Q.1. Explain why the electron gain enthalpy of fluorine is less negative than that of chlorine.
Ans. In fluorine, the new electron to be added goes to 2p-subshell while in chlorine, the added electron goes to 3p-subshell. Since the 2p-subshell is relatively small as compared to 3p-subshell, the added electron in small 2p-subshell experiences strong interelectronic repulsions in comparison to that in 3p-subshell in Cl. As a result, the incoming electron does not feel much attraction from the nucleus and therefore, the electron gain enthalpy of F is less negative than that of Cl.
Q.2. All transition elements are d-block elements, but all d-block elements are not transition elements. Explain.
Q.3. Identify the group and valency of the element having atomic number 119. Also predict the outermost electronic configuration and write the general formula of its oxide.
Ans. The present set up of the Long Form of the Periodic Table can accommodate at the maximum 118 elements. After this, according to Aufbau principle, 8.y-orbital should be filled. Therefore, the outer electronic configuration of element having atomic number 119 will be 8s1. Since it has one electron in the outermost s-orbital, its valency will be 1 and it should belong to group 1 along with alkali metals. The general formula of its oxide will be M20, where M represents the element.
Q.4. Ionisation enthalpies of elements of second period are given below: Ionisation enthalpy/ k cal mol–1 : 520, 899, 801, 1086, 1402, 1314, 1681, 2080.
Match the correct enthalpy with the elements and complete the graph given in Figure. Also write symbols of elements with their atomic number.
Ans. To match the correct enthalpy with the elements and to complete the graph, the following points are taken into consideration. As we move from left to right across a period, the ionization enthalpy keeps on increasing due to increased nuclear charge and simultaneous decrease in atomic radius. However, there are some exceptions given below:
Q.5. Among the elements B, Al, C and Si,
(i) which element has the highest first ionisation enthalpy?
(ii) which element has the most metallic character?
Justify your answer in each case.
Ans. Arranging the elements into different groups and periods:
Q.6. Write four characteristic properties of p-block elements.
The four important characteristic properties of p-block elements are the following:
(i) p-Block elements include both metals and non-metals but the number of non-metals is much higher than that of metals. Further, the metallic character increases from top to bottom within a group and non-metallic character increases from left to right along a period in this block.
(ii) Their ionization enthalpies are relatively higher as compared to s-block elements.
(iii) They mostly form covalent compounds.
(iv) Some of them show more than one (variable) oxidation states in their compounds. Their oxidizing character increases from left to right in a period and reducing character increases from top to bottom in a group.
Q.7. Choose the correct order of atomic radii of fluorine and neon (in pm) out of the options given below and justify your answer.
(i) 72, 160
(ii) 160, 160
(iii) 72, 72
(iv) 160, 72
Solution: Atomic radius decreases as we move from left to right in a period in the periodic table. Fluorine has the smallest atomic radius. As we move to neon in the same period, the atomic radius increases as it has van der Waals radius and van der Waals radii are bigger than covalent radii.
Q.8. Illustrate by taking examples of transition elements and non-transition elements that oxidation states of elements are largely based on electronic configuration.
Q.9. Nitrogen has positive electron gain enthalpy whereas oxygen has negative. However, oxygen has lower ionisation enthalpy than nitrogen. Explain.
Q.10. First member of each group of representative elements (i.e., s and p-block elements) shows anomalous behaviour. Illustrate with two examples.
Ans. First member of each group of s- and p-block elements shows anomalous behaviour due to the following reasons:
(i) Small size
(ii) High ionization enthalpy
(iii) High electronegativity
(iv) Absence of d-orbitals
Examples: Li in the first group shows different properties from the rest of elements like covalent nature of its compounds, formation of nitrides.
Similarly, beryllium, the first element of second group differs from its own group in the following ways:
Q.11. p-Block elements form acidic, basic and amphoteric oxides. Explain each property by giving two examples and also write the reactions of these oxides with water.
Ans. In p-block, when we move from left to right in a period, the acidic character of the oxides increases due to increase in electronegativity. For example,
On moving down the group, acidic character decreases and basic character increases, e.g.,
(a) Nature of oxides of 13 group elements
(b) Nature of oxides of 15 group elements
Among the oxides of same element, higher the oxidation state of the element, stronger is the acid. For example, SO3 is a stronger acid than SO2.
B2O3 is weakly acidic and on dissolution in water, it forms orthoboric acid. Orthoboric acid does not act as a protonic acid (it does not ionise) but acts as a weak Lewis acid.
