Periodicity In Properties MCQ Quiz in मल्याळम - Objective Question with Answer for Periodicity In Properties - സൗജന്യ PDF ഡൗൺലോഡ് ചെയ്യുക

Concept:

Ionization enthalpy is the energy required to remove the most loosely bound electron from a gaseous atom in its ground state. The general trend of ionization enthalpy along the period and group is influenced by several factors:

• Across a Period: Ionization enthalpy increases from left to right across a period due to increasing nuclear charge, which pulls the electrons closer to the nucleus, making them harder to remove.

• Down a Group: Ionization enthalpy decreases as we move down a group due to an increase in atomic size and shielding effect, making it easier to remove electrons from the outermost shell.

• Shielding Effect: The inner electrons shield the outer electrons from the full attraction of the nucleus, making the ionization energy lower.

• Atomic Radius: As the atomic radius increases, the distance between the nucleus and the outermost electron increases, lowering the ionization enthalpy.

• Electronic Configuration: Elements with stable electronic configurations (such as noble gases) have much higher ionization enthalpies compared to elements with an unstable configuration.

Explanation: 

• Sodium (Na) has the lowest ionization enthalpy because it is an alkali metal with a single electron in its outermost shell, making it easy to remove.

• Magnesium (Mg) has a higher ionization enthalpy than Na and Al due to a smaller atomic radius , complete 3s orbital ,and a higher nuclear charge, making electron removal more difficult.

• Aluminum (Al) follows Mg, but its ionization energy is slightly lower than expected because of its electron configuration (3p electron is easier to remove than a 3s electron).

• Silicon (Si), Phosphorus (P), and Sulfur (S) have progressively higher ionization enthalpies as their nuclear charge increases across the period.

• Chlorine (Cl) has the highest ionization enthalpy because it is a halogen with a high effective nuclear charge and a small atomic radius.

Conclusion:

The correct order of first ionization enthalpy for Na, Mg, Al, Si, P, S, and Cl is: Na < Al < Mg < Si < S < P < Cl

Concept:

Isoelectronic species are atoms and ions that have the same number of electrons but different nuclear charges (atomic numbers). The size of an isoelectronic species is influenced by the effective nuclear charge (Z_eff), which is the net positive charge experienced by electrons in the valence shell. The greater the effective nuclear charge, the more strongly the electrons are pulled towards the nucleus, resulting in a smaller atomic or ionic radius.

Explanation:

• Fluoride ion (F^–): Has 9 protons and 10 electrons.

• Neon atom (Ne): Has 10 protons and 10 electrons.

• Sodium ion (Na⁺): Has 11 protons and 10 electrons.

The increasing nuclear charge (number of protons) among these species causes their sizes to decrease as you go from F^– to Ne to Na⁺.

Order of Effective Nuclear Charge (Z_eff):

The nuclear charge increases from F^– to Ne to Na⁺, resulting in the following order of effective nuclear charge and atomic or ionic sizes:

• F^– (smallest Z_eff, largest size)

• Ne (intermediate Z_eff, intermediate size)

• Na⁺ (largest Z_eff, smallest size)

Order of Sizes

Therefore, the correct order of the sizes of the isoelectronic species from largest to smallest is:

F^– > Ne > Na⁺

Conclusion:

The order of sizes for isoelectronic species F^– > Ne > Na⁺.

CONCEPT:

Quantum Numbers and Electronic Configuration in Lu³⁺

• Quantum numbers are used to describe the position and energy of electrons in an atom.
• The principal quantum number (n) describes the energy level of the electron and the distance from the nucleus.
• The azimuthal quantum number (l) describes the shape of the orbital and can have integer values from 0 to (n-1).

EXPLANATION:

To determine the principal (n) and azimuthal (l) quantum numbers of the valence electrons in tripositive Lutetium (Lu³⁺), we need to first determine the electronic configuration of neutral Lutetium (Lu).

Lu: [Xe] 4f¹⁴ 5d¹ 6s²

Removing three electrons from neutral Lutetium involves the removal of two 6s electrons and one 5d electron:

Lu³⁺: [Xe] 4f¹⁴

• Neutral Lutetium has an atomic number of 71. Its electron configuration is:
• In the tripositive state (Lu³⁺), three electrons are removed. The electrons are generally removed from the outermost orbitals first.
• In Lu³⁺, the valence electrons are in the 4f subshell.
• The principal quantum number (n) for the 4f subshell is 4, and the azimuthal quantum number (l) for the f subshell is 3.

