Class 10 Science Chapter 5

The Modern periodic table

In the modern periodic table, the elements are arranged in rows and columns. These rows and columns are known as periods and groups respectively. The table consists of 7 periods and 18 groups.

Do You Know:

In the modern periodic table, hydrogen is placed above alkali metals because of resemblance with their electronic configurations. However, it is never regarded as an alkali metal. This makes hydrogen a unique element.

If you look at the modern periodic table, you will find that all elements in the same group contain the same number of valence electrons. Let us see the following activity to understand better.

Activity 1: Look at group two of the modern periodic table. Write the name of the first three elements followed by their electronic configurations.

What similarity do you observe in their electronic configurations? How many valence electrons are present in these elements?

The first three elements of group two are beryllium, magnesium, and calcium. All these elements contain the same number of valence electrons. The number of valence electrons present in these elements is 2. On the other hand, the number of shells increases as we go down the group.

Again, if you look at periods in the modern periodic table, you will find that all elements in the same period contain the same valence shell. Let us see the following activity to understand better.

Activity 2: Look at the elements of the third period of the modern periodic table. Write the electronic configuration of each element and calculate the number of valence electrons present in these elements.

What do you observe from the given activity? Do these elements contain the same number of shells? How many valence electrons are present in these elements?

You will find that elements such as sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon are present in that period. The valence shell in all these elements is the same, but they do not have the same number of valence electrons.

Thus, the number of electrons in the valence shell increases by one unit as the atomic number increases by one unit on moving from left to right in a period.

Let us calculate the number of elements that are present in the first, second, third, and fourth periods.

The maximum number of electrons that a shell can hold can be calculated using the formula 2n². Here, n represents the number of shells from the nucleus. For example, n is equal to 1, 2, and 3 for K, L, and M shells respectively. Hence, the maximum number of electrons that each of these shells can hold can be calculated by substituting the value of n in the given formula.

Number of electrons that K shell can accommodate = 2n²

= 2

Hence, K shell can accommodate only 2 electrons and only two elements are present in the first period.

Similarly, the second and third shell (L and M respectively) can accommodate 8 and 18 electrons respectively. Since the outermost shell can contain only 8 electrons, there are only 8 elements in both the periods.

Important Note:

The position of an element in the Modern Periodic Table tells us about its chemical reactivity. The valence electrons determine the kind and the number of bonds formed by an element.

IUPAC Nomenclature for Elements with Atomic Number  > 100

• Latin word roots for various digits are listed in the given table.

Notation for IUPAC Nomenclature of Elements

Digit	Name	Abbreviation
0	nil	n
1	un	u
2	bi	b
3	tri	t
4	quad	q
5	pent	p
6	hex	h
7	sept	s
8	oct	o
9	enn	e

• Latin words for various digits of the atomic number are written together in the order of digits, which make up the atomic number, and at the end, ‘ium’ is added.
• Nomenclature of elements with the atomic number above 100 is listed below.

Nomenclature of Elements with Atomic Number Above 100

Atomic number	Name	Symbol	IUPAC Official Name	IUPAC Symbol
101	Unnilunium	Unu	Mendelevium	Md
102	Unnilbium	Unb	Nobelium	No
103	Unniltrium	Unt	Lawrencium	Lr
104	Unnilquadium	Unq	Rutherfordium	Rf
105	Unnilpentium	Unp	Dubnium	Db
106	Unnilhexium	Unh	Seaborgium	Sg
107	Unnilseptium	Uns	Bohrium	Bh
108	Unniloctium	Uno	Hassnium	Hs
109	Unnilennium	Une	Meitnerium	Mt
110	Ununnilium	Uun	Darmstadtium	Ds
111	Unununnium	Uuu	Rontgenium	Rg
112	Ununbium	Uub
113	Ununtrium	Uut
114	Ununquadium	Uuq
115	Ununpentium	Uup
116	Ununhexium	Uuh
117	Ununseptium	Uus
118	Ununoctium	Uuo

Do you know that Mendeleev’s periodic table was a huge success? Let us now study the achievements of Mendeleev’s periodic table.

Achievements of Mendeleev’s periodic table:

• Mendeleev left some gaps in his periodic table. He had predicted that some elements were yet to be discovered. He left these gaps deliberately so that these undiscovered elements could get a place in his periodic table.

• Mendeleev named the undiscovered elements using the Sanskrit word Eka (meaning one) as a prefix, with the name of the preceding element in the same group. For example, gallium was not discovered in Mendeleev’s time. Therefore, he left a gap for it in his periodic table and named it Eka-aluminium. He also predicted the properties of these undiscovered elements based on their positions in the periodic table. A comparison of the properties of gallium as predicted by Mendeleev and its actual properties is given in the following table.

Comparison of the properties of Eka-aluminium and gallium

In the periodic table, Ea is used as the symbol for Eka-aluminium. It can be clearly observed from the table that Mendeleev’s predictions were almost exact. This extraordinary achievement made Mendeleev’s periodic table very popular.

• Noble gases were not discovered at the time when Mendeleev gave the periodic table. These were discovered in recent times as they are very inert and are present in very low concentrations in the atmosphere. When these gases were finally discovered, they got a place in his periodic table as a separate column. The accommodation of these gases in the periodic table did not disturb the positions of other elements. This underlined the strength of Mendeleev’s periodic table.

• Although Mendeleev arranged the elements in the increasing order of their atomic masses, there were instances where he placed an element with a slightly higher atomic mass before an element with a slightly lower atomic mass.

