Periodicity in Periodic Properties
Periodicity refers to the recurrence of similar properties of elements after certain regular intervals in increasing atomic number order. The repetition of similar electronic configurations of their atoms in the outermost energy shell (or valence shell) after certain regular intervals is the cause of periodicity in element properties. All alkali metals, for example, have a strong tendency to lose a single electron to achieve a stable noble gas configuration.
Physical properties, such as atomic size, ionisation enthalpy, electron affinity, electronegativity, valency, and so on, are, however, directly related to atoms' electronic configuration. Let's take a look at some of the most important periodic table periodic properties and variations of properties.
Variation of Ionisation Enthalpy in the Periodic Table
Variation Along a Period
In general, the ionisation enthalpy increases as the atomic number in a period increases. Atomic size and nuclear charge can explain the overall increase over time. We know that moving across a period from left to right increases the nuclear charge, and the atomic size decreases over time while the principal energy level remains constant. The valence electrons are tightly held by the nucleus as the nuclear charge increases and the atomic size decreases. As a result, more and more energy is required to remove the electron. Thus the ionisation enthalpy rises.
Exception: These are due to half-filled and completely-filled configurations which have extra stability. In Be, the electron removed during ionisation is a 2s-electron whereas the electron removed during ionisation of B is a 2p-electron. Thus, the ionisation enthalpy of B < Be. As a result, the 2p-electron of B is not tightly held by the nucleus due to the 2s-electron of Be. Hence, the ionisation enthalpy of B is less than that of Be.
Variation Down a Group
Moving from top to bottom within a group, there is a gradual decrease in ionisation enthalpy. The net effect of the following factors can explain the decrease in ionisation enthalpy down a group. The nuclear charge increases as one moves downwards in a group. The addition of the principal energy shell causes a gradual increase in atomic size (n).
As the number of inner electrons increases, the shielding effect on the outermost electron increases. The effect of increasing atomic size and shielding is much greater than the effect of increasing nuclear charge as a result. As we move down the group, the electron becomes less and less tightly bound to the nucleus. As a result, the ionisation enthalpies in a group gradually decrease.
Variation of Electron Affinity in the Periodic Table
Variation Along a Period
Moving from left to right across a period increases effective nuclear charge and thus size. So, the added electron would be closer to the nucleus on average, i.e. strongly attracted by the nucleus. As a result, the value of electron affinity increases as we move from left to right along a period.
Exception: Due to their stable half-filled configuration, the electron gain enthalpies of elements in groups 2 and 15 are low, whereas those of elements in group 18 are zero due to their complete valence shell.
Variation Down the Group
Moving down a group increases the size of the atom, which increases the distance between the nucleus and the incoming electron, or decreases the force of attraction between the nucleus and the incoming electron. As a result of this electron affinity decreases down the group.
Exception: The electron affinity of elements in the second period (O, F) is less negative than that of elements in the third period (S, Cl). This is because, in the case of elements of the second period, the incoming electron enters the smaller n=2 quantum level and, thus, experiences the greater force of repulsion by the other electrons present in the level.
Whereas, in elements of the third period, it enters the larger n=3 quantum level and thus experiences much less force of repulsion than in elements of the second period. As a result, chlorine has the periodic table's highest negative electron affinity.
Nature of Oxides in Groups and Periods
An element's chemical reactivity is best demonstrated by its reactions to oxygen. Elements at extremes of a period easily combine with oxygen to form oxides. The oxide formed by the element on the far left (most metallic) is the most basic (Na2O), whereas the oxide formed by the element on the far right (most non-metallic) is the most acidic (e.g., Cl2O7).
The amphoteric (e.g., Al2O3) or neutral oxides of the elements in the centre (e.g., CO, etc.) exhibit acidic and basic properties. They react acidically with bases and basically with acids. Neutral oxides, on the other hand, have neither acidic nor basic properties.
In general, as one moves from left to right across a period, the basic character of the oxides decreases while the acidic character increases. Moving across the third period, for example, reveals that Na2O is strongly basic, MgO is less basic, Al2O3 is amphoteric, SiO2 is weakly acidic, P2O5 is acidic, overline SO3 is strongly acidic and Cl2O7 is extremely acidic.
Conclusion
The physical and chemical properties of the element are the periodic function of their atomic number. The recurrence of similar properties after a certain regular interval is called periodicity. The energy required to remove an electron from the atom/ion in the gaseous state is called ionisation enthalpy whereas the energy released when an electron is added to the atom/ion in the gaseous state is called electron affinity.
Moving from left to right across a period increases effective nuclear charge and thus size. So, the added electron would be closer to the nucleus on average, i.e. strongly attracted by the nucleus. As a result, the value of electron affinity increases as we move from left to right along a period.
FAQs on Variation Down a Group and Along a Period of Periodic Properties for JEE
1. Write short notes on variation in the melting point of the chlorides of alkali metals.
The amount of energy required to separate one mole of a solid ionic compound into its gaseous ions is defined as lattice energy. The higher the lattice energy, the higher the melting point of the alkali metals halide and the lower its water solubility. For the same halide ion, the melting point of lithium halides is lower than that of sodium halides. It continues to fall as we move down the group from Na to Cs. The low melting point of LiCl (887 K) in comparison to NaCl is most likely because LiCl is covalent and NaCl is ionic.
As a result, the correct order is NaCl > KCl > CsCl > LiCl.
2. Which of the s-block elements is liquid from the following given elements?
Na, Mg, Ca, Cs, Fr
The liquid elements of s-block are Cs and Fr (at temperatures ranging from 25° to 40°C). The S – Block Elements are the elements of groups 1 and 2. The element in which the last electron enters the s-orbital is referred to as the s-block element. The hardness of the s-block elements is significantly lower than that of the others. Cs and Fr have much lower melting points. This is because of their large size and relatively weak intermetallic bonds. They can melt at normal room temperature or slightly higher temperatures.
