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16.3: Ionic, Metallic, and Network Condensed Phases

Created by: CK-12

Lesson Objectives

The student will:

  • describe the forces holding molecules together in ionic compounds, metallic solids, and network solids.
  • describe the metallic bond and explain some of the characteristics that are due to metallic bonding.
  • describe ionic or network solid structures and explain some of the characteristics due to this type of solid bonding.
  • identify the type of solid given the characteristics of a solid, such as conductivity of solid and liquid phase, solubility in water, and malleability.


  • alloy
  • conductor
  • metallic bond


The intermolecular forces of attraction include London dispersion forces, dipole-dipole interactions, and hydrogen bonding. These forces can be found to varying degrees in all phases of matter. Depending on the strength of the forces and the temperature of the molecules, these forces affect the properties of the matter. For example, solids held together only by London dispersion forces are poor conductors of electricity and are not malleable. These forces also play a large role in solubility. The general rule is “like dissolves like.” This means that polar solvents will dissolve polar or ionic substances, but not non-polar substances. Non-polar solutes, however, can be readily dissolved in non-polar solvents.

Ionic Solids and Liquids

Ionics solids are held together by the electrostatic attraction between oppositely charged ions. The ions are formed into various types of crystal lattice structures depending on the comparative sizes of the ions and the charges on each. These ionic charges are full charges (+1, +2, -1, -2 and so on), so they are considerably stronger than either polar attractions or hydrogen bonds. This will cause the melting points for ionic substances to be quite high compared to the substances we have been considering. For example, the melting point of sodium chloride, \mathrm{NaCl}, is 801^\circ\mathrm{C}, and the melting point of calcium sulfate, \mathrm{CaSO}_4, is \mathrm{1460}^\circ\mathrm{C}.

In solid state, the ions in ionic solids are held firmly in position. There is no space large enough for the ions to move through, even if they could escape the forces of attraction. Since the ions cannot move, ionic solids are non-conductors of electricity. When the solid is melted to a liquid, however, the ions are free to migrate. Therefore, ionic liquids are good conductors of electric current. A conductor is a substance capable of transmitting electricity and/or heat. Ionic solids are usually quite soluble in water, and the water solutions of ionic solids are good conductors of electricity because of the freedom of the ions to migrate through the solution.

Metallic Solids

Of the 81 elements which can be clearly classified as metals, all of them except mercury are solids at room temperature. Any model that explains the bonding in metallic solids must account for the properties of metals. Metals typically 1) are excellent conductors of heat and electricity in both solid and liquid phase, 2) are malleable, 3) are white and shiny, 4) are not soluble in any common solvent, polar or non-polar, and 5) have a wide-range of melting points that are mostly higher than the melting points of polar solids.

The simplest proposed model that explains metallic behavior is the metallic bond, where a regular pattern of cations is surrounded by a “sea of electrons.” The metal ions occupy positions in a lattice structure while the mobile, free-moving sea of valence electrons occupy all of the overlapping valence shell area, as illustrated in the figure below. The metallic bond consists of non-directed bonds in which a “sea of electrons” surrounds all the bonded atoms. All the atoms are bonded in a single bond that includes the entire piece of metal.

For an animation showing the form of a metallic bond, (2a) see http://www.youtube.com/watch?v=ijw8OBt4btM (1:54).

For purposes of comparison, consider the covalent bonding in the trigonal planar molecule shown below. The central atom in this molecule contains three pairs of electrons. The electrons take positions as far away from each other as possible due to electrostatic repulsion. Therefore, the pairs of electrons maintain positions at angles of 120^\circ from each other. The atoms that share these electron pairs in the covalent bond must be placed so that the shared electrons are in the overlapped orbitals of both atoms. Therefore, these bonded atoms may not move with respect to each other. The atoms hold their relative positions because of the directional bonds. Neither the bonding electrons nor the atoms are free to move with respect to each others.

