Network atomic solids semiconductor

Even though silicon is metallic in appearance, it is not generally classified as a metal. The electrical conductivity of silicon is so much less than that of ordinary metals it is called a semiconductor. Silicon is an example of a network solid (see Figure 20-1)—it has the same atomic arrangement that occurs in diamond. Each silicon atom is surrounded by, and covalently bonded to, four other silicon atoms. Thus, the silicon crystal can be regarded as one giant molecule. 

All metals conduct electricity on account of the mobility of the electrons that bind the atoms together. Ionic, molecular, and network solids are typically electrical insulators or semiconductors (see Sections 3.f3 and 3.14), but there are notable exceptions, such as high-temperature superconductors, which are ionic or ceramic solids (see Box 5.2), and there is currently considerable interest in the electrical conductivity ol some organic polymers (see Box 19.1). 

We learn about solids in which the atoms are held together by extended networks of covalent bonds. We learn how the electronic structure and properties of semiconductors differ from those of metals. 

COVALENT-NETWORK SOLIDS (SECTION 12.7) Covalent-network solids consist of atoms held together in large networks by covalent bonds. These solids are much harder and have higher melting points than molecular solids. Important examples include diamond, where the carbons are tetrahedrally coordinated to each other, and graphite, where the sp -hybridized carbon atoms form hexagonal layers. Semiconductors are solids that do conduct electricity, but to a far lesser extent than metals. Insulators do not conduct electricity at all. 

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