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Variation in Physical Properties of Elements Across the Periodic Table
Elements of the periodic table show variation in
physical properties along
a period and down a group. These physical properties are:
This is the energy required to remove an electron from the outermost shell of a free gaseous atom to form a free gaseous ion. The energy required to remove the first electron is called the first ionization energy; to remove the second, the energy is the second ionization energy, etc.
The greater the ease in the removal of electron, the lower the ionization
energy, and vice versa.
Along a particular period, i.e., from left to right, the ionization energies of elements increases, i.e., it becomes more difficult to ionize an atom. E.g., in period 3, the ionization energies of the elements is
in the order:
Na < Mg < Al < Si < P < S < Cl < Ar
The reason for this order is that: as we go from left to right along a period, there is increase in the nuclear charge, (resulting from the successive increase in the number of protons), but there is no corresponding increase in number of shells (i.e. the number of shells along a period is constant, e.g. for period 3, it is three). Hence, the force of attraction between the protons and the outward electrons becomes progressively stronger, thereby making it more difficult for an electron to be removed from the outermost shells of the atoms of the elements.
Along a particular group, due to increase in the number of shells down the group, the ionization energies of the elements decrease. This is because the electrons in the outermost shells become progressively farther away and shielded from the nucleus. Even with increase in the number of protons, the force of attraction between the protons and the outermost electrons are recorded to be weaker down a group and stronger up a group.
Hence, it becomes easier to remove the outermost electrons down a group.
For example, in group IA, the ionization energies are in the order:
K < Na < Li < H
Atomic Sizes or Radii
Along a period, atomic radii decrease due to increase in nuclear charge without a corresponding increase in shells. The outermost electrons become more pulled to the nucleus, resulting in the atom shrieking and thus, smaller.
For example, in period 3, an atom of sodium is the largest in size, while an atom of Argon is the smallest.
Down a group, due to increase in shells, the outward electrons are progressively shielded from the nucleus, thereby reducing their pull to the nucleus. The result is that the sizes or radii of atoms of elements in the same group increase downward in the group.
The atomic size or radius of an element can be deduced by dividing the distance between the nuclei of two atoms brought in covalency by 2.
When an atom losses an electron to form
a positive ion, its radius decreases.
This is because the nuclear charge (number of protons) now acts on fewer electrons, thereby pulling them closer to the nucleus. However, when an atom gains an electron, the nuclear charge now acts on more electrons, pulling them less closely to the nucleus (this is further compounded by the increase in repulsion between the electrons) – its radius increases.
From the above explanation, notice that cations are smaller than anions. For example, in period 3 - while neutral sodium atom is of larger size or radius than chlorine atom, sodium ion Na+ is of smaller radius than chlorine ion, Cl-.
In a particular group, ionic radii is of the same trend as the atomic radii. I.e., down a group, ionic radii as well as atomic radii increase.
For example, in group IA, the order is ;
K+> Na+> Li+> H+.
This is due to progressive increase in shells.
This is the energy given off when an electron is added to the outermost shell of an atom. The greater the electron affinity, the easier it is for the atom to gain an electron.
Along a period from left to right, electron affinity increases. I.e., it is easier for a chlorine atom to gain an electron than for a sodium atom.
This is due to increase in nuclear charge and smaller atomic radii from left to right of a period.
In a group, electron affinity decreases downwards and increases upward – the farther the outermost shell is from the nucleus, the lower the electron affinity, and vice versa. For example, in group VIIA, the order of electron affinity is Cl > Br > I. Notice that electron affinity of fluorine is unexpectedly lower than that of chlorine.
The observation is due to the relatively small size of the fluorine atom, which results in relatively high repulsion between the electrons. Hence, the order of electron affinity of the halogens is Cl > F> Br > I.
This is the tendency for atoms to acquire electrons and become negative ions (anion). The variation in electronegativity across a period and along a group is in the same order as that of electron affinity, except that in group VIIA, fluorine is now more electronegative than chlorine, though it has less electron affinity than chlorine.
The reason for the trend in electronegativity is same as discussed above for electron affinity.
This is the tendency for atoms to lose electrons and become positive ions (cation). Notice that the higher the ionization energy, the lower the electropositivity, and vice-versa. Hence, electropositivity decreases along a period from left to right , and increases down a group. Notice also that if an atom A has a higher electropositivity than another B, then atom B will have a higher electronegativity than A.
For example, in period 3, Na is more electropositive than Cl, therefore Cl is more electronegative than Na. In group IA, Fr is more electropositive than Na, therefore, Na is more electronegative than Fr.
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