The model of the metallic bond, however, provides for mobile electrons that are free to move throughout the entire piece of metal, thus providing the means for the metal to be an excellent conductor of heat and electricity. Extra electrons pushed onto one side of the metal can easily move to the other side through the valence electron shells. The metal ions are not directionally bounded to their immediate neighbors, which allows them to be pushed to new positions without causing the bond to break. As long as an atom or ion is not separated from the piece of metal, its position can be significantly changed while remaining bounded to the other atoms. This malleability allows metal cubes to be pounded into flat sheets without breaking the bond.

The freedom of the electrons on the surface of a piece of metal also allows the metal to absorb and emit many frequencies of light, which accounts for the white, shiny appearance of many metals. The metals on the far left of the periodic table have the fewest electrons in the valence shells. As a result, the valence electrons in these metals would be least crowded, have the most freedom, and present the most complete metal character. The metals of families IA and IIA are excellent conductors, exceptionally malleable (soft enough to be cut with a spoon), and white and shiny in color.

Other elements can be introduced into a metallic crystal relatively easily to produce substances known as alloys. An alloy is a substance that contains a mixture of elements and has metallic properties. A fairly well-known alloy is brass, which is an alloy composed of approximately two-thirds copper atoms and one-third zinc atoms. Sterling silver is an alloy composed of about 93% silver and 7% copper. Iron is a metal that is commonly alloyed with carbon to produce steel. Carbon forms directional bonds with some of the iron atoms to make steel less malleable than pure iron. Steel with less than 0.2% carbon remains somewhat malleable and is used for nails and cables. In comparison, steel with around 0.6% carbon is harder and is used for railroad rails and structural steel beams, while steel with around 1.5% carbon is very hard and is used for tools and cutlery.

Network Solids

In some solids, all the atoms in the entire structure are bounded with covalent chemical bonds. These solids are a single giant molecule and are called network solids. When considering the strength of the various bonds and attractions that hold particles together in the solid state, the strongest is the covalent bond. Therefore, network solids have the highest melting points of all solids, as melting a network solid requires enough molecular motion to disrupt the covalent bonds. Network solids are not soluble in any common solvent. Some examples of network solids are graphite, diamond, mica, and asbestos. The structures of graphite and diamond are shown below.

Most network solids are non-conductors, although graphite is an exception and is a good conductor of electricity. The solid structure in graphite involves large, two-dimensional molecules of covalently bonded carbon atoms. The carbon atoms form flat sheets (like a sheet of paper) bounded in the fashion shown above. Layers of these sheets are laid on top of each other, and the sheets are held together by much weaker London dispersion forces. The sheets are extremely strong in the two dimensions bounded with covalent bonds, but the forces holding the sheets together are weak and easily broken. The flat sheets slide over each other readily, making graphite a good lubricant for metal parts.

The mineral mica is also bounded in this two-dimensional network style. Mica is found in nature and appears as a rock, but you can slide your fingernail between sheets and pull off large flats sheets of the rock. One type of mica, called muscovite mica, is transparent enough that you can see through several sheets. This material has been used to make small windows in furnaces that are transparent but won't melt.

Diamonds are giant molecules of carbon atoms bounded three-dimensionally in tetrahedral units. Every carbon atom in the structure is covalently bounded to four other carbon atoms. Diamonds are the hardest substance known and have one of the highest melting points of all substances.

Some forms of asbestos are one-dimensional network solids in which the atoms are bounded in a chain. The result is a fibrous molecule that can be woven into fabric (see example in Figure below). Due to its high melting point, asbestos fabric was used to make heat resistant materials (for example, fireman’s gloves, furnace padding, clutch plates) for many years until it was determined that asbestos fibers are hazardous if inhaled.

Asbestos (chrysotile)

Asbestos in the form of chrysotile. You can see fibers of asbestos in the upper left hand corner of the photo.

