Electronic Configurations Of Elements MCQ Quiz in తెలుగు - Objective Question with Answer for Electronic Configurations Of Elements - ముఫ్త్ [PDF] డౌన్‌లోడ్ కరెన్

Correct answer: 1)

Concept:

Electron gain Enthalpy: Electron gain enthalpy is the enthalpy change when an electron is added to a neutral gaseous atom (X) to convert it into a negative ion.

• Negative Electron Gain Enthalpy: Negative electron gain enthalpy means negative values as the energy gets released, the group 17 elements (halogen atoms) gain stability by gaining electrons. As the halogens have a very strong affinity to reach to the stable nearest noble gas state, the halogens have a higher negative electron gain enthalpy.

• Positive Electron Gain Enthalpy: Positive electron gain enthalpy is the process when the element shows a certain reluctance during accepting a new electron (generally the second atom). As the noble gases have a high positive electron gain enthalpy due to stable full-filled orbitals, it places the extra gained electron into the higher maximum energy levels and leads to a highly reactive and unstable electronic configuration. Due to the addition of one electron, the atoms then get negatively charged, and thus the addition of the next electron often gets disrupted by electrostatic repulsion. For such types of reactions require a further supply of high energy, causing the electron gain enthalpy of the second electron positive in nature.

Explanation:

Fully-filled and half-filled orbitals have less tendency to gain an electron due to their stability. 

Elements having less than one electron than the nearby noble gas configuration have a  strong tendency to gain electrons.

Elements having 1 or 2 electrons in the outer shell have the tendency to lose electrons to gain the nearest noble gas configuration.

Thus, among the given element's configuration, D has a strong tendency to gain an electron than B. A is the noble gas (Ne) configuration and C has a strong tendency to lose electrons. 

Conclusion:

 Therefore, the correct order of increasing tendency to gain electron A < C < B < D 

Additional Information

Concept:

Actinoids - 

1. Actinides or actinoids are the elements of f-block in periodic table having an atomic number between 90 to 103.
2. They are following the element Actinium and that's why called actinoids or actinides.
3. All actinoids are radioactive in nature.
4. General electectronic configuration of actinoids are [Rn] 5f^1-14 6d^0-1 7s².
5. Examples of actinoids are - Uranium, Neptunium, Thorium, etc.

Explanation:

Actinides or actinoids are the elements of f-block in periodic table having an atomic number between 90 to 103.

⇒ Terbium (Z = 65),

• Terbium(Tb) having atomic number 65 is not lying in the range of actinoid's atomic number.
• Although, terbium itself belongs to the f-block of the periodic table but it is a member of the lanthanide series of f-block.
• Thus, terbium is a lanthanoid, not an actinoid.

The rest of the given options are between atomic no. 90 to 103 and therefore are actinoids.

Hence, only terbium(Z=65) is not an actinoid.

∴ the correct answer is option 1

\(\dfrac {X}{2} \rightarrow \dfrac {X^+}{2} +e \, \dfrac {1}{2}\) .....(i)

\(e+\dfrac {X}{2} \rightarrow \dfrac {X^-}{2} \, \dfrac {1}{2} E.A (-ve)\).....(ii)

(i)+(ii)

\(x \rightarrow \dfrac {1}{2} X^++\dfrac {1}{2} X^- \, \dfrac {1}{2} (I.E-E.A)=410 \, \text{kJ}\)

\(I.E-E.A=820 \, \text{kJ}\)

Now \(\dfrac {1}{2} x^-\rightarrow \dfrac {1}{2} x^++2e^-\)....(iii)

\(\Delta H=735\)

Now equation (iii) can be achieved by (i) + reverse (ii) and we will get

\(\dfrac {1}{2} I.E +\dfrac {1}{2} E.A=735\)

\(I.E +E.A =1470\)....(iv)

\(2E.A=650\)

\(E.A=325 \, \text{kJ/mol}\)

CONCEPT:

Stability of Half-Filled and Fully Filled Electronic Configurations

• Half-filled and fully filled electronic configurations are more stable than other configurations due to exchange energy.
• Exchange energy is the energy released when electrons with parallel spins exchange their positions in degenerate orbitals (orbitals with the same energy).
• This energy is significant in stabilizing the atom or ion.

EXPLANATION:

• In an atom, electrons tend to occupy orbitals in a way that maximizes the number of parallel spins, leading to half-filled or fully filled configurations.
• The stability provided by exchange energy can be understood by considering the following points:
  • Electrons with parallel spins repel each other less than those with opposite spins.
  • This reduced repulsion leads to a lower overall energy of the atom or ion.
  • As a result, configurations with half-filled or fully filled orbitals are more stable due to the maximized exchange energy.
• For example, in the case of the p-orbital:
  • A p³ configuration (half-filled) and a p⁶ configuration (fully filled) are particularly stable.

Therefore, the half-filled and fully filled electronic configurations are more stable than partly filled configurations due to exchange energy.

If the atom or molecule is negatively charged, as an additional electron occupies an outer orbital, there is increased electron-electron repulsion (and hence, increased shielding) which pushes the electrons further apart. Because the electrons now outnumber the protons in the ion, the protons can not pull the extra electrons as tightly toward the nucleus; this results in decreased \(Z_{eff}\). Therefore the ionic radii increase whereas it is vice versa when it is positively charged.

