Periodicity In Properties MCQ Quiz - Objective Question with Answer for Periodicity In Properties - Download Free PDF

CONCEPT:

Ionic Radii

• The ionic radius is the measure of an atom's ion in a crystal lattice. It is half the distance between two ions that are barely touching each other.
• The size of an ion is affected by its charge. As the number of electrons increases (for anions), the ion becomes larger. Conversely, as the number of electrons decreases (for cations), the ion becomes smaller.

EXPLANATION:

•  C4- :
  • With the lowest nuclear charge ( Z = 6 ), C4-  has the largest ionic radius in this series.
• N3- :
  • Nitrogen ( Z = 7 ) has a slightly higher nuclear charge than carbon, so N3- has a smaller ionic radius than C4-  .
• O2- :
  • Oxygen ( Z = 8 ) has a higher nuclear charge than nitrogen, so O2-  has a smaller ionic radius than N3- .
• F- :
  • Fluorine ( Z = 9 ) has the highest nuclear charge in this series, so F^- has the smallest ionic radius.

Increasing Order of Ionic Radii: F- < O2- < N3- < C4- .

CONCEPT:

Trends in Ionization Energy

• Ionization energy generally increases across a period from left to right due to increasing nuclear charge.
• Ionization energy generally decreases down a group due to the increasing distance of the outer electron from the nucleus and increased electron shielding.

Explanation:-

C > Si > Ge > Sn < Pb

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• Statement (I): The first ionization energy of Pb is greater than that of Sn
  • True. Pb has a slightly higher ionization energy compared to Sn due to relativistic effects and poor shielding by 4f and 5d electrons, making it more difficult to remove the outermost electron.
• Statement (II): The first ionization energy of Ge is greater than that of Si
  • False. The first ionization energy of Si is greater than that of Ge. This is because Ge is lower in the group, and thus, it has a larger atomic radius and more shielding, making it easier to remove an electron compared to Si.

Conclusion:-

The correct answer is: Statement I is true but Statement II is false (Option 1).

CONCEPT:

Electron Gain Enthalpy

• Electron gain enthalpy (also known as electron affinity) is the energy change that occurs when an electron is added to a neutral atom in the gaseous state to form a negative ion. Generally, the more negative the electron gain enthalpy, the greater the tendency of the atom to accept an electron.
• The trend in electron gain enthalpy is influenced by several factors such as atomic size, nuclear charge, and electron configuration.

ELECTRON GAIN ENTHALPY VALUES:

• Fluorine (F): –333 kJ/mol
• Bromine (Br): –325 kJ/mol
• Sulfur (S): –200 kJ/mol
• Argon (Ar): +96 kJ/mol

SEQUENCE-BASED ON ELECTRON GAIN ENTHALPY:

• The most negative electron gain enthalpy indicates the greatest tendency to accept an electron.
• Given the values, the sequence from most negative to least negative (or even positive) is:
  • C. Fluorine (F): –333 kJ/mol
  • B. Bromine (Br): –325 kJ/mol
  • D. Sulfur (S): –200 kJ/mol
  • A. Argon (Ar): +96 kJ/mol

Therefore, the correct answer is A > D > B > C.

CONCEPT:

Classification of Molecules Based on the Octet Rule

• Molecules obeying the octet rule: These are molecules where all atoms have eight electrons in their valence shell.
  • Examples: CCl₄, CO₂
• Molecules with incomplete octet: These molecules have atoms that do not have eight electrons in their valence shell.
  • Examples: BCl₃, AlCl₃
• Molecules with incomplete octet with an odd electron: These are molecules with an odd number of electrons, resulting in at least one atom not having a complete octet.
  • Examples: NO, NO₂
• Molecules with expanded octet: These are molecules where the central atom has more than eight electrons in its valence shell.
  • Examples: H₂SO₄, PCl₅

EXPLANATION:​

• (A) Molecules obeying the octet rule → (IV) CCl₄, CO₂
• (B) Molecules with incomplete octet → (II) BCl₃, AlCl₃
• (C) Molecules with incomplete octet with an odd electron → (I) NO, NO₂
• (D) Molecules with expanded octet → (III) H₂SO₄, PCl₅

Therefore, the correct answer is: A-IV, B-II, C-I, D-III

Explanation:-

Statement (I): An element in the extreme left of the periodic table forms acidic oxides.

