CBSE Class 11 Chemistry Classification of Elements and Periodicity in Properties Worksheet Set 04

Periodic Properties Worksheet

Question 1. Name the factors that affect the ionisation enthalpy of an element.
Answer:
1. Size of the atom: the larger the atom, the farther the outer electron is from the nucleus and the lower the ionisation enthalpy.
2. Nuclear charge: the greater the nuclear charge, the higher the ionisation enthalpy.
3. Screening (shielding) effect: more inner electrons shield the outer electron, so the ionisation enthalpy is lower.
4. Penetration of the electron: within a shell the penetration is s > p > d > f, so an s electron is held more strongly than a p electron of the same shell.
5. Electronic configuration: filled and half-filled sub-shells are stable, so more energy is needed to remove an electron from them.
In simple words: Smaller size, bigger nuclear charge, less shielding and a stable configuration all raise the ionisation enthalpy.

Teacher's Note:
1. Learn the five factors as size, charge, shielding, penetration and configuration.
2. Give one example for the configuration factor, such as N higher than O.

Question 2. Explain why the ionisation enthalpy (a) increases from left to right across a period and (b) decreases down a group.
Answer:
(a) Across a period: the nuclear charge rises by one unit at each step, the extra electron enters the same shell, and the atomic size decreases. The effective nuclear charge increases, so the outer electron is held more tightly and more energy is needed to remove it. (Period 2: Li 520 to Ne 2081 kJ mol⁻¹.)
(b) Down a group: a new shell is added at every step, so the outer electron is farther from the nucleus and is shielded by more inner electrons. The attraction falls and less energy is needed. (Group 1: Li 520, Na 496, K 419, Rb 403, Cs 376 kJ mol⁻¹.)
In simple words: Across a period the pull gets stronger; down a group the electron gets farther away.

Teacher's Note:
1. Mention both nuclear charge and size in part (a).
2. Mention the new shell and shielding in part (b).

Question 3. Among the elements Li, K, Ca, S and Kr, which is expected to have the lowest first ionisation enthalpy and which the highest?
Answer:
First ionisation enthalpies (kJ mol⁻¹): K 419, Li 520, Ca 590, S 1000 and Kr 1351.
Lowest: K. It is a large alkali metal with a single 4s electron far from the nucleus.
Highest: Kr. It is a noble gas with a complete outer shell (4s² 4p⁶) at the right-hand end of its period.
In simple words: The alkali metal has the lowest value and the noble gas the highest.

Teacher's Note:
1. Use position in the table: left and down means low; right and up means high.
2. Ca is in the same period as K but has a higher nuclear charge, so Ca > K.

Question 4. Would you expect the first ionisation enthalpies of two isotopes of the same element to be the same or different? Justify your answer.
Answer:
The ionisation enthalpy depends on the electronic configuration and the nuclear charge (number of protons).
Isotopes of an element have the same number of protons and electrons, so they have the same electronic configuration and the same nuclear charge.
They differ only in the number of neutrons, which does not affect the attraction on the outer electron. Therefore their first ionisation enthalpies are the same.
In simple words: Isotopes differ only in neutrons, and neutrons do not attract electrons.

Teacher's Note:
1. Neutrons are neutral and have no effect on ionisation enthalpy.
2. The same argument shows that isotopes have the same chemical properties.

Question 5. Explain: (a) The first ionisation enthalpy of Mg is higher than that of Al. (b) The first ionisation enthalpy of P is higher than that of S.
Answer:
(a) Mg is 1s² 2s² 2p⁶ 3s² and Al is 1s² 2s² 2p⁶ 3s² 3p¹. The electron removed from Mg comes from a filled 3s sub-shell, which is more stable and penetrates closer to the nucleus. The electron removed from Al is a 3p electron, which is less strongly attracted. So 738 kJ mol⁻¹ (Mg) > 577 kJ mol⁻¹ (Al).
(b) P is 3s² 3p³ with a half-filled 3p sub-shell, which is extra stable. S is 3s² 3p⁴, and one of the 3p orbitals holds a pair of electrons, so the paired electron is removed more easily. So 1012 kJ mol⁻¹ (P) > 1000 kJ mol⁻¹ (S).
In simple words: Filled and half-filled sub-shells hold their electrons more tightly.

Teacher's Note:
1. These are the exceptions to the general rise of ionisation enthalpy across a period.
2. The same reason explains Be > B and N > O.

Question 6. Why is the third ionisation enthalpy of magnesium (7733 kJ mol⁻¹) so much higher than its first (738) and second (1451) ionisation enthalpies?
Answer:
Mg is 1s² 2s² 2p⁶ 3s². The first two electrons are removed from the 3s sub-shell, which is the valence shell. This gives Mg⁺ and then Mg²⁺.
Mg²⁺ has the stable neon configuration 1s² 2s² 2p⁶. The third electron has to be removed from this completely filled inner shell, which is close to the nucleus and strongly held.
Therefore the energy needed jumps sharply (738 → 1451 → 7733 kJ mol⁻¹).
In simple words: The first two electrons are valence electrons; the third comes from a stable neon core.

Teacher's Note:
1. A large jump in successive ionisation enthalpies marks the end of the valence electrons.
2. Here the jump after the second electron shows that Mg is in group 2.

