Classification of Elements and Periodicity in Properties Set-3

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

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Questions in this Quiz

Q1: Electron Gain Enthalpy ($\Delta_{eg} H$) is defined as the enthalpy change accompanying the process where a neutral gaseous atom ($\text{X}$) forms:

$\text{X}^{+}(\text{g}) + e^{-}$
$\text{X}^{-}(\text{g})$
$\text{X}^{2+}(\text{g}) + 2e^{-}$
$\text{X}(\text{l})$

Q2: Group 17 elements (Halogens) have highly negative electron gain enthalpies because:

They are located on the left side of the periodic table.
They can attain a stable noble gas electronic configuration by picking up one electron
They have high ionization enthalpies.
They are large atoms.

Q3: Why do Noble Gases have large **positive** electron gain enthalpies?

The added electron must enter the next higher principal quantum level ($n+1$), leading to a very unstable electronic configuration
They are monatomic.
They have small atomic radii.
Their ionization enthalpy is very low.

Q4: Among $\text{P, S, Cl, F}$, which element will have the **most** negative electron gain enthalpy ($\Delta_{eg} H$)?

$\text{F}$
$\text{P}$
$\text{S}$
$\text{Cl}$ (Due to the larger $3p$ orbital accommodating the electron with less repulsion than $2p$ of $\text{F}$)

Q5: Why is the electron gain enthalpy of Oxygen ($\text{O}$) less negative than that of Sulfur ($\text{S}$)?

$\text{S}$ has a much smaller atomic radius than $\text{O}$.
The added electron in $\text{O}$ goes to the small $n=2$ quantum level and suffers significant electron-electron repulsion
$\text{O}$ is located lower in the group than $\text{S}$.
$\text{O}$ has a half-filled $p$-orbital.

Q6: What is the fundamental difference between Electron Gain Enthalpy and Electronegativity?

Electronegativity is a measureable quantity, while $\Delta_{eg} H$ is not.
$\Delta_{eg} H$ measures energy change for an isolated gaseous atom forming an anion, while electronegativity is a measure of an atom's ability to attract **shared** electrons in a compound
Both measure the tendency to gain electrons, but $\Delta_{eg} H$ is qualitative.
$\Delta_{eg} H$ is expressed on the Pauling scale, while electronegativity is in $\text{kJ mol}^{-1}$.

Q7: Electronegativity generally increases across a period because:

The atomic radius decreases and the attraction between the valence electrons and the nucleus increases
The size of the atom increases.
Metallic properties increase.
Ionization enthalpy decreases.

Q8: Which scale, used to quantify electronegativity, arbitrarily assigned a value of $4.0$ to Fluorine?

Mulliken-Jaffe scale
Allred-Rochow scale
Pauling scale
Dobereiner scale

Q9: Which of the following sets of elements exhibits the strongest tendency to form an anion (i.e., highest non-metallic character)?

$\text{Na, Mg, Al}$
$\text{Si, P, S}$
$\text{C, N, O}$
$\text{F, Cl, Br}$ (Group 17, Halogens)

Q10: The second electron gain enthalpy of Oxygen ($\text{O}^{-} \to \text{O}^{2-}$):

Must be highly negative.
Must be positive (endothermic), requiring energy to overcome repulsion between the incoming electron and the existing negative charge
Must be zero.
Is approximately equal to the first electron gain enthalpy.

Q11: Considering the elements $\text{F, Cl, O, and N}$, the correct order of their non-metallic character is:

$\text{Cl} > \text{F} > \text{O} > \text{N}$
$\text{F} > \text{N} > \text{O} > \text{Cl}$
$\text{F} > \text{O} > \text{N} > \text{Si}$
$\text{F} > \text{O} > \text{N} > \text{Cl}$ (Decreases down group, increases across period)

Q12: The least reactive element among $\text{I, II, III, IV, V, VI}$ (from Problem 3.31), which has a large positive $\Delta_{eg} H$ and very high $\Delta_i H_1$, is likely to be:

Element $\text{I}$
Element $\text{III}$
Element $\text{V}$ (Noble Gas: $\Delta_i H_1 = 2372, \Delta_{eg} H = +48$)
Element $\text{IV}$

Q13: The most reactive non-metal among $\text{I, II, III, IV, V, VI}$ (from Problem 3.31), characterized by high $\Delta_i H_1$ and highly negative $\Delta_{eg} H$, is likely to be:

Element $\text{III}$ (Halogen: $\Delta_i H_1 = 1681, \Delta_{eg} H = -328$)
Element $\text{IV}$
Element $\text{VI}$
Element $\text{I}$

Q14: Based on the values in Table 3.7, which element pair shows the greatest deviation from the expected trend that $\Delta_{eg} H$ becomes less negative down a group?

$\text{H}$ and $\text{Li}$
$\text{O}$ and $\text{S}$
$\text{Cl}$ and $\text{Br}$
$\text{Na}$ and $\text{K}$

Q15: The valence of representative elements is usually equal to the number of electrons in the outermost orbitals, or:

Eight minus the number of outermost electrons
Half the atomic number
The principal quantum number ($n$)
The group number divided by two

Q16: An element $\text{M}$ from Group 2 (Alkaline Earth Metal) combines with an element $\text{N}$ from Group 15 (Valence 3). Predict the formula of the stable binary compound formed.

