There is a periodic fluctuation in the electronic configuration of the elements as the atomic number increases. Therefore, the physical and chemical properties of the elements vary in a periodic manner.
Elements with a similar valence shell electronic configuration are placed in the same group, one below the other. Chemical properties depend on the electronic configuration of the valence shell.
Because all the elements in a given group have similar valence shell electronic configurations, they have similar chemical properties. Physical properties depend on the size of atoms. Because the sizes of atoms change gradually from top to bottom in a group.
Therefore, the elements show a gradation of physical properties within the same group. In the period of the periodic table, the number of electrons in the valence shell increases gradually from left to right. Their chemical and physical properties also differ in the same way.
In this section, you will learn about the variation of physical properties of certain elements within a group and across a period.
Figure 4.3 shows electronic configuration of Li, Be and Mg. Which atom has more shells, Be or Mg? Which atom has more electrons between the nucleus and the valence electrons, Be or Mg?
Electrons present in the inner shells cut off attractive force between the nucleus and the valence electrons.
The reduction in force of attraction between nucleus and the valence electrons by the electrons present in the inner sub-shells is called shielding effect.
Which atom has greater shielding effect, Be or Mg?
Which atom, Li or Be has greater number of shells? Which atom, Li or Be has greater number of electrons between nucleus and valence electrons?
As you move from left to right in a period the number of electrons in the inner shells remains constant, therefore, shielding effect remains constant.
As you move from top to bottom in a group the number of electronic shells increase. So the number of electrons in the inner shell also increase. As a result shielding effect increases.
All the physical and chemical properties of elements depend on the electronic configuration of their atoms. We now consider some properties of atoms that are affected by electronic configuration: atomic size, ionization energy, electron affinity and electronegativity. They usually increase and decrease repeatedly throughout the periodic table. That is, they show consistent changes or trends within a group or a period. These tendencies are correlated with their behaviour.
Choose the elements whose atoms you expect to have greater shielding effect.
(a) Be or Mg
(b) C or Si
Look at the periodic table and find the relative position of given elements in the periodic table. Apply the trend of increasing shielding effect in a group.
(a) Mg atoms will have greater shielding effect.
(b) Si atoms will have greater shielding effect.
Choose the element whose atoms you expect to have smaller shielding effect.
(a) F or Cl
(b) Li or Na
(c) B or Al
The size of an atom depends on its electronic configuration. Atomic size is the average distance between the atomic nucleus and the electronic outer shell.
Figure 4.4 shows the atomic radii of the main group elements. Figure 4.4 shows the variation of atomic radii within a period and within a group. You can see two general trends in atomic radii.
The atomic radius decreases in each period as you move across the period. This is because as you move from one element in the sequence to the next, to the right of it:
Example: Going from lithium to beryllium, the atomic size decreases. You can understand this from the electronic configuration of the valence shell of Li (2s¹) and Be (2s²). Moving from Li to Be, the number of shells does not change, but the atomic number increases from 3 to 4. Therefore, the strength of the nucleus on the valence shell electron increases, and the atomic radius decreases.
Atomic radius increases in each main group as you move down the element group. This is because the size of an atom is determined by the size of its valence shell.
As you move down the group to the next lower element, the atom has an additional shell of electrons. This increases the radius of the atom.
Example: Going from Li to Na, the atomic radius increases. Consider the electron configuration of Li (1s², 2s¹) and Na (1s², 2s², 2p⁶, 3s¹). A new electron shell is added, increasing the size of the atom.
Choose the element whose atom you expect to have larger atomic radius in each of the following pairs.
(a) Mg, Al
(b) C, Si
Remember that the larger atom in any:
(a) Period lies further to the left in the periodic table.
(b) Group lies closer to the bottom in the periodic table.
(c) Check the periodic table and choose the element.
(a) The larger atom is Mg
(b) The larger atom is Si
Using the periodic table but without looking at the figure 4.4, choose the element whose atom you expect to have smaller atomic radius in each of the following pairs.
(a) O or S
(b) O or F
Ionization energy is an important property of atoms that explains cation formation.
"Ionization energy is defined as the minimum amount of energy required to remove the outermost electron from an isolated gaseous atom".
M(g) + ionization energy → M⁺(g) + e⁻
Ionization energy is a measure of the extent to which the nucleus attracts the outermost electron.
Figure 4.5 shows the ionization energies of the main group elements. Values are given in units of kJ mole⁻¹ or kJ/mole.
The value of the ionization energy decreases from top to bottom in the group. This is because:
Which atom has a greater shielding effect, Li or Na?
As you move from left to right in the period:
Which atom has the higher ionization energy, Li or Be?
Electron affinity explains the anion formation. Electron affinity is defined as the amount of energy released when an electron adds up in the valence shell of an isolated atom to form a uni negative gaseous ion.
X(g) + e⁻ → X⁻(g) + electron affinity
Figure 4.6 shows electron affinities of main group elements. Factors affecting electron affinity are:
As you move from left to right through a period, electron affinity generally increases. This is due to an increase in nuclear charge and a decrease in atomic radius, which binds the extra electron more tightly to the nucleus. But the shielding effect remains constant in each cycle. Therefore, the alkali metals have the lowest and the halogens the highest electron affinities in each period.
Electron affinity decreases from top to bottom in a group. This is due to an increase in the shielding effect. Due to the increased shielding effect and increase in atomic radius, the added electron binds less tightly to the nucleus. As a result, less energy is released.
There are several exceptions to the general trend of election affinity values. You will learn reasons for it in higher grade.
Electronegativity is the ability of an atom to attract electrons toward itself in a chemical bond. Figure 4.7 shows as a scale of electronegativities of the elements devised by Linus Pauling. The American chemist Linus Pauling devised a method for calculating the relative electronegativities of elements.