Chapter 10 - The s-Block Elements

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:

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.

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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.

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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).

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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.

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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.

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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.

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In the next section, we will see chemical properties of alkali metals.

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Chapter 3.12 - Periodic Trends in Chemical Reactivity

In the previous section 3.11, we completed a discussion on the periodic trends in valence. In this section, we will see Anomalous properties of 2nd period elements. We will also see the periodic trends in chemical reactivity

1. Consider the table 3.8 below:

Table 3.8

• It shows the elements of two periods: Period 2 and Period 3
• We see only the ‘first two elements’ of groups 1, 2, 13, 14, 15, 16, 17 and 18
• Note that, these groups constitute the s and p blocks of the periodic table
2. Take out the first element from each of those groups (except group 18):
    ♦ Li from group 1
    ♦ Be from group 2
    ♦ B from group 13
    ♦ C from group 14
    ♦ N from group 15
    ♦ O from group 16
    ♦ F from group 17
3. Let us analyze each element from the above set:
• Consider the first entry of the above set: Li of group 1
    ♦ We have learnt that all elements of a group will have similar properties
    ♦ But Li differs in many respects from the other members of the group 1    
• Consider the next entry of the above set: Be of group 2
    ♦ We have learnt that all elements of a group will have similar properties
    ♦ But Be differs in many respects from the other members of the group 2
• Consider the next entry of the above set: B of group 13
    ♦ We have learnt that all elements of a group will have similar properties
    ♦ But B differs in many respects from the other members of the group 13
- - - -
- - - -
• Consider the last entry of the above set: F of group 17
    ♦ We have learnt that all elements of a group will have similar properties
    ♦ But B differs in many respects from the other members of the group 13
4. It is clear that, all first elements differ from their group members
■ Let us see some examples:
• Lithium mostly form covalent compounds
    ♦ But all other members of the group 1 forms ionic compounds
• Beryllium also mostly form covalent compounds
    ♦ But all other members of the group 2 forms ionic compounds
5. Here scientist noticed 3 interesting points:
(i) Behavior of Li is more similar to the second element of the following group
• Let us elaborate on this:
• Li belongs to group 1
    ♦ So the ‘following group’ is group 2 
    ♦ The ‘second element’ of the group 2 is Mg
• So Li is more similar to Mg
• This is indicated by the yellow arrow in the table below:

Table 3.9

(ii) Behavior of Be is more similar to the second element of the following group
• Let us elaborate on this:
• Be belongs to group 2
    ♦ So the ‘following group’ is group 3 
    ♦ The ‘second element’ of the group 3 is Al
• So Be is more similar to Al
• This is indicated by the pink arrow in the table above
(iii) In the same way, B is more similar to Si
• Let us elaborate on this:
• B belongs to group 13
    ♦ So the ‘following group’ is group 14 
    ♦ The ‘second element’ of the group 14 is Si
• So B is more similar to Si
• This is indicated by the blue arrow in the table above
6. The diagonal relationship:
• We see that, Li and Mg are diagonal elements
• Also, Be and Al are diagonal elements
• Again, B and Si are diagonal elements
■ So this type of similarity is commonly referred to as diagonal relationship in the periodic properties

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• Let us see the reason for this type of a similarity. We will write it in steps:

1. Imagine that, we move along the period 2
• Let us move from Li to Be
2. We know that, the size of the atom decreases
• So Be will be smaller than Li
3. Now imagine that, from Be, we move downwards along the group 2
• Let us move from Be to Mg
4. We know that, the size of the atom increases
• Mg will be larger than Be
5. So we have two effects:
(i) ‘Decreasing size’ when moving horizontally towards the right from Li to Be
(ii) ‘Increasing size’ when moving downwards from Be to Mg
6. These two effects are opposite to each other and so cancel out
• The result of this ‘cancelling effect’ can be written in 3 steps:
(i) Imagine that we move in a ‘L-shaped’ path from Li to Mg (through Be)
(ii) We will find that the size decreases and then ‘increases back’ to nearly the original size
(iii) That means, the atomic size of Li and Mg are nearly the same
• We find that, it is indeed true:
    ♦ Metallic radii of Li, Be and Mg are 152, 111 and 160 pm respectively
    ♦ Note that, 160 is close to 152
7. Consider the next set: Be and Al
• Let us move in a ‘L-shaped’ path
    ♦ We start from Be and reach Al (through B)
    ♦ When B is reached, the size becomes small
    ♦ But when Al is reached, the size ‘increases back’ to nearly the size of Be
8. A similar analysis can be written in the case of the next set 'B and Si' also:
• Let us move in a ‘L-shaped’ path
    ♦ We start from B and reach Si (through C)
    ♦ When C is reached, the size becomes small
    ♦ But when Si is reached, the size ‘increases back’ to nearly the size of B
9. We will see more about 'diagonal relationship' when we discuss in detail about s and p block elements in later chapters