$\text{MN}$
$\text{MN}_2$
$\text{M}_3\text{N}_2$ (Valences $2$ and $3$)
$\text{M}_2\text{N}_3$

Q17: The oxide formed by the element on the extreme right of a period, such as $\text{Cl}_2\text{O}_7$, is typically:

Neutral
Amphoteric
Most acidic
Most basic

Q18: The anomalous behavior of the first member of groups 1, 2, and 13-17 (e.g., $\text{Li}, \text{Be}, \text{B}$) is primarily attributed to their:

Large size and low electronegativity
Small size, large charge/radius ratio, and high electronegativity
High number of available valence orbitals
Low ionization enthalpy

Q19: What is the maximum covalency typically exhibited by the first member of a $p$-block group (e.g., Boron), compared to subsequent members like Aluminium? (Hard Level)

4 (due to only four valence orbitals: $2s$ and $2p$)
6 (due to availability of $d$-orbitals)
8 (due to expanded octet)
5 (due to ability to form $p\pi-p\pi$ bonds)

Q20: An atom has the electronic configuration $\text{1s}^2 \text{2s}^2 \text{2p}^6 \text{3s}^2 \text{3p}^6 \text{3d}^3 \text{4s}^2$. It will be placed in which group?

$\text{Fifth}$ (5th)
$\text{Third}$ (3rd)
$\text{Fourth}$ (4th)
$\text{Sixth}$ (6th)

Q21: The correct order of first ionization energy enthalpy ($\text{IE}_1$) for the given four elements ($\text{C, O, N, F}$) is:

$\text{C} < \text{N} < \text{O} < \text{F}$
$\text{F} < \text{N} < \text{O} < \text{C}$
$\text{C} < \text{O} < \text{F} < \text{N}$
$\text{C} < \text{O} < \text{N} < \text{F}$

Q22: The correct order of atomic radii in Group $13$ element ($\text{B, Al, Ga, In, Tl}$) is:

$\text{B} < \text{Al} < \text{Ga} < \text{In} < \text{Tl}$
$\text{Tl} > \text{In} > \text{Al} > \text{Ga} > \text{B}$
$\text{B} < \text{Ga} < \text{Al} < \text{In} < \text{Tl}$
$\text{B} < \text{Al} < \text{In} < \text{Ga} < \text{Tl}$

Q23: The species $\text{Ar, K}^+, \text{Ca}^{2+}$ contain the same number of electrons. In which order do their radii increase?

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

Q24: Which of the following order of ionic radius is correctly represented?

$\text{H}^+ > \text{H} > \text{H}^-$
$\text{Na}^+ < \text{F}^- < \text{O}^{2-}$
$\text{O}^{2-} < \text{F}^- < \text{Na}^+$
$\text{None of the above}$

Q25: Which one of the following arrangements represents the correct order of least negative to most negative electron gain enthalpy ($\Delta_{\text{eg}} \text{H}$) of $\text{C, Ca, Al, F, and O}$?

$\text{Ca} < \text{Al} < \text{C} < \text{O} < \text{F}$
$\text{F} > \text{O} > \text{C} > \text{Al} > \text{Ca}$
$\text{Ca} < \text{C} < \text{O} < \text{Al} < \text{F}$
$\text{Al} < \text{Ca} < \text{C} < \text{O} < \text{F}$

Q26: Identify the wrong statement in the following:

Among isoelectronic species, **smaller the positive charge** on the cation, **smaller** is the ionic radius.
Among isoelectronic species, greater the negative charge on the anion, greater is the atomic/ionic radius.
Atomic radius of the element increases as one moves down the first group of the periodic table.
Atomic radius of the element decreases as one moves across from left to right in the second period of the periodic table.

Q27: Among the element $\text{Ca, Mg, P, Cl}$, the order of increasing atomic radii is:

$\text{Cl} < \text{P} < \text{Ca} < \text{Mg}$
$\text{Ca} < \text{Mg} < \text{P} < \text{Cl}$
$\text{Cl} < \text{P} < \text{Mg} < \text{Ca}$
$\text{Mg} < \text{Ca} < \text{Cl} < \text{P}$

Q28: Identify the correct order of increasing ionic size for the following isoelectronic species: $\text{Ca}^{2+}, \text{K}^+, \text{Ar}, \text{Cl}^-, \text{S}^{2-}$.

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

Q29: Which electronic configuration indicates that there will be an abnormally high difference in the second ($\text{IE}_2$) and third ($\text{IE}_3$) ionization enthalpy of the element?

$\text{1s}^2 \text{2s}^2 \text{2p}^6 \text{3s}^2$
$\text{1s}^2 \text{2s}^2 \text{2p}^6 \text{3s}^1$
$\text{1s}^2 \text{2s}^2 \text{2p}^6 \text{3s}^2 \text{3p}^1$
$\text{1s}^2 \text{2s}^2 \text{2p}^6 \text{3s}^2 \text{3p}^2$

Q30: The first ionization enthalpy (in $\text{eV}$) for Beryllium ($\text{Be}$) and Boron ($\text{B}$) respectively, are:

$8.29, 9.32$
$9.32, 9.32$
$10.00, 10.00$
$9.32, 8.29$

Q31: Which one of the following ions will be the smallest in size?

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