Question

Which of the following orders is incorrect? (a) \(\mathrm{NH}_{3}<\mathrm{PH}_{3}<\mathrm{AsH}_{3} \rightarrow\) acidic nature. (b) \(\mathrm{Li}<\mathrm{Be}<\mathrm{B}<\mathrm{C} \rightarrow\) first ionization energy (c) \(\mathrm{Al}_{2} \mathrm{O}_{3}>\mathrm{MgO}<\mathrm{Na}_{2} \mathrm{O}<\mathrm{K}_{2} \mathrm{O} \rightarrow\) basic nature. (d) \(\mathrm{Li}^{+}<\mathrm{Na}^{+}<\mathrm{K}^{+}<\mathrm{Cs}^{+} \rightarrow\) ionic radius

Short answer

Option (c) is incorrect in the order given.

Step-by-step solution

Understand the Concepts

First, let's understand the properties involved: acidic nature, ionization energy, basic nature, and ionic radius. Acidic nature generally decreases as we move down a group in the periodic table, ionization energy usually increases across a period from left to right, basic nature typically increases as we move down a group, and ionic radius increases down a group and across periods as the atomic number increases.

Evaluate Option (a)

For option (a), which is about acidic nature, moving down the group from \(\mathrm{NH}_3\) to \(\mathrm{AsH}_3\), the acidic nature generally decreases. Hence, the order \(\mathrm{NH}_3 < \mathrm{PH}_3 < \mathrm{AsH}_3\) is correct for the increase in acidic nature.

Evaluate Option (b)

For option (b), first ionization energy should increase across a period. Therefore, for \(\mathrm{Li}, \mathrm{Be}, \mathrm{B}, \mathrm{C}\) in Period 2 of the periodic table, the ionization energy generally increases from \(\mathrm{Li}\) to \(\mathrm{C}\). Hence, the order \(\mathrm{Li} < \mathrm{Be} < \mathrm{B} < \mathrm{C}\) is correct.

Evaluate Option (c)

For option (c), as we move down the group, the basic nature of oxides increases. Hence, for the increasing basic nature, the order should typically be \(\mathrm{Na}_2\mathrm{O} > \mathrm{MgO} \). Since this is not the case here, the statement appears incorrect by given trend compared to standard principles.

Evaluate Option (d)

For option (d), the ionic radius increases down the group. Therefore, for \(\mathrm{Li}^+, \mathrm{Na}^+, \mathrm{K}^+, \mathrm{Cs}^+\), the ionic radius should increase from \(\mathrm{Li}^+\) to \(\mathrm{Cs}^+\). The given order \(\mathrm{Li}^+ < \mathrm{Na}^+ < \mathrm{K}^+ < \mathrm{Cs}^+\) is correct.

Key concepts

Acidic Nature

Acidic nature is an important concept in chemistry, which refers to the ability of a compound to donate protons (H+) in a reaction. In the periodic table, as we move down a group, elements tend to show a decrease in acidic strength. This is because larger atoms have a weaker hold on their valence electrons, making it less likely for them to donate a proton.
For instance, when considering hydrides like \(\mathrm{NH}_3\), \(\mathrm{PH}_3\), and \(\mathrm{AsH}_3\)\, nitrogen at the top of the group forms compounds that are more acidic compared to arsenic at the bottom. This is why the general order of decreasing acidic nature down the group is as follows: \(\mathrm{NH}_3 \) \(< \mathrm{PH}_3 < \mathrm{AsH}_3 \)\.
Understanding these trends is crucial because they help predict the behavior of compounds in various chemical reactions, especially those involving acid-base equilibria.

Ionization Energy

Ionization energy is the energy required to remove an electron from a neutral atom in the gaseous state. In the periodic table, ionization energy generally increases across a period from left to right due to the increasing nuclear charge. Electrons are held more tightly as a result of the stronger attraction between the nucleus and the electrons.
For Period 2 elements like \(\mathrm{Li}, \mathrm{Be}, \mathrm{B}, \mathrm{C}\), we observe that ionization energy increases from lithium \(\mathrm{Li}\) to carbon \(\mathrm{C}\). This is due to the increasing positive charge in the nucleus, which pulls the electrons closer, requiring more energy to remove one.
Therefore, the order \(\mathrm{Li} < \mathrm{Be} < \mathrm{B} < \mathrm{C}\) is correct in demonstrating the trend of increasing ionization energy across the period. Recognizing these patterns helps predict how elements will react with each other.

Basic Nature

In contrast to acidic nature, basic nature refers to the tendency of a compound to accept protons. In the periodic table, the basic nature of oxides often increases down a group. This is primarily because larger cations can stabilize the negative charge formed by hydroxide ions, which is a typical measure of basicity.
For alkaline earth metals (Group 2), oxides like \(\mathrm{Na}_2\mathrm{O}, \mathrm{MgO}\), and \(\mathrm{K}_2\mathrm{O}\) demonstrate increasing basicity from magnesium to potassium. The order given in the exercise indicates a deviation from this trend, pointing out a misunderstanding in the usual progression.
Recognizing the actual order, such as \(\mathrm{Na}_2\mathrm{O} > \mathrm{MgO}\), reinforces understanding of how elements behave chemically, especially when forming basic compounds.

Ionic Radius

The ionic radius is a measure of an ion's size and is influenced by the atom's electron cloud's size. As you move down a group in the periodic table, the ionic radius generally increases. This happens because each subsequent element has an additional electron shell, making the ion larger.
Consider alkali metals like \(\mathrm{Li}^+, \mathrm{Na}^+, \mathrm{K}^+, \mathrm{Cs}^+\). These ions show a typical increase in ionic size from lithium to cesium. Thus, the sequence \(\mathrm{Li}^+ < \mathrm{Na}^+ < \mathrm{K}^+ < \mathrm{Cs}^+\) is accurate.
Understanding ionic radii trends is essential, as it affects properties like lattice energy and solubility in water. This knowledge aids in predicting the physical and chemical properties of elements and their compounds.