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Chemical reactivity

So far in this chapter, we have seen the periodic trends in the following items:

A. Physical properties:
  (i) atomic radii
 (ii) Ionic radii
(iii) Ionization enthalpy
(iv) Electronegativity
B. Chemical properties:
 (i) Valence
(ii) Oxidation state
■ By now, we know that, the periodic trends is a result of this:
Systematic arrangement of elements on the basis of electronic configuration
• We have one more item to learn:
The chemical reactivity
• The four physical properties and the two chemical properties that we have seen, can be used to explain chemical reactivity
• We will write it in steps:
1. Recall the 4 horizontal arrows that we saw in fig.3.14 in an earlier section 3.9
• For convenience, that fig.3.14 is shown again below:

2. Out of the four arrows, two are important for our present discussion:
(i) The arrow for ionization enthalpy
(ii) The arrow for electron gain enthalpy
3. Consider the arrow for ionization enthalpy:
■ It indicates that:
In a period, the ionization enthalpy is least for the element at the extreme left end
• That means:
In a period, the element at the extreme left end can lose it’s electron in the most easiest way
• It follows that:
In a period, the element at the extreme left end can become a cation in the most easiest way
• It follows that:
In a period, the element at the extreme left end will have the highest chemical reactivity
    ♦ Chemical Reactivity is the ‘readiness to  take part in chemical reactions’
4. Consider the arrow for electron gain enthalpy
■ It indicates that:
In a period, the electron gain enthalpy is 'highest negative' for the element at the extreme right end
• That means:
In a period, the element at the extreme right end can gain an electron in the most easiest way
• It follows that:
In a period, the element at the extreme right end can become an anion in the most easiest way
• It follows that:
In a period, the element at the extreme right end will have the highest chemical reactivity
5. So we have two extremes:
• One at the left end and the other at the right end
    ♦ Both have the highest reactivities
6. The 'variation of reactivity' can be written in 3 steps:
  (i) The reactivity decreases (initially) as we go from left to right across a period
 (ii) It reaches the 'lowest reactivity' at the center of the periodic table
(iii) After passing the center, it again increases and reaches the 'highest reactivity' at the extreme right end
7. We can write a note on oxidation and reduction also:
(i) Oxidation:
• 'Losing electrons' is oxidation
    ♦ That is., 'elements which lose electrons' are said to be oxidised 
• The elements at the left easily loses electrons
    ♦ So we can say: Elements at the left are easily oxidised
(ii) Reduction:
• 'Gaining electrons' is reduction
    ♦ That is., 'elements which gain electrons' are said to be reduced
• The elements at the right easily gains electrons
    ♦ So we can say: Elements at the right are easily reduced
8. We can write a note on metallic and non-metallic character also:
(i) Metallic character:
• Metals are electron donors
    ♦ That is., metals lose electrons easily 
• Elements at the left lose electrons easily
    ♦ So we can say: Elements at the left are more metallic
(ii) Non-metallic character:
• Non-Metals are electron acceptors
    ♦ That is., non-metals accept electrons easily 
• Elements at the right accepts electrons easily
    ♦ So we can say: Elements at the right are more non-metallic

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So we have completed a discussion on ‘periodic trends (along periods) in reactivity’. In the next section, we will see some examples to demonstrate those trends. We will also see ‘periodic trends (along groups) in reactivity’

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