Classification of Elements and Periodicity in Properties Set-2

Test your knowledge on Classification of Elements and Periodicity in Properties from Chemistry, Class 11.

Minutes

Questions

1 / -0

Marking Scheme

Don't have an account? Sign up for free to save your progress and track your history.

Questions in this Quiz

Q1: How is the covalent radius of a non-metallic element like Chlorine calculated?

Half the distance between two atoms when they are bound by a single covalent bond
The full internuclear distance between the atoms
Half the internuclear distance in a metallic crystal
Half the van der Waals radius

Q2: The metallic radius of Copper ($\text{Cu}$) is assigned a value of $128\text{ pm}$. This is based on the internuclear distance of $256\text{ pm}$ between adjacent $\text{Cu}$ atoms in the solid crystal, where the metallic radius is:

Twice the internuclear distance
Half the internuclear distance
Equal to the internuclear distance
$256\text{ pm}$

Q3: Why does atomic size generally decrease across a period (left to right)?

Effective nuclear charge decreases.
The principal quantum number ($n$) increases.
Effective nuclear charge increases, increasing attraction of electrons to the nucleus within the same valence shell
The number of core electrons increases drastically.

Q4: Why does the atomic radius increase as we descend a group?

The nuclear charge decreases down the group.
The valence electrons are placed farther from the nucleus due to increased principal quantum number ($n$) and shielding
The electron gain enthalpy becomes more negative.
All inner energy levels are completely filled.

Q5: When comparing the size of an anion ($\text{X}^-$) to its parent atom ($\text{X}$), the anion is:

Smaller, due to increased effective nuclear charge
Larger, because the added electron increases repulsion and decreases effective nuclear charge
Smaller, because it has more protons
The same, as the nuclear charge is constant

Q6: An atom ($\text{X}$) forms a cation ($\text{X}^+$). Which statement correctly describes the relationship between their sizes?

The cation is larger because the remaining electrons are less attracted by the nucleus.
The cation is smaller because it has fewer electrons while the nuclear charge remains the same
The cation is the same size because only valence electrons are involved.
The cation's radius is half the atomic radius.

Q7: Consider the isoelectronic species $\text{O}^{2-}, \text{F}^{-}, \text{Na}^{+}, \text{Mg}^{2+}$. Which ion has the largest radius?

$\text{Mg}^{2+}$
$\text{Na}^{+}$
$\text{F}^{-}$
$\text{O}^{2-}$ (Anion with greater negative charge has larger radius due to net electron repulsion)

Q8: Which isoelectronic species will have the smallest radius among $\text{N}^{3-}, \text{O}^{2-}, \text{F}^{-}$?

$\text{N}^{3-}$
$\text{O}^{2-}$
$\text{F}^{-}$
All are the same size

Q9: Arrange the following isoelectronic species in the order of increasing ionic radii: $\text{Mg}^{2+}, \text{Na}^{+}, \text{F}^{-}$.

$\text{Mg}^{2+} < \text{Na}^{+} < \text{F}^{-}$
$\text{F}^{-} < \text{Na}^{+} < \text{Mg}^{2+}$
$\text{Na}^{+} < \text{Mg}^{2+} < \text{F}^{-}$
$\text{Mg}^{2+} < \text{F}^{-} < \text{Na}^{+}$

Q10: Why are the atomic radii of noble gases not typically compared with the covalent radii of other elements?

They are monoatomic, and their non-bonded (van der Waals) radii values are typically very large
They are highly reactive.
Their atomic numbers are too high.
They are located in the $p$-block.

Q11: In the sequence $\text{Mg}, \text{Mg}^{2+}, \text{Al}, \text{Al}^{3+}$, identify the species with the largest size.

$\text{Mg}$ (Largest parent atom)
$\text{Mg}^{2+}$
$\text{Al}$
$\text{Al}^{3+}$

Q12: The ionic radius of Fluoride ion ($\text{F}^{-}$) is $136\text{ pm}$, while the atomic radius of Fluorine ($\text{F}$) is $64\text{ pm}$. This drastic increase in size is primarily due to:

Decrease in nuclear charge
Increased electron-electron repulsion and decrease in effective nuclear charge
The loss of the valence shell
The electron entering the $n=3$ quantum level

Q13: Which of the following elements, $\text{Li}, \text{Na}, \text{K}, \text{Rb}, \text{Cs}$, has the largest atomic radius?

$\text{Rb}$
$\text{K}$
$\text{Cs}$
$\text{Li}$

Q14: Considering elements $\text{Na, Mg, Al, Si, P, S, Cl}$ (Period 3), which element has the smallest atomic radius?

$\text{Na}$
$\text{Si}$
$\text{Cl}$
$\text{Mg}$

Q15: The species $\text{Ar}, \text{K}^{+}, \text{Ca}^{2+}$ are isoelectronic. The correct order of their radii (from smallest to largest) is:

$\text{Ar} < \text{K}^{+} < \text{Ca}^{2+}$
$\text{Ca}^{2+} < \text{K}^{+} < \text{Ar}$
$\text{K}^{+} < \text{Ar} < \text{Ca}^{2+}$
$\text{Ar} < \text{Ca}^{2+} < \text{K}^{+}$

Q16: Ionization Enthalpy ($\Delta_i H$) is the energy required to remove an electron from an atom in which specific state?

