NCERT
Solutions for Class 11 Chemistry Chapter 3 Classification of Elements and
Periodicity in Properties
These
Solutions are part of NCERT Solutions for Class 11 Chemistry. Here we have
given NCERT Solutions for Class 11 Chemistry Chapter 3 Classification of
Elements and Periodicity in Properties.
Question
1.
What is the basic theme of organisation
in periodic table ?
Answer:
The basic theme of organisation in periodic table is to group the elements with
same characteristics (physical and chemical) together so that it may become
rather easy to follow these characteristics. Since these mainly depend upon the
valence shell electronic configurations, the elements placed in a group have
therefore same valence shell configuration of their atoms. In fact, the valence
shell electronic configuration of the elements placed in a group gets repeated
after definite gaps of atomic numbers (8, 8, 18, 18, 32).
Question
2.
Which important property did Mendeleev
use to classify the elements in periodic table and did he stick to that ?
Answer:
The property used by Mendeleev is known as Mendeleev’s Periodic Law. According
to the law,
Physical
and chemical properties of the elements are a periodic function of their mass
numbers.
Mendeleev
did stick to it and classified elements into groups and periods. This
classification was perhaps the first systematic way in order to study the
characteristics of the elements. No doubt, later on many defects were pointed
out and scientists challenged the basis of the classification.
Question
3.
What is the basic difference in
approach between Mendeleev’s periodic law and Modern periodic law ?
Answer:
The atomic mass of the elements is the basis of periodicity according to
Mendeleev’s periodic law. On the other hand, the atomic number of the elements
constitutes the same according to Modern periodic law. For more details,
·
Modern Periodic Law
Dmitri
Mendeleev, a Russian Chemist, was the first to make a very significant
contribution in the formation of the periodic table. He studied the chemical
properties of the large number of the elements and in 1869 and came out with a
statement that the chemical properties of the elements are a periodic function
of their atomic masses. When Mendeleev came to know about the work of Lother
Meyer, he modified his statement and gave a law called Mendeleev-Lother Meyer
Periodic Law or simply Mendeleev’s Periodic Law.
Modern Periodic Law
It states that In the year 1913, an English physicist, Henry Moseley, a young
physicist from
Actually,
when Moseley plotted a graph between and atomic number of the different metals,
a straight line was obtained. But it was not the case when a graph was plotted
between (4v) and atomic mass of the metals. These studies led Moseley to
believe that atomic number and not the atomic mass is the fundamental property
of an element. Therefore, atomic numbers must form the basis of the
classification of the elements in the periodic table. In the light of the above
observations. Moseley gave the modern periodic law which states that
Question
4.
On the basis of quantum numbers,
justify that the sixth period of the periodic table should have 32 elements.
Answer:
The sixth period corresponds to sixth shell. The sub-shells present in this
shell are 6s, 4f, 5p and 6d. The maximum number of electrons which can be
present in these sub-shell are 2 + 14 + 6 + 10 = 32. Since the number of
elements in a period correspond to the number of electrons in the shells,
therefore, sixth period should have a maximum of 32 elements.
Question
5.
In terms of period and group where will
you locate the element with Z = 114 ?
Answer:
We know that the gaps of atomic numbers in a group are 8, 8, 18, 18, 32. The
element which proceeds the element with Z = 114 in the same group must have
atomic number equal to 114 – 32 = 82. This represents lead (Pb) which is
present in 6th period and group 14 of the p-block. This means that the element
with Z = 114 must also belong to the group 14 of p-block and it must be a
member of 7th period.
Question
6.
Write the atomic number of the element
present in the third period and seventeenth group of the periodic table.
Answer:
The element is chlorine (Cl) with atomic number (Z) = 17.
Question
7.
Which element do you think would have
been named by :
(i)
(ii) Seaborg’s group ?
Answer:
(i) Lawrencium
(Lr) with atomic number (Z) = 103
(ii) Seaborgium
(Sg) with atomic number (Z) = 106.
Question
8.
Why do elements in a group have similar
physical and chemical properties ?
