Question

Why is the first ionization energy of aluminum less than the first ionization energy of magnesium?

Short answer

Aluminum has a lower first ionization energy than magnesium because its outer electron is in a 3p orbital, which is higher energy and further from the nucleus compared to the 3s orbital in magnesium, and is also more shielded, thus reducing the effective nuclear charge.

Step-by-step solution

Understanding Atomic Structure and Periodic Trends

First, understand that the ionization energy is the energy required to remove the most loosely held electron from an atom in its ground state. Magnesium (Mg) and Aluminum (Al) are both in period 3 of the periodic table. Magnesium has an atomic number of 12, resulting in an electron configuration of 1s2 2s2 2p6 3s2, while Aluminum with an atomic number of 13 has an electron configuration of 1s2 2s2 2p6 3s2 3p1.

Analyzing Electron Configurations

The electron to be removed during ionization of magnesium is in the 3s orbital, which is closer to the nucleus and more tightly bound than the electrons in the 3p orbital of aluminum. Since the 3p subshell, where aluminum has its outermost electron, is higher in energy and further away from the nucleus than the 3s subshell in magnesium, it is easier to remove the outer electron from aluminum.

Considering Effective Nuclear Charge

Although aluminum has one more proton in the nucleus than magnesium, resulting in a higher nuclear charge, the effective nuclear charge experienced by the outer electron in aluminum is less due to the additional shielding by the same inner electrons. Therefore, the attraction between the nucleus and the outer electron in aluminum is weaker than that in magnesium.

Comparing Ionization Energies

Since it is easier to remove an electron from a 3p orbital than a 3s orbital, and due to the lower effective nuclear charge experienced by the outer electron in aluminum, the first ionization energy of aluminum is less than that of magnesium.

Key concepts

Atomic Structure

Exploring atomic structure is foundational for understanding why elements behave the way they do. Atoms consist of a nucleus made of protons and neutrons, surrounded by electrons in different energy levels, or shells. Electrons within these shells are organized into subshells and orbitals, where they adhere to specific rules regarding their arrangement, known as electron configurations.

As seen in the exercise, magnesium and aluminum both occupy period 3 in the periodic table, reflecting their similar energy levels. However, the atomic structure reveals distinct electron configurations that influence their ionization energies. The heightened difficulty of removing an electron from magnesium's 3s orbital compared to aluminum's 3p orbital is a direct consequence of the principles governing atomic arrangement.

Periodic Trends

Periodic trends are patterns observed throughout the periodic table that provide insights into an element's chemical and physical properties. One key trend is ionization energy—the energy required to remove an electron from an atom. Generally, ionization energy increases across a period from left to right and decreases down a group.

This trend is linked to atomic size and the effective nuclear charge each electron experiences. In the context of the exercise, we see that despite being neighbors on the periodic table, aluminum has a lower ionization energy than magnesium, which may initially seem counterintuitive. Yet, our understanding of periodic trends offers an explanation: the positioning of electrons in different orbitals affects their relative energy levels and the ease with which they can be removed.

Electron Configurations

Electron configurations depict the distribution of electrons in an atom's orbitals, which is crucial for predicting how atoms will interact during chemical reactions. The electron configuration for magnesium (Mg), \(1s^2 2s^2 2p^6 3s^2\), indicates a full 3s subshell, making the electrons more firmly held. In contrast, aluminum (Al) has the configuration \(1s^2 2s^2 2p^6 3s^2 3p^1\), with a single electron in the higher energy 3p orbital.

As electrons in p orbitals are farther from the nucleus than those in s orbitals, they are less strongly attracted to the nucleus and have higher energy, making them easier to ionize. The exercise demonstrates the practical application of this concept; aluminum's lone 3p electron is easier to remove than magnesium's 3s electrons.

Effective Nuclear Charge

Effective nuclear charge (ENC) is the net positive charge experienced by an electron in a multi-electron atom. The ENC takes into account not just the total charge of the nucleus but also the shielding effect of electrons located in inner shells. This determines the strength of the attraction between the nucleus and an electron in a given orbital.

In our exercise example, aluminum has a higher actual nuclear charge due to the additional proton compared to magnesium. However, because both elements are in the same period and thus have the same inner-shell electron configuration, the increase in nuclear charge does not significantly enhance the proton-electron attraction for aluminum's outermost electron. The interplay between increased nuclear charge and consistent shielding results in a lower ENC for the outer electrons in aluminum, rationalizing the lower ionization energy relative to magnesium.