Question

Which outer electron configurations would you expect to correspond to a noble gas? To a metalloid? $$\text { a. } n s^{2} \quad \text { b. } n s^{2} n p^{6} \quad \text { c. } n s^{2} n p^{5} \quad \text { d. } n s^{2} n p^{2}$$

Short answer

The noble gas configuration would correspond to option b, 'ns^2 np^6', as it represents a full valence shell. Metalloid configurations that can be expected are option c, 'ns^2 np^5', and option d, 'ns^2 np^2' as they are close to completing their valence shell.

Step-by-step solution

Identify Noble Gas Configuration

The noble gases have full valence shells, which correspond to a filled 's' and 'p' subshells in their outermost energy level. The general electron configuration for a noble gas is therefore represented as 'ns^2 np^6', where 'n' stands for the period number or energy level of the valence shell.

Choose Noble Gas Configuration

Option b, 'ns^2 np^6', represents a filled s and p subshell within the same energy level, which is characteristic of noble gases. This is the complete set of eight valence electrons, often referred to as an octet, which makes the atom stable.

Identify Metalloid Configuration

Metalloids have properties that are intermediate between metals and nonmetals. Their outer electron configurations generally do not have a full valence shell, as in the case of noble gases, but are relatively close to completing their valence shell. These could include 'ns^2 np^2', 'ns^2 np^3', or 'ns^2 np^5', which show varying degrees of reactivity.

Choose Metalloid Configuration

Among the given options, both 'ns^2 np^5' and 'ns^2 np^2' (options c and d) are plausible configurations for metalloids. Option c, 'ns^2 np^5', is one electron short of a full valence shell and can gain one electron to achieve stability, this is typical for metalloids like arsenic. Option d, 'ns^2 np^2', needs to either gain or lose electrons to reach a stable state and is also seen in metalloids such as silicon.

Key concepts

Noble Gases

When trying to understand the properties of elements, knowing their electron configuration can be extremely helpful. Particularly for noble gases, which are a group of elements with very low reactivity, their electron configurations are quite distinctive. Noble gases have completely filled outer shell electron configurations, making them very stable and largely unreactive.

Noble gases, such as helium (He), neon (Ne), and argon (Ar), have their outermost 's' and 'p' orbitals fully occupied. This is generally represented as 'ns² np⁶', where 'n' refers to the principal quantum number that denotes the energy level of the valence electrons. Having 'np⁶' indicates that there are six electrons in the 'p' orbital and 'ns²' indicates two electrons in the 's' orbital, completing an octet which gives utmost stability to the noble gases.

Elements strive to achieve this configuration because it represents a state of minimal energy, a principle known as the octet rule. The octet rule is a chemical rule of thumb that reflects the observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, thereby achieving noble gas configuration and gaining stability.

Metalloids

Metalloids are fascinating elements that possess a mix of metallic and non-metallic properties, which makes them very useful in various technological applications, especially in semiconductors. Their electron configurations are not as straightforward as those of noble gases.

The configurations for metalloids can vary, as they don't have completely filled valence shells. Typically, their outer electron layers are close to having a full set of electrons which would either complete the octet for p-block elements or achieve a filled subshell for s-block elements. For example, 'ns² np⁵' indicates one electron is missing from completing the 'p' orbital octet, and 'ns² np²' shows a need for either gaining or losing electrons to reach a stable state.

Metalloids like silicon (Si) and germanium (Ge) are used in electronics because they can act as conductors or insulators depending on the circumstances, such as the presence of impurities (doping) or temperature changes. Their versatility is largely due to their intriguing electron configurations, which allow for the formation of covalent bonds with other elements to achieve a more stable electronic structure.

Valence Electrons

The term 'valence electrons' refers to the electrons in the outermost shell of an atom that are involved in forming chemical bonds. They are significant because their number and arrangement determine the chemical behavior and reactivity of an element.

The number of valence electrons is crucial when it comes to predicting how an element will interact with others. For instance, elements with a single valence electron, like the alkali metals, are highly reactive because they tend to lose that electron easily to attain a noble gas configuration. On the other hand, elements with a nearly complete set of valence electrons, such as halogens, tend to be reactive as well, but in their case, they are eager to gain an extra electron to achieve stability.

To visualize the valence electrons, the electron configuration is written out, indicating which orbitals are filled or partially filled in the outermost shell. For example, oxygen has the electron configuration '2s² 2p⁴', which means it has six valence electrons (the sum of the s and p orbital electrons in the outer shell), thereby needing two more to complete its octet.