Answer:
The elements in a group have the same valence shell electronic configurations
of their atoms. However, they differ in the atomic sizes which tend to
increase. Therefore, the elements in a group have similar chemical properties
while their physical characteristics show a little variation.
Question
9.
What do atomic radius and ionic radius
really mean to you ?
Answer:
For the definition and details of atomic radius,
Variation
of Atomic Radius in the Periodic Table
We shall
discuss the variation or change of the atomic radius along the period from left
to the right and also down the group.
Variation in a period. Along a period, the atomic radii of the elements
generally decrease from left to the right.
The atomic radii (metallic radii for Li, Be and B while covalent radii for the
remaining elements) of the elements present in the second period are given :
The trend
in atomic radii have also been shown in the figure 3.5. From the values, it is
clear that the alkali metal (Li) placed on the extreme left has the maximum
atomic radius while the halogen (F) on the extreme right has the minimum value.
Explanation. In moving from left to the right in a period, the nuclear charge
gradually increases by one unit and at the same time one electron is also being
added in the electron shell. Due to increased nuclear charge from left to the
right, the electrons are also getting attracted more and more towards the
nucleus. Consequently, the atomic size is expected to decrease as shown in case
of the elements of second period.
Ionic
Radii
Ionic compounds are crystalline solids and the ionic radii are related to the
ions present in them. The ions formed by the loss of one or more electrons from
the neutral atom is known as cation (positive ion) while the one formed by the
gain of electrons by the neutral atom is called anion (negative ion). Ionic
radius may be defined as :
the
effective distance from the centre of the nucleus of the ion upto which it
exerts its influence on the electron cloud.
Since the
electron cloud may extend itself to a very large distance from the nucleus it
may not be possible to determine the ionic radius experimentally. However, some
theoretical methods have been given to calculate the radius of the ion and the
values as given by Pauling are most acceptable.
But it is
quite easy to determine the inter-nuclear distance between two oppositely
charged ions (say Na+ and
Cl– ions) in the crystal lattice by
X-ray studies. If the absolute radius of one of them is known, that of the other
can be obtained by subtracting the same from the inter-nuclear distance.
Similarly
for the details of ionic radius, In a group, the atomic radii increase
downwards.
For
example, in sodium chloride crystals, the inter-nuclear distance is 276 pm and
the radius of the Na+ ion
is 95 pm (Pauling method). Therefore, the radius of CL– ion is 276 – 95 = 181 pm.
It may be
noted that the value of the ionic radius is linked with the atomic radius and
it varies accordingly. Thus, the ionic radius always increases down the group
and decreases along the period provided the ions involved have the same
magnitude of charge.
This is on
account of the increase in the number of electron shells and also due to
increase in the magnitude of shielding or screening effect. In a period, the
trend is the reverse.
The atomic
radii decrease from left to the right because the electrons are filled in the
same shell and no new shell is formed. The ionic radii of the elements also
follow the same trend.
The atomic
radii of the elements influence the other periodic properties such as
ionisation enthalpy, electron gain enthalpy and electronegativity.
Question
10.
How do atomic radii vary in a period
and in a group ? How do you explain the variation ?
Answer:
Variation in a group. The
atomic radii of the elements in every group of the periodic table increase as
we move downwards. Covalent Radii of the alkali metals present in group 1 are
given. Since all of them are metallic in nature, their metallic radii have also
been given for reference.
The trend
in the atomic radii are also shown in the figure 3.6. From the values, it is
quite clear that the atomic radius of lithium (Li) is the minimum while that of
cesium (Cs) is the maximum. The value of the last element francium (Fr) is not
known since being as radioactive in nature, it is unstable.
Explanation. On moving down a group, there is
an increase in the principal quantum number and thus, increase in the number of
electron shells. Therefore, the atomic size is expected to increase. But at the
same time, there is an increase in the atomic number or nuclear charge also. As
a result, the atomic size must decrease. However, the effect of increase in the
electron shells is more pronounced than the effect of increase in nuclear
charge. Consequently, the atomic size or atomic radius increases down a group.
This is well supported by the values given in Table 3.5 for the alkali metals.
Question
11.