This statement is false. Elements in the extreme left of the periodic table are alkali metals, which form basic oxides, not acidic oxides. For example, sodium oxide (Na₂O) and potassium oxide (K₂O) are basic in nature.

Statement (II): Acid is formed during the reaction between water and oxide of a reactive element present in the extreme right of the periodic table.

This statement is true. Elements in the extreme right of the periodic table are non-metals. Non-metal oxides are acidic and form acids in reaction with water. For example, sulfur trioxide (SO₃) reacts with water to form sulfuric acid (H₂SO₄).

Considering the above analysis:

Correct Option: 1) Statement-I is false but Statement-II is true.

Concept: 

Iso-electronic species - 

• The atoms or ions or molecules having the same no. of valence electrons are called the iso-electronic species.
• These species have some similarities in their properties due to the same electronic structure, but some properties like ionic or atomic radii may vary due to different nuclear charges.
• Example of iso-electronic species - O2- and F-, Mg2+and Na+, etc.

Explanation:

In the case of iso-electronic species, the variation of ionic radii is considered with respect to the nuclear charge present on the ion.

→ With an increase in nuclear charge or the number of protons in the nucleus, the value of ionic radii decreases.

Reason - In the case of isoelectronic species,

• The number of electrons remains the same but the value of nuclear charge changes with a change in the number of protons or atomic number.
• Thus, with the increase in the nuclear charge, the last electrons hold firmly to the nucleus and the value of ionic radii thus decreases with an increase in the nuclear charge.

Thus, for given isoelectronic species Na+, Mg2+, F– and O2– the order will be - Mg2+ < Na+ < F– < O2–  because nuclear charge varies in order O2– < F– < Na+ < Mg2+.

Conclusion: Therefore, the correct order of increasing the length of their radii is  Mg2+ < Na+ < F– < O2–

Hence, the correct answer is option 2.

Concept:

• Electron affinity is one of the periodic properties of the elements.
• Along the group from top to bottom electron affinity decreases.
• Along the group from left to right electron affinity increases.

Explanation:

• In the case of halogens, the regular trend in electron affinity is not observed between Cl and F.
• Cl-atom has more electron affinity than F-atom.
• The reason is Cl-atom's size is larger than F-atom.
• So, F-atom experiences more repulsions between the electrons.
• In Cl-atom the charge distribution is more compared to fluorine.
• So, when an extra electron is added to chlorine, then more amount of energy is released compared to fluorine and it attains more stability.

Thus, the order of electron affinity among halogens follows the order of Cl > F > Br > I.

CONCEPT:

Ionic Radii

• The ionic radius is the measure of an atom's ion in a crystal lattice. It is half the distance between two ions that are barely touching each other.
• The size of an ion is affected by its charge. As the number of electrons increases (for anions), the ion becomes larger. Conversely, as the number of electrons decreases (for cations), the ion becomes smaller.

EXPLANATION:

•  C4- :
  • With the lowest nuclear charge ( Z = 6 ), C4-  has the largest ionic radius in this series.
• N3- :
  • Nitrogen ( Z = 7 ) has a slightly higher nuclear charge than carbon, so N3- has a smaller ionic radius than C4-  .
• O2- :
  • Oxygen ( Z = 8 ) has a higher nuclear charge than nitrogen, so O2-  has a smaller ionic radius than N3- .
• F- :
  • Fluorine ( Z = 9 ) has the highest nuclear charge in this series, so F^- has the smallest ionic radius.

Increasing Order of Ionic Radii: F- < O2- < N3- < C4- .

The correct answer is:+2.2

Concept:-

• Atomic number (Z): This is the number of protons in the nucleus of an atom, which determines the chemical properties of an element and its place in the periodic table. Sodium (Na) has an atomic number of 11, meaning it has 11 protons.
• Shell and subshell distribution: Electrons in an atom are organized in shells and subshells around the nucleus. Sodium has a configuration of 1s² 2s² 2p⁶ 3s¹.
• Shielding effect and penetration: We use the term shielding to describe the phenomenon that occurs when electrons are obscured from the influence of the nuclear positive charge due to the presence of other electrons. Some orbitals have greater penetrating power towards the nucleus hence they experience a greater effective nuclear charge.

