Trends represent the range of positive or negative changes occurring in a specific area of interest. Periodic tables have been edited a number of times based upon many reasons. Each year many new elements and newly researched thing have been added. Some of the common periodic trends are electronegativity, atomic radii, ionic radius, metallic character, ionization energy, electron affinity, and chemical reactivity. The main reason behind the periodic trends is the change in the atomic structure of the chemical elements within their respective period. Periodic laws make certain that the chemical elements should be arranged in a periodic table on the basis of their atomic structures and properties. Some of the basic trends of the periodic table are as explained below.
Atomic radius is the distance between the centre of the nucleus of an atom to the outermost energy level. It is not easy to measure the size of an atom. This is because the possibilities of locating the outermost valence electron far away from the centre is difficult. Depending on the current position of an atom, we can use various types of atomic radius; they are the covalent radius, crystal/metallic radius, and van der Waal’s radius.
1. Covalent radius
Covalent radius denoted as rcov refers to the measurement ( unit) of the atomic radius resulting in the formation of a single covalent bond. The commonly used units of measurements are the pm (picometres) and Å ( Angstrom). One angstrom is equivalent to a hundred picometres. The mathematical relationship between any 2 atoms can be represented as follows:
R (AB) = r (A) + r (B).
From the above relationship, we can introduce radii for a single bond, double bonds as well as triple bonds.
2. Van der Waal’s Radius
Van der Waal’s radius refers to half of the internuclear distance between adjacent atoms of the two neighbouring molecules in the solid-state. The radius can be calculated as follows:
3. Crystal/metallic radius
Metallic radius refers to half of the distance between the two adjacent metal ions in the metallic structure. It mostly depends on the nature of the atom as well as its environment. Some of the external factors such as temperature and pressure exerted on the metal are also responsible for the phenomenon. As the effective nuclear charge increase across the period, the metallic radii decrease. Metallic radius for A-A can be calculated as:
d = rA+rA rA=
In summary, rcovalent< rmetallic < rvan der Waal
Atomic and Ionic Radii of the Elements Of The Periodic Table
Atomic radius decreases as you move from left to right of the period. The progressive decrease in size is the result of a gradual increase in positive nuclear charge because, as we move from left to right of the table, the number of protons increases progressively. The same trend is seen with the number of electrons, but all the additional electrons occupy the same energy level. As the number of protons increases, so does the positive nuclear charge. Similarly, there is an increase in the atomic number as you move down the group. This is due to the increase in the principal quantum number. There is a general trend in an increase in the first ionization energy from left to right of the period. As discussed earlier, atomic radii decrease progressively as you move from LHS to RHS of the periodic table. The result is that the outermost electrons of the elements on the RHS nearer to the nucleus in comparison with those elements found on the LHS. Consequently, the force of attraction exerted by the nucleus on the outermost electrons increases as you move from LHS to RHS of the period. On the other, the ionic radius increases as you move down the group. This is due to the increase in the principal quantum number. In most cases, metals form ions by losing an electron(s) to form cations while non-metals gain an electron(s) to form anions.
Ionization energy refers to the minimum energy required to overcome the attraction exerted by the nuclear charge and remove an electron from a gaseous atom. Removal of an electron leads to the formation of cations, M+. The amount of energy needed to remove the 1st outermost valence electron is referred to as 1st ionization energy. Second ionization energy is the amount of energy required to remove a subsequent electron from a gaseous ion. Generally, the 2nd ionization energy is greater than 1st ionization energy. This is because it is easier to remove an electron from a group 1A member for it to form a cation. On the other hand, it is difficult to remove an electron in the outermost energy level. As you move down the group, 1st ionization energy decreases due to the increase of the size of atoms down the group thus making the valence electron to be far away from the nucleus.
The 1st ionization energy of the Alkali metals increases as one moves from LHS and RHS of the period. The nuclear charge increases and the shielding effect are constant as you move across the period. A greater attraction of the nucleus for the electron leads to an increase in ionization energy. As you move from left to right of the period, the electron gain enthalpy tends to be more negative because of the decrease in the atomic size that leads to an increase in the nuclear charge. For this reason, the addition of an extra electron is favoured since there is an increased force between the incoming electron and the nucleus.
Trends in electronegativity
“Electronegativity of an element is the tendency of the atoms in an element to attract electrons when they are chemically combined with atoms of another element”.
Electro-negativities have been calculated for the elements and are expressed in arbitrary units on the Pauling electronegativity scale. The numerical scale is based on some factors, including the ionization energies of elements. Electronegativity generally decreases as one moves down the group. On the other hand, as one moves from left to right of the periodic table, the electronegativity of representative elements increases. By contrast, the non-metallic elements at the far right excluding the noble gases have high electronegativity. You will realize that the electronegativity of transition elements is irregular and they are not tabulated. Electronegativity of Cs which is the least electronegative element is approximately 0.7 while that of fluorine is four since fluorine has a strong tendency to attract electrons during the chemical reaction.
Electron Gain Enthalpy
Electron gain enthalpy is the process through which an electron is added to a neutral gaseous atom to form an anion. The process is crucial because it provides a measure of the ease with which an atom adds an electron to form a negatively charged ion as. This can be written as
Y (g) + e- Y-(g)
When an electron is added to an atom of an element, two forces act on it. These are an attraction by the positively charged nucleus and the repulsion force by the electrons in the atom. In most cases, the addition of the first electron releases energy but the addition of subsequent electrons requires a supply of energy. As mentioned earlier, non-metals gain an electron(s) with the aim of attaining stability. For instance, group 7A members such as chlorine require only a single electron for them to attain stability. Group 6A members such as Sulphur require two electrons for them to form a stable ion. Depending on the element, the process of adding an electron(s) to atoms can either be endothermic or exothermic. In most cases, energy is released during the gaining of electrons, and the electron enthalpy is always negative. For this reason, group 7A elements such as chlorine have a high affinity because they can attain stability by gaining electrons. Noble gases, on the other hand, have exhibited a large positive electron gain enthalpies since the electron has to enter the next higher principal quantum level leading to a very unstable configuration.
As the atomic size increases down the group, electron gain enthalpies decrease.
Nomenclature of elements with atomic number greater than 100