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Chapter 12: Problem 142

Would you expect an ionic solid or a network covalent solid to have the higher melting point?

Short Answer

Expert verified
Generally, a network covalent solid is expected to have a higher melting point than an ionic solid.

Step by step solution

01

Understanding the Characteristics of Ionic Solids

Ionic solids are composed of positive and negative ions held together by strong electrostatic forces of attraction, known as ionic bonds. Due to these strong bonding forces, ionic solids generally have high melting points.
02

Understanding the Characteristics of Network Covalent Solids

Network covalent solids are formed by covalent bonds, where atoms share electrons to achieve stability. The bonding in these solids is also very strong since the network structure extends throughout the entire solid, leading to substantial melting points.
03

Comparing Ionic and Network Covalent Solids

While both types have quite high melting points due to strong bonds, the comparison can depend on the specific substances being compared. In general, however, it's commonly accepted that network covalent solids oftentimes have higher melting points due to the extensive, strong bonding structure throughout the entire solid.

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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 fascinating due to the way they are structured and bonded. These solids consist of a lattice framework comprised of charged ions—cations and anions. The electrostatic attraction, or ionic bonding, between these oppositely charged ions is extremely robust. This intense bonding is the primary reason ionic solids usually exhibit high melting points.

A few key aspects to consider include:
  • Lattice Strength: The strength of the lattice is reinforced by the multitude of ionic bonds, requiring substantial energy to disrupt and melt the solid.
  • Composition: The size and charge of ions can influence the melting point; smaller ions with higher charges form stronger attractions and thus have higher melting points.
Examples of ionic solids with high melting points include table salt (NaCl) and magnesium oxide (MgO). These solids are typically crystalline and brittle, illustrating the rigid nature of the ionic framework.
Network Covalent Solids
Network covalent solids, sometimes called covalent network crystals, offer an intriguing chemistry and physical performance. These solids form when atoms are bonded by covalent bonds into a continuous and extensive structural network. This structure contributes to their remarkably high melting points.

The defining characteristics of network covalent solids include:
  • Intense Bonding: Covalent bonds involve the sharing of electrons between atoms, creating a sturdy link that contributes to high thermal stability.
  • Uniform Structure: The network is continuous, meaning that the extreme strength is uniform throughout the entire solid, unlike in other substances where molecular interactions vary.
Diamond and quartz (SiO₂) exemplify network covalent solids, known for their incredible hardness and resilience. Because the bonds in these solids extend throughout their entire grid, melting requires breaking numerous strong bonds, leading to very high melting points.
Bonding Structures
The bonding structures in ionic and network covalent solids demonstrate the profound impact that atomic arrangements and interactions have on physical properties like melting point.

Comparing Bonding Structures:
  • Ionic Bonding: In ionic solids, the electrostatic forces between charged ions form a rigid and stable structure. Nevertheless, despite high melting points, the actual temperature can vary based on ion types and lattice arrangement.
  • Covalent Network Bonding: Often surpassing ionic solids in thermal stability, network covalent solids owe their high melting points to the seamless connectivity of covalent bonds across the entire solid.
Despite both having strong bonds, network covalent solids generally tend to have higher melting points than ionic solids. This is attributed to their extensive and consistent bonding framework which requires immense energy to break. For example, while sodium chloride melts at 801°C, diamond—a network covalent solid—remains solid at temperatures well exceeding 3500°C.

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Most popular questions from this chapter

The normal melting point of copper is \(1357 \mathrm{K}\), and \(\Delta \mathrm{H}_{\text {fus }}\) of \(\mathrm{Cu}\) is \(13.05 \mathrm{kJ} \mathrm{mol}^{-1}\). (a) How much heat, in kilojoules, is evolved when a \(3.78 \mathrm{kg}\) sample of molten Cu freezes? (b) How much heat, in kilojoules, must be absorbed at 1357 K to melt a bar of copper that is \(75 \mathrm{cm} \times\) \(15 \mathrm{cm} \times 12 \mathrm{cm} ?\) (Assume \(d=8.92 \mathrm{g} / \mathrm{cm}^{3}\) for \(\mathrm{Cu}\).)

Germanium has a cubic unit cell with a side edge of \(565 \mathrm{pm} .\) The density of germanium is \(5.36 \mathrm{g} / \mathrm{cm}^{3}\) What is the crystal system adopted by germanium?

In an ionic crystal lattice each cation will be attracted by anions next to it and repulsed by cations near it. Consequently the coulomb potential leading to the lattice energy depends on the type of crystal. To get the total lattice energy you must sum all of the electrostatic interactions on a given ion. The general form of the electrostatic potential is $$V=\frac{Q_{1} Q_{2} e^{2}}{d_{12}}$$ where \(Q_{1}\) and \(Q_{2}\) are the charges on ions 1 and \(2, d_{12}\) is the distance between them in the crystal lattice. and \(e\) is the charge on the electron. (a) Consider the linear "crystal" shown below. The distance between the centers of adjacent spheres is \(R .\) Assume that the blue sphere and the green spheres are cations and that the red spheres are anions. Show that the total electrostatic energy is $$V=-\frac{Q^{2} e^{2}}{d} \times \ln 2$$ (b) In general, the electrostatic potential in a crystal can be written as $$V=-k_{M} \frac{Q^{2} e^{2}}{R}$$ where \(k_{M}\) is a geometric constant, called the Madelung constant, for a particular crystal system under consideration. Now consider the NaCl crystal structure and let \(R\) be the distance between the centers of sodium and chloride ions. Show that by considering three layers of nearest neighbors to a central chloride ion, \(k_{M}\) is given by $$k_{M}=\left(6-\frac{12}{\sqrt{2}}+\frac{8}{\sqrt{3}}-\frac{6}{\sqrt{4}} \cdots\right)$$ (c) Carry out the same calculation for the CsCl structure. Are the Madelung constants the same?

Without doing calculations, indicate how you would expect the lattice energies of \(\mathrm{LiCl}(\mathrm{s}), \mathrm{KCl}(\mathrm{s}), \mathrm{RbCl}(\mathrm{s}),\) and \(\mathrm{CsCl}(\mathrm{s})\) to compare with the value of \(-787 \mathrm{kJmol}^{-1}\) determined for \(\mathrm{NaCl}(\mathrm{s})\) on page \(543 .\) [ Hint: Assume that the enthalpies of sublimation of the alkali metals are comparable in value. What atomic properties from Chapter 9 should you compare?

The vapor pressure of trichloromethane (chloroform) is 40.0 Torr at \(-7.1^{\circ} \mathrm{C}\). Its enthalpy of vaporization is \(29.2 \mathrm{kJ} \mathrm{mol}^{-1} .\) Calculate its normal boiling point.
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