Explanation:-

Conclusion:-

So, The effective nuclear charge for Na (11) is  +2.2

CONCEPT:

Isoelectronic Species and Ionic Radii

• Isoelectronic Species: Isoelectronic species are atoms, ions, or molecules that have the same number of electrons. For example, O^2–, F^–, Na⁺, and Mg²⁺ all have 10 electrons.
• Nuclear Charge: The nuclear charge is the charge of the nucleus, which is equal to the number of protons in the nucleus.
• Ionic Radii:
  • The size of an ion generally decreases with increasing positive charge (cations are smaller than their parent atoms) and increases with increasing negative charge (anions are larger than their parent atoms).
  • For isoelectronic species, the size decreases with increasing nuclear charge because the electrons are drawn closer to the nucleus.

Explanation:-

• Given species: O^2–, F^–, Na⁺, and Mg²⁺
• Isoelectronic: All four species have 10 electrons, so they are isoelectronic. (A) is correct.
• Nuclear Charge:
  • O^2– has 8 protons.
  • F^– has 9 protons.
  • Na⁺ has 11 protons.
  • Mg²⁺ has 12 protons.
  Statement (B) is incorrect since they do not have the same nuclear charge.
• Ionic Radii:
  • O^2– has the largest radius because it has the smallest nuclear charge for the same number of electrons. (C) is correct.
  • Mg²⁺ has the smallest radius because it has the largest nuclear charge for the same number of electrons. (D) is correct.

CONCLUSION:

The correct answer is  (A), (C) and (D) only

The correct answer is 
Concept:-

• Nuclear Charge: The more protons in an atom's nucleus, the greater is the nuclear charge. This greater charge pulls the electrons closer to the nucleus, making the ion smaller.
• Electron Configuration: The arrangement of electrons in an atom also affects its size. If an atom loses electrons and is transformed into a cation, it will become smaller.
• Shielding Effect: This refers to the ability of inner-shell electrons to minimize the attractive force exerted by the nucleus on the outer-shell electrons.

Explanation:-

Ca²⁺ (Calcium Ion): when it loses two electrons to form a Ca²⁺ ion, there are 20 protons in the nucleus and only 18 electrons. The increased positive charge pulls the remaining electrons closer to the nucleus, effectively reducing the size of the ion.

Mg²⁺ (Magnesium Ion): When two electrons are lost forming Mg²⁺ ion, there are 12 protons but only 10 electrons. These 12 protons can exert a stronger pull on fewer electrons (only 10), causing the ion to be smaller than it was in its neutral state.

Zn²⁺ (Zinc Ion): When it loses two electrons to form a Zn²⁺ ion, there are now 30 protons but only 28 electrons. The imbalance in favor of the protons causes the ion to be smaller.

The strength of the positive charge due to protons in the nucleus directly impacts the extent of pull on the electrons.

Here, Zn²⁺ has the maximum number of protons (30), followed by Ca²⁺ (20), and then Mg²⁺ (12).

However, we must also consider the number of electrons each ion has. To this effect, Mg²⁺ has fewer electrons (10) than do Ca²⁺ (18) and Zn²⁺ (28), and thus its fewer electrons are drawn more closely to the nucleus, resulting in a smaller ionic radius.

So, in terms of the number of protons and electrons, Mg²⁺ (Magnesium Ion) gets shrunk to a greater extent than do Ca²⁺ and Zn²⁺.

Therefore, Mg²⁺ has the smallest ionic radius among the three given ions.

Concept:

• Isoelectronic species contain the same number of valence electrons.
• In the case of isoelectronic species, the anion with a high charge is larger in size. and the cation with a more positive charge has the smallest size.

Explanation:

The given ionic species are Na+, F−, Mg2+, and O2−.

• All these ions have the same number of valence electrons which is 10.
• They are isoelectronic species.

