In the previous section, we completed a discussion on hydrogen. In this chapter, we will see s-block elements.
Some basics about group 1 elements can be written in 11 steps:
1. The s-block of the periodic table contains two groups: Group 1 and Group 2.
2. Elements of the group 1 are:
lithium, sodium, potassium, rubidium, caesium and francium.
• In all these elements, the last electron is situated in the outer most s-orbital.
• All these elements are metals.
• These elements form hydroxides on reaction with water. The hydroxides thus formed are strongly alkaline in nature. So these elements are collectively known as alkali metals.
3. Elements of the group 2 are:
beryllium, magnesium, calcium, strontium, barium and radium.
• In all these elements, the last electron is situated in the outer most s-orbital.
• All these elements are metals.
• These elements (except beryllium) form oxides and hydroxides. The oxides and hydroxides thus formed are strongly alkaline in nature.
• The oxides of these elements are found in the earth’s crust. So these elements are collectively known as alkaline earth metals.
[Crust is the outer most layer of the earth. It’s average thickness is about 40 km. Compared to the inner layers of the earth, this thickness is very small. So we say that, the earth’s crust is thin]
4. Abundance of group 1 elements:
• Among the group 1 elements, sodium and potassium are abundant in nature.
♦ Sodium is the sixth most abundant metal in the earth’s crust.
♦ Potassium is the seventh most abundant metal in the earth’s crust.
• Lithium, rubidium and caesium have much lower abundances.
• Francium is highly radioactive. Even the longest-lived isotope 223Fr has a half life of only 21 minutes.
5. Among the group 2 elements, calcium and magnesium are abundant in nature.
♦ Calcium is the fifth most abundant metal in the earth’s crust.
♦ Magnesium is the eighth most abundant metal in the earth’s crust.
• Strontium and barium have much lower abundances.
Beryllium is rare.
• Radium is the rarest among the group 2 elements. Only 10-10 per cent of igneous rocks is radium. [Igneous rocks are formed when the molten rock (magma) cools and hardens]
6. We can write the electronic configuration of any group 1 element, if we know the electronic configuration of the noble gas coming just before.
• The general electronic configuration is:
[noble gas]ns1.
• Here n is the period in which the group 1 element is situated.
7. We can write the electronic configuration of any group 2 element, if we know the electronic configuration of the noble gas coming just before.
• The general electronic configuration is:
[noble gas]ns2.
• Here n is the period in which the group 2 element is situated.
8. We expect all members of a group to exhibit the same chemical properties. We expect so because, all members of a group will be having the same number of electrons in the outer most shell.
• But consider lithium. It is the first member of group 1. Lithium has properties which are different from those of other members of that group.
• Properties of lithium resemble the properties of magnesium.
• Magnesium is the second member of group 2. So lithium and magnesium occupy diagonal positions as shown by the yellow double headed arrow in fig.10.1 below:
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| Fig.10.1 |
9. Similarly, consider beryllium. It is the first member of group 2. Beryllium has properties which are different from those of other members of that group.
• Properties of beryllium resemble the properties of aluminium.
• Aluminium is the second member of group 3. So beryllium and aluminium occupy diagonal positions as shown by the magenta double headed arrow in fig.10.2 above.
10. This type of diagonal similarity is known as diagonal relationship in periodic table.
There can be two reasons for diagonal relationship:
(i) Similarity in ionic sizes.
(ii) Similarity in $\frac{\text{charge}}{\text{radius}}$ ratios
• Importance of $\frac{\text{charge}}{\text{radius}}$ ratio can be demonstrated using an example. It can be written in 3 steps:
(i) Mg2+ and Ca2+ have the same charge.
(ii) Mg2+ have a smaller radius than Ca2+.
(iii) So the $\frac{\text{charge}}{\text{radius}}$ ratio will be greater for Mg2+.
11. Consider the following ions obtained from the s-block:
♦ Monovalent sodium ion (Na+)
♦ Monovalent potassium ion (K+)
♦ Divalent magnesium ion (Mg2+)
♦ Divalent calcium ion (Ca2+)
• Biological fluids in plants and animals contain large proportions of these ions. These above ions perform two main functions:
(i) Maintenance of ion balance in the body.
(ii) Conduction of nerve impulses.
Electronic configuration of Alkali metals
• This can be written in 3 steps:
1. Consider any alkali metal. It will be having one valence electron. This electron will be in the outer most s-orbital.
• Some examples are:
♦ potassium: [noble gas]4s1
♦ caesium: [noble gas]6s1
2. This valence electron is loosely held. That means, alkali metals readily lose this electron. Consequently, alkali metals are the most electropositive metals.
