I.Very Short Answer Type Questions ( Classification of Elements )
Question 1. State the Modem Periodic Law.
Answer: Modem Periodic Law states that physical and chemical properties of the elements are a periodic function of their atomic numbers.
Question 2. Why is ionization enthalpy of nitrogen greater than that of oxygen?
Answer: Nitrogen has exactly half filled p-orbitals.
Question 3. Why are electron gain enthalpies of Be and Mg positive?
Answer: They have fully filled s-orbitals and hence have no tendency to accept an additional electron. That’s why energy is needed if an extra electron is to be added. Therefore, electron gain enthalpies of Be and Mg are positive.
Question 4. Give four examples of species which are isoelectronic with ca2+.
Answer: Ar, K+, CT, S2-, or P3- are isoelectronic with ca2+.
Question 5. Which two elements of the following belong to the same period?
Al, Si, Ba and O
Answer: Al and Si.
Question 6. Explain why chlorine can be converted into chloride ion more easily as compared to fluoride ion from fluorine ?
Answer: Electron gain enthalpy of Cl is more negative than that of F.
Question 7. What are horizontal rows and vertical columns of the periodic table called?
Answer: Horizontal rows are called periods and vertical columns are called groups.
Question 8. Which has a larger radius?
(i)Mg or Ca (ii) S or Cl
Answer: (i) Ca (ii) S.
Question 9. What are representative elements?
Answer: The elements of group 1 (alkali metals), group 2,(alkaline earth metals) and group 13 to 17 constitute the representative elements. They are elements of s-block and p-block.
Question 10. Give general electronic configuration off-block elements?
Answer: General electronic configuration of f-block elements =(n – 2) f1-14 (n -1) d0-1 ns2.
Question 11. What are inner transition metals? Why are they called rare earth metals?
Answer: Lanthanoids (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements.
Question 12. Define ionisation enthalpy.
Answer: It is the energy required to remove an electron from an isolated gaseous atom in its ground state. M (g) + I.E. àM+ (g) + e–
Question 13. The electronic configuration of an element is Is 2s 2p 3s 3p 4s . Locate the element in the periodic table.
- As the principal quantum number for the valence shell is 4, the element is present in the 4th period.
- Since the last electron has been filled in 4s sub-shell (or orbital), the element belongs to s-block.
- As there is only one electron in the valence s-sub-shell, the element is present in group I.
II. Short Answer Type Questions
Question 1. What is the cause of periodicity in properties of the elements? Explain with two examples.
Answer: The cause of periodicity in properties is the repetition of similar outer electronic configuration after certain regular intervals.
For example, all the elements of group LA i.e., alkali metals, have similar outer electronic configuration as ns1.
Where n refer to the number of outermost principal shell.
In a similar manner all the halogens i.e., elements of group VILA have similar other electronic configuration i.e., ns2 ns5 and hence possess similar properties.
Question 2. Show by a chemical reaction with water that Na20 is a basic oxide and Cl207 is an acidic oxide.
Answer: Na2 0reacts with water to form sodium oxide which turns red litmus blue.
Na20 +H20——-> 2NaOH
Therefore, Na20 is a basic oxide
In contrast,Cl207 reacts with water to form perchloric acid which turns blue litmus red.
Therefore, Cl207 is an acidic oxide.
Question 3. What do you understand by ‘Representative elements’? Name the groups whose elements are called representative elements.
Answer: The elements of s and p-block are collectively called representative or main group elements. These include elements of group I (alkali metals), group 2 (alkaline earth metals). .
Question 4. Name different blocks of elements in the periodic table. Give general electronic configuration of each block.
Answer: Elements in the long form of the periodic table have been divided into four blocks i.e., s, p, d and f. This division is based upon the name of the orbital which receives the last electron. General electronic configuration of s-block elements: ns1 – 2 where n = 2 – 7
p-block elements: ns2 np1- 6 where n = 2 – 6
d-block elements: (n -1) d1 -10 ns0- 2 where n = 4 – 7
f-block elements: (n – 2)f0-14(n -1) d0-1 ns2 where n = 6 – 7
Question 5. Elements A, B, C and D Iwoe atomic numbers 12,19, 29, and 36 respectively. On the basis of electronic configuration, write to which group of the periodic table each element belongs.
