Question

Which of these elements are unlikely to form covalent bonds: \(\mathrm{S}, \mathrm{H}, \mathrm{K}, \mathrm{Ar},\) Si? Explain your choices.

Short answer

Among the elements S, H, K, Ar, and Si, argon (Ar) is the element that is unlikely to form covalent bonds, since it has a complete outer electron shell and does not need to share/gain electrons. The other nonmetals (S, H, and Si) are likely to form covalent bonds to achieve a stable electron configuration.

Step-by-step solution

Identify the nonmetal elements

First, we must identify the nonmetal elements among S, H, K, Ar, and Si. In the periodic table, nonmetals are found to the right of the metalloids and include elements such as hydrogen, oxygen, and nitrogen. In this list, S (sulfur), H (hydrogen), Ar (argon), and Si (silicon) are nonmetals. K (potassium) is a metal.

Determine their electron configuration

To determine which nonmetals are more likely to form covalent bonds, let's analyze their electron configurations. Nonmetals tend to form covalent bonds by completing their outer electron shell, either by sharing or gaining electrons from other nonmetals. Sulfur (S): 1s2, 2s2, 2p6, 3s2, 3p4 Hydrogen (H): 1s1 Argon (Ar): 1s2, 2s2, 2p6, 3s2, 3p6 Silicon (Si): 1s2, 2s2, 2p6, 3s2, 3p2

Assess their likelihood to form covalent bonds

Now let's assess their likelihood to form covalent bonds based on their electron configurations. Sulfur (S): Sulfur has 6 valence electrons (3s2, 3p4) and needs two more electrons to complete its outer shell. It is likely to form covalent bonds. Hydrogen (H): Hydrogen has 1 valence electron (1s1) and needs one more electron to complete its outer shell. It is also likely to form covalent bonds. Argon (Ar): Argon has a full outer electron shell (1s2, 2s2, 2p6, 3s2, 3p6). Since it has a stable electron configuration, it is a noble gas and does not need to form covalent bonds. Silicon (Si): Silicon has 4 valence electrons (3s2, 3p2) and needs four more electrons to complete its outer shell. It is likely to form covalent bonds. Among the elements S, H, K, Ar, and Si, argon (Ar) is the element that is unlikely to form covalent bonds, since it has a complete outer electron shell and does not need to share/gain electrons. The other nonmetals are likely to form covalent bonds to achieve a stable electron configuration.

Key concepts

Electron Configuration

Electron configuration describes the distribution of electrons in an atom's electron shells. Electrons orbit the nucleus in "shells" or "energy levels." These shells are filled sequentially, starting from the closest to the nucleus. The configuration follows the "Aufbau principle," which describes the order that electrons fill orbitals:

\[1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p\]

Each number represents the energy level (n), the letter shows the type of orbital (s, p, d, f) and the superscript indicates the number of electrons in those orbitals. Electron configuration is crucial in determining chemical properties and bonding behavior.

For example:

• **Sulfur (S)**: Its electron configuration is \(1s^2, 2s^2, 2p^6, 3s^2, 3p^4\). Here, electrons occupy up to the 3rd energy level.
• **Hydrogen (H)**: Its simple configuration is \(1s^1\), indicating one electron in the first shell.
• **Argon (Ar)**: With a configuration of \(1s^2, 2s^2, 2p^6, 3s^2, 3p^6\), Argon's outer shell is complete, making it "noble."

This completion of shells indicates full stability for Ar, which explains its inert nature.

Valence Electrons

Valence electrons are the outermost electrons of an atom and determine how an atom will react chemically. These electrons participate in forming chemical bonds, such as covalent bonds. When atoms have incomplete outer shells, they are more likely to interact with other atoms to achieve stability by sharing, losing, or gaining electrons.

For example, when we consider:

• **Sulfur (S)**: It has 6 valence electrons \( (3s^2, 3p^4) \) and needs two more electrons to fill its outer shell, so it is likely to form covalent bonds by sharing electrons with other atoms.
• **Hydrogen (H)**: With just 1 valence electron \((1s^1)\), it needs one more to fill its shell. It can easily form a covalent bond by sharing this electron.
• **Silicon (Si)**: Possessing 4 valence electrons \((3s^2, 3p^2)\), it can form up to four covalent bonds.

In each of these cases, the number of valence electrons is crucial in predicting the type and number of bonds the element can form. Atoms are most stable when their valence shell is full, often resembling the nearest noble gas configuration.

Noble Gases

Noble gases are elements located in Group 18 of the periodic table. They include helium, neon, argon, krypton, xenon, and radon. These gases are known for their remarkable stability which arises from having a complete set of valence electrons. Their outer electron shell is full, making them generally unreactive.

Noble gases adhere to the "octet rule," which states that atoms are most stable with eight electrons in their outermost orbit. This fully populated valence shell results in minimal chemical reactivity. Thus, noble gases rarely form compounds. For example:

• **Argon (Ar)**: Its electron configuration is \(1s^2, 2s^2, 2p^6, 3s^2, 3p^6\), with a filled outer shell, explaining its lack of chemistry under normal conditions.

Understanding noble gases is key to grasping why elements like Argon do not form covalent bonds. Their stability set them apart as reference points for electron configurations of other elements seeking stability through bonding.