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Chemical elements in s-
Group 1 2 18
1 1
2 3
3 11
4 19
5 37
6 55
7 87

The s-block is one of four blocks of elements in the periodic table. The element of s- group have a common property. The electron in their most outward electron shell are in the s-orbital.[1] Elements in the s- are in the first two periodic table groups.[2] The elements in group one are called the alkali metals. The elements in group two are called the alkaline earth metals.

The modern periodic law says that "The properties of elements are periodic function of their atomic number." This means that some properties of elements are repeated as the atomic number of the elements gets larger. These repeating properties have been used to separate the elements into four s. These s are s-, p-, d-, and f-.

Group I
Element atomic
Hydrogen 1 1
Lithium 3 2,1
Sodium 11 2,8,1
Potassium 19 2,8,8,1
Rubidium 37 2,8,18,8,1
Caesium 55 2,8,18,18,8,1
Francium 87 2,8,18,32,18,8,1
Group II
Element atomic
Beryllium 4 2,2
Magnesium 12 2,8,2
Calcium 20 2,8,8,2
Strontium 38 2,8,18,8,2
Barium 56 2,8,18,18,8,2
Radium 88 2,8,18,32,18,8,2
Element atomic
Helium 2 2

Properties of s- elements[change | change source]

All of the s- elements are metals (except Hydrogen). In general, they are shiny, silvery, good conductors of heat and electricity. They lose their valence electrons easily. In fact, they lose their trademark s orbital valence electrons so easily that the s- elements are some of the most reactive elements on the periodic table.

The elements in group 1, known collectively as the alkali metals (except hydrogen), always lose their one valence electron to make a +1 ion. These metals are characterized by being silvery, very soft, not very dense and having low melting points. These metals react extremely vigorously with water and even oxygen to produce energy and flammable hydrogen gas. They are kept in mineral oil to reduce the chance of an unwanted reaction or worse, an explosion.

The elements in group 2, known as the alkaline earth metals (except helium), always lose their two valence electrons to make a +2 ion. Like the alkali metals, the alkaline earth metals are silvery, shiny and relatively soft. Some of the elements in this column also react vigorously with water and must be stored carefully.

S- elements are famous for being ingredients in fireworks. The ionic forms of potassium, strontium and barium make appearances in firework displays as the brilliant purples, reds and greens.

Francium is considered to be the most rare naturally occurring element on earth. It is estimated that there is only ever one natural atom of Francium present on earth at a time. Francium has a very unstable nucleus and undergoes nuclear decay rapidly.

Chemical properties of alkali metals

1.Alkali metals react with dry hydrogen to form hydrides.

  • These hydrides are ionic in nature
  • These hydrides of alkali metals react with water to form corresponding hydroxides and hydrogen gas. - LiH+ H2O->LiOH+H2
  • These hydrides are strong reducing agents and their reducing nature increases down the group.
  • Alkali metals also form complex hydrides such as LiAlH4 which is a good reducing agent. Alkali metal hydrides do not exist in water and this reaction with any other agent is carried out in protic solvent.
  • Fused alkali metal hydrides on electrolysis produce H2 gas at anode.

2.Formation of oxides and hydroxides.

  • These are most reactive metals and have strong affinity towards O2 ,they form oxides on surface. They are kept under kerosene or paraffin oil to protect them from air.
  • When burnt in air (O2) ,li forms Li20 , Na forms Na 2O2 and other alkali metals form superoxides.

3. They are purely metallic , as they lose the electrons from the outermost shell readily. They are highly reactive metals and they have low ionization energy .

4. Beryllium is amphoteric in nature .

Diagonal relationship[change | change source]

Group 1 2 13 14
Period 2 Li Be B C
Period 3 Na Mg Al Si

The first element in group one, Lithium, and the first in group two, Beryllium, behave differently to other members of their groups. Their behaviour is like the second element of the next group. So lithium is similar to magnesium, and beryllium is similar to aluminum.

In the periodic table this is known as a 'diagonal relationship'. The diagonal relationship is because of similarities in ionic sizes and charge/radius ratio of the element. The similarity between lithium and magnesium is because of their similar sizes:

Radii, Li=152pm Mg=160pm

Lithium[change | change source]

Lithium has many different behaviours to other elements in group one. This difference caused by:

  1. the small size of the lithium atom and its ion.
  2. the higher polarization power of li
    (i.e. charge size ratio). This means increased covalent character of its compounds which is responsible for their solubility in organic solvents
  3. high ionisation enthalpy and high electronegative character of lithium as compared to other alkali metals
  4. non availability of d-orbitals in its valence shell
  5. strong intermetallic bonding

Some of the ways in which lithium behaves differently from other members of are:

  1. Lithium is harder than sodium and potassium which are so soft that they can be cut by a knife.
  2. The melting and boiling points of lithium are higher.
  3. Lithium forms monoxide with oxygen, other alkali form peroxide and superoxide.
  4. Lithium combines with nitrogen to form nitrides, while other alkali metals do not.
  5. Lithium Chloride is deliquescent and crystallizes as a hydrate LiCl.2H2O. Other alkali metal chlorides do not form hydrates.

References[change | change source]

  1. "Electron Configuration for all the elements in the Periodic Table". Retrieved 22 October 2016.
  2. "Archived copy" (PDF). Archived from the original (PDF) on 2013-01-23. Retrieved 2012-02-16.{{cite web}}: CS1 maint: archived copy as title (link)

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