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13.3: Solids

Created by: CK-12

Lesson Objectives

  • Describe a solid according to the kinetic-molecular theory.
  • Understand that a solid also has a vapor pressure, and describe the relationship between the vapor pressure of a solid and sublimation.
  • Describe the features of the seven basic crystal systems.
  • Define a unit cell.
  • List the four classes of crystalline solids and describe the properties of each.
  • Describe an amorphous solid.

Lesson Vocabulary

  • amorphous solid
  • crystal
  • deposition
  • melting point
  • sublimation
  • unit cell

Check Your Understanding

Recalling Prior Knowledge

  • How does the kinetic-molecular theory treat the condensed states of matter differently than it treats gases?
  • What happens to the particles of a liquid as they undergo a change of state into a gas?

So far, we have studied gases and liquids. In this lesson, you will gain an understanding of the nature of solids, focusing on the many different ways in which particles can be arranged within ordered solid crystals.

Properties of Solids

Solids are similar to liquids in that both are condensed states, with particles that are far closer together than those of a gas. However, while liquids are fluid, solids are not. The particles of most solids are packed tightly together in an orderly arrangement. The motion of individual atoms, ions, or molecules in a solid is restricted to vibrational motion about a fixed point. Solids are almost completely incompressible, and for most substances, their solid form is the densest of the three states of matter.

As a solid is heated, the average kinetic energy of its particles still increases, but due to their relatively fixed positions, this manifests itself as stronger and more rapid vibrations. Eventually, the organization of the particles within the solid structure begins to break down and the solid starts to melt. The melting point is the temperature at which a solid changes into a liquid. At its melting point, the disruptive vibrations of the particles in the solid overcome the attractive forces operating within the solid. As with boiling points, the melting point of a solid is dependent on the strength of those attractive forces. Sodium chloride (NaCl) is an ionic compound, so it consists of a multitude of strong ionic bonds. Sodium chloride melts at 801°C. Ice (solid H2O) is a molecular compound whose molecules are held together by hydrogen bonds. Though hydrogen bonds are the strongest of the intermolecular forces, they are still much weaker than ionic bonds. The melting point of ice is 0°C.

The melting point of a solid is the same as the freezing point of the corresponding liquid. At that temperature, the solid and liquid states of the substance are in equilibrium. For water, this equilibrium occurs at 0°C.

\mathrm{H_2O}(s) \rightleftharpoons \mathrm{H_2O}(l)

You can view a cartoon with a simple explanation of the properties of solids at http://www.abpischools.org.uk/page/modules/solids-liquids-gases/slg2.cfm?age=Age%20range%207-11&subject=Science. Page 2 is specifically about the properties of solids.

Vapor Pressure of a Solid

In the last lesson, you learned about the vapor pressure of a liquid and its dependence upon temperature. Solids also have a vapor pressure, though it is generally much less than that of a liquid. A snow bank will gradually disappear even if the temperature stays below 0°C. The snow does not melt but, instead, passes directly from the solid state to the vapor state. Sublimation is the change of state from a solid to a gas without passing through the liquid state.

Iodine is an example of a substance for which sublimation can be readily observed at room temperature, as seen below (Figure below). Although the vapor pressure of solid iodine at room temperature is actually quite low (< 1 mmHg), its vapor is a distinctive purple color and has a very strong scent, making it easy to detect.

Solid iodine sublimes readily, forming a purple vapor.

The following video shows the sublimation of iodine in a beaker. The beaker is covered with an evaporating dish filled with ice water, and the dark iodine crystals are deposited on the outside of that dish. Deposition is the change of state from a gas to a solid.

http://www.youtube.com/watch?v=4fAOI6BeMZY (3:52)

Carbon dioxide is another substance that sublimes at atmospheric pressures. Carbon dioxide in the solid state is known as dry ice. Dry ice is very cold (−78°C), so it is used as a coolant for goods such as ice cream that must remain frozen during shipment. Because the dry ice sublimes rather than melting, there is no liquid mess associated with its change of state as the dry ice warms.

Crystalline Solids

The majority of solids are crystalline in nature. A crystal is a substance in which the particles are arranged in an orderly, repeating, three-dimensional pattern. Particles of a solid crystal may be ions, atoms, or molecules, depending on the type of substance. The three-dimensional arrangement of a solid crystal is referred to as the crystal lattice. The general nature of crystal lattices of ionic compounds was introduced in the chapter, Ionic and Metallic Bonding. Different arrangements of the particles within a crystal cause them to adopt several different shapes.

