Question

Which of the following has maximum ionization energy? (a) \(\mathrm{Ca} \rightarrow \mathrm{Ca}^{2+}+2 \mathrm{e}^{-}\) (b) \(\mathrm{Mg} \rightarrow \mathrm{Mg}^{2+}+2 \mathrm{e}^{-}\) (c) \(\mathrm{Ba} \rightarrow \mathrm{Ba}^{+}+\mathrm{e}^{-}\) (d) \(\mathrm{Be} \rightarrow \mathrm{Be}^{+}+\mathrm{e}^{-}\)

Short answer

(d) Be has the highest ionization energy.

Step-by-step solution

Understand Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It is generally measured in kilojoules per mole (kJ/mol). The higher the ionization energy, the more difficult it is to remove an electron.

Evaluate Trends in the Periodic Table

Ionization energy increases across a period from left to right and decreases down a group. Thus, elements in the top-right of the periodic table generally have higher ionization energies compared to those in the bottom-left.

Examine Each Element

(a) Ca is in group 2 and period 4, meaning its ionization energy is relatively lower. (b) Mg is in group 2 and period 3. Being above Ca in the periodic table, Mg has a higher ionization energy. (c) Ba is in group 2 and period 6, lower than both Ca and Mg in terms of ionization energy because it is further down the group. (d) Be is in group 2 and period 2. Since it's at the top of the group, it has the highest ionization energy among these options.

Compare Ionization Processes

Evaluate which ionization process requires the highest energy. Be goes from a neutral atom to a singly charged ion (removal of one electron), while Mg and Ca go through multiple removals to become doubly charged. The first ionization involves less energy than subsequent ionizations. Since Be+ forms from Be with the least electron shielding and greatest nuclear attraction, it typically requires higher energy.

Find the Answer

Based on the periodic trend and the given reactions, (B) for Mg and (A) for Ca involve removal of two electrons, while (D) for Be involves removal of only one electron, with high nuclear attraction and the least electron shielding.

Key concepts

Trends in the Periodic Table

The periodic table displays numerous trends in element properties, one of the most significant being ionization energy. As you move from left to right across a period, the ionization energy tends to increase. This pattern occurs because the added electrons are drawn closer to the nucleus within the same energy level, strengthening the nuclear attraction. Moreover, as you go down a group, ionization energy typically decreases. This happens because the outer electrons are further from the nucleus and experience greater shielding from inner electron shells, making it easier to remove an electron.

Understanding these trends is crucial when comparing ionization energies among elements, as position in the periodic table heavily influences the energy required to remove an electron.

Electron Removal

Electron removal, or ionization, involves taking an electron away from an atom or ion, usually in the gaseous state. Removing an electron requires overcoming the attraction between the negatively charged electron and the positively charged nucleus. The energy required for this process is termed ionization energy.

During ionization, different elements require different amounts of energy. Factors influencing this include nuclear charge, electron distance from the nucleus, and electron shielding. For instance, elements with tightly held electrons, closer to the nucleus, will have higher ionization energies, making them less prone to losing electrons.

Nuclear Charge Influence

The influence of nuclear charge is a key factor in determining ionization energy. The nuclear charge refers to the number of protons in the nucleus of an atom. A higher nuclear charge means a stronger attraction force exerted on the electrons. As a result, more energy is needed to overcome this force, leading to a higher ionization energy.

As you move across a period in the periodic table, the nuclear charge increases without adding more shielding, which leads to a progressively tighter hold on the electrons, thus increasing ionization energy. Conversely, within a group, although the nuclear charge is higher, added electron shells lead to increased shielding, lowering the ionization energy compared to elements higher up in the group.

Electron Shielding

Electron shielding, also known as electron screening, occurs when inner-shell electrons reduce the effective nuclear charge felt by outer-shell electrons. This effect is notable as you move down a group in the periodic table.

With each added electron shell, the outer electrons are less strongly attracted by the nucleus due to increased interference by intervening electron shells. This added distance and shielding make it easier to remove an electron, thereby lowering the ionization energy.

In the context of ionization energy comparisons, remember that with less electron shielding, as seen in higher periods of the periodic table, more energy is required to remove the electrons.

Group and Period Comparison

Analyzing elements through group and period comparisons helps understand trends in ionization energies. Let's consider elements from the same group: Be, Mg, Ca, and Ba, which all belong to group 2 of the periodic table.

When compared across a period, group 2 elements higher in a period (like Be) usually have higher ionization energies than those lower down (like Ba). Be exhibits high ionization energy, attributed to its relatively low atomic radius and effective nuclear charge being greater than that of elements further down the group.

Understanding these periodic trends and group dynamics allows us to predict and explain differences in ionization energy effectively, guiding us in chemical reactivity predictions.