Answer: Electronic configuration of A (Z = 12)
=1s2 2s2 2p6 3s2
period = 3, Element’s name = Mg block = s, Group = II Electronic configuration of B (Z = 19)
Element’s name = K (potassium)
=1s2 2s2 2p6 3s2 3p6 4s1 n = 4, period = 4 Block = s, Group = I Electronic configuration of C (Z = 29)
=1s2 2s2 2p6 3s2 3p6 3d10 4s1 n – 4, period = 4 Block = d Electronic configuration of D (Z = 36)
=1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 period = 4 Block = p-Block group = 18
Question 6. Define the term ionization enthalpy? How does it vary along a period and along a group?
Answer: Ionization Enthalpy. The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom so as to convert it into a gaseous cation is called its ionization enthalpy or energy. It is represented by A; H.
This process may be represented as M (g) + ∆iH -> M+ (g) + e–(g)
where M (g) is isolated gaseous atom.
M+ (g) is the resultant cation (a position ion)
Variation along a period. Moving from left to right in a period, the ionization enthalpy increases with atomic number.
Variation within a group. The ionization enthalpies keep on decreasing regularly as we move down a group from one element to the other.
Question 7. Discuss briefly the various factors on which ionization enthalpy depends.
- Atomic size. With the increase in the atomic size, the number of electron shells
increases. Therefore, the force that binds the electrons with the nucleus decreases. Thus, the ionization enthalpy decreases with increase in atomic size.
- Nuclear charge. As the magnitude of the positive charge on the nucleus of an atom increases, the attraction with the electrons also increases. Therefore, the ionization enthalpy increases with the increase in the magnitude of the nuclear charge.
- Screening or shielding effect. Greater the magnitude of the screening effect, less will be the value of ionization enthalpy or potential.
Question 8. What are Dobereiner’s triads? Name two such triads.
Answer: Dobereiner arranged certain elements with similar properties in groups of three in such a way that the atomic mass of the middle element was nearly the same as the average atomic masses of the first and third elements.
Triad: lithium sodium potassium
Atomic mass: 7 23 39
Atomic mass of Na =(39+7)/2= 23
Triad: Chlorine Bromine Iodine
Atomic mass: 35.5 80 127
Atomic mass of Br =127 + 35.5/2 = 81.25
Question 9. Give the electronic configuration of the transition elements. Write their four important characteristics.
Answer: The d-block elements are known as transition elements.
Electronic configuration =(n – 1) d1-10 ns1 -2
Characteristics of d-block elements:
- They show variable oxidation states.
- Their compounds are generally paramagnetic in nature.
- Most of the transition elements form coloured compounds.
- They are all metals with high melting and boiling points.
Question 10. What is screening or shielding effect? How does it influence the ionization enthalpy ?
Answer: In a multielectron atom, the electrons present in the inner shells shield the electrons in the valence shell from the attraction of the nucleus or they act as a screen between the nucleus and these electrons. This is known as shedding effect or screening effect. As the screening effect increases, the effective nuclear charge decreases. Consequently, the force of attraction by the nucleus for the valence shell electrons decreases and hence the ionization enthalpy decreases.
Question 11. Define electron gain enthalpy. What are its units?
Answer: The energy which is released by an atom in gaining an electron from outside atom or ion to form negative ion (or anion) is called electron gain enthalpy (∆egH).
Unit of electron gain enthalpy is kJ/ mol.
In some cases, like in noble gas, atoms do not have any attraction to gain an electron. In that case energy has to be supplied.
Ne (g) + e– —> Ne– (g)
∆egH = + 116 kJ mol -1
III. Long Answer Type Questions ( Classification of Elements )
Question 1. Discuss the main features of long form of the periodic table. What are the advantages of long . form of periodic table?
Answer: Main features of long form of periodic table:
- Groups. The vertical columns in the periodic table are known as groups. There are 18 groups in the long form of periodic table.
Each group having the same electronic configuration in the outermost shell.
- Periods. There are 7 periods in the long form of periodic table.
It is denoted by n which means highest principal quantum number.
- Lanthanoids. Group of 14 elements in the sixth period. They are placed after Lanthanum.
- Actinides. Group of 14 elements in the seventh period after actinium. Both Lanthanoids and actinoids are placed in separate panel at the bottom of the periodic table.
Advantages of long form of periodic table: ( Classification of Elements )
- It gives a suitable link between the position of element and its electronic configuration.