 O^2-

• The anion O^2- has the largest size.
• The reason for its larger size is due to the decrease in its effective nuclear charge.
• In its nucleus, there are 8 protons only.
• The number of protons is less than the number of electrons.

Mg²⁺

• The size of Mg²⁺ is smaller due to the increased effective nuclear charge between the nucleus and the valence electrons.
• There are 10 electrons and 12 protons in it.

Thus, the overall size order of the given isoelectronic species

Mg²⁺ < Na²⁺ < F^- < O^2-

The correct answer is Si < Be < Mg < Na

Concept:-

• Atomic Size: Larger atoms have less attraction between their nuclei (which possess a positive charge) and their valence electrons (which possess a negative charge).
• Effective Nuclear Charge: As you move across a period in the periodic table (from left to right), the number of protons in the nucleus (nuclear charge) increases, and the valence electrons become more tightly held due to increased attraction to the nucleus. This phenomenon, known as an effective nuclear charge, increases across a period and decreases down a group.

Explanation:-

• This lower attraction makes it easier for larger atoms to lose electrons, which is, in essence, the metallic property.
• As you move down a group in the periodic table, atomic size increases due to the addition of new energy levels or electron shells, leading to an increase in metallic character.
• The increase in effective nuclear charge moving across the period makes it more difficult for elements to lose electrons, leading to a decrease in metallic character.
• Now, let's arrange the given elements - Si (Silicon), Be (Beryllium), Mg (Magnesium), and Na (Sodium) - in order of increasing metallic character.
• These elements are located in the periodic table as follows:

Be (Group 2, Period 2)
Mg (Group 2, Period 3)
Si (Group 14, Period 3)
Na (Group 1, Period 3)

• Following the rule that metallic character decreases as one moves to the right across a period, Si would have the least metallic character among the given elements.

As one moves down the periodic table, the metallic character increases because the atomic size becomes larger, and the electron shielding effect overpowers the increased nuclear charge.

So, Be and Mg are both in Group 2, but Mg is below Be, thus Mg has a larger metallic character than Be.

Finally, among these elements, Na is the most metallic. It is in Group 1 in Period 3, from the left side of the periodic table, and it possesses one valance electron, which is easily lost, a characteristic of high metallic behavior.

Si < Be < Mg < Na

Conclusion:-

So, the order of the given elements in increasing order of metallic character is:

Si < Be < Mg < Na

Concept:

Electron affinity:

• The electron affinity of an element is defined as the ease with which it can accept an electron.
• Chlorine the group 17th element has the highest electron affinity also called halogen represented by X.
• Electron affinity increases going from left to right across a period.
• Electron affinity decreases going down the group.
• Electron Affinity= 1/ Atomic size

Explanation:

• Although Fluorine has the highest electronegativity, Chlorine has the highest electron affinity and this is because of the considerable repulsion in the tightly packed 2p subshell of Fluorine.
• As fluorine and chlorine are the elements of the halogens family. The tendency to accept one electron to attain the noble gas configuration. is more therefore the electron affinity is higher in the halogen family.

There the correct order of the electron affinity is N < O < F < Cl

Correct answers: 1 and 3)

Concept:

Elements of group 13 of the periodic table: B (non-metal), Al(Metal), Ga(Metal), In(Metal), Th(Metal).

Substances that have free electrons are good conductors of electricity.

Explanation:

• The general electronic configuration of group 13 elements is ns² np¹ since they have 3 electrons in the valence shell.
• The element belonging to the 3rd period and group-13 of the periodic table is aluminum which is solid metal and a good conductor of electricity. Aluminum has strong conductivity for electricity and heat because it has three electrons in the free state.
• Boron has a high melting point. This is because of the icosahedral structure. In the boron family, gallium has the lowest melting point.
• The elements of group 13 have higher densities than group 2 elements. This is because they have smaller sizes, and hence small volumes. The density increases from B to Tl.
• Both Al and Gallium are amphoteric in nature. It means they react with both acids and bases.

Conclusion:

Therefore, Al belongs to the 3rd period and group-13 of the periodic table and it is a solid metal as well as a good conductor of electricity.

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