What do you understand by isoelectronic
species ? Name a species that will be isoelectronic with each of the following
atoms or ions.
(i) F– (ii) Ar (in) Mg2+ (iv) Rb+.
Answer:
Isoelectronic species are those species (atoms/ions) which have same number of
electrons (iso-same ; electronic- electrons). The isoelectronic species are
listed :
(i) Na+
(ii) K+
(iii) Na+
(iv) Sr2+.
Question
12.
Consider the following species :
N3-, O2-, F–, Na+, Mg2+, Al3+ .
(a) What is common in them ?
(b) Arrange them in order of increasing
radii.
Answer:
(a) All of them
are isoelectronic in nature and have 10 electrons each.
(b) In
isoelectronic species, greater the nuclear charge, lesser will be the atomic or
ionic radius. Based upon that, the increasing order of atomic radii are :
Al3+ < Mg2+ < Na+ < F– < O2- < N3-.
Question
13.
Explain why are cations smaller and
anions larger in radii than their parent atoms.
Answer:
Cations (positive ions) are formed by the loss of electrons from the parent
atoms. Therefore, they have smaller radii than the parent elements. On the
other hand, anions are formed when the parent atoms take up one or more
electrons. They have therefore, bigger radii than the parent atoms. For
illustration,
Question
14.
What is the significance of the terms—
‘isolated gaseous atom’ and ‘ground state’ while defining ionization enthalpy
and electron gain enthalpy ?
Answer:
(a) Significance of term ‘isolated
gaseous atom’ : The three states of matter as we know differ
in the interparticle spaces. These are the maximum in the solid state while
least in the gaseous state.
The atoms
in the gaseous state are far separated in the sense that they do not have any
mutual attractive and repulsive interactions. These are therefore, regarded as
isolated atoms.
When an
atom in the gaseous state is isolated, its electron releasing tendency and
electron accepting tendency are both absolute in nature.
It means
that their values of ionisation enthalpy and of electron gain enthalpy are not
influenced by the presence of the other atoms. It is not possible to express
these when the atoms are in the liquid or solid state due to the presence of
inter atomic forces.
(b)
Significance of ground state : The ground state means that in a particular atom, the
electrons associated are in the lowest energy state i.e.. they neither lose nor
gain electrons. This represents the normal energy state of an atom. Both
ionisation enthalpy and electron gain enthalpy are generally expressed with
respect to the ground state of an atom only.
Question
15.
The energy of an electron in the ground
state of the hydrogen atom is – 2.18 × 10-18 J. Calculate the ionisation
enthalpy of atomic hydrogen in terms of J mol-1.
Answer:
The ionisation enthalpy is for 1 mole atoms. Therefore, ground state energy of
one mole atoms may be expressed as:
E(groun
state) = (-2.18
× 10-18J) × (6.022 × lO23 mol-1) = – 1.312 × 106Jmol-1
Ionisation enthalpy = Ex –
Egroun
state
= 0 – (-1.312 × 106 J
mol-1) = 1-312 × 106 J mol-1
Question
16.
Among the second period elements, the
actual ionization enthalpies are in the order :
Li<B<Be<C<O<N<F< Ne.
Explain why
(i) Be has higher ∆iH1 than B ?
(ii) O has lower ∆i8H1 than N and F ?
Answer:
(i)We know that the
cations or positive ions are formed when the neutral atoms (generally metal
atoms) lose electrons. They can do so only in the gaseous state since they are
isolated and interatomic forces are the minimum. Now energy is needed to
overcome the force of attraction between the nucleus and electrons so that the
latter may be released. This is known as ionization enthalpy. It may be defined
as :
The minimum amount of energy which is needed to remove the most loosely bound
electron from a neutral isolated gaseous atom in its ground state to form a
cation also in the gaseous state.
Atom (g) → Positive ion (g) + e– ; IE (∆iH1)
(Cation)
Na (g) → Na+ (g)
+ e– ; IE (∆iH1)
The energy required can also be expressed in the form of ionization potential
which is the minimum amount of potential needed to remove the most loosly held
electron from an isolated neutral gaseous atom in its ground state.