Al2O3 is amphoteric in nature. It is insoluble in water but dissolves in alkalies and reacts with acids.
Tl2O is as basic as NaOH due to its lower oxidation state (+1).
Tl2O + 2HCl → 2TlCl + H2O
P4O10 on reaction with water gives orthophosphoric acid
Cl2O7 is strongly acidic in nature and on dissolution in water, it gives perchloric acid.
Q.12. How would you explain the fact that first ionisation enthalpy of sodium is lower than that of magnesium but its second ionisation enthalpy is higher than that of magnesium?
Ans. Electronic configuration of Na is 1s22s22p63s1. After losing one electron from its outermost shell, sodium easily attains stable electronic configuration (1s22s22p6 ), while magnesium does not lose its electron easily due to presence of two electrons in s-orbital (ls22s22p63s2). Hence first ionisation energy of sodium is less than magnesium.
When one electron is removed from Na and Mg, their configurations become ls22s22p6 and 1s22s22p63s1 respectively. Now it is easier to remove one electron from 3s of Mg+ than 2p6 of Na+. Hence, second ionisation energy of Mg is less than Na.
Q.13. What do you understand by exothermic reaction and endothermic reaction? Give one example of each type.
Ans. Exothermic reaction: The reaction which is accompanied with evolution of heat is known as exothermic reaction.
CaO + CO2 → CaCO3, ΔH= -178.3 kJ mol-1
Endothermic reaction: The reaction which is accompanied with absorption of energy is known as endothermic reaction, e.g.,
2NH3 → N2 + 3H2; ΔH = +91.8 kJ mol-1
Q.14. Arrange the elements N, P, O and S in the order of
(i) increasing first ionisation enthalpy.
(ii) increasing non metallic character.
Give reason for the arrangement assigned.
(i) S < P < O < N
Ionisation enthalpy increases from left to right in a period and decreases down the group. N has higher ionisation enthalpy than O due to extra stability of half-filled orbitals. Similarly, P has higher ionisation enthalpy than S due to half-filled orbitals.
(ii) P < S < N < O
Non-metallic character decreases down the group and increases from left to right in a period.
Q.15. Explain the deviation in ionisation enthalpy of some elements from the general trend by using Figure.
Ans. Ionisation enthalpy increases in a period with increase in atomic number. The graph shows few exceptions and not the linear relationship. Ionisation enthalpy of Be is greater than that of B due to filled s-orbital in Be (Be – l.s2 1s2, B – 1s2 2s2 2pl).
Ionisation enthalpy of N is greater than that of O due to half-filled/?-orbitals in nitrogen.
(N – Is2 2s 22p3 O – ls2 2s2 2p4).
Q.16. Explain the following:
(a) Electronegativity of elements increase on moving from left to right in the periodic table.
(b) Ionisation enthalpy decrease in a group from top to bottom?
(a) The electronegativity generally increases on moving across a period from left to right (e.g., from Li to F in the second period). This is due to decrease in atomic size and increase in effective nuclear charge. As a result of increase in effective nuclear charge, the attraction for the outer electron and the nucleus increases in a period and therefore, electronegativity also increases.
(b) On moving down a group there is a gradual decrease in ionisation enthalpy. The decrease in ionisation enthalpy down a group can be explained in terms of net effect of the following factors:
(i) In going from top to bottom in a group, the nuclear charge increases.
(ii) There is a gradual increase in atomic size due to an additional main energy shell (n).
(iii) There is increase in shielding effect on the outermost electron due to increase in the number of inner electrons.
The effect of increase in atomic size and the shielding effect is much more than the effect of increase in nuclear charge. As a result, the electron becomes less and less firmly held to the nucleus as we move down the group. Hence, there is a gradual decrease in the ionization enthalpies in a group.
Q.17. How does the metallic and non metallic character vary on moving from left to right in a period?
Ans.As we move from left to right in a period, the number of valence electrons increases by one at each succeeding element but number of shells remains same. Due to this, effective nuclear charge increases. More is the effective nuclear charge, more is the attraction between the nucleus and electron.
Hence, the tendency of the element to lose electrons decreases. This results in decrease in metallic character. Furthermore, the tendency of an element to gain electrons increases with increase in effective nuclear charge, so non- metallic character increases on moving from left to right in a period.
Q.18. The radius of Na+ cation is less than that of Na atom. Give reason.
Ans. The radius of Na+ is less than Na atom because Na+ is formed by losing one energy shell.