Isolated gaseous atom in its ground state
Liquid atom
Solid crystal lattice
Gaseous ion in an excited state

Q17: The second ionization enthalpy ($\Delta_i H_2$) is always higher than the first ionization enthalpy ($\Delta_i H_1$) because:

It is easier to remove an electron from a positively charged ion than from a neutral atom.
Energy is always required to remove electrons from an atom.
It is more difficult to remove an electron from a positively charged ion than from a neutral atom
Effective nuclear charge decreases after the first removal.

Q18: On a plot of first ionization enthalpies ($\Delta_i H$) versus atomic number ($Z$), where do the maxima (highest values) occur?

At the noble gases
At the alkali metals
At the halogens
At the alkaline earth metals

Q19: Why does ionization enthalpy generally increase across a period?

Increase in shielding outweighs the increase in nuclear charge.
Successive electrons are added to orbitals in the same principal quantum level, and increasing nuclear charge outweighs shielding
The valence electrons are increasingly farther from the nucleus.
The principal quantum number ($n$) decreases.

Q20: The first ionization enthalpy of Boron ($\text{Z}=5$) is slightly less than that of Beryllium ($\text{Z}=4$) because:

The $2s$ electron in $\text{Be}$ is more shielded than the $2p$ electron in $\text{B}$.
The $2p$ electron in $\text{B}$ is more shielded by inner core electrons than the $2s$ electron in $\text{Be}$, making it easier to remove
$\text{Be}$ has a higher atomic number.
$\text{B}$ has a completely filled $2s$ orbital.

Q21: The "anomaly" of Oxygen ($\text{O}$) having a smaller first ionization enthalpy than Nitrogen ($\text{N}$) arises because:

Oxygen has a lower nuclear charge.
Oxygen has increased electron-electron repulsion due to paired electrons in a $2p$-orbital, making removal easier
Nitrogen has a completely filled $2p$ subshell.
Oxygen is more electronegative than Nitrogen.

Q22: The first ionization enthalpy values for $\text{Na}, \text{Mg}$, and $\text{Si}$ are $496, 737$, and $786\text{ kJ mol}^{-1}$ respectively. Based on shielding and orbital configuration, the first $\Delta_i H$ value for $\text{Al}$ should be closest to: (Hard Level)

$760\text{ kJ mol}^{-1}$
$575\text{ kJ mol}^{-1}$ (Lower than $\text{Mg}$ due to effective shielding of $3p$ electron by $3s$ electrons)
$490\text{ kJ mol}^{-1}$
$737\text{ kJ mol}^{-1}$

Q23: For an element with the electronic configuration $\text{1s}^2 \text{2s}^2 \text{2p}^6 \text{3s}^1$, which ionization enthalpy will show a significantly large jump in value?

$\Delta_i H_1$
$\Delta_i H_2$ (Removal breaks into the stable noble gas core, $[\text{Ne}]$)
$\Delta_i H_3$
$\Delta_i H_4$

Q24: The valence electron in Lithium ($\text{Li}$) is shielded from the nucleus by the inner core of $1s$ electrons. The reduction in the effective nuclear charge experienced by this valence electron is known as:

Electronegativity
Ionization Potential
Shielding or Screening
Penetration Effect

Q25: The ionization enthalpies generally decrease down a group primarily because:

Effective nuclear charge increases significantly.
The outermost electron is increasingly farther from the nucleus, and increased shielding outweighs the increasing nuclear charge
The atomic radius decreases.
The number of valence electrons increases.

Q26: An element $\text{I}$ has $\Delta_i H_1 = 520\text{ kJ mol}^{-1}$ and a very high $\Delta_i H_2 = 7300\text{ kJ mol}^{-1}$. Which statement is true about element $\text{I}$?

It is the least reactive element.
It is an Alkali Metal (Group 1)
It forms a stable binary halide $\text{MX}_2$.
It is a Noble Gas.

Q27: An element $\text{VI}$ has $\Delta_i H_1 = 738\text{ kJ mol}^{-1}$ and $\Delta_i H_2 = 1451\text{ kJ mol}^{-1}$. This small difference between $\Delta_i H_1$ and $\Delta_i H_2$ suggests that element $\text{VI}$ is likely to be:

A noble gas
An alkali metal
An alkaline earth metal (forms stable $\text{M}^{2+}$ ion)
A halogen

Q28: What is the correct order of the first ionization enthalpy for the second period elements $\text{Li, Be, B}$?

$\text{Li} < \text{Be} < \text{B}$
$\text{Li} < \text{B} < \text{Be}$
$\text{B} < \text{Be} < \text{Li}$
$\text{Be} < \text{B} < \text{Li}$

Q29: When considering the elements $\text{B, Al, Mg,}$ and $\text{K}$, what is the correct order of their metallic character (from least metallic to most metallic)?

$\text{B} < \text{Al} < \text{Mg} < \text{K}$
$\text{K} < \text{Mg} < \text{Al} < \text{B}$
$\text{B} < \text{Mg} < \text{Al} < \text{K}$
$\text{Al} < \text{B} < \text{Mg} < \text{K}$

Q30: Which statement regarding ionization enthalpy is incorrect?

The removal of an electron from orbitals bearing lower $n$ value is easier than from an orbital having higher $n$ value
Ionization enthalpy increases for each successive electron removed.
A big jump in ionization enthalpy marks the end of valence electrons.
The greatest increase in ionization enthalpy occurs upon removal of an electron from the core noble gas configuration.