(ii) ∆iH1 values of the there elements present in second
period are given :
N(1402 kJ mol-1) ;
O(1314 kJ mol-1) ;
F(1681 kJ mol-1).
• ∆iH1 of O is expected to more than that of N but is
actually lesser because the electronic configuration of N is more symmetrical
as well as stable in comparison to O. For electronic configuration.
• ∆iH1 of O is less than that of F because the ionization
enthalpy in general increases along a period.
Question
17.
∆iH1 value of Mg is more as compared
to that of Na while its ∆iH2 value is less. Explain.
Answer:
The electronic configuration of the two elements and their first and second
ionization enthalpies are.,given :
∆iH1 value of Mg is more than that of Na due to greater
symmetry and smaller size. But ∆iH2 value
of Na is higher because Na+ ion
has the configuration of noble gas element neon while Mg+ ion does not have a symmetrical
configuration.
Question
18.
What are the various factors due to
which the ionisation enthalpy of the main group elements tends to decrease down
the group ?
Answer:
Variation down a group. The ionization enthalpies of the elements decrease on
moving from top to the bottom in any group. The trend is quite evident from the
data given in the table 3.10. The ∆iH1 values
of the members belonging to the alkali metals of group 1 have been given
separately in the fig 3.10.
The decrease in ionization enthalpies down any group is because of the
following factors :
(i) There is an
increase in the number of the main energy shells (n) in moving from one element
to the other.
(ii) There is also
an increase in the magnitude of the screening effect due to the gradual
increase in the number of inner electrons.
Although the nuclear charge qtlso increases down the group which is likely to
result in increased ionization enthalpy but its effect is less pronounced than
the two factors listed above. The net result is the decrease in ionization
enthalpies in a group in the periodic table.
Question
19.
The first ionisation enthalpy values
(kJ mol-1) of group 13 elements are :
How would you account for their
deviation from the general trend ?
Answer:
The decrease in ∆iH1 value from B to Al is quite
expected because of the bigger size of Al atom. But the element Ga has ten 3d
electrons present in the 3d sub-shell which do not screen as much as is done by
s and p electrons. Therefore, there is an unexpected increase in the magnitude
of effective nuclear charge resulting in increased 3 Classification of Elements
and Periodicity in Properties value. The same explanation can be offered
in moving from In to T1. The latter has fourteen 4f electrons with very poor
shielding effect. This also results in unexpected increase in the effective
nuclear charge resulting in inflated ∆iH1 value.
Question
20.
Which of the following pairs of
elements would have a more negative electron gain enthalpy ?
(i) O or F
(ii) F or Cl
Answer:
(i) O or F. The
negative electron gain enthalpy of F(∆egH = – 328 kJ mol-1) is more than that of O(∆egH = – 141 kJ mol-1). This is on account of smaller size of the F atom and its
greater urge to have a noble gas configuration (only one electron is needed) as
compared to oxygen (two electrons are needed).
(ii) F or Cl. The
negative electron gain enthalpy of Cl (∆egH = – 349 kJ mol-1) is more than that of F(∆egH = – 328 kJ mol-1). Actually, the size of F atom (atomic radius = 72 pm) is quite
small as compared to Cl atom (atomic radius = 99 pm). This results in greater
crowding of electrons in small space around the nucleus in case of F and the
attraction for outside electron is less as compared to Cl in which the atomic
size is large and the crowding of electrons is less.
Question
21.
Would you expect the second electron
gain enthalpy of O as positive, more negative or less negative than the first ?
Justify your answer.
Answer:
For oxygen atom ; O(g) + e– →
0–(g) ; (∆egH) = – 141 kJ mol-1
O–(g) e– → 02-(g) ; (∆egH) = + 780 kJ mol-1
The first electron gain enthalpy of oxygen is negative because energy is
released when a gaseous atom accepts an electron to form monovalent anion. The
second electron gain enthalpy is positive because energy is needed to overcome
the force of repulsion between monovalent anion and the second incoming
electron.
Question
22.
What is basic difference between the
terms electron gain enthalpy and electronegativity ?
Answer:
Difference between electron gain enthalpy and electronegativity.