The following web site has data and explanatory reasons for the trends in melting and boiling points of some period 3 elements. http://www.creative-chemistry.org.uk/alevel/module1/trends8.htm

Amorphous Solids

Many important solids do not have the regular, repeating arrangement of atoms or molecules present in crystalline solids. Solids with irregular, unpredictable molecular organization are called amorphous solids. There are many solids that can form either crystalline or amorphous solids, depending on how rapidly the liquid is cooled. Very rapid cooling frequently results in an amorphous solid, whereas slow cooling produces crystalline solids. Amorphous solids have been described as appearing to have their molecules frozen in place before they have time to settle into an organized pattern. Examples of amorphous solids are glass, paper, plastics, cement, and rubber.

Amorphous solids are called solids because they maintain their shape and volume. Some researchers insist that certain amorphous solids will flow under pressure, which is a characteristic of liquids. Some antique windows have been found to be thicker at the bottom than at the top. Some chemists claim this is because the glass flowed downward very slowly over a hundred years due to the force of gravity. Other chemists claim that these antique glasses are of different thicknesses due to flaws in the glass making process of a hundred years ago, although this opinion doesn't explain why the thicker part of the glass was always at the bottom and never at the top.

Crystalline solids melt at sharply defined temperatures. In comparison, glass and some other amorphous solids only soften as they are heated. As a result, some authors refer to amorphous solids as “not true solids,” others call them “super-cooled liquids,” and still others insist that amorphous solids are absolutely solids.

Lesson Summary

  • Ionic solids are a type of solid in which the intermolecular forces of attraction is the electrostatic attractions due to the opposite charges of the ions.
  • In one type of solid formed by metallic atoms, a sea of electrons exerts a force of attraction on the positive ions (metallic bond).
  • Network solids have every atom in the structure attached to other atoms in the structure by covalent chemical bonds.
  • Amorphous solids are solids that cooled so rapidly, the molecules did not get into the tight, organized solid pattern. Due to their disorganized structure, amorphous solids have some properties more like liquids.

Further Reading / Supplemental Links

The learner.org website allows users to view streaming videos of the Annenberg series of chemistry videos. You are required to register before you can watch the videos, but there is no charge to register. The website has a video that apply to this lesson called “Metals” that details the value of accuracy and precision. In the video, malleability, ductility, and conductivity are examined, along with the methods for extracting metals from ores and blending alloys.

The following sites provide more information about crystals, glasses, and amorphous materials.

Review Questions

  1. Identify the most important type of inter-particle force present in the following solids that is responsible for binding the particles into a solid.
    1. \mathrm{Kr}
    2. \mathrm{BaCl}_2
    3. \mathrm{Mg}
    4. \mathrm{NaNO}_3
    5. \mathrm{BCl}_3
    6. diamond
    7. \mathrm{NH}_3
  2. Predict which substance in the following pairs would have the stronger force of attraction between molecules and justify your answer.
    1. \mathrm{NaI} or \mathrm{I}_2
    2. solid argon or solid sodium
    3. \mathrm{HF} or \mathrm{HBr}
  3. An unknown solid is not soluble in water or \mathrm{CCl}_4. The solid conducts electricity and has a melting point of 800^\circ\mathrm{C}. Identify the most likely attractive forces holding the particles in the solid state.
  4. An unknown solid is soluble in water but not in \mathrm{CCl}_4. The solid does not conduct electricity but its liquid does. The solid shatters when hammered and has a melting point of 1430^\circ\mathrm{C}. Identify the most likely attractive forces holding the particles in the solid state.
  5. Why would you expect ionic solids to have higher melting points that polar solids?
  6. Why does the melting point of water decrease with increasing surrounding pressure?
  7. When a drop of liquid is placed on a surface, the more spherical the drop remains is an indication of the strength of the intermolecular forces of attraction. The four drops in the sketch below represent mineral spirits (a non-polar molecule), acetone (a polar molecule), water (a hydrogen bond forming molecule), and mercury (a metallic liquid). Identify each of these liquids as one of the drops (A, B, C, or D).

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