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Chapter 12: Problem 103

Selected chlorides have the following melting points: NaCl \(\left(801^{\circ} \mathrm{C}\right), \mathrm{MgCl}_{2}\left(714^{\circ} \mathrm{C}\right), \mathrm{PCl}_{3}\left(-94^{\circ} \mathrm{C}\right), \mathrm{SCl}_{2}\left(-121^{\circ} \mathrm{C}\right)\) $$ \begin{array}{l}{\text { (a) For each compound, indicate what type its solid }} \\ {\text { form is (molecular, metallic, ionic, or covalent-network). }} \\\ {\text { (b) Predict which of the following compounds has a }} \\ {\text { higher melting point: } \mathrm{CaCl}_{2} \text { or } \mathrm{SiCl}_{4} \text { . }}\end{array} $$

Short Answer

Expert verified
The types of solids for the given compounds are as follows: \(NaCl\) and \(MgCl_2\) are ionic solids, while \(PCl_3\) and \(SCl_2\) are molecular solids. Comparing the melting points of \(CaCl_2\) and \(SiCl_4\), since \(CaCl_2\) forms a stronger ionic bond, it is predicted to have a higher melting point than \(SiCl_4\), which has a covalent bond.

Step by step solution

01

NaCl - Sodium Chloride

Na is a metal (alkali metal in group 1) and Cl is a non-metal (halogen in group 17), so the bond between them is ionic. Thus, NaCl is an ionic solid.
02

MgCl2 - Magnesium Chloride

Mg is a metal (alkaline earth metal in group 2) and Cl is a non-metal (halogen in group 17), so the bond between them is ionic. Hence, MgCl2 is an ionic solid.
03

PCl3 - Phosphorus Trichloride

P and Cl both are non-metals (P in group 15 and Cl in group 17) which form a covalent bond. As PCl3 consists of discrete molecules instead of a continuous network, it is a molecular solid.
04

SCl2 - Sulfur Dichloride

S and Cl are both non-metals (S in group 16 and Cl in group 17) which form a covalent bond. Like PCl3, SCl2 also consists of discrete molecules, so it is a molecular solid. Step 2: Predicting which compound has a higher melting point between CaCl2 and SiCl4
05

Electronegativity of elements in CaCl2 and SiCl4

Ca is a metal, situated in group 2 (alkaline earth metals). Cl is a non-metal, in group 17 (halogens). Their electronegativity difference is significant, and they form an ionic bond. Si is a metalloid situated in group 14, and Cl is a non-metal in group 17 (halogens). They form a covalent bond due to their closer electronegativities.
06

