STRUCTURE ATOM AND PERIODIC TRENDS ( CH 7 )

PERIODIC TRENDS

In chemistry, periodic trends are the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of the periodic table of elements.

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 radius tends to decrease as one progresses across a period from left to right because the effective nuclear charge increases, thereby attracting the orbiting electrons and lessening the radius. The atomic radius usually increases while going down a group due to the addition of a new energy level (shell). However, diagonally, the number of electrons has a larger effect than the sizeable radius. For example, lithium (145 picometer) has a smaller atomic radius than magnesium (150 picometer).^{[citation needed]} Atomic radius decreases from left to right across a period, and also increases from top to bottom down a group.

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 one, the nth ionization energy is the energy required to remove the atom's nth electron, after the (n−1) electrons before it have been removed. Trend-wise, ionization energy tend 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.^{[citation needed]} 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.

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

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.^[1] Consequently, atoms with a more negative electron affinity value are considered to have a lower 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.in this case, class 7a which has the highest electron affinity than the other groups. Reactivity elements of group 7 down from top to bottom in a group - fluorine is the most reactive and most non reactive iodine.

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 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.

In the group 13 elements electronegativity increases from aluminium to thallium. In group 14 electronegativity of lead is higher than that of tin.

Metallic properties

Metallic property decreases across a period with increase in number of valence electrons as well as a decrease in atomic radius, and it increases down the group with increase in number of shells and atomic radius

Non-Metallic properties

Non-metallic property increases across a period and decreases down the group due to the same reason.