Periodic trends of chemical elements
Major periodic trends include: electronegativity, ionization energy, electron affinity, atomic radius, melting point, metallic character, and ionic radius which fully explained below as you read on. Periodic trends, arising from the arrangement of the periodic table, provide chemists with an invaluable tool to quickly predict an element's properties. These trends exist because of the similar atomic structure of the elements within their respective group families or periods, and because of the periodic nature of the elements.
This principle
was discovered after number of investigations done by scientists in nineteenth
century such as Lothar Meyer and Dmitri
Mendeleev. Initially, no theoretical explanation for the Periodic
Law was available and it was used only as an empirical principle. But, with the
development of electronic theory of atomic structure, it became possible to
understand the theoretical basis for the Periodic Law. From the modern periodic
table, it is evident that the periodic recurrence of elements with
similar physical and chemical properties, when the elements are listed in order
of increasing atomic number, results directly from the periodic recurrence of
similar electronic configurations in the outer shells of respective atoms.
Discovery of
Periodic Law constitutes one of the most singularly important events in the
history of chemical science. Almost every chemist makes extensive and continued
use of Periodic Law. Periodic Law also led to the development of the periodic
table, which is widely used nowadays.
Atomic radius
Main article:
Atomic radius
The atomic
radius is the distance from the atomic nucleus to the outermost stable electron
orbital in an atom that is at equilibrium. The atomic radii tend to decrease
across a period from left to right. The atomic radius usually increases while
going down a group due to the addition of a new energy level (shell). However,
atomic radii tend to increase diagonally, since the number of electrons has a
larger effect than the sizeable nucleus. For example, lithium (145 picometer)
has a smaller atomic radius than magnesium (150 picometer).
Atomic radius
can be further specified as:
Covalent
radius: half the distance between two atoms of a diatomic compound, singly
bonded.
Van der Waals
radius: half the distance between the nuclei of atoms of different molecules in
a lattice of covalent molecules.
Metallic
radius: half the distance between two adjacent nuclei of atoms in a metallic
lattice.
Ionic radius:
half the distance between two nuclei
Ionization energy
The ionization
potential is the minimum amount of energy required to remove one electron from
each atom in a mole of atoms in the gaseous state. The first ionization energy
is the energy required to remove two, the ionization energy is the energy
required to remove the atom's nth electron, after the (n−1) electrons before it
has been removed. Trend-wise, ionization energy tends to increase while one
progresses across a period because the greater number of protons (higher
nuclear charge) attract the orbiting electrons more strongly, thereby
increasing the energy required to remove one of the electrons. Ionization
energy and ionization potentials are completely different. The potential is an
intensive property and it is measured by "volt"; whereas the energy
is an extensive property expressed by "eV" or "kJ/mole".
As one
progresses down a group on the periodic table, the ionization energy will
likely decrease since the valence electrons are farther away from the nucleus
and experience a weaker attraction to the nucleus's positive charge. There will
be an increase of ionization energy from left to right of a given period and a
decrease from top to bottom. As a rule, it requires far less energy to remove
an outer-shell electron than an inner-shell electron. As a result, the ionization
energies for a given element will increase steadily within a given shell, and
when starting on the next shell down will show a drastic jump in ionization
energy. Simply put, the lower the principal quantum number, the higher the
ionization energy for the electrons within that shell. The exceptions are the
elements in the boron and oxygen family, which require slightly less energy
than the general trend.
Electron affinity
The electron
affinity of an atom can be described either as the energy gained by an atom
when an electron is added to it, or conversely as the energy required to detach
an electron from a singly charged anion. The sign of the electron affinity can
be quite confusing, as atoms that become more stable with the addition of an
electron (and so are considered to have a higher electron affinity) show a
decrease in potential energy; i.e. the energy gained by the atom appears to be
negative. For atoms that become less stable upon gaining an electron, potential
energy increases, which implies that the atom gains energy. In such a case, the
atom's electron affinity value is positive.[2] Consequently, atoms with a more
negative electron affinity value are considered to have a higher electron
affinity (they are more receptive to gaining electrons), and vice versa.
However, in the reverse scenario where electron affinity is defined as the
energy required to detach an electron from an anion, the energy value obtained
will be of the same magnitude but have the opposite sign. This is because those
atoms with a high electron affinity are less inclined to give up an electron,
and so take more energy to remove the electron from the atom. In this case, the
atom with the more positive energy value has the higher electron affinity. As
one progresses from left to right across a period, the electron affinity will
increase.
Although it
may seem that Fluorine should have the greatest electron affinity, the small
size of fluorine generates enough repulsion that Chlorine has the greatest
electron affinity.
Electronegativity
Electronegativity
is a measure of the ability of an atom or molecule to attract pairs of
electrons in the context of a chemical bond. The type of bond formed is largely
determined by the difference in electronegativity between the atoms involved,
using the Pauling scale. Trend-wise, as one moves from left to right across a
period in the periodic table, the electronegativity increases due to the
stronger attraction that the atoms obtain as the nuclear charge increases.
Moving down in a group, the electronegativity decreases due to the longer
distance between the nucleus and the valence electron shell, thereby decreasing
the attraction, making the atom have less of an attraction for electrons or
protons.
However, in
the group 13 elements electronegativity increases from aluminium to thallium,
and in group 14 electronegativity of lead is lower than that of tin.
Valence
electrons
Main article:
Valence electrons
Valence
electrons are the electrons in the outermost electron shell of an isolated atom
of an element. Sometimes, it is also regarded as the basis of Modern Periodic
Table. In a period, the number of valence electrons increases (mostly for light
metal/elements) as we move from left to right side. However, in a group this
periodic trend is constant, that is the number of valence electrons remains the
same.
Valency
Valency in the
periodic table across a period First Increases then Decreases.In the periodic
table going from top to bottom remains same.
However, this
periodic trend is sparsely followed for heavier elements (elements with atomic
number greater than 20), especially for lanthanide and actinide series.
It is also
important to consider the core electrons when speaking about the valence
electrons.
Metallic and
non-metallic properties
Metallic
properties increase down groups as decreasing attraction between the nuclei and
the outermost electrons causes the outermost electrons to be loosely bound and
thus able to conduct heat and electricity. Across the period, increasing
attraction between the nuclei and the outermost electrons causes metallic
character to decrease.
Non-metallic
property increases across a period and decreases down the group due to the same
reason.due to decrease in nuclear attractive force
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