Comparing melting points of CaCl2 and SiCl4

CaCl2 forms an ionic bond which is stronger than the covalent bond in SiCl4. Ionic solids typically have higher melting points because they require more energy to break the bonds. Therefore, we predict that CaCl2 would have a higher melting point than SiCl4.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Ionic Solids
Ionic solids are a type of chemical compound where the ions are held together by strong electrostatic forces known as ionic bonds. These solids form when metals bond with non-metals, where the metal loses electrons to become a positively charged ion (cation), and the non-metal gains electrons to become a negatively charged ion (anion).
The attraction between ions of opposite charges results in a rigid lattice structure which contributes to the solid's strength and stability.
  • High Melting Points: Ionic solids have high melting points due to the strong forces of attraction between the oppositely charged ions. For example, the melting point of sodium chloride (NaCl) is 801掳C.
  • Conductivity: While ionic solids do not conduct electricity in solid form, they do conduct when melted or dissolved in water, as the ions are free to move and carry charge.
  • Form: These compounds are usually crystalline in form, often appearing as clear or white crystals that are hard and brittle.
Molecular Solids
Molecular solids are composed of molecules held together by weaker intermolecular forces rather than by ionic or covalent bonds. These solids are typically formed from non-metallic elements and compounds.
Unlike ionic solids, their structures consist of discrete molecules.
  • Low Melting Points: Due to the weak forces holding the molecules together, molecular solids generally have low melting points. For instance, phosphorus trichloride (PCl3) has a melting point of -94掳C.
  • Soft and Insulating: They are usually soft and do not conduct electricity, as there are no free ions or electrons to carry a charge.
  • Examples: Common examples include ice (solid water) and dry ice (solid carbon dioxide).
These unique properties make molecular solids useful in a variety of applications where low melting and boiling points are required.
Melting Points
The melting point of a substance is the temperature at which it changes from a solid to a liquid. It is a key physical property that provides insight into the strength of the forces holding a solid together.
Various factors determine the melting point of a compound, including the type of bonding and the mass of the atoms involved.
  • Ionic vs. Molecular Solids: Ionic solids, like sodium chloride (NaCl), possess high melting points because of the strong ionic bonds. In contrast, molecular solids like sulfur dichloride (SCl2), have much lower melting points due to weaker molecular attractions.
  • Predicting Melting Points: When comparing substances, ionic compounds generally have higher melting points than covalent compounds. For example, calcium chloride (CaCl2) is expected to have a higher melting point than silicon tetrachloride (SiCl4) due to its ionic character.
  • Temperature Indicators: High melting points typically indicate that a lot of heat energy is needed to overcome the interactions holding the particles together in the solid state.
Understanding melting points can help in determining a material's state and usability in different temperature conditions.
Electronegativity Differences
Electronegativity refers to the ability of an atom to attract electrons towards itself within a chemical bond. It is an important factor in determining the type and strength of bonds between atoms. The difference in electronegativity between two atoms can indicate the likely nature of the bond between them.
  • Ionic Bonds: A large difference in electronegativity (typically greater than 1.7) often results in the formation of ionic bonds, as seen in magnesium chloride (MgCl2).
  • Covalent Bonds: Smaller differences, such as those found in phosphorus trichloride (PCl3), often lead to covalent bonding, where electrons are shared rather than transferred.
  • Polar vs. Non-polar: Intermediate differences can result in polar covalent bonds, where electrons are shared unequally, creating dipoles.
Electronegativity differences not only help predict the type of bond that will form, but also the physical properties of the resulting compound, such as solubility, boiling point, and conductivity.

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Most popular questions from this chapter

Both covalent-network solids and ionic solids can have melting points well in excess of room temperature, and both can be poor conductors of electricity in their pure form. However, in other ways their properties are quite different. $$ \begin{array}{l}{\text { (a) Which type of solid is more likely to dissolve in water? }} \\ {\text { (b) Which type of solid can become a considerably better }} \\ {\text { conductor of electricity via chemical substitution? }}\end{array} $$

Besides the cubic unit cell, which other unit cell(s) has edge lengths that are all equal to each other? (a) Orthorhombic, (b) hexagonal, (c) rhombohedral, (a) triclinic, (e) both rhombohedral and triclinic.

Indicate whether each statement is true or false: $$ \begin{array}{l}{\text { (a) Substitutional alloys tend to be more ductile than inter- }} \\ {\text { stitial alloys. }} \\ {\text { (b) Interstitial alloys tend to form between elements with }} \\ {\text { similar ionic radii. }} \\ {\text { (c) Nonmetallic elements are never found in alloys. }}\end{array} $$

Indicate whether each statement is true or false: $$ \begin{array}{l}{\text { (a) The band gap of a semiconductor decreases as the parti- }} \\ {\text { cle size decreases in the } 1-10 \text { -nm range. }} \\\ {\text { (b) The light that is emitted from a semiconductor, upon }} \\\ {\text { external stimulation, becomes longer in wavelength as }} \\ {\text { the particle size of the semiconductor decreases. }}\end{array} $$

For each of the following groups, which metal would you expect to have the highest melting point: (a) gold, rhenium, or cesium; (b) rubidium, molybdenum, or indium; (c) ruthenium, strontium, or cadmium?
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