Short & Long Question Answers: Classification of Elements & Periodicity in Properties | Chemistry Class 11 - NEET PDF Download

Short Answer Type Questions

Q1: Write four characteristic properties of p-block elements.

Ans: The four most important properties of p-block elements are as follows:

(a) Both metals and nonmetals are present in p-block elements, but the number of nonmetals is much greater than that of metals. Furthermore, within a group, the metallic character increases from top to bottom, while the non-metallic character increases from left to right along a period in this block.

(b) Their ionisation enthalpies are higher than those of s-block elements.

(c) They mostly combine to form covalent compounds.

(d) Some of them have compounds with multiple (variable) oxidation states. In a period, their oxidising character increases from left to right, while their reducing character increases from top to bottom.

Q2:  All transition elements are d-block elements, but all d-block elements are not transition elements. Explain.

Ans: All transition metals are d-block elements, but not all d-block elements are transition elements because all d-block elements that do not have completely filled d- orbitals are not counted as transition elements, making such elements exceptional. Zn, Cd, and Hg are a few examples.

Q3: 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?

Ans: Among the elements, B, Al, C and Si

(i) The element that has the highest first ionisation enthalpy is C.

(ii) The element that has the most metallic character is Al.

Q4: Explain why the electron gain enthalpy of fluorine is less negative than that of chlorine.

Ans: This is due to the small size of the fluorine atom. As a result of the strong interelectronic repulsions in fluorine’s relatively small 2p orbitals, the incoming electron does not experience much attraction.

Q5: Nitrogen has positive electron gain enthalpy whereas oxygen has negative. However, oxygen has lower ionisation enthalpy than nitrogen. Explain.

Ans: The outermost electronic configuration of nitrogen is 2s² 2p_x¹, 2p_y¹, 2p_z¹ whereas that of oxygen is 2s² 2p_x², 2p_y¹, 2p_z¹. Since oxygen acquires a stable configuration, i.e., 2p³, by removing one electron from the 2p-orbital, it has a lower ionisation enthalpy than nitrogen. In the case of nitrogen, however, due to its stable configuration, it is difficult to remove one of the three 2p-electrons.

Q6: First member of each group of representative elements (i.e., s and p-block elements) shows anomalous behaviour. Illustrate with two examples.

Ans: The first member of each group of representative elements (i.e., the s- and p- block elements) exhibits anomalous behaviour due to:

(i) small size

(ii) higher ionisation enthalpy

(iii) higher electronegativity

(iv) the absence of d- orbitals.

For example, in s – block elements, lithium behaves differently than the other alkali metals.

Q7:  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:  Due to their various properties, p – block elements produce acidic, basic, and amphoteric oxides:

• The higher an element’s electronegativity, the more acidic its oxide.
  For example- Boron has an electronegativity of -2, carbon has an electronegativity of 2.5, and nitrogen has an electronegativity of 3. As a result, the order of acidic character of B, C, and N oxides is B₂O₃ < CO₂ < N₂O₃
• If the ionisation enthalpy of an element is high, it will form acidic oxide; if it is intermediate, it will form amphoteric oxide; and if it is low, it will form basic oxides.
  For example, the ionisation enthalpy of boron is 800 while that of carbon is 1086.5, implying that carbon oxide is more acidic than boron oxide.
• The oxides of the first element in each group in the p – block are more acidic than the oxides of other elements. As we move down the group, the acidic character decreases, followed by elements that form amphoteric oxides and then basic oxides.
  For example- In the Boron family, B forms a weak acidic oxide, while Al, Ga, and In form amphoteric oxides, and Tl forms a strong basic oxide.

Reactions of some of the oxides with water:

• Acidic Oxides:
  B₂O₃ + 3H₂O → 2H₃BO₃
• Basic Oxides:
  Tl₂O + H₂O → 2TlOH
• Amphoteric Oxides are insoluble in water and thus reacts with acid and base:
  Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O
  Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Q8:  Arrange the elements N, P, O and S in the order of-

(i) increasing first ionisation enthalpy.

(i) increasing non-metallic character.

Give reason for the arrangement assigned.

Ans:  (i) Due to the extra stable exactly half-filled 2p-orbitals, the ionisation enthalpy of nitrogen (1s²,2s²,2p³) is greater than that of oxygen (1s²,2s²,2p⁴). Likewise, the ionisation enthalpy of phosphorous (1s²,2s²,2p⁶,3s²,3p³) is greater than that of sulphur (s²,2s²,2p⁶,3s²,3p³)

Ionisation enthalpy decreases with decreasing atomic size as one moves down the group.

As a result, the increasing order of first ionisation enthalpy is S < P < O < N

(ii) Nonmetallic character increases across a period (left to right), but decreases as one moves down the group.

As a result, the increasing non-metallic order is P < S < N < O.

Long Answer Type Questions

Q1: Discuss the factors affecting electron gain enthalpy and the trend in its variation in the periodic table.

Ans:  Factors influencing electron gain enthalpy includes-

(i) Nuclear charge: As the nuclear charge increases, the electron gain enthalpy becomes more negative. If the nuclear charge is high, there is a greater attraction for the incoming electron.

(ii) Atomic size: As the atom’s size increases, so does the distance between the nucleus and the incoming electron, resulting in less attraction. As a result, as the size of the element’s atom increases, the electron gain enthalpy becomes less negative.

(iii) Electronic configuration: Elements with stable electronic configurations of half-filled and completely filled valence subshells have a very low tendency to accept additional electrons, resulting in less negative electron gain enthalpies.

Variations in electron gain enthalpies in the periodic table

In general, electron gain enthalpy becomes more negative from left to right in a period and less negative from top to bottom in a group.