Na – 1s2 2s2 2p6 3s1
Na+– 1s22s2 2p6
Q.19. Among alkali metals which element do you expect to be least electronegative and why?
Ans. Cs is the most electropositive or least electronegative element among the alkali metals since electronegative character decreases in a group.
Q.20. Match the correct atomic radius with the element.
Ans. All the given elements are of same period and along period, atomic radii decrease because effective nuclear charge increases. Thus, the order of atomic radii is
O < N < C < B < Be or, Be = 111 pm, O = 66 pm, C = 77 pm, B = 88 pm, N = 74 pm.
Q.21. Match the correct ionisation enthalpies and electron gain enthalpies of the following elements.
(i) → (b);
(ii) → (a);
(iii) → (d);
(iv) → (c)
Q.22. Electronic configuration of some elements is given in Column I and their electron gain enthalpies are given in Column II. Match the electronic configuration with electron gain enthalpy.
Ans. (i → D), (ii → A), (iii → B), (iv → C)
(i) This electronic configuration corresponds to the noble gas, i.e., neon. Since, noble gases have +ΔegH values, therefore, electronic configuration (A) corresponds to the ΔegH = +48 kJ mol-1.
(ii) This electronic configuration corresponds to the alkali metal, i.e., potassium. Alkali metals have small negative ΔegH values, hence, electronic configuration (B) corresponds to ΔegH = -53 kJ mol-1.
(iii) This electronic configuration corresponds to the halogen, i.e., fluorine. Since, halogens have high negative ΔegH values, therefore, electronic configuration (C) corresponds to ΔegH = -328 kJ mol-1.
(iv) This electronic configuration corresponds to the chalcogen, i.e., oxygen. Since, chalcogens have ΔegH values less negative than those of halogens, therefore, electronic configuration (D) corresponds to ΔegH = -141 kJ mol-1.
In the following questions a statement of Assertion (A) followed by a statement of reason (R) is given. Choose the correct option out of the choices given below each question.
Q.23. Assertion (A) : Generally, ionisation enthalpy increases from left to right in a period.
Reason (R) : When successive electrons are added to the orbitals in the same principal quantum level, the shielding effect of inner core of electrons does not increase very much to compensate for the increased attraction of the electron to the nucleus.
(i) Assertion is correct statement and reason is wrong statement.
(ii) Assertion and reason both are correct statements and reason is correct explanation of assertion.
(iii) Assertion and reason both are wrong statements.
(iv) Assertion is wrong statement and reason is correct statement.
Ionisation enthalpy increases from left to right across the period due to decrease in atomic size.
The electrons present within the subshell has almost same effective nuclear charge.
Q.24. Assertion (A) : Boron has a smaller first ionisation enthalpy than beryllium.
Reason (R) : The penetration of a 2s electron to the nucleus is more than the 2p electron hence 2p electron is more shielded by the inner core of electrons than the 2s electrons.
(i) Assertion and reason both are correct statements but reason is not correct explanation for assertion.
(ii) Assertion is correct statement but reason is wrong statement.
(iii) Assertion and reason both are correct statements and reason is correct explanation for assertion.
(iv) Assertion and reason both are wrong statements.
In beryllium, the electron removed during the ionization is an s-electron whereas the electron removed during ionization of boron is a p-electron. The penetration of a 2s-electron to the nucleus is more than that of a 2p-electron; hence the 2p electron of boron is more shielded from the nucleus by the inner core of electrons than the 2s electrons of beryllium. Therefore, it is easier to remove the 2p-electron from boron compared to the removal of a 2s-electron from beryllium. Thus, boron has a smaller first ionization enthalpy than beryllium.
Q.25. Assertion (A) : Electron gain enthalpy becomes less negative as we go down a group.
Reason (R) : Size of the atom increases on going down the group and the added electron would be farther from the nucleus.
(i) Assertion and reason both are correct statements but reason is not correct explanation for assertion.
(ii) Assertion and reason both are correct statements and reason is correct explanation for assertion.
(iii) Assertion and reason both are wrong statements.
(iv) Assertion is wrong statement but reason is correct statement.
On moving down the group, the electron gain enthalpy becomes less negative because on moving down the group the atomic size increases and the added electron lies away from the nucleus.
Q.26. Discuss the factors affecting electron gain enthalpy and the trend in its variation in the periodic table.
