Atomic Radius: Atomic radius generally decreases across a period. It happens because within a period, outer electrons are added in the same valence shell. With increase in number of protons, the effective nuclear charge increases. This results in increased attraction of electrons towards the nucleus. Due to this, the atomic radii decrease from left to right across a period. When we move down a group, a new shell is added at subsequent period. Due to this, valence electrons become farther from the nucleus. So, atomic radius increases as we move down a group.
Ionic Radius: A cation is smaller than its parent atom because it has less number of electrons while its nuclear charge remains the same. On the other hand, an anion is larger than its parents because of addition of electrons results in increased repulsion among the electrons and a decrease in effective nuclear charge.
Isoelectronic Species: Atoms or ions with same number of electrons are called isoelectronic species. For example, O2-, F-, Na+ and Mg2+ have the same number of electrons (10). Because of their different nuclear charges, their radii would be different. The cation with the greater positive charge will have smaller radius due to greater attraction of electrons to the nucleus. Anion with the greater negative charge will have the larger radius because of larger repulsion among electrons.
Ionisation Enthalpy: Energy required to remove an electron from an isolated gaseous atom (X) in its grounded state is called ionisation enthalpy. Energy is always required to remove an electron from an atom. Hence, ionisation enthalpy is always positive. The second ionisation enthalpy is more than the first ionisation enthalpy because it is more difficult to remove an electron from an ion than from an atom.
This graph shows first ionization enthalpy of elements of the second period. The minimum value is for alkali metal, while the maximum value is for noble gas. Same trend can be observed in each period with minima appearing at alkali metal and maxima appearing at noble gas (as shown in next graph).
The increasing trend in ionization enthalpy can be attributed to increased effective nuclear charge on subsequent addition of electron in a period. This is in tune with decrease in atomic radius as we move across a period. When we move down a group, a new shell is added which results in reduction of ionization enthalpy. This explains increase in reactivity through a group and decrease in reactivity across a period.
In the first graph, you will notice that ionization enthalpy of Boron (Z = 5) is less than that of Beryllium (Z = 4), in spite of Boron having a greater nuclear charge. To understand this, let us look at electronic configurations of these elements.
Be (4): 1s2 2s2
B (5): 1s2 2s2 2p1
In case of Beryllium, the outermost electron is in s orbital, while the outermost electron in Boron is in p orbital. Electrons in s orbital are more attracted towards the nucleus compared to electrons in p orbital. Due to this, ionization enthalpy of beryllium is less than that of Boron. This happens because a fully filled inner orbital effectively shields the outermost electron from nuclear charge.
Electron Gain Enthalpy: The enthalpy change on addition of electron to a neutral gaseous atom (X) is called electron gain enthalpy. The process of adding an electron to an atom can be endothermic or exothermic, and it depends on the element. So, electron gain enthalpy can be either negative or positive. For example; halogens have very high negative electron gain enthalpies because they just need to gain one electron to attain noble gas configuration. But noble gases have very high positive electron gain enthalpies because electron has to enter the next higher principal quantum level; resulting in a very unstable electronic configuration. Electron gain enthalpies have large negative values towards the upper right side of the periodic table preceding the noble gases.
Compared to ionization enthalpy, the variation in electron gain enthalpies is less systematic. Generally, electron gain enthalpy becomes more negative from left to right in a period. This happens because increase in effective nuclear charge makes it easier to add an electron to a smaller atom. Electron enthalpy generally becomes less negative when we move down a group.
But electron gain enthalpy of O or F is less negative than that of succeeding elements. This is because the extra electron goes to the smaller n = 2 quantum level and suffers significant repulsion from other electrons present in this level. For S or Cl, the n = 3 quantum level, electron-electron repulsion is much less.
Electronegativity: A qualitative measure of ability of an atom (in a chemical compound) to attract shared electrons is called electronegativity. Electronegativity increases across a period and decreases down a group. This can be correlated with atomic radii. So, electronegativity increases with decrease in atomic radius. Metals are less electronegative, while non-metals are more electronegative. Halogens show the highest electronegativity.
Valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals and/or equal to eight minus the number of outermost electrons. Now-a-days, the term oxidation state is frequently used for valence. Oxidation state of an element in a particular compound is defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in that compound.
To understand this, let us take examples of two compunds of oxygen, i.e. OF2 and Na2. The descending order of negativity of these elements is F > O > Na.
The outermost electronic configuration of F, O and Na are as follows:
F: 2s2 2p5
O: 2s2 2p4
F can easily gain 1 electron to attain noble gas configuration. Since there are two atoms of fluorine in OF2, hence oxygen shares its two electrons, one with each atom of fluorine to make this compound. In this compound, the oxidation state of F is -1, while the oxidation state of O is +2.
Na can easily give up 1 electron to attain noble gas configuration. So, oxidation state of Na is +1 and that of O is -2 in Na2O.
This example shows that oxidation state of oxygen can be either -2 or +2.
The first elment of each of the groups 1 and 2 and groups 13-17 differs in many respects from other elements of their respective groups. For example; unlike other alkali metals lithium forms covalent compounds. Similarly, beryllium forms covalent compounds while other members of these groups generally make ionic compounds. Behavior of lithium and beryllium is more similar to the second element of the following group, i.e. Mg and Al respectively. This type of similarity is called diagonal relationship in periodic properties.
This happens because of their sall size, large charge/radius ratio and high electronegativity. Moreover, the first member of group has only four valence orbitals (2s and 2p) available for bonding, while second member of the groups have nine valence orbitals (3s, 3p, 3d). Due to this, the maximum covalency of the first member of each group is 4. On the other hand, other members of the group can expand their valence shell to accommodate more than 4 pairs of electrons. Moreover, first member of p-block elements show greater ability to form multiple bonds to itself, e.g. C=C, C≡C, C=N, C≡N, N=O, etc.
Periodic Trend in Chemical Reactivity: Elements on the extreme left and right of periodic table are highly reactive, while those in middle are moderately reactive.
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