  For example, cobalt, whose atomic mass is 58.9, was placed before nickel whose atomic mass is 58.7. This was done to maintain consistency in the properties of the elements present in a group i.e., to group the elements with similar properties together.

Though, Mendeleev’s periodic table was a huge success at that time, it still had many limitations. Now, let us discuss the limitations of Mendeleev’s periodic table.

Limitations of Mendeleev’s periodic table:

1. Hydrogen’s position was not justified in Mendeleev’s periodic table. He positioned hydrogen in the first column above alkali metals. He did so because hydrogen and alkali metals have similar properties. For example, hydrogen reacts with halogens, oxygen, and sulphur to form compounds whose formulae are similar to those of alkali metals.

It is because all these elements have one thing in common. You must have observed that these elements contain the same number of shells, and the last electron enters shell M as shown in the given table:

Elements of the third period

Do you know the valency of the elements present in the third period?

Valency is defined as the number of electrons an atom requires to lose, gain, or share in order to complete its valence shell to attain the stable noble gas configuration. Valencies of the elements can also be determined by the number of electrons present in the outermost shell known as the valence shell.

Therefore, on moving across a period, from left to right, the valency first increases from 1 to 4, and then decreases from 4 to 0.

Do you know how the valency of an element, present in the same group, changes on moving down the group?

The number of valence electrons present in atoms of elements belonging to a particular group remains the same. As a result, valency also remains the same. Hence, the valency of group 1 elements is one, group 2 elements is two, and so on.

For example, the valency of the first four elements of group 1 of the periodic table is given as follows. The valency of all these four elements is 1.

Valency of the first four elements of group 1

Atomic size is the radius of an atom in its neutral state i.e. the distance of the nucleus from its valence shell in an isolated gaseous atom.

Do you know how atomic size changes as we move from top to bottom or across the periodic table. Let us see.

Changes in atomic size on moving down a group

The atomic radii of the elements present in the third period are given in Pico meters.

Changes in atomic size on moving across a period

Do you know that the metallic character of elements can also be determined with the help of the periodic table?

Metals are electropositive in nature as they lose electrons to form positive ions called cations. For example, Na loses one electron to form Na⁺ ion.

We have learnt that there is a decrease in the atomic size, as we move across a period. As a result, it becomes difficult for an atom to lose electrons as we move from left to right across the periodic table. Hence, metallic character decreases on moving across a period from left to right.

On the other hand, atomic size increases when we move down a group. This makes it easier for an element to lose electrons. Hence, metallic character increases on moving down a group.

Therefore, the elements present in the 1^st group are the most metallic in their respective periods as they contain only one electron, and have the largest atomic size in their respective periods. 

Changes in metallic character on moving down a group

Do you know that with the help of a periodic table, we can also predict the nature of the oxides formed by the various elements? The oxides formed by metals are basic in nature, whereas the oxides of non-metals are acidic in nature. Hence, the elements present on the left hand side of the periodic table will form basic oxides, whereas the elements present on the right hand side of the periodic table will form acidic oxides.

Elements show periodicity because of their valence shell configuration. All elements showing periodicity in properties have the same number of electrons in the last or valence shell. The properties that will be discussed here are:

• Atomic radius
• Ionisation potential
• Electron affinity
• Electronegativity
• Metallic character

Atomic Radius: The atomic radius is usually considered as the distance from the centre of the nucleus to the outermost shell i.e., to a point where the electron density is effectively zero.

Across the period i.e., from left to right: Atomic radius decreases

Down the group i.e., from top to bottom: Atomic radius increases

Reason: Across the period, the effective nuclear charge increases. This is due to the fact that the number of electrons increases (in the same subshell), and so does the number of protons in the nucleus. This pulls the valence shell of electrons in an atom towards itself, thus decreasing the atomic radius. But as we move down the group, the number of shells keeps on increasing along with the number of protons. The space required to accommodate the extra shells becomes the dominating factor and therefore the atomic size increases.

Ionisation Potential: It is the energy required to remove one mole of electrons from the valence shell of one mole of isolated gaseous atoms.

Across the period i.e., from left to right: Ionisation potential increases

Down the group i.e., from top to bottom: Ionisation potential decreases

Reason: Across the period, the effective nuclear charge increases. This causes the atomic radius to decrease, thus getting the valence shell closer to the nucleus. This makes it difficult to remove electrons. But as we move down the group, the number of shells keeps on increasing along with the number of electrons. The distance from the nucleus coupled with the interference of the electron between the nucleus and the valence shell renders the valence electrons weakly bound to the nucleus.

Electronegativity: The tendency of an atom of an element to attract a shared pair of electrons towards itself when bonded with another element in a compound is called electronegativity.

Across the period i.e., from left to right: Electronegativity increases

Down the group i.e., from top to bottom: Electronegativity decreases

Reason: Across the period, the effective nuclear charge increases, thus decreasing the atomic radius. This favours the increase in electronegativity of elements across the period. But as we move down the group, the number of shells keeps on increasing and therefore the atomic size increases and the electronegativity decreases.

Metallic character: It is defined as the tendency of an atom to lose electrons.

Across the period i.e., from left to right: Metallic character decreases

Down the group i.e., from top to bottom: Metallic character increases

Reason: Across the period, the effective nuclear charge increases, thus decreasing its atomic radius. This favours the increase of electronegativity and therefore the tendency to lose electrons is low. This accounts for the decrease in the metallic character along a period. But as we move down the group, the number of orbits keeps on increasing and therefore the atomic size increases. This means that the electronegativity decreases. This enhances the loss of electrons and therefore the metallic character increases down a group.