Question

Three particles have the same electron configuration. One is a cation of an alkali metal, one is an anion of the halide in the third period, and the third particle is an atom of a noble gas. What are the identities of the three particles (including charges)? Which particle should have the smallest atomic/ionic radius, which should have the largest, and why?

Short answer

The three particles are Na+ (cation of an alkali metal), Cl- (anion of the third period halide), and Ar (noble gas). The Na+ cation should have the smallest radius due to loss of electrons, Cl- anion should have the largest radius due to gain of electrons, and Ar should be in between.

Step-by-step solution

Determine the Noble Gas

Since all three particles have the same electron configuration and one is a noble gas, determine the noble gas that has a complete octet in its outermost shell for the third period. This noble gas is Argon (Ar), which has an electron configuration of [Ne]3s^2 3p^6.

Identify the Cation of an Alkali Metal

The alkali metal cation must have lost one electron to have the same electron configuration as Argon. The alkali metal in the third period is Sodium (Na). Upon losing an electron, it becomes Na+ with the electron configuration [Ne]3s^2 3p^6, which is identical to that of Argon.

Identify the Halide Anion

The halide anion in the third period would have gained an electron to have the same electron configuration as Argon. The halide in the third period is Chlorine (Cl). Upon gaining an electron, it becomes Cl- with the electron configuration [Ne]3s^2 3p^6, which is identical to that of Argon.

Compare Atomic/Ionic Radii

To determine which particle has the smallest or largest radius, recall that cations are smaller than their neutral atoms due to the loss of electrons and increased nuclear charge exerted on the remaining electrons. Anions are larger than their neutral atoms because the addition of electrons results in greater electron-electron repulsion. Noble gases, without net charge, will generally have a radius between that of an analogous cation and anion. Therefore, the cation (Na+) should have the smallest radius, the anion (Cl-) should have the largest radius, and the noble gas (Ar) should have a radius in between.

Key concepts

Atomic/Ionic Radius

Understanding the size of an atom or ion is crucial when exploring the world of chemistry. The atomic radius refers to the distance from the nucleus to the outermost shell of electrons in an atom. When an atom gains or loses electrons, becoming an ion, we refer to its size as the ionic radius.

Cations, which are positively charged ions, typically have a smaller radius than their atom counterparts due to the loss of one or more electrons. This loss leads to a decrease in electron-electron repulsion and allows the nucleus to pull the remaining electrons closer, thereby shrinking the ion. For instance, when a sodium atom (Na) loses an electron to become Na+, its radius decreases.

In contrast, anions have a larger radius than their respective atoms. Upon gaining extra electrons, the increased electron-electron repulsion within the same shell causes the ion to expand. Chlorine (Cl), for example, gains an electron to become Cl-, resulting in an increased radius due to this repulsion.

Noble gases are a special case as they exist in a state with a complete octet and do not typically form ions. Their size stands between that of the cation and anion of similar electron configurations. The size trends among ions and atoms are valuable for predicting the behavior of substances in different chemical environments.

Periodic Table Trends

The periodic table exhibits distinct trends that can help us predict the properties of elements, such as atomic and ionic radius, electronegativity, and ionization energy. As we move across a period from left to right, the atomic number increases, adding more protons to the nucleus and more electrons to the outer shell. This increased nuclear charge pulls electrons closer to the nucleus, decreasing the atomic or ionic radius.

Conversely, as we descend down a group, elements have additional electron shells, and the increased distance between the nucleus and the outermost electrons leads to a larger atomic or ionic radius. This trend is why sodium, located on the left side of the third period, has a significantly larger atomic radius than argon, which is on the far right of the same period.

Impact of Nuclear Charge and Shielding

The effective nuclear charge increases across a period, contributing to a smaller atomic radius, while electron shielding by inner orbitals in atoms down a group leads to an increased radius. Understanding these periodic trends not only helps when comparing different elements, but also offers insight into the chemical reactivity and bonding potential of the elements.

Cation and Anion Identification

The first step in identifying a cation or anion is to examine an element's electron configuration. Cations are positively charged ions, formed when an atom loses electrons. This often happens with metals which have relatively low ionization energies and a tendency to lose valence electrons to achieve a stable electron configuration resembling that of a noble gas. For example, the alkali metal sodium becomes a cation (Na+) after losing one electron.

Anions, on the other hand, are negatively charged ions formed when an atom gains electrons. This is typical of nonmetals, like halogens, that have a high affinity for electrons and tend to gain them to complete their outer electron shell. Chlorine, for instance, becomes an anion (Cl-) by gaining an extra electron, filling its outermost shell to resemble argon.

Using Electronic Configurations

Electron configurations are a powerful tool for predicting the charge and type of ion an element is likely to form. By understanding whether an element is more likely to give up or gain electrons, we can determine whether it will form a cation or an anion. This understanding is crucial for predicting how elements will interact in chemical reactions and is a fundamental concept in chemistry education.