Factors affecting electron gain enthalpy:
(i) Nuclear charge: The electron gain enthalpy become more negative as the nuclear charge increases. This is due to greater attraction for the incoming electron if nuclear charge is high.
(ii) Size of the atom: With the increase in size of the atom, the distance between the nucleus and the incoming electron increases and this results in lesser attraction. Consequently, the electron gain enthalpy become less negative with increase in size of the atom of the element.
(iii) Electronic configuration: The elements having stable electronic configurations of half filled and completely filled valence subshells show very small tendency to accept additional electron and thus electron gain enthalpies are less negative.
Variation of electron gain enthalpies in periodic table:
Electron gain enthalpy, in general, becomes more negative from left to right in a period and becomes less negative as we go from top to bottom in a group.
(a) Variation down a group: On moving down a group, the size and nuclear charge increases. But the effect of increase in atomic size is much more pronounced than that of nuclear charge and thus the additional electron feels less attraction by the large atom. Consequently, electron gain enthalpy becomes less negative. This is clear from decrease of electron gain enthalpy in going from chlorine to bromine and to iodine.
(b) Variation along a period: On moving across a period, the size of the atom decreases and nuclear charge increases. Both these factors result in greater attraction for the incoming electron, therefore, electron gain enthalpy, in general, becomes more negative in a period from left to right. However, certain irregularities are observed in the general trend. These are mainly due to the stable electronic configurations of certain atoms.
Important Trends in Electron Gain Enthalpies:
There are some important features of electron gain enthalpies of elements.
(i) Halogens have the highest negative electron gain enthalpies.
(ii) Electron gain enthalpy values of noble gases are positive while those of Be, Mg, N and P are almost zero.
(iii) Electron gain enthalpy of fluorine is unexpectedly less negative than that of chlorine.
Q.27. Define ionisation enthalpy. Discuss the factors affecting ionisation enthalpy of the elements and its trends in the periodic table.
Ans. Ionisation enthalpy: It is defined as the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom.
M(g) →M+(g) + e–I1 = First ionisation enthalpy
Similarly, second and third electrons are also removed by providing successive ionisation enthalpies.
Factors on which Ionisation Enthalpy Depends:
(i) Size of the atom: The larger the atomic size, smaller is the value of ionisation enthalpy. In a larger atom, the outer electrons are far away from the nucleus and thus force of attraction with which they are attracted by the nucleus is less and hence can be easily removed.
Ionisation enthalpy ∝ 1/ Atomic size
(ii) Screening effect: Higher the screening effect, the lesser is the value of ionisation enthalpy as the screening effect reduces the force of attraction towards nucleus and hence the outer electrons can be easily removed.
Ionisation enthalpy ∝ 1/ Screening effect
(iii) Nuclear charge: As the nuclear charge increases among atoms having same number of energy shells, the ionisation enthalpy increases because the force of attraction towards nucleus increases.
Ionisation enthalpy ∝ Nuclear charge
(iv) Half filled and fully filled orbitals: The atoms having half filled and fully filled orbitals are comparatively more stable, hence more energy is required to remove the electron from such atoms. The ionization enthalpy is rather higher than the expected value in case of such an atom.
Ionisation enthalpy ∝ Stable electronic configuration
(v) Shape of orbital: The s-orbital is more close to nucleus than the p-orbital of the same orbit. Thus, it is easier to remove electron from a p-orbital in comparison to s-orbital. In general, the ionisation enthalpy follows the following order
(s>p> d>f) orbitals of the same orbit.
Variation of ionisation enthalpy in the periodic table
ln general, the ionisation energy decreases down the group due to increase in atomic size. On the other hand, the ionisation energy increases across the period from left to right, again due to decrease in atomic size from left to right.
Q.28. Justify the given statement with suitable examples— “the Properties of the elements are a periodic function of their atomic numbers”.
Ans. The physical and chemical properties of the elements are periodic functions of their atomic numbers. Periodic Law is essentially the consequence of the periodic variation in electronic configurations, which indeed determine the physical and chemical properties of elements and their compounds. Elements having similar outer electronic configurations in their atoms are arranged in vertical columns, referred to as groups or families.
Q.29. Write down the outermost electronic configuration of alkali metals. How will you justify their placement in group 1 of the periodic table?
Since all the elements have one electron in their valence shell, hence, they are placed in first group of the periodic table.
Q.30. Write the drawbacks in Mendeleev’s periodic table that led to its modification.