Crystal Systems

Crystals are classified into general categories based on their shapes. A crystal is defined by its faces, which intersect with one another at specific angles, which are characteristic of the given substance. The seven crystal systems are shown below (Table below), along with an example of each. The edge lengths of a crystal are represented by the letters a, b, and c. The angles at which the faces intersect are represented by the Greek letters α, β, and γ. Each of the seven crystal systems differs in terms of the angles between the faces and in the number of edges of equal length on each face.

Seven Basic Crystal Systems and an Example of Each
Crystal System Diagram Example

Cubic

a = b = c; α = β = γ = 90°

Pyrite

Tetragonal

a = b ≠ c; α = β = γ = 90°

Wulfenite

Orthorhombic

a ≠ b ≠ c; α = β = γ = 90°

Aragonite

Monoclinic

a ≠ b ≠ c; α ≠ 90° = β = γ

Azurite

Rhombohedral

a = b = c; α = β = γ ≠ 90°

Calcite

Triclinic

a ≠ b ≠ c; α ≠ β ≠ γ ≠ 90°

Microcline

Hexagonal

a = b ≠ c; α = β = 90°, γ = 120°

Beryl

Unit Cells

A unit cell is the smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire crystal. A crystal can be thought of as the same unit cell repeated over and over in three dimensions. Illustrated below (Figure below) is the relationship of a unit cell to the entire crystal lattice.

A unit cell is the smallest repeating portion of a crystal lattice.

Unit cells occur in many different varieties. As one example, the cubic crystal system is composed of three different types of unit cells: (1) simple cubic, (2) face-centered cubic, and (3) body-centered cubic. These are shown in three different ways below (Figure below).

Three unit cells of the cubic crystal system. Each sphere represents an atom or an ion. In the simple cubic system, the atoms or ions are at the corners of the unit cell only. In the face-centered unit cell, there are also atoms or ions in the center of each of the six faces of the unit cell. In the body-centered unit cell, there is one atom or ion in the center of the unit cell in addition to the corner atoms or ions.

Classes of Crystalline Solids

Crystalline substances can also be described by the types of particles in them and the types of chemical bonding that take place between the particles. There are four types of crystals: (1) ionic, (2) metallic, (3) covalent network, and (4) molecular. Properties and several examples of each type are listed and described below (Table below).

Crystalline Solids – Melting and Boiling Points
Type of Crystalline Solid Examples (formulas) Normal Melting Point (°C) Normal Boiling Point (°C)
Ionic

NaCl

CaF2

801

1418

1413

2533

Metallic

Hg

Na

Au

W

−39

371

1064

3410

630

883

2856

5660

Covalent network

B

C (diamond)

SiO2

2076

3500

1600

3927

3930

2230

Molecular

H2

I2

NH3

H2O

−259

114

-78

0

−253

184

-33

100

  1. Ionic crystals—The ionic crystal structure, discussed in the chapter Ionic and Metallic Bonding, consists of alternating positively charged cations and negatively charged anions. The ions may either be monatomic or polyatomic. Generally, ionic crystals form from a combination of metal cations and Group 16 or 17 nonmetal anions, although nonmetallic polyatomic ions are also common components of ionic crystals. Ionic crystals are hard and brittle and have high melting points. Ionic compounds do not conduct electricity as solids, but they do conduct when molten or dissolved in water.
  2. Metallic crystals—Metallic crystals, also discussed in the chapter Ionic and Metallic Bonding, consist of metal cations surrounded by a “sea” of mobile valence electrons. These electrons, also referred to as delocalized electrons, do not belong to any one atom and are capable of moving through the entire crystal. As a result, metals are good conductors of electricity. As seen in the table above (Table above), metallic crystals can have a wide range of melting points.
  3. Covalent network crystals—A covalent network crystal consists of atoms at the lattice points of the crystal, with each atom being covalently bonded to its nearest neighbor atoms. The covalently bonded network is three-dimensional and contains a very large number of atoms. Network solids include diamond (Figure below), quartz, many metalloids, and oxides of transition metals and metalloids. Network solids are hard and brittle, with extremely high melting and boiling points. Being composed of atoms rather than ions, they do not conduct electricity well in any state.
  4. Molecular crystals—Molecular crystals typically consist of molecules at the lattice points of the crystal, held together by relatively weak intermolecular forces. The intermolecular forces may be dispersion forces in the case of nonpolar substances or dipole-dipole forces in the case of polar substances. Some molecular crystals, such as ice, have molecules held together by hydrogen bonds. When one of the noble gases is cooled and solidified, the lattice points are individual atoms rather than molecules. However, because the atoms are held together by dispersion forces and not by covalent or metallic bonds, the properties of such a crystal are most similar to the crystals of molecular substances. In all cases, the intermolecular forces holding the particles together are far weaker than either ionic or covalent bonds. As a result, the melting and boiling points of molecular crystals are much lower. Lacking ions or free electrons, molecular crystals are poor electrical conductors.

