The elements of Group 1 and Group 2 of the modern periodic table are called S-block elements. Two types of s block elements are possible, i.e., the elements with one electron (s1) or the elements with two electrons (s2) in their s-subshell.
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S-block comprises 14 elements, namely hydrogen (H), lithium (Li), helium (He), sodium (Na), beryllium (Be), potassium (K), magnesium (Mg), rubidium (Rb), calcium (Ca), caesium (Cs), strontium (Sr), francium (Fr), barium (Ba), and radium (Ra).
S-block elements having only one electron in their s-orbital are called group one or alkali metals, whereas the s block elements having two electrons filling their s-orbital are called group two or alkaline earth metals.
The electrons present in an atom occupy various sub-orbitals of available energy levels in the order of increasing energy. The last electron of an atom may find itself in either of the s, p, d and f subshells. Accordingly, the elements of the atom having their last valence electron present in the s-suborbital are called s block elements.
The electronic configuration of S-block elements is explained below.
The alkali elements in the s-block consist of a single valence electron in their outermost shell. This outermost electron is loosely held, which makes these metals highly electropositive. Due to this, they are not available in a free state in nature. The general electronic configurations of s block elements – group 1 are as shown in the table below:
| Element | Symbol | Electronic Configuration |
|---|---|---|
| Lithium | Li | 1s 2 2s 1 |
| Sodium | Na | 1s 2 2s 2 2p 6 3s 1 |
| Potassium | K | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 |
| Rubidium | Rb | 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1 |
| Caesium | Cs | [Xe]6s 1 |
| Francium | Fr | [Rn]7s 1 |
The electronic configurations of elements included in group 2 of S block elements are shown below:
| Elements | Symbols | Electronic Configuration |
| Beryllium | Be | [He]2s 2 |
| Magnesium | Mg | [Ne]3s 2 |
| Calcium | Ca | [Ar]4s 2 |
| Strontium | Sr | [Kr]5s 2 |
| Barium | Ba | [Xe]6s 2 |
| Radium | Ra | [Rn]7s 2 |
Both alkali and alkaline earth elements show a regular gradation in their properties among their respective group elements. But the first member of both S block elements, namely, Lithium and Beryllium, differ much from the rest of their members, but at the same time, they resemble more with the diagonal element present in the next column.
The anomaly of these S-block elements is due to the following:
Greater polarization of s block elements makes the first element more covalent and differentiates them from the rest, which is ionic.
The similarity in size and charge density makes them resemble the element diagonally placed in the next group (diagonal relationship).
It is observed that the physical and chemical properties of these s block elements change in a particular trend as the atomic number of the elements increases. Changes in the various properties of the group are as mentioned below:
When the s block elements of the modern periodic table are observed, it is seen that the size of the alkali metals is larger compared to other elements in a particular period. As the atomic number increases, the total number of electrons increases along with the addition of shells.
On moving down the group, the atomic number increases. As a result, the atomic and ionic radius of the alkali metals increases.
As we go down the group, the size of the atoms increases, due to which the attraction between the nucleus and the electrons in the outermost shell decreases. As a result, the ionization enthalpy decreases. The ionization enthalpy of the alkali metals is comparatively lesser than other elements.
As the ionic sizes of the elements increase, the hydration enthalpy decreases. The smaller the size of the ion, the hydration enthalpy is high as the atom has the capacity to accommodate a larger number of water molecules around it due to the high charge/radius ratio, and hence gets hydrated.
A diagonal relationship in s block elements exists between adjacent elements, which are located in the second and third periods of the periodic table. For example, the lithium of group 1A and the second period shows similarities with the properties of magnesium, which are located in the 2nd group and 3rd period.
Similarly, the properties of beryllium, which are located in the 2nd group and 2nd periods, show a likeness with the properties of aluminium which is located in the third period and third group. The two elements which show similarities in their properties can be called diagonal pairs or diagonal neighbours.
The properties of s block elements vary significantly when compared to the other elements of the sub-group they belong to. The diagonal neighbours show a lot of similarities. Such a relationship is exhibited as you move left to the right and down the group; the periodic table has opposing factors.
For example, the electronegativity of the s block elements increases as we go across the period and decreases as we go down the group. Therefore, when it is moved diagonally, the opposite tendencies cancel out, and the value of electronegativity almost remains the same.
Lithium and beryllium, the first members of the s-block family, differ much from the rest of their members. The anomaly of these elements is due to their,
Greater polarization makes the first element form covalent compounds compared to the rest, which are ionic.
Sodium and magnesium form oxides with lower oxidation numbers, while heavier atoms form oxides with higher oxidation numbers. Sodium forms oxide and peroxide, whereas oxygen has an oxidation number of -2 and -1, respectively.
In magnesium oxide, oxygen is in a -2 state. But caesium forms super-oxides where the oxidation state of oxygen is – 0.5. Similarly, the heavier barium form peroxide, having an oxidation state of oxygen as -1.
Solubility depends on two factors:
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Covalent compounds like beryllium sulphate have a higher enthalpy of dissociation than ionic barium sulphate. But the smaller entities like beryllium have higher charge density, resulting in higher solvation, and hence, the release of hydration enthalpy is larger than the dissociating energy.
So, BeSO4 is more soluble than ionic BaSO4.
S-block elements are strong electropositive elements with low reduction potential, indicating their strong reducing ability compared to others. So, substances having lower reducing ability than them will not be able to reduce them. Reducing the ability of an atom is related to the ease of releasing electrons for reduction. Decreasing ionization energy down the column suggests that caesium is a stronger reducing agent than Lithium.
But, reducing ability (oxidation potential) depends on the combined energy difference of three processes:
Lithium, being the smallest ion, its hydration enthalpy is very high than caesium and compensates more than its higher ionization enthalpy. Thus, lithium has the highest reducing ability (highest oxidation potential or lowest reduction potential = -3.04V) compared to caesium.
S block elements, or their halides on exposure to flame, undergo electronic transitions in the visible region of the light spectrum. Hence, they induce characteristic colour into the flame. The colours are as follows:
| Metals | Lithium | Sodium | Potassium | Rubidium | Caesium |
| Flame colour | Crimson red | Yellow | Violet | Red violet | Blue |
| Metals | Beryllium | Magnesium | Calcium | Strontium | Barium |
| Flame colour | – | – | Brick red | Crimson red | Apple green |