3. Consider the atom of any alkali metal. It will form a monovalent ion (M+).
• Such an ion will combine with other -ve ions. So we will never find alkali metals in free state in nature.
Atomic and ionic radii of Alkali metals
• This can be written in 4 steps:
1. Consider the alkali metal from any period. Atom of that alkali metal will be the largest among all atoms in that period.
• For example,
♦ Na atom is the largest atom in the third period.
♦ Cs atom is the largest atom in the sixth period.
2. Consider the atom (M) of any alkali metal.
• The monovalent ion (M+) will be smaller than M.
♦ For example, Na+ ion is smaller than Na atom.
3. The atomic radius increases as we move down the group 1.
• For example:
♦ Atomic radius of Na is larger than that of Li.
♦ Atomic radius of K is larger than that of Na.
♦ So on . . .
4. The ionic radius also increases as we move down the group 1.
• For example:
♦ Ionic radius of Na+ is larger than that of Li+
♦ Ionic radius of K+ is larger than that of Na+.
♦ So on . . .
• We already know the reason for such pattern. We saw them in the chapter on periodic trends (Details here).
Ionization Enthalpy of Alkali metals
• This can be written in 2 steps:
1. Consider the alkali metal from any period. Its ionization enthalpy will be the lowest among all elements in that period.
2. Within the group, the ionization enthalpies decrease as we go down.
• So Cs will be having the lowest ionization enthalpy.
• We do not consider the ionization enthalpy of Fr because it is a very rare element. Also it is radioactive with a very small half life of 21 minutes. So it is difficult to determine it’s ionization enthalpy and other chemical properties.
Hydration Enthalpy of Alkali metals
• This can be written in 2 steps:
1. As we go down group 1, hydration enthalpy decreases.
Li+ > Na+ > K+ > Rb+ > Cs+
2. Li+ has the maximum degree of hydration. So lithium salts are mostly hydrated. LiCl.2H2O is an example.
Physical properties of Alkali metals
• This can be written in 6 steps:1. All alkali metals are:
♦ silvery white in color.
♦ soft.
♦ light in weight.
2. We have seen that, the alkali metal atoms are the largest in the corresponding periods. So they occupy larger volumes.
• But they are lower in mass. So the density (= mass/volume) of alkali metals will be low.
3. We know that, ionic bonding occurs between a +ve ion (derived from a metallic atom) and a -ve ion (derived from a non-metallic atom).
• But in a sample of a metal, there are only metallic atoms. Then how can -ve ions form? With out -ve ions, how do the atoms in that sample stick together?
• The answer is:
Atoms in a sample of metal are kept together by metallic bonding.
• This type of bonding can be explained in 5 steps:
(i) We know that in metals, the outermost electrons are loosely held.
(ii) So the outer most electrons get delocalized. Those electrons do not belong to any particular atom.
• We can say that, each atom in the sample of a metal consists of two parts:
♦ The inner core.
♦ The delocalized electron.
(iii) The delocalized outer most electrons form a “sea” of electrons.
(iv) The inner cores are strongly attracted towards this “sea”. This is because, the inner cores are +ve ly charged while the "sea" is -ve ly charged.
♦ So the “sea” binds the inner cores together.
(v) When we say the word “sea”, the blue surface of the sea comes to our minds. We consider it as an area. But the sea has depth. It has volume. In the same way, the “sea” in our present case also has volume. It extends to the whole volume of the metal sample. So all the metal cores in the sample are attracted towards the "sea".
(v) When the metal cores are kept together in this way, we call it metallic bonding.
4. The alkali metals have only one outer most electron. So there will be only lesser number of electrons in the “sea”. Consequently, the bonding between the cores will be weak.
• We can say that:
In the sample of an alkali metal, the atoms are kept together by weak metallic bonding.
• That is why the alkali metals have low melting and boiling points.
5.When alkali metals are heated with an oxidizing flame, the outer most electron absorbs energy and jumps to a higher energy level.
• This electron releases energy and returns to the ground state. The released energy is in the form of visible light.
♦ If the metal is lithium, then the visible light will be crimson red in color.
♦ If the metal is sodium, then the visible light will be yellow in color.
♦ If the metal is potassium, then the visible light will be violet in color.
♦ If the metal is rubidium, then the visible light will be red violet in color.
♦ If the metal is caesium, then the visible light will be blue in color.
• So alkali metals can be detected by flame test, flame photometry or atomic emission spectroscopy.
6. We know that alkali metals have low ionization enthalpy. So the outermost electron can be easily removed from the core.
• Some times, light energy is sufficient to remove the electrons. These electrons can be collected to form an electric current.
• Caesium and potassium are widely used as electrodes in photoelectric cells.
In the next section, we
will see chemical properties of alkali metals.
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