Diamond is a network solid and consists of carbon atoms covalently bonded to one another in a repeating three-dimensional pattern. Each carbon atom makes four single covalent bonds in a tetrahedral geometry.

Amorphous Solids

Unlike a crystalline solid, an amorphous solid is a solid that lacks an ordered internal structure. Some examples of amorphous solids include rubber, plastic, and gels. Glass is a very important amorphous solid that is made by cooling a mixture of materials in such a way that it does not crystallize. Glass is sometimes referred to as a supercooled liquid rather than a solid. If you have ever watched a glassblower in action, you have noticed that he takes advantage of the fact that amorphous solids do not have a distinct melting point like crystalline solids do. Instead, as glass is heated, it slowly softens and can be shaped into all sorts of interesting forms. When a glass object shatters, it does so in a very irregular way. In contrast, a crystalline solid always breaks along specific planes as dictated by its crystal system.

Lesson Summary

  • Intermolecular forces keep the particles of liquids close together, though liquids are still able to flow. Surface tension is a property of liquids that gives objects the ability to float on its surface.
  • Molecules of a liquid evaporate even when the temperature of the liquid is below its boiling point. Molecules with high kinetic energies escape the surface of the liquid, leaving the remaining liquid cooler than it was before.
  • The vapor pressure of a liquid is determined by the strength of the intermolecular forces. Vapor pressure increases as temperature increases due to an increased rate of evaporation. The relationship of vapor pressure to temperature is shown on a vapor pressure curve.
  • Boiling of a liquid occurs when the vapor pressure of the liquid is equal to the external pressure.

Lesson Review Questions

Reviewing Concepts

  1. How are the particles arranged in most solids?
  2. Describe what happens when a solid is heated to its melting point. What phases are in equilibrium at the melting point?
  3. Wet clothes can be hung on a clothesline in sub-freezing temperatures, and they will still dry. Explain.
  4. What accounts for the differences between the seven basic crystal systems?
  5. Explain the difference between a crystal lattice and a unit cell.
  6. What are three kinds of unit cells that would be classified as having a cubic crystal system?

Problems

  1. Use the table above (Table above) to identify which crystal system or systems fit the following criteria.
    1. The lengths of all edges are the same.
    2. All angles between faces are 90°.
    3. No faces intersect at 90° angles, and no two edges are the same length.
  2. From the given descriptions, identify each of the following as an ionic crystal, a metallic crystal, a covalent network crystal, or a molecular crystal.
    1. Substance A melts at 125°C and does not conduct electricity as either a solid or a liquid.
    2. Substance B is very hard and melts at 3440°C. It does not conduct electricity either as a solid or a liquid.
    3. Substance C is hard and brittle and melts at 1720°C. It does not conduct electricity as a solid, but it does conduct as a liquid.
    4. Substance D is malleable and conducts electricity well as a solid. It melts at 1135°C.
  3. Identify whether the solid form of each of the following substances is an ionic crystal, a metallic crystal, a covalent network crystal, or a molecular crystal.
    1. Mn
    2. MgBr2
    3. Xe
    4. NH4Cl
    5. CO2
    6. As
    7. Pd
    8. SiC
  4. Use the figure above (Figure above) to list the three kinds of cubic unit cells in order from least densely packed to most densely packed.

Further Reading / Supplemental Links

References

For the table above (Table above),

Points to Consider

As part of the discussions about the three states of matter, you have seen how matter undergoes changes from one state to another.

  • Under what conditions of temperature and pressure do the various changes of state occur?
  • What is a phase diagram, and how can it be used to understand the nature of a given substance?

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Date Created:

Aug 02, 2012

Last Modified:

Dec 24, 2014
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CK.SCI.ENG.SE.1.Chemistry-Intermediate.13.3

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