Physical Properties and Crystal Types

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Physical Properties and Crystal Types Physical Properties and Crystal Types From personal experience you known that solids exhibit a wide range of physical properties. Some, such as diamond, are very hard; others, such as naphthalene (moth balls) or ice, are soft, by comparison, and are easily crushed. Some solids, such as salt crystals or iron, have high melting points, whereas others, such as candle wax, melt at low temperatures. Some conduct electricity well, but others are nonconducting. Physical properties such as these depend on the kinds and strengths of the attractive forces that hold the particles together in the solid. Even though exact predictions cannot be made some generalizations do exist. In making these generalizations, we can divide crystals into several types according to the kinds of particles located at sites in the lattice and the kinds of attractions that exist between the particles. Ionic Crystals Ionic crystals are hard, have high melting points, and are brittle. When they melt, the resulting liquids conduct electricity well. These properties reflect the strong attractive forces between ions of opposite charge as well as the repulsions that occur when ions of like charge are placed near each other. They are brittle and tend to shatter into smaller crystals when stressed. When a crystal is hammered or stressed, ions with like charges are forced into close proximity. The crystal then literaly self destructs because of electrostatic repulsion. Molecular Crystals Molecular crystals are solids in which the lattice sites are occupied either by atoms - as in solid argon or krypton - or by molecules - as in solid CO2, SO2, or H2O. Such solids tend to be soft and have low melting points because the particles in the solid experience relatively weak intermolecular attractions. The crystals are soft because little effort is needed to separate the particles or to move them past each other. The solid melts at low temperatures because the particles need little kinetic energy to break away from the solid. If the crystals contain only individual atoms, as in solid argon or krypton, or if they are composed of non-polar molecules, as in naphthalene, the only attractions between the molecules are the London forces. In crystals containing polar molecules, such as sulphur dioxide, the major forces that hold the particles together are dipole-dipole attractions. In crystals such as water the primary forces of attractions are due to hydrogen bonding. Covalent Crystals Covalent crystals are solids in which the lattice points are occupied by atoms that are covalently bonded to other atoms at neighbouring lattice sites. The result is a crystal that is essentially one gigantic molecule. These solids are sometimes called network solids because of the interlocking network of covalent bonds extending throughout the crystal in all directions. A typical example is the diamond, the structure of which is illustrated below. The above structure is diamond. Notice that each carbon atom is covalently bonded to four others at the corners of a tetrahedron. This structure extends throughout an entire diamond crystal. (In diamond, of course, all the atoms are identical. They are different shades here only to make it easier to visualize the structure.) Covalent crystals tend to be hard and to have very high melting points because of the strong attractions between covalently bonded atoms. They do not conduct electricity because the electrons are bound too tightly to the bonds. Other examples of covalent crystals are quartz (SiO2 - typical grains of sand) and silicon carbide (SiC - a common abrasive used in sandpaper). Metallic crystals Metallic crystals have properties that are quite different from those of the other three types of crystals above. They conduct heat and electricity well, and they have the lustre that we characteristically associate with metals. A number of different models have been developed to describe metals. The simplest one views the crystal as having positive ions at the lattice positions which are surrounded by electrons in a cloud that spreads throughout the entire solid. The electrons in this cloud belong to no single positive ion, but rather to the crystal as a whole. Because the electrons aren't localized on any one atom, they are free to move easily, which accounts for the electrical conductivity of metals. By their movement, the electrons can also transmit kinetic energy rapidly thorough the solid, so metals are also good conductors of heat. Because the electrons are free to move easily, even the lattice points are moveable or deformable and this helps explain the malleability of metals. Even though the lattice points and electrons are free to move there are some strong attractive forces at work and these help explain ductility in some metals. This model also explains the lustre of metals. When light shines on the metal, the loosely held electrons vibrate easily and readily reemit the light with essentially the same frequency and intensity. It is not possible to make simple generalizations about the melting points of metals. Some have very high melting points, like tungsten, whereas others like mercury have quite low melting points. To some degree the melting point depends on the charge of the positive ions in the metallic crystals. The ions of Group IA tend to exist as cations with a +1 charge, and they are only weakly attracted to the "electron sea" that surrounds them. Atoms of Group IIA metals, however, tend to form cations with a +2 charge. These more highly charged ions are attached more strongly to the surrounding electron sea, so the Group IIA metals have higher melting points than their neighbours in Group IA. For example magnesium melts at 650oC but sodium melts at 98oC. Tungsten which has a very high melting point, must have very strong attractions between their atoms, which suggests that there probably is some covalent bonding in them as well. Crystal Types Type  Particles Occupying  Lattice Sites   Type of  Attractive Force Typical  Examples  Typical  Properties Ionic Positive & negative ions Attractions between  opposite ions  NaCl,  NaNO3,  CaCl2 Hard; high melting points;  non-conductors of electricity as solids  but good conductors when   melted   Molecular Atoms or Molecules Dipole-dipole attractions  London forces  Hydrogen bonding HCl  SO2 N2, Ar ,  CH4 H2O    Soft, low melting  points;  nonconductors of electricity  in both solid and  melted form Covalent (network) Atoms  Covalent bonds   between atoms Diamond,  SiC, SiO2 Very hard; very high melting pts. nonconductors of electricity   Metallic  Atoms Positive Attractions  between ions and  an electron cloud  that extends throughout  the crystal in both  solid and liquid Cu Ag Fe Na Hg Au Range from very hard to soft; melting points range from high  to low; conduct electricity and  heat well; have a characteristic

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