs
Exception
The group 13 elements are an exception and thus the electronegativity increases from aluminum to thallium. Also, in group 14, the electronegativity of tin is higher than lead.
5. Electron Affinity
Electron affinity can be defined as the tendency of an atom to accept an electron or an electron pair. This is a characteristic feature of nonmetals as they gain electrons to become anions. Across a period, the electron affinity increases as nuclear charge increases.
Down the group, it decreases, as the nuclear charge decreases. Fluorine has the highest electronegativity and noble gasses are not included in this. The reason being they have a full valence shell and thus can neither gain nor lose electrons.
6. Shielding Effect
It can be defined as the repelling of an outer electron by the inner electrons. It can also be used to explain how many nuclei can control the outer electrons.
The effective nuclear charge decreases down the group due to the increased shielding effect. Across a period, the effective nuclear charge increases as the nuclear charge increases.
| Language | Gujarati |
| No. of Pages | 1 |
| PDF Size | 0.05 MB |
| Category | Education |
| Source/Credits | – |
Periodic Table PDF Free Download