Question

When atoms lose more than one electron, the ionization energy to remove the second electron is always more than the ionization energy to remove the first. Similarly, the ionization energy to remove the third electron is more than the second and so on. However, the increase in ionization energy upon the removal of subsequent electrons is not necessarily uniform. For example, consider the first three ionization energies of magnesium: \(\begin{array}{ll}\text { First ionization energy } & 738 \mathrm{~kJ} / \mathrm{mol} \\ \text { Second ionization energy } & 1450 \mathrm{~kJ} / \mathrm{mol} \\ \text { Third ionization energy } & 7730 \mathrm{~kJ} / \mathrm{mol}\end{array}\) The second ionization energy is roughly twice the first ionization energy, but then the third ionization energy is over five times the second. Use the electron configuration of magnesium to explain why this is so. Would you expect the same behavior in sodium? Why or why not?

Short answer

The jump in ionization energies is due to electrons being removed from increasingly stable and lower energy levels. Magnesium's first and second electrons are in the 3s subshell, while the third electron removed is from the 2p subshell. The same behavior in sodium wouldn't be observed for the second ionization energy because only one electron is in the outermost shell, and the next one to be removed would be from a complete 2p subshell, similar to the jump in magnesium from the second to third ionization energies.

Step-by-step solution

Understanding Magnesium's Electron Configuration

Identify the electron configuration of magnesium. Magnesium has an atomic number of 12, meaning it has 12 protons and, when neutral, 12 electrons. The electron configuration for magnesium is 1s2 2s2 2p6 3s2. This configuration shows that magnesium has two electrons in its outermost shell (the 3s subshell).

Analyzing the First Ionization Energy

The first ionization energy involves removing the first electron from the outermost shell, which is the 3s subshell in magnesium. This electron is relatively easy to remove because it is in the highest energy level and is shielded by the inner electrons.

Analyzing the Second Ionization Energy

The second ionization energy concerns the removal of the second electron from the 3s subshell. Since removing the first electron results in a +1 charge, the second electron will experience a greater effective nuclear charge and thus be harder to remove, requiring more energy. That's why the second ionization energy is significantly greater than the first.

Explaining the Large Increase to the Third Ionization Energy

The third electron to be removed comes from the 2p subshell, after the 3s electrons have been removed. Removing an electron from a filled or half-filled subshell of a lower energy level requires much more energy than removing one from a higher energy level. This is because electrons in the 2p subshell are closer to the nucleus and more strongly attracted to it due to a higher effective nuclear charge, hence the large increase to the third ionization energy.

Comparing to Sodium's Electron Configuration

Sodium has the electron configuration of 1s2 2s2 2p6 3s1. Since sodium only has one electron in its outermost shell, you would expect only one relatively low ionization energy to remove this electron. Any subsequent ionization energy would be much higher, as it would involve removing an electron from the 2p shell, similar to the jump from the second to third ionization energy in magnesium.

Key concepts

Electron Configuration

Understanding an atom's electron configuration is crucial for explaining trends in its chemical properties, including ionization energy. Electron configuration describes the distribution of electrons in an atom's orbitals. For instance, magnesium, with an atomic number of 12, has the electron configuration of 1s2 2s2 2p6 3s2. This notation shows that magnesium is structured with full inner shells and two electrons in its outermost (3s) shell. These outermost electrons are the most easily removed because they are farther from the nucleus and experience the least electrostatic pull from the positive charge of the nucleus due to shielding by inner electrons. This concept lays the foundation for understanding why the first ionization energy is the lowest.

Effective Nuclear Charge

The effective nuclear charge (Zeff) is the net positive charge experienced by an electron in a multi-electron atom. This charge is not the full charge of the nucleus because the shielding effect of inner shell electrons reduces it. For magnesium's second electron in the 3s shell, after the first electron is removed, the Zeff increases. The remaining electron experiences a stronger pull from the nucleus, as there's one less electron to repel it and contribute to shielding. Consequently, the ionization energy rises because it takes more energy to remove an electron that is more strongly bound. This principle explains why each successive electron removal requires more energy, particularly when entering a new shell, as with magnesium's drastic increase from the second to the third ionization energy.

Atomic Structure

Atomic structure is fundamental in predicting an element's ionization energies. The structure is composed of the nucleus, containing protons and neutrons, surrounded by electrons in defined orbitals. The energy required to remove an electron depends on its position in these orbitals. For instance, magnesium's third electron removed is from the 2p subshell. This subshell is closer to the nucleus, has less shielding than the 3s, and electrons here are held more tightly due to a lower principal quantum number. Thus, a significant increase in energy is necessary to overcome the attraction between these inner-shell electrons and the nucleus, leading to the sharply higher third ionization energy observed.

Periodic Trends

Periodic trends are patterns within the periodic table that illustrate variations in elemental properties, like ionization energy, as you move across a period or down a group. Ionization energy generally increases across a period as more protons in the nucleus create a higher charge, attracting electrons more strongly. Conversely, as you go down a group, ionization energy typically decreases since electrons are further from the nucleus in higher energy levels, reducing the electrostatic pull. Magnesium and sodium, although in the same period, have different electron configurations, leading to their differing ionization energies. The trends illustrate how atomic structure intricately connects with the elements' periodic placement and their chemical behavior.