(a) Downward variation within a group: Moving down a group increases the size and nuclear charge. However, the effect of increasing atomic size is much more pronounced than that of increasing nuclear charge, so the additional electron feels less attraction by the large atom. As a result, the electron gain enthalpy decreases. This is evident from the decrease in electron gain enthalpy when transitioning from chlorine to bromine and then to iodine.

(b) Periodic variation: As one moves across a period, the size of the atom decreases and the nuclear charge increases. Because both of these factors increase the attraction for the incoming electron, electron gain enthalpy becomes more negative in a period from left to right. However, there are some anomalies in the overall trend. These are primarily due to certain atoms’ stable electronic configurations.

Important Trends in Electron Gain Enthalpies

The electron gain enthalpies of elements have some important characteristics. They are as follows:

(i) The negative electron gain enthalpies of halogens are the highest.

(ii) Noble gases have positive electron gain enthalpy values, whereas Be, Mg, N, and P have nearly zero.

(iii) Fluorine’s electron gain enthalpy is unexpectedly less negative than chlorine’s.

Q2: Define ionisation enthalpy. Discuss the factors affecting ionisation enthalpy of the elements and its trends in the periodic table.

Ans: Ionisation Enthalpy: The amount of energy required to remove an e from an isolated gaseous atom in its gaseous state is defined as an element’s ionisation enthalpy.

The following factors influence ionisation enthalpy:

• Atom size: The larger the atomic size, the lower the value of ionisation enthalpy. The outer e^– are far away from the nucleus in large atoms, so the force of attraction with which they are attracted by the nucleus is less and thus they can be easily removed.
  Ionization enthalpy ∝ 1/atomic size
• Screening Effect: Because the screening effect reduces the force of attraction towards the nucleus, the outer e can be easily removed.
  Ionization enthalpy ∝ 1/Screening effect
• Nuclear charge: The Ionisation enthalpy increases as nuclear charge increases among atoms with the same number of energy shells because of the force of attraction towards the nucleus increases.
  Ionisation enthalpy ∝ nuclear charge
• Half-filled and fully-filled orbitals: Because atoms with half-filled and fully-filled orbitals are more stable, it takes more energy to remove an electron from such atoms. In the case of such an atom, the ionisation enthalpy is somewhat higher than expected. Ionisation enthalpy ∝ stable electronic configuration
• Orbital shape: The s-orbital of the same orbit is closer to the nucleus than the p-orbital. As a result, removing an electron from a p-orbital is easier than from an s-orbital. The shape for orbitals: s > p > d > f

Variation of ionisation enthalpy in the periodic table

In general, as atomic size increases, the ionisation energy decreases down the group. The ionisation energy, on the other hand, increases across the period from left to right, because of a decrease in atomic size from left to right

Q3: Write down the outermost electronic configuration of alkali metals. How will you justify their placement in group 1 of the periodic table?

Ans: Alkali metals’ outermost electronic configuration is ns¹.

All elements of group IA (or I), i.e., alkali metals, have the same outer electronic configuration, ns¹, where n denotes the number of principal shells. These electronic configurations are shown in the table below.

As a result of their similarity in electronic configuration and properties, all of these elements are placed in group 1 of the periodic table.

Q4. Write the drawbacks in Mendeleev’s periodic table that led to its modification.

Ans: Drawbacks in Mendeleev’s periodic table are:

1. Hydrogen’s position: Hydrogen is assigned to group I. It does, however, resemble elements from Group I (alkali metals) as well as elements from Group VIIA (halogens). As a result, the position of hydrogen in the periodic table is incorrect.
2. Anomalous pairs: The increasing order of atomic masses was not followed in certain pairs of elements. Mendeleev arranged the elements in these cases based on similarities in their properties rather than the increasing order of their atomic masses. Argon (Ar, atomic mass 39.9), for example, is placed before potassium (K, atomic mass 39.1). Likewise, cobalt (Co, atomic mass 58.9) comes before nickel (Ni, atomic mass 58.6), and tellurium (Te, atomic mass 127.6) comes before iodine (I, atomic mass 126.9). These positions were not justified.
3. Isotopes are atoms of the same element that have different atomic masses but the same atomic number. As a result, according to Mendeleev’s classification, these should be classified differently based on their atomic masses. For example, hydrogen isotopes with atomic masses 1, 2, and 3 should be placed in three different locations. Isotopes, on the other hand, do not have their own spot in the periodic table.
4. Several gaps in the periodic table were left because he believed that several elements were yet to be discovered, for example, gallium was not discovered at the time.
5. Position of lanthanoids (or lanthanides) and actinoids (or actinides): The fourteen elements that follow lanthanum (known as lanthanoids, atomic numbers 58-71) and the fourteen elements that follow actinium (known as actinoids, atomic numbers 58-71) are not included separately.

Q5: In what manner is the long form of periodic table better than Mendeleev’s periodic table? Explain with examples.

Answer. Due to the following reasons, the long-form periodic table is considered more letter than the Mendeleev’s table:

1. All elements in the long-form periodic table are arranged in increasing order of atomic numbers, whereas the table is arranged in increasing order of atomic masses.
2. The position of hydrogen in the long-form periodic table has been justified, whereas there is no such justification in Mendeleev’s periodic table.
3. The long-form periodic table considers the filling of electrons in s,p,d, and subshells, whereas the table considers the atomic numbers of the elements.
4. The periodic table is divided into four blocks: s, p, d, and f, whereas the periodic table has no blocks.
5. Long-form periodic table groups are not further subdivided into subgroups, whereas each group in Mendeleev’s periodic table has subgroups A and B.
6. Long-form periodic tables are simple and easy to reproduce, whereas Mendeleev’s periodic tables are quite difficult to reproduce.

​