Therefore, the correct set of quantum numbers for the valence electrons in tripositive Lutetium is n = 4 and l = 3

So, the correc answer is Option 4: n = 4 and l = 3

Concept:

Electron gain enthalpy is the energy change that occurs when an electron is added to a neutral atom in the gas phase to form a negative ion. More negative values indicate a higher tendency to gain an electron. Typically, electron gain enthalpy becomes more negative across a period from left to right due to increasing nuclear charge and becomes less negative down a group owing to increasing atomic size and electron-electron repulsions.

Explanation:

• Sulfur (S): Has a moderately high electron affinity, resulting in a fairly negative electron gain enthalpy. However, it is less negative than those of fluorine and chlorine.

• Phosphorus (P): Has a half-filled orbital which gives it some stability. Adding an electron would disrupt this stability, resulting in a less favorable (less negative) electron gain enthalpy.

• Fluorine (F): Has a very high nuclear charge and small atomic radius, resulting in the most negative electron gain enthalpy among these elements.

• Chlorine (Cl): Has a high nuclear charge and a relatively small atomic radius, resulting in a highly negative electron gain enthalpy, almost as negative as fluorine’s.

Conclusion:

Based on the explanations, the element with the least negative electron gain enthalpy among the given options is: Phosphorus(P).

Concept:

The oxidation state of a transition metal refers to the charge it would have if all bonds to different atoms were completely ionic. Transition metals can exhibit a variety of oxidation states, often including very high ones, due to the involvement of d-electrons.

Explanation:

Consider the highest known oxidation states of the given transition metals:

• Palladium (Pd): The highest oxidation state of palladium is +4.

• Osmium (Os): Osmium can reach an oxidation state of +8, as seen in compounds like osmium tetroxide (OsO₄).

• Chromium (Cr): The highest oxidation state of chromium is +6, seen in compounds such as potassium dichromate (K₂Cr₂O₄).

• Manganese (Mn): Manganese can exhibit an oxidation state of +7, as seen in potassium permanganate (KMnO₄).

Conclusion:

The transition metal that exhibits the highest oxidation state is: Osmium (Os) with an oxidation state of +8.

Concept:

The electronic configuration of an element reveals the arrangement of electrons in its atomic orbitals. This information helps to classify the element as a metal, non-metal, metalloid, or inert gas.

• Metals: Metals typically have few electrons in their outermost shell(ns¹ or ns² or (n-1)d^1-10 ns^1-2) and tend to lose electrons to attain a stable configuration. This gives them characteristic properties such as high electrical conductivity, malleability, ductility, and luster. Examples include alkali metals, alkaline earth metals, and transition metals.
• Non-metals: Non-metals usually have more electrons in their outer shells(ns² np^3-6) and tend to gain or share electrons to achieve stability. They are generally poor conductors of heat and electricity, and they often have lower melting and boiling points compared to metals. Common non-metals include carbon, nitrogen, oxygen, etc.
• Metalloids: Metalloids have properties intermediate between metals and non-metals. They can exhibit mixed characteristics and are often semiconductors. Examples include elements like silicon and germanium.
• Inert Gases (Noble Gases): Inert gases have a complete outer electron shell(ns² np⁶), making them very stable and mostly unreactive. These elements are found in Group 18 of the periodic table and include helium, neon, argon, etc.

Explanation:

The electronic configuration provided is: 1s², 2s², 2p⁶, 3s². We can sum the electrons to determine the atomic number:

2 (1s) + 2 (2s) + 6 (2p) + 2 (3s) = 12

This corresponds to an element with atomic number 12, which is Magnesium (Mg).

Alternate Method Given electronic configuration is comparable to ns² electronic configuration which belongs to metallic group.

Magnesium (Mg) belongs to Group 2 of the periodic table and is classified as an alkaline earth metal. Metals are characterized by their ability to lose electrons and form positive ions, their high electrical conductivity, and their malleability and ductility.

Conclusion:

The element with the electronic configuration 1s², 2s², 2p⁶, 3s² is Metal.

Concept:

The ionization potential (IP) is the energy required to remove an electron from a gaseous atom or ion. The second ionization potential refers to the energy required to remove a second electron after the first has already been removed. It usually requires more energy than the first ionization potential because it involves removing an electron from a positively charged ion.

