When metals combine with each other, the bonding is usually described as metallic bonding you could've guessed that. In this model, each metal atom donates one or more of its valence electrons to make an electron sea that surrounds all of the atoms, holding the substance together by the attraction between the metal cations and the negatively charged electrons. Since the electrons in the electron sea can move freely, metals conduct electricity very easily, unlike molecules, where the electrons are more localized.
Metal atoms can move past each other more easily than those in ionic compounds which are held in fixed positions by the attractions between cations and anions , allowing the metal to be hammered into sheets or drawn into wire.
Different metals can be combined very easily to make alloys , which can have much different physical properties from their constituent metals. Steel is an alloy of iron and carbon, which is much harder than iron itself; chromium, vanadium, nickel, and other metals are also often added to iron to make steels of various types.
Brass is an alloy of copper and zinc which is used in plumbing fixtures, electrical parts, and musical instruments. Bronze is an alloy of copper and tin, which is much harder than copper; when bronze was discovered by ancient civilizations, it marked a significant step forward from the use of less durable stone tools.
Key: metal metalloid nonmetal. Non-metals can gain electrons completely to give anions and hence can form ionic bonds. But they can also share electron pairs with other non-metal atoms and can form covalent bonds too. There are two types of atomic bonds - ionic bonds and covalent bonds. They differ in their structure and properties. Covalent bonds consist of pairs of electrons shared by two atoms, and bind the atoms in a fixed orientation. Pure silicon is necessary in semiconductor electronic devices such as transistors, computer chips, and solar cells.
Like some metals, passivation of silicon occurs due the formation of a very thin film of oxide primarily silicon dioxide, SiO 2. Silicon dioxide is soluble in hot aqueous base; thus, strong bases destroy the passivation. Removal of the passivation layer allows the base to dissolve the silicon, forming hydrogen gas and silicate anions. For example:. Silicon reacts with halogens at high temperatures, forming volatile tetrahalides, such as SiF 4.
Unlike carbon, silicon does not readily form double or triple bonds. Silicon compounds of the general formula SiX 4 , where X is a highly electronegative group, can act as Lewis acids to form six-coordinate silicon. Antimony reacts readily with stoichiometric amounts of fluorine, chlorine, bromine, or iodine, yielding trihalides or, with excess fluorine or chlorine, forming the pentahalides SbF 5 and SbCl 5.
Depending on the stoichiometry, it forms antimony III sulfide, Sb 2 S 3 , or antimony V sulfide when heated with sulfur. As expected, the metallic nature of the element is greater than that of arsenic, which lies immediately above it in group These nonpolar molecules contain boron with sp 2 hybridization and a trigonal planar molecular geometry.
The fluoride and chloride compounds are colorless gasses, the bromide is a liquid, and the iodide is a white crystalline solid. Except for boron trifluoride, the boron trihalides readily hydrolyze in water to form boric acid and the corresponding hydrohalic acid. Boron trichloride reacts according to the equation:. Boron trifluoride reacts with hydrofluoric acid, to yield a solution of fluoroboric acid, HBF 4 :. In this reaction, the BF 3 molecule acts as the Lewis acid electron pair acceptor and accepts a pair of electrons from a fluoride ion:.
All the tetrahalides of silicon, SiX 4 , have been prepared. Silicon tetrachloride can be prepared by direct chlorination at elevated temperatures or by heating silicon dioxide with chlorine and carbon:.
It is possible to prepare silicon tetrafluoride by the reaction of silicon dioxide with hydrofluoric acid:. Hydrofluoric acid is the only common acid that will react with silicon dioxide or silicates.
This reaction occurs because the silicon-fluorine bond is the only bond that silicon forms that is stronger than the silicon-oxygen bond. For this reason, it is possible to store all common acids, other than hydrofluoric acid, in glass containers. Except for silicon tetrafluoride, silicon halides are extremely sensitive to water. Upon exposure to water, SiCl 4 reacts rapidly with hydroxide groups, replacing all four chlorine atoms to produce unstable orthosilicic acid, Si OH 4 or H 4 SiO 4 , which slowly decomposes into SiO 2.
Boric oxide is necessary for the production of heat-resistant borosilicate glass, like that shown in Figure 4 and certain optical glasses. Boric oxide dissolves in hot water to form boric acid, B OH 3 :.
Figure 4. Laboratory glassware, such as Pyrex and Kimax, is made of borosilicate glass because it does not break when heated. The inclusion of borates in the glass helps to mediate the effects of thermal expansion and contraction. This reduces the likelihood of thermal shock, which causes silicate glass to crack upon rapid heating or cooling.
Figure 5. The boron atom in B OH 3 is sp 2 hybridized and is located at the center of an equilateral triangle with oxygen atoms at the corners. In solid B OH 3 , hydrogen bonding holds these triangular units together. Boric acid, shown in Figure 5, is a very weak acid that does not act as a proton donor but rather as a Lewis acid, accepting an unshared pair of electrons from the Lewis base OH — :. Complete water loss, at still higher temperatures, results in boric oxide.
Borates are salts of the oxyacids of boron. Borates result from the reactions of a base with an oxyacid or from the fusion of boric acid or boric oxide with a metal oxide or hydroxide.
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