Ans. Drawbacks of Mendeleev’s periodic table:
1. Position of hydrogen: Hydrogen is placed in group I. However, it resembles the elements of group I (alkali metals) as well as the elements of group VILA, (halogens). Therefore, the position of hydrogen in the periodic table is not correctly defined.
2. Anomalous pairs: In certain pairs of elements, the increasing order of atomic masses was not obeyed. In these cases, Mendeleev placed elements according to similarities in their properties and not in increasing order of their atomic masses. For example, argon (Ar, atomic mass 39.9) is placed before potassium (K, atomic mass 39.1). Similarly, cobalt (Co, atomic mass 58.9) is placed before nickel (Ni, atomic mass 58.6) and tellurium (Te, atomic mass 127.6) is placed before iodine (I, atomic mass 126.9). These positions were not justified.
3. Position of isotopes: Isotopes are the atoms of the same element having different atomic masses but same atomic number. Therefore, according to Mendeleev’s classification, these should be placed at different places depending upon their atomic masses. For example, isotopes of hydrogen with atomic masses 1,2 and 3 should be placed at three places. However, isotopes have not been given separate places in the periodic table.
4. Some similar elements are separated and dissimilar elements are grouped together: In the Mendeleev’s periodic table, some similar elements were placed in different groups while some dissimilar elements had been grouped together. For example, copper and mercury resembled in their properties but they had been placed in different groups. At the same time, elements of group IA such as Li, Na and K were grouped with copper (Cu), silver (Ag) and gold (Au), though their properties are quite different.
5. Cause of periodicity: Mendeleev did not explain the cause of periodicity among the elements.
6. Position of lanthanoids (or lanthanides) and actinoids (or actinides): The fourteen elements following lanthanum (known as lanthanoids, from atomic number 58-71) and the fourteen elements following actinium (known as actinoids, from atomic number 90 – 103) have not been given separate places in Mendeleev’s table.
In order to cover more elements, Mendeleev modified his periodic table.
Q.31. In what manner is the long form of periodic table better than Mendeleev’s periodic table? Explain with examples.
Ans. Superiority of the Long form of the Table over Mendeleev’s Table:
(i) This table is based on a more fundamental property, i.e., atomic number.
(ii) It correlates the position of the elements with their electronic configurations more clearly.
(iii) The completion of each period is more logical. In a period as the atomic number increases, the energy shells are gradually filled up until an inert gas configuration is reached. It eliminates the even and odd series of IV, V, VI and VII periods of Mendeleev’s periodic table.
(iv) The position of VIIIth group is appropriate in this table. All the transition elements have been brought in the middle as the properties of transition elements are intermediate between s- and p-block elements.
(v) Due to separation of two subgroups, dissimilar elements do not fall together. One vertical column accommodates elements with same electronic configuration thereby showing same properties.
(vi) In this table, a complete separation between metals and non-metals has been achieved. The non-metals are present in the upper right comer of the periodic table.
(vii) There is a gradual change in properties of the elements with increase in their atomic numbers, i.e., periodicity of properties can be easily visualized. The same properties o recurrence in properties occur after the intervals of 2, 8, 8, 18, 18 and 32 elements which indicates the capacity of various periods of the table.
(viii) This arrangement of elements is easy to remember and reproduce.
Q.32. Discuss and compare the trend in ionisation enthalpy of the elements of group 1 with those of group 17 elements.
Ans. The ionization enthalpies decrease regularly as we move down a group from one element to the other. This is evident from the values of the first ionisation enthalpies of the elements of group 1 (alkali metals) and group 17 elements as given in table and figure.
Given trend can be easily explaihed on the basis of increasing atomic size and screening effect as follows
(i) On moving down the group, the atomic size increases gradually due to the addition of one new principal energy shell at each succeeding element. Hence, the distance of the valence electrons from the nucleus increases.
Consequently, the force of attraction by the nucleus for the valence electrons decreases and hence the ionisation enthalpy should decrease.
(ii) With the addition of new shells, the shielding or the screening effect increases. As a result, the force of attraction of the nucleus for the valence electrons further decreases and hence the ionisation enthalpy should decrease.
(iii) Nuclear charge increases with increase in atomic number. As a result, the force of attraction by the nucleus for the valence electrons should increase and accordingly the ionisation enthalpy should increase.
The combined effect of the increase in the atomic size and the screening effect more than compensates the effect of the increased nuclear charges. Consequently, the valence electrons become less and less firmly held by the nuclear and hence the ionisation enthalpies gradually decrease as move down the group.