Problem 68 - A Complete Resource Book in Chemistry for JEE Main

Question

The second ionization energies of the \(\mathrm{C}, \mathrm{N}, \mathrm{O}\) and \(\mathrm{F}\) atoms are such that (a) \(\mathrm{O}>\mathrm{N}>\mathrm{F}>\mathrm{C}\) (b) \(\mathrm{F}>\mathrm{O}>\mathrm{N}>\mathrm{C}\) (c) \(\mathrm{O}>\mathrm{O}>\mathrm{N}>\mathrm{F}\) (d) \(\mathrm{O}>\mathrm{F}>\mathrm{N}>\mathrm{C}\)

Step-by-Step Solution

Answer

(d) O > F > N > C.

Step 1: Understand Second Ionization Energy

The second ionization energy is the energy required to remove an electron from a positively charged ion after one electron has already been removed. It generally increases across a period due to increasing nuclear charge and decreases down a group as effective nuclear charge felt by valence electrons decreases.

Step 2: Analyze Periodic Trends

In a period, the second ionization energy increases from left to right. Thus, for elements in the same period, one would expect the order of second ionization energy to generally be F > O > N > C.

Step 3: Consider Electron Configuration

The electron configuration plays a significant role, especially for O and N. After removing the first electron, O forms a half-filled stable configuration, making it harder (more energy) to remove the next electron compared to N.

Step 4: Order of Second Ionization Energies

Using the observations and the periodic trends, we can establish the order of second ionization energies as O > F > N > C because O reaches a more stable configuration before F, and N's configuration is relatively stable but not as strong as O or F after ionization.

Key Concepts

Periodic Trends

Periodic trends refer to patterns observed in the properties of elements as you move across a period (row) or down a group (column) in the periodic table. These trends arise because of the predictable change in atomic structure, such as the increase in proton number and electron shells.

Across a period from left to right, elements have a higher nuclear charge due to an increased number of protons. This increased charge causes the electrons in the outer shell to be pulled closer to the nucleus. As a result, various properties such as electronegativity and ionization energy typically increase.

• Ionization energy is higher because more energy is needed to remove an electron that's more tightly bound to the nucleus.
• This trend is part of why elements like fluorine (F) have high second ionization energies compared to carbon (C) despite being in the same period.

Conversely, going down a group, elements have more energy levels, which increases the atomic radius and typically reduces ionization energies because the outer electrons are further from the nucleus and more shielded by inner electrons.

Electron Configuration

Electron configuration refers to the distribution of electrons in an atom's orbitals. Each element has a unique electron configuration, which influences chemical behavior and reactivity.

When considering ionization energies, it's important to consider how electrons fill these orbitals. Electrons fill orbitals from lower to higher energy levels while obeying the Pauli exclusion principle and Hund's rule.

• The stability of an electron configuration affects how easily an atom loses electrons. Half-filled or fully filled subshells are more stable configurations.
• For example, when an oxygen atom loses one electron, it achieves a half-filled p subshell, which is particularly stable.

Thus, even though fluorine may have a higher initial ionization energy than nitrogen, oxygen's more stable half-filled configuration after losing an electron makes it harder to further ionize beyond the first electron removal compared to nitrogen.

Ionization Energy

Ionization energy is the energy necessary to remove an electron from an atom or ion. This energy indicates how strongly an atom holds onto its electrons and generally reflects the reactivity of the element.

The second ionization energy refers specifically to removing an electron from an already positively charged ion (one that has lost an initial electron). This energy is often higher than the first because:

• The remaining electrons are closer to the nucleus and have a stronger interaction with the positive charge.
• Elements that achieve a stable electron configuration after losing the first electron will require significantly more energy to lose the second electron.

In relation to the periodic trends discussed earlier, second ionization energies typically increase as you move from left to right across a period due to the higher effective nuclear charge felt by the remaining electrons. Meanwhile, elements like oxygen which reach a more stable configuration quickly will have higher second ionization energies compared to others like nitrogen and carbon, which do not reach such stable configurations as early.