The increasing order of the ionic radii of the given isoelectronic species is \(Ca^{2+},K^{+},Cl^{-},S^{2-}\)

Bicarbonates of \(Li\) and \(Mg\) are more soluble in water than carbonates whereas carbonates of alkali metals are more soluble.

Hydroxides and nitrates of both Li and Mg decompose on heating to give oxide.

\(LiNO_3 \xrightarrow {\Delta} Li_2O +NO_2+O_2\)

\(Mg(NO_3)_2 \xrightarrow {\Delta} MgO+NO_2+O_2\)

Magnesium always form basic carbonates while lithium form normal carbonates, it does not form basic carbonates.

CONCEPT:

Electronic Configuration and Position of Elements

• The position of an element in the periodic table can be determined from its outer electronic configuration. The principal quantum number (n) determines the period, and the type of orbital (s, p, d, f) along with the number of electrons determines the group.
• For an element with the configuration ns₂np₂ (n = 6), the outer electrons suggest the element belongs to group 14 and period 6, as the highest energy level is 6p₂.
• An element with the configuration (n-1)d₂ns₂ (n = 4) belongs to group 4 and period 4, as the d and s orbitals fill accordingly.
• For an element with the configuration (n-2)f₇(n-1)d₁ns₂ (n = 6), it indicates a lanthanide series element, belonging to group 3 and period 6.

EXPLANATION:

• The element 'A' has the electronic configuration [Xe]4f₁₄5d₁₀6s₂6p₂, which corresponds to group 14 and period 6. This is the correct configuration for an element in the p-block of the periodic table, specifically in group 14 and period 6.
• Thus, statement 1 ("The element ‘A’ belongs to 3rd period and 16th group.") is incorrect as element A is in period 6 and group 14.
• Statement 2 ("The element ‘A’ belongs to 6th period and 14th group.") is correct as explained by the electronic configuration of element A.
• For element 'B', the electronic configuration is [Ar]3d₂4s₂, which belongs to group 4 and period 4. This is the configuration of a transition metal in the d-block, confirming the position in group 4 and period 4.
• For element 'C', the configuration is [Xe]4f₇5d₁6s₂, which corresponds to the lanthanide series, group 3, and period 6.
• Statement 3 ("The element ‘C’ belongs to 6th period and is a lanthanide element.") is correct based on the configuration and position of element C in the lanthanide series.
• Statement 4 ("All A, B, C elements are metals.") is also correct since all elements A, B, and C are metals (element A is a p-block metal, B is a transition metal, and C is a lanthanide metal).

Therefore, the correct statements are: 2, 3, and 4.

(B) \(Be\) and \(Mg\) have \(ns^2\) configuration (stable configuration). Thus they have a higher first ionization enthalpy than \(B\) and \(Al\).

(C) \(\text{Due to lanthanide contraction,}\) the atomic and ionic radii of the Niobium and Tantalum are almost the same.

(D) \(r_{metallic}\)>\(r_{covalent}\) (covalent bond formation involves the overlapping of orbitals).

All options are correct.

CONCEPT:

Properties of Halogens

• Halogens are the elements in Group 17 of the periodic table, and their properties change predictably as you move down the group.
• As you move down the group, the atomic number of the halogens increases, and this affects several of their properties.

EXPLANATION:

• Ionisation Enthalpy:
  • Ionisation enthalpy decreases as you move down the group because the outer electrons are further from the nucleus and are more easily removed.
• Ionic Radius:
  • The ionic radius increases as you move down the group. This is because each successive element has an additional electron shell, making the ion larger.
• Enthalpy of Vaporisation:
  • Enthalpy of vaporisation generally increases as you move down the group due to the increase in molecular size and the strength of the van der Waals forces between the molecules.
• More than one of the above:
  • Both the ionic radius and the enthalpy of vaporisation increase as you move down the group.

Therefore, the correct answer is: 4) More than one of the above.

CONCEPT:

Ionisation Enthalpy

• Ionisation enthalpy is the energy required to remove an electron from an atom or ion in its gaseous state.
• The ionisation enthalpy generally increases across a period due to increasing nuclear charge and decreases down a group due to increasing atomic size and shielding effect.

Electron Affinity

• Electron affinity is the amount of energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion.
• Electron affinity generally becomes more negative across a period and less negative down a group.

Atomic Size

• Atomic size refers to the distance from the nucleus of an atom to the outermost shell of electrons.
• Atomic size generally decreases across a period due to increasing nuclear charge and increases down a group due to the addition of electron shells.

EXPLANATION:

• Na < Al < Si < Mg : Ionisation enthalpy
  • This order is incorrect. The correct order should be Na < Mg < Al < Si. Magnesium has a higher ionisation enthalpy than Aluminium due to its filled s-orbital.
•  I < Br < Cl < F : Electron affinity
  • This order is correct. Electron affinity becomes more negative as we move from iodine to fluorine across the period.
• Na+ < F- < O2- < Ne : Atomic size
  • This order is incorrect. The correct order should be Na+ < Ne < F- < O2-. Ne is smaller than F- and O2- because it has a higher nuclear charge with the same number of electrons.

Therefore, the correct answer is Option 4: More than one of the above.