Both electronegativity and electron gain enthalpy are the electron attracting
tendencies of the elements but they differ in many respects. The important
points of distinction are listed.
Question
23.
How will you react to the statement
that the electronegativity of N on Pauling scale is 3-0 in all the nitrogen
compounds ?
Answer:
On Pauling scale, the electronegativity of nitrogen, (3-0) indicates that it is
sufficiently electronegative. This is quite expected as well due to its very
small size and its requirement to have only three electrons to achieve the
noble gas configuration. But it is not correct to say that the
electronegativity of nitrogen in all the compounds is 3. It depends upon its
state of hydridisation in a particular compound. It may be noted that greater
the percentage of ^-character, more will be the electronegativity of the
element. Thus, the electronegativity of nitrogen increases in moving from sp3
hybridised orbitals to sp hydridised orbitals i.e., as sp3 < sp2 < sp.
Question
24.
Describe the theory associated with the
radius of an atom as it :
(a) gains an electron.
(b) loses an electron.
Answer:
(a) When an atom
gains an electron. It changes into an anion (M + e– → M–). In
doing so, there is an increase in the size. Actually, there is no change in the
atomic number or nuclear charge but there is an increase in the electrons by
one. As a result, electrons experience less attraction towards the nucleus and
the anionic radius or size increases. For example, the radius of F atom (72 pm)
is less than that of F– ion
(136 pm).
(b) When an atom loses an electron.
It changes into a cation (M → M+ + e–).
In doing so, there is a decrease in size. In this case also, the nuclear charge
is the same but the electrons have decreased by one. As a result, the electrons
experience more attraction towards the nucleus and the cationic radius
decreases. For example, radius of Na atom (157 pm) is less than that of Na+ ion (95 pm).
Question
25.
What are the major differences between
metals and non-metals ?
Answer:
Metals and non-metals as the names suggest are quite different from one
another. They differ in physical as well as chemical characteristics.
Distinction based upon physical
properties. The main points of distinction are given :
Distinction based upon chemical
properties. The main points of distinction are given :
Question
26.
Would you expect the ionization
enthalpies of two isotopes of the same element to be same or different ?
Justify your answer.
Answer:
The ionization enthalpies are related to the magnitude of nuclear charge as
well as the electronic configuration of the elements. Since the isotopes of an
element have same nuclear charge as well as electronic configuration, they have
the same ionization enthalpies.
Question
27.
Use periodic table to answer the
following questions :
(a) Identify the element with five
electrons in the outer sub-shell.
(b) Identify the element that would
tend to lose two electrons.
(c) Identify the element that would
tend to gain two electrons.
Answer:
(a) Element
belongs to nitrogen family (group 15) e.g. nitrogen
(b) Eletnent
belongs to alkaline earth family (group 2) e.g. magnesium
(c) Element belongs
to oxygen family (group 16) e.g. oxygen
Question
28.
The increasing order of reactivity
among group 1 elements is Li < Na < K < Rb < Cs whereas among the
group 17 elements, it is F > Cl > Br >
Answer:
In group 1 elements (alkali metals) the reactivity of the metals is mainly due
to the electron releasing tendency of their atoms, which is related to
ionization enthalpy. Since ionization enthalpy decreases down the group, the
reactivity of alkali metals increases.
In group 17 elements (halogens-non-metals), the reactivity is linked with
electron accepting tendency of the members of the family. It is linked with
electronegativity and electron gain enthalpy. Since both of them decrease down
the group, the reactivity therefore decreases.
Thus, we conclude that the order of reactivity in two cases is different since
in one case it is due to electron releasing tendency and it is because of
electron accepting tendency in the other case.
Question
29.
Write the general electronic
configuration of s, p, d and /block elements.
Answer:
s-Block elements : ns1-2 (n
varies from 2 to 7)
p-Block elements : ns2 np1-6 (n varies from 2 to 7)
d-Block elements : (n – 1) d1-10 ns1-2 (n varies from 4 to 7)
f-Block elements : (n – 2)f1-14 (n – 1) d0-1 ns2 (n
may have value either 6 or 7)
Question
30.