- On the basis of atomic numbers it easier to remember all the elements.
- The elements in the same group have similar properties due to their outer-most (valence shell) configuration. Thus it gives is a logical classification.
- Justified positions are provided to transition and inner transition elements. ‘
- It makes the study of elements systematic and simple.
Question 2. Discuss the main characteristics of four blocks of elements in the periodic table? Give their general electronic configuration.
Answer: s-block elements:
- They are highly reactive elements and thus occurs in combined state. On moving down the group their reactivity increases.
- They have good reducing characters.
- They generally form electropositive ion by losing 1 or 2 electrons, that’s why they are electro positive in nature.
- They are good conductors of heat and electricity.
p-block elements: ( Classification of Elements )
- Most of the p-block elements show variable oxidation states.
- They include both metals and non-metals.
- They are generally covalent in nature.
- As move from left to right the non-metallic character of the element increases.
- On moving down the group metallic character increases.
d-block elements: ( Classification of Elements )
- d-block elements show variable oxidation states.
- They are generally paramagnetic in nature.
- Their compounds are generally coloured. Those which form complex compounds.
- Most of the elements and their compounds acts as catalyst.
f-block elements: ( Classification of Elements )
- They are generally heavy metals having high melting and boiling points.
- Their compounds are generally coloured.
- Variable oxidation states are generally shown by these elements.
- Most of Activities are radioactive.
General electronic configuration: ( Classification of Elements )
p-block —ns2 np1- 6
d-block —(n -1) d1 -10 ns0-2
f-block —(n – 2)f0-14 (n -1) d0- 1 ns2
Question 3. Define electron gain enthalpy. What are its units? Discuss the factors which influence the electron gain enthalpy.
Answer: Electron gain enthalpy is the energy released when an isolated gaseous atom is converted into a negative ion by adding an extra electron.
Electron gain enthalpy is denoted by the sign∆egH.
The process may be represented by
M(g) + e– ———————>M– (g)
neutral gaseous atom anion
∆ H=∆eg H
electron gain enthalpy is negative or positive it depends upon the nature of the element. For example. For halogens it is highly negative, because they can acquire the noble gas configuration by accepting an extra electron.
In contrast. For noble gases have positive electron gain enthalpy because energy has to be supplied to the element.
Factors on which electron gain enthalpy depends:
- Atomic size. As the size of an atom increases, the distance between its nucleus and the incoming electron also increases. Therefore, the force of attraction between the nucleus and the incoming electron decreases and hence the electron gain enthalpy becomes less negative.
- Nuclear charge. As the nuclear charge increases force of attraction for the incoming electron increases and thus electron gain enthalpy becomes more negative.
- Symmetry of electronic configuration. Elements having symmetrical configuration (Either half filled or fully filled orbitals in the same sub shell)
having no attraction for electron because by accepting electron their configuration becomes less stable. In that case energy has to be supplied to accept electron. Thus electron gain enthalpy will be positive.
Question 4. Discuss the factors that influence the magnitude of ionization enthalpy. What are the general trends of variation of ionization enthalpy in the periodic table? Explain.
Answer: Factors affecting Ionization enthalpy.
- Atomic size. With the increase in atomic size, the number of electron shells increases and thus the force of attraction between the electrons and the nucleus decreases. Therefore the ionization enthalpy decreases.
- Nuclear charge. As the nuclear charge increases the attraction for the electron also increases that’s why ionization enthalpy increases.
- Screening or shielding effect. In a multi-electron atom, the electron present in the inner shells shield the electrons in the valence shell as a result these electrons
experience less attraction from the nucleus. This leads to lesser ionization enthalpy.
- Variation along a period. On moving from left to right in a period the nuclear charge increases and the atomic size decreases as a result ionization enthalpies are expected to increase.
- Variation within a group. On moving down the group as the atomic size of the elements increases that’s why ionization enthalpy decreases down the group.
Question 5. (a) How does atomic radius vary in group in the periodic table?
(i) Radius of cation is less than that of the atom.
(ii) Radius of anion is more than that of the atom.
(iii) In iso-electronic ion, the ionic radii decreases with increase in atomic number.
Answer: (a) Variation of atomic radius in a group:
On moving down the group there is an increase in the principal quantum number and therefore no. of electron shells increases and thus the atomic size increases. Thus the atomic radii of the element increases.
