Question

Classify the solid states in terms of crystal types of the elements in the third period of the periodic table. Predict the trends in their melting points and boiling points.

Short answer

Third-period elements include Na, Mg, Al (metallic solids); Si (covalent network); P, S, Cl (molecular solids); Ar (noble gas). Melting/boiling points generally decrease across the period due to weaker bonds.

Step-by-step solution

Identify Third Period Elements

The third period of the periodic table consists of the elements: Sodium (Na), Magnesium (Mg), Aluminum (Al), Silicon (Si), Phosphorus (P), Sulfur (S), Chlorine (Cl), and Argon (Ar).

Determine Crystal Types

Each element can exist in a particular crystal structure. Sodium (Na) and Magnesium (Mg) have metallic crystal structures. Aluminum (Al) is also a metal and exhibits a face-centered cubic crystal structure. Silicon (Si) has a covalent network structure, forming a diamond-like lattice. Phosphorus can form covalent networks in certain allotropes, while sulfur forms molecular crystals as S8 rings. Chlorine and Argon exist as molecular solids at low temperatures, with Cl as diatomic molecules and Ar as a monatomic noble gas.

Analyze Trends in Melting Points

Melting points vary widely among these elements based on their bonding and structure. As a rule of thumb, metals like Na, Mg, and Al have relatively high melting points due to metallic bonding, with Aluminum having the highest among them. Silicon has an even higher melting point due to its strong covalent network structure. Phosphorus, Sulfur, and Chlorine have lower melting points, decreasing in the order listed. Argon has the lowest melting point, reflecting its status as a noble gas.

Analyze Trends in Boiling Points

Boiling points generally increase with stronger intermolecular forces or bonding types. Similar to the melting points, the trend is that metals have higher boiling points due to metallic bonding, with Aluminum higher than Sodium and Magnesium. Silicon as a covalent network solid has a very high boiling point. The molecular structures of Phosphorus, Sulfur, and Chlorine contribute to lower boiling points than the metals and silicon, with Argon having the lowest boiling point due to weak Van der Waals forces.

Key concepts

Crystal Structures

Crystal structures refer to the orderly and repeating arrangements of atoms or molecules in a solid. Different elements form various types of crystal structures based on their chemical and physical properties. For example, metals like sodium (Na), magnesium (Mg), and aluminum (Al) often form metallic crystal structures. These are characterized by the presence of closely packed atoms, often in arrangements such as face-centered cubic (FCC) or body-centered cubic (BCC).
Silicon (Si), known for its covalent network crystal structure, adopts a diamond-like lattice. This structure is a strongly bonded 3D network, resulting in high stability and hardness.
On the other hand, phosphorus, depending on its allotrope, can form covalent network crystals as well. Sulfur typically forms molecular crystals composed of S8 ring structures, while elements like chlorine (Cl) and argon (Ar) exist as molecular solids at low temperatures, with individual Cl molecules and Ar atoms weakly bonded.

Periodic Table

The periodic table is an essential tool in chemistry, used to identify elements and predict their properties. Along a period, such as the third period, there is a transition from metals to non-metals. This is due to the increasing number of electrons and protons as you move across.
The third period specifically includes a variety of elements from sodium (in Group 1) to argon (a noble gas in Group 18). Moving across the period, the elements transition from exhibiting metallic characteristics (Na, Mg, Al) to more covalent or molecular characteristics (Si, P, S) before reaching the noble gases (Cl, Ar).
This progression affects the type of bonding and crystal structures formed by these elements. As we move from metals to non-metals, metallic bonding gradually gives way to covalent and then to weak Van der Waals forces in simple diatomic (Cl) or monatomic (Ar) forms.

Melting Points

Melting points are impacted by the types of bonding and structures present in elements. Metals like sodium, magnesium, and aluminum typically have higher melting points, thanks to their robust metallic bonding. Among these, aluminum, with its face-centered cubic (FCC) structure, generally has the highest melting point due to tightly packed atoms and strong metallic bonds.
Silicon, with its diamond-like covalent network crystal, exceeds even aluminum in melting point. This is due to the extensive and strong covalent bonding throughout its structure. On the opposite end of the spectrum, non-metals such as phosphorus, sulfur, and chlorine have lower melting points. Their molecular structure and weaker intermolecular forces lead to easier breakage of bonds upon heating.
Argon, as a noble gas, has the lowest melting point among these elements. Its simple atomic structure and weak Van der Waals forces make it easier to transition from solid to liquid.

Boiling Points

Boiling points, much like melting points, depend heavily on bonding strength and atomic arrangements. Within the metals (Na, Mg, Al), aluminum usually has the highest boiling point, supported by strong metallic bonds in its FCC arrangement.
Silicon, once again, stands out due to its covalent network bonding, which requires significant energy to overcome, resulting in a high boiling point. In contrast, non-metals like phosphorus, sulfur, and chlorine have lower boiling points because of their molecular nature.

• Phosphorus and sulfur exhibit weaker forces between molecules, hence lower boiling points compared to metals and covalent solids.
• Chlorine, existing in a diatomic form (Cl2), follows this pattern with relatively low boiling points.

Argon, with minimal intermolecular forces as a monatomic gas, exhibits the lowest boiling point of all third-period elements.

Metallic Bonding

Metallic bonding is a key factor in understanding the properties and behaviors of metal elements. It involves the sharing of freely moving electrons across a lattice of positive ions. This "sea of electrons" allows for the conductivity and malleability seen in metals.

• In sodium (Na) and magnesium (Mg), this electron delocalization is relatively spread out, resulting in unique conductivity and moderate melting and boiling points.
• Aluminum (Al), with its highly packed FCC structure and dense electron cloud, exhibits greater strength in metallic bonding. This leads to higher melting and boiling points than Na or Mg.

Furthermore, these electron interactions create a cohesive force that keeps the metallic structure intact, making metals like Na, Mg, and Al strong and stable in their solid state.