Assign the position of the elements
having outer electronic configuration
(i) ns2p4 for n = 3
(ii) (n – 1 )d2ns2 for n = 4 and
(iii) (n – 2) f7 (n – 1 )d1 ns2 for n = 6 in the periodic table.
Answer:
(i) The element is
present in third period (n = 3) and belongs to group 16(10 + 2 + 4 = 16).
The element is sulphur : [1s2 2s2 2p6 3s2 3p4].
(ii) The element
is present in fourth period (.n = 4) and belongs to group 4 (2 + 2 = 4).
The element is titanium : [Ne] 3d2 4s2.
(iii) The element
is present in sixth period (n = 6) and belongs to group 3.
The element is gadolinium (Z = 64): [Xe] 4f7 5d1 6s2.
Question
31.
The first (∆iH1) and the second (∆iH2) ionization enthalpies (in kJ mol-1) and the (∆egH) electron gain enthalpy (in kJ mol1) of a few elements are given below :
Which of the above elements is likely
to be :
(a) the least reactive element.
(b) the most reactive metal.
(c) the most reactive non-metal.
(d) the least reactive non-metal.
(e) the metal which can form a stable
binary halide of the formula MX2 (X = halogen)
(f) the metal which can form a
predominantly stable covalent halide of the formula MX (X = halogen) ?
Answer:
The nature and reactivity of the elements are guided by ionisation enthalpies
and electron gain enthalpies. Based upon
the available data :
(a) The least reactive element is ‘V’. It
has very high ∆iH
(2372 kJ mol-1)
and ∆egH is (+ 48 kJ mol-1) which indicate that it is a noble gas
element. The value of ∆egH
suggests it to be helium (He).
(b)
The most reactive metal is ‘II’. It has a very low ∆iH equal to 419 kJ mol-1 and its negative ∆eg value is also quite small (- 48
kJ mol-1). These values correspond to the
alkali metal potassium (K).
In general, alkali metals are the most reactive metals in the periodic table.
(c)
The most reactive non-metal is ‘III’. Since the element has very high ionization enthalpies
(∆iH1 as well as ∆iH2)
and very high negative electron gain enthalpy, it is expected to be a highly
reactive non-metal. The values correspond to non-metal fluorine (F).
(d)
The least reactive non-metal is ‘IV’. The character is revealed by the moderate values of
∆iH1 and ∆iH2.
The value of electron gain enthalpy indicates that the element is iodine (I).
(e)
The metal forming stable binary halide (MX2) is ‘VI’. The formula MX2 suggests that the metal belongs
to the family of alkaline earth (group 2). The ∆iH1 and ∆iH2 values
suggest it to be Mg (However, ∆egH does not agree).
(f)
The metal which can form a predominantly stable covalent halide of the formula.
MX (halogen) is ‘I’. The
metal is monovalent in nature and data suggests it to be lithium (Li) because ∆iH1 ] value is very small as compared to ∆iH2 value. In fact, in alkali metal lithium (Li), the
monovalent cation (Li+)
has the electronic configuration of the noble gas helium (He) The halide is
LiX. The covalent character of LiX is due to the reason that the electron cloud
of Li+ ion attracts the electron cloud
of X-ion towards itself. As a result, the halide is of covalent nature.
Question
32.
Predict the formulas of the stable
binary compounds that would be formed by the combination of the following
pairs of elements :
(a) Lithium and oxygen
(b) Magnesium and nitrogen
(c) Aluminium and iodine
(d) Silicon and oxy gen
(e) Phosphorus and fluorine
(f) Element 71 and fluorine
Answer:
(a) Li2O (Lithium oxide)
(b) Mg3N2 (Magnesium nitride).
(c) AlI3 (Aluminium iodide)
(d) SiO2 (Silicon dioxide)
(e) PF3 (Phosphorus trifluoride) or PF5 (Phosphorus pentafluoride)
(f) The element
with Z.
= 71 is lutetium (Lu) with electronic configuration [Xe] 4f145d1 6s2. With fluorine, it will form a binary compound of
formula LuF3.
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Elements and Periodicity in Properties, help you. If you have any query
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