(b) (i) Radius of cation is less than that of the atom:
Since the cation is formed by losing of one or more electrons.
Na —> Na++ e–
Thus the radius of Na+will be less than the Na.
(ii) Radius of anion is more than that of the atom.
Since the anion is formed by gaining one or more electron. Therefore, the atomic radius is larger than the corresponding atom.
(iii) In iso-electronic ions, atoms have same number of electrons but different magnitude of nuclear charges. As the nuclear charge increases ionic radius decreases.
For example. N3-, O2-, F– have same No. of electrons = 10 but different ionic radii = 171, 140, 136 respectively.
IV. Multiple Choice Questions ( Classification of Elements )
Question 1. The highest ionization energy is exhibited by
(a) halogens (b) alkaline earth metals
(c) transition metals (d) noble gases
Question 2. Which of the following oxides is neutral?
(a) Sn02 (b) CO (c) Al203 (d) Na20
Question 3. Which of the following is arranged in order of increasing radius?
(a) K+ (aq) < Na+ (aq) < Li+ (aq) (b) K+ (aq) > Na+ (aq) > Zn2+ (aq)
(c) K+ (aq) > Li+ (aq) > Na+ (aq) (d) Li+ (aq) < Na+ (aq) < K+ (aq)
Question 4. What is the electronic configuration of the elements of group 14?
(a) ns2 np4 (b) ns2 np6 (c) ns2 np2 (d) ns2
Question 5. Among the following elements, which has the least electron affinity?
(a) Phosphorous (b) Oxygen (c) Sulphur (d) Nitrogen
Question 6. In halogens, which of the following, increases from iodine to fluorine?
(a) Bond length (b) Electronegativity
(c) The ionization energy of the element (d) Oxidizing power
Question 7. Diagonal relationships are shown by
(a) Be and A1 (b) Mg and A1 (c) Li and Mg (d) B and P
Question 8. Which of the following species are not known?
(a) AgOH (b) PbI4 (c) PI5 (d) SH6
(e) All of the above
Question 9. Which one of the following is isoelectronic with Ne?
(a) N3- (b) Mg2+ (c) Al3+ (d)all of the above
Question 10.Which element has smallest size?
(a) B (b) N (c) Al (d) P
Answer: 1. (b) 2. (b) 3. (d) 4. (c) 5. (d)
6. (b) (c) and (d)7. (a) and (c)8. (e)9. (d) 10. (b)
V. Hots Questions ( Classification of Elements )
Question 1. Arrange the following as stated: (i) N2, 02, F2, Cl2(Increasing order of bond dissociation energy) (ii) F, Cl, Br, I (Increasing order of electron gain enthalpy) (iii) F2, N2, Cl2, O2(Increasing order of bond length).
Answer: (i) F2 < Cl2 < 02 < N2
(ii) I < Br < F < Cl
(iii) N2 < 02 < F2 < Cl2
Question 2. The first ionisation enthalpy of magnesium is higher than that of sodium. On the other hand, the second ionisation enthalpy of sodium is very much higher than that of magnesium. Explain. ( Classification of Elements )
Answer: The 1st ionisation enthalpy of magnesium is higher than that of Na due to higher nuclear charge and slightly smaller atomic radius of Mg than Na. After the loss of first electron, Na+ formed has the electronic configuration of neon (2, 8). The higher stability of the completely filled noble gas configuration leads to very high second ionisation enthalpy for sodium. On the other hand, Ma+ formed after losing first electron still has one more electron in its outermost (3s) orbital. As a result, the second ionisation enthalpy of magnesium is much smaller than that of sodium.
Question 3. Give reasons: ( Classification of Elements )
(i) IE1 of sodium is lower than that of magnesium whereas IE2 of sodium is higher than that of magnesium.
(ii) Noble gases have positive value of electron gain enthalpy.
Answer: (i) The effective nuclear charge of magnesium is higher than that of sodium. For these reasons, the energy required to remove an electron from magnesium is more than the energy required in sodium. Hence, the first ionization enthalpy of sodium is lower than that of magnesium.
However, the second ionization enthalpy of sodium is higher than that of magnesium.
This is because after losing an electron, sodium attains the stable noble gas configuration. On the other hand, magnesium, after losing on electron still has one electron.
(ii)Because of stable configuration.