Question 7. The first ionisation enthalpy of carbon is greater than that of boron, but the second ionisation enthalpy of boron is greater than that of carbon. Explain.
Answer:
First: C (1s² 2s² 2p²) has a higher nuclear charge and a smaller size than B (1s² 2s² 2p¹), so ΔiH₁ of C (1086) is greater than that of B (801 kJ mol⁻¹).
Second: after losing one electron, B⁺ is 1s² 2s², which has a stable, filled 2s sub-shell, while C⁺ is 1s² 2s² 2p¹. Removing an electron from the filled 2s of B⁺ is harder than removing the 2p electron of C⁺.
So ΔiH₂ of B (2427) is greater than that of C (2353 kJ mol⁻¹).
In simple words: B⁺ has a filled 2s sub-shell, so it resists a second ionisation more than C⁺.

Teacher's Note:
1. Write the configuration of the M⁺ ion before comparing second ionisation enthalpies.
2. A filled s sub-shell gives extra stability.

Question 8. Out of Na⁺ and Ne, which has the higher ionisation enthalpy? Explain.
Answer:
Na⁺ and Ne are isoelectronic (10 electrons, 1s² 2s² 2p⁶).
The nuclear charge of Na⁺ is +11 and that of Ne is +10. The same 10 electrons are held by a larger nuclear charge in Na⁺, so they are held more tightly and Na⁺ is smaller.
Hence Na⁺ has the higher ionisation enthalpy (4562 kJ mol⁻¹ for Na⁺ against 2081 kJ mol⁻¹ for Ne).
In simple words: Same electrons, more protons: Na⁺ holds them harder.

Teacher's Note:
1. For isoelectronic species, ionisation enthalpy rises with nuclear charge.
2. Na⁺ → Na²⁺ is the second ionisation of sodium.

Question 9. Noble gases have positive (zero or near zero) electron gain enthalpy values. Explain.
Answer:
The outermost shell of a noble gas is completely filled (ns²np⁶, and 1s² for He).
An incoming electron has to enter the next higher shell, which is farther from the nucleus and is screened by the filled shell, so it is not attracted.
Energy has to be supplied to add the electron, so the electron gain enthalpy is positive (for example He +48, Ne +116 and Ar +96 kJ mol⁻¹).
In simple words: The shell is full, so the new electron has nowhere stable to go.

Teacher's Note:
1. Positive ΔegH means energy is absorbed.
2. Compare with the halogens, whose ΔegH is the most negative because one electron completes the octet.

Question 10. Why are the electron gain enthalpies of Be and Mg positive?
Answer:
Be (2s²) and Mg (3s²) have completely filled s orbitals in their valence shells.
An extra electron would have to go into the higher-energy p orbital of the valence shell, which is poorly attracted and less stable.
Energy has to be supplied, so the electron gain enthalpy is positive.
In simple words: A filled s sub-shell does not welcome another electron.

Teacher's Note:
1. Nitrogen shows a similar pattern because of its half-filled 2p³ sub-shell.
2. Filled and half-filled sub-shells lower the tendency to gain electrons.

Teacher's Note:
1. The words 'isolated gaseous atom' and 'shared pair' are the key distinction.
2. Cl has a more negative electron gain enthalpy than F, but F is more electronegative than Cl.

Question 12. Fluorine is the most electronegative element. Explain.
Answer:
Electronegativity increases across a period and decreases down a group.
Fluorine is at the top of group 17, so it has the smallest size of all the halogens. The shared pair in a bond is therefore very close to the fluorine nucleus and is strongly attracted.
The noble gases (group 18) have complete shells, so the most electronegative elements of a period are the halogens. Thus F, being at the top right (excluding the noble gases), has the highest electronegativity, 4.0 on the Pauling scale.
In simple words: Fluorine is the smallest atom that is short of one electron, so it pulls bonding electrons the hardest.

Teacher's Note:
1. Quote the value 4.0 on the Pauling scale.
2. Remember that noble gases are not given electronegativity values in the usual tables.

Question 13. The increasing order of reactivity among group 1 elements is Li < Na < K < Rb < Cs, whereas that among group 17 elements is F > Cl > Br > I. Explain.
Answer:
Group 1: each element has one valence electron and reacts by losing it. Ionisation enthalpy decreases down the group (Li 520 to Cs 376 kJ mol⁻¹), so the tendency to lose the electron increases and reactivity rises: Li < Na < K < Rb < Cs.
Group 17: each element has seven valence electrons and reacts by gaining one. The tendency to accept an electron falls down the group because of the increasing size and shielding, so reactivity falls: F > Cl > Br > I.
(Fluorine's reactivity is also helped by its low F–F bond dissociation enthalpy.)
In simple words: Metals react by losing electrons (easier down a group); non-metals react by gaining them (harder down a group).

Teacher's Note:
1. Link the reactivity of metals to ionisation enthalpy and that of non-metals to electron gain enthalpy.
2. Do not write that 'F has the most negative electron gain enthalpy'; Cl does.

Teacher's Note:
1. Metals lie on the left and in the middle of the periodic table, non-metals on the upper right.
2. Metalloids such as Si and Ge show intermediate behaviour.