Question

Consider the elements: \(\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}, \mathrm{Si}, \mathrm{P}\) a. Which element has the highest second ionization energy? b. Which element has the smallest atomic radius? c. Which element is least metallic? d. Which element is diamagnetic?

Short answer

a. Mg has the highest second ionization energy. b. P has the smallest atomic radius. c. Si is the least metallic. d. Mg (after losing 2 electrons) is diamagnetic.

Step-by-step solution

Understanding Second Ionization Energy

Second ionization energy refers to the energy required to remove the second electron from an atom after the first has been removed. Generally, second ionization energy increases across a period from left to right due to the increased effective nuclear charge. Among the given elements, magnesium (Mg) would require more energy to remove a second electron because its first electron is removed from a completed 3s subshell, and the second electron removal implies breaking into a stable noble gas configuration.

Comparing Atomic Radii

Atomic radius generally decreases across a period because of the increasing effective nuclear charge, which pulls electrons closer to the nucleus. Among the given elements, phosphorus (P) has the smallest atomic radius because it is furthest to the right in the third period of the periodic table.

Identifying Metallic Character

Metallic character decreases across a period from left to right as the elements become more nonmetallic. Silicon (Si) is the element among the given ones that is the least metallic, being a metalloid; it is towards the right side, bordering the metal-nonmetal divide in the periodic table.

Determining Diamagnetic vs Paramagnetic Elements

An element is considered diamagnetic if it has no unpaired electrons and therefore is not attracted to a magnetic field. In the elements listed, (Neon) Ne would be diamagnetic because it has a full outer electron shell with paired electrons. However, among the given elements, Na, Mg, Al, Si, and P, Magnesium (Mg) will be diamagnetic after losing its two 3s electrons through ionization, which leaves it with a complete octet of 2s and 2p electrons, all paired.

Key concepts

Ionization Energy

Ionization energy is a critical concept that refers to the amount of energy needed to remove an electron from an atom in its gaseous state. The first ionization energy is the energy required to remove the first electron, while higher ionization energies pertain to the removal of subsequent electrons.

• Ionization energy generally increases across a period from left to right on the periodic table, and the reason for this is the increasing effective nuclear charge. This increase in nuclear charge attracts the electrons more strongly, making them harder to remove.
• As an atom loses electrons, it becomes a positive ion, and the ion becomes increasingly stable. Thus, higher ionization energies are often associated with greater stability in the resulting ion.

Using these principles, we can determine that the process of ionizing magnesium (Mg) twice—as in the case where the second electron is removed—is especially energy-intensive since it disrupts a stable noble gas electron configuration.

Understanding the trend for ionization energy not only helps in predicting the reactivity of elements but also their electron configuration stability post-ionization.

Atomic Radius

The atomic radius is one of the most fundamental periodic properties, indicating the size of an atom.

• Across a period, the atomic radius decreases because the increasing number of protons in the nucleus exerts a stronger pull on the electron cloud, drawing it closer in.
• On the other hand, down a group, the atomic radius increases due to the addition of electron shells, which outweighs the effect of increased nuclear charge.

In the given set of elements, phosphorus (P) stands out with the smallest atomic radius since it is furthest to the right within its period.

A more in-depth understanding of atomic radii is essential not just for visualizing atomic structure but also for grasping interactions between atoms, such as bonding and van der Waals forces.

Metallic Character

Metallic character describes an element's ability to lose electrons and form positive ions or cations.

• Metallic properties decline across a period from the left (metallic side) to the right (nonmetallic side) of the periodic table, which coincides with the trend of increasing ionization energy.
• The element's ability to conduct electricity and heat, luster, and malleability all contribute to its metallic character.

Among the elements in question, silicon (Si) possesses the least metallic character, existing on the borderline between metals and nonmetals, characterized as a metalloid.

Understanding metallic character is vital not only for classifying elements but also for predicting their physical and chemical behavior in various applications.

Diamagnetism and Paramagnetism

Diamagnetism and paramagnetism relate to the magnetic properties of an element based on its electronic structure.

• An element is considered diamagnetic when it has all of its electrons paired and thus does not have a net magnetic field. Such elements are weakly repelled by a magnetic field.
• Paramagnetism occurs in elements with unpaired electrons, which align with external magnetic fields and are thus attracted to them.

While none of the given elements are inherently diamagnetic in their neutral state, magnesium (Mg) becomes diamagnetic once it loses its two 3s electrons, leaving a filled shell of paired electrons.

Understanding diamagnetism and paramagnetism is not only crucial for studying magnetic properties but also for applications in fields such as materials science and chemistry where magnetic behavior plays a role.