Explanation:

The second ionization potential follows trends in the periodic table and depends on the additional stability gained or lost when an electron is removed. 

• Potassium (K): After losing the first electron, potassium forms a K⁺ ion which has a noble gas configuration (Ar). Removing another electron from a stable, noble gas configuration requires significantly more energy.

• Calcium (Ca): After losing the first electron, calcium forms a Ca⁺ ion. The removal of the second electron, leading to Ca²⁺, still results in a relatively stable noble gas configuration (Ar), but requires less energy compared to potassium due to it not starting from a noble gas configuration originally.

• Barium (Ba): Barium, similar to calcium, after losing the first electron forms a Ba⁺ ion. However, due to its larger atomic size and lower effective nuclear charge compared to calcium, it requires even less energy for the second ionization than calcium.

Conclusion:

The decreasing order of the second ionization potential of K, Ca, and Ba is: K > Ca > Ba

Concept:

Pauling's scale of electronegativity is a useful tool in chemistry. Electronegativity refers to the ability of an atom to attract shared electrons in a chemical bond. These values are particularly useful in predicting certain properties of compounds.

Explanation:

Pauling's electronegativity: \(\Delta \chi = \sqrt{ \left(E(A-B) - \frac{E(A-A) + E(B-B)}{2} \right)}\)

Where

•  \(\)\(\Delta \chi\) is the difference in electronegativity between atom A and atom B.
• E(A-B) is the bond energy of the A-B bond.
• E(A-A) is the bond energy of the A-A bond.
• E(B-B)  is the bond energy of the B-B bond.

Using Pauling's electronegativity values:

• 1) Polarity of bonds in molecules: Electronegativity differences between atoms in a molecule help in predicting whether a bond is nonpolar covalent, polar covalent, or ionic. A larger difference in electronegativity between two atoms typically results in a more polar bond.
• 2) Position of elements in Electromotive series: This is typically determined by standard reduction potentials rather than electronegativity values.
• 3) Coordination number: Coordination number is more related to the size and geometry of ions, and the specific bonding characteristics in coordination compounds, rather than electronegativity.
• 4) None of these

Conclusion:

Pauling's electronegativity values are primarily useful in predicting the polarity of bonds in molecules. Therefore, the correct answer is: Polarity of bonds in molecules

Concept:

In the periodic table, the melting points (mp) and boiling points (bp) of elements can vary across different groups. These trends are influenced by the type of bonding, atomic size, and structure of the elements.

Explanation:

Comparing the melting and boiling point trends down the groups of the periodic table:

• Group 1 (Alkali Metals): Generally, mp/bp decrease down the group.
• Group 2 (Alkaline Earth Metals): Variation is not straightforward, but generally mp tends to decrease slightly, while bp shows irregular trends.
• Group 17 (Halogens): Both mp/bp increase down the group due to increasing molecular size and stronger van der Waals forces.
• Group 13 (Boron Group): Show varied trends, usually not a straightforward increase or decrease.

Among these groups, the halogens (Group 17) show a clear trend where both melting points and boiling points increase down the group. This is primarily due to the increase in molecular size, which leads to stronger van der Waals forces that require more energy to overcome.

Conclusion:

Therefore, the correct group where melting points and boiling points increase down the group is: Group 17.

Concept:

Electron affinity is the amount of energy released when an atom in the gaseous phase accepts an electron to form a negative ion. Generally, electron affinity becomes more negative across a period (from left to right) and less negative down a group (from top to bottom) in the periodic table.

Explanation:

To determine which element has the maximum electron affinity, let us consider the known values for the listed elements:

• Fluorine (F): Has a very high electron affinity, but the small size of the atom results in significant electron-electron repulsion in the valence shell, slightly reducing its electron affinity compared to chlorine.
• Chlorine (Cl): Chlorine has the highest electron affinity among the halogens and indeed among all the elements. Chlorine's larger atomic size compared to fluorine reduces electron-electron repulsion while still having a strong effective nuclear charge.
• Bromine (Br): Has a high electron affinity, but it is lower than that of chlorine due to increased atomic size, which leads to less effective nuclear charge experienced by the added electron.
• Iodine (I): Has an even lower electron affinity than bromine due to even larger atomic size and increased shielding effect, making the added electron less bound to the nucleus.

Conclusion:

The element with the maximum electron affinity is: Cl