# Structure and General Properties of the Metalloids (2023)

### Learning Outcomes

• Describe the general preparation, properties, and uses of the metalloids
• Describe the preparation, properties, and compounds of boron and silicon

A series of six elements called the metalloids separate the metals from the nonmetals in the periodic table. The metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. These elements look metallic; however, they do not conduct electricity as well as metals so they are semiconductors. They are semiconductors because their electrons are more tightly bound to their nuclei than are those of metallic conductors. Their chemical behavior falls between that of metals and nonmetals. For example, the pure metalloids form covalent crystals like the nonmetals, but like the metals, they generally do not form monatomic anions. This intermediate behavior is in part due to their intermediate electronegativity values. In this section, we will briefly discuss the chemical behavior of metalloids and deal with two of these elements—boron and silicon—in more detail.

The metalloid boron exhibits many similarities to its neighbor carbon and its diagonal neighbor silicon. All three elements form covalent compounds. However, boron has one distinct difference in that its 2s22p1 outer electron structure gives it one less valence electron than it has valence orbitals. Although boron exhibits an oxidation state of 3+ in most of its stable compounds, this electron deficiency provides boron with the ability to form other, sometimes fractional, oxidation states, which occur, for example, in the boron hydrides.

Silicon has the valence shell electron configuration 3s23p2, and it commonly forms tetrahedral structures in which it is sp3 hybridized with a formal oxidation state of 4+. The major differences between the chemistry of carbon and silicon result from the relative strength of the carbon-carbon bond, carbon’s ability to form stable bonds to itself, and the presence of the empty 3d valence-shell orbitals in silicon. Silicon’s empty d orbitals and boron’s empty p orbital enable tetrahedral silicon compounds and trigonal planar boron compounds to act as Lewis acids. Carbon, on the other hand, has no available valence shell orbitals; tetrahedral carbon compounds cannot act as Lewis acids. Germanium is very similar to silicon in its chemical behavior.

Arsenic and antimony generally form compounds in which an oxidation state of 3+ or 5+ is exhibited; however, arsenic can form arsenides with an oxidation state of 3-. These elements tarnish only slightly in dry air but readily oxidize when warmed.

Tellurium combines directly with most elements. The most stable tellurium compounds are the tellurides—salts of Te2- formed with active metals and lanthanides—and compounds with oxygen, fluorine, and chlorine, in which tellurium normally exhibits an oxidation state 2+ or 4+. Although tellurium(VI) compounds are known (for example, TeF6), there is a marked resistance to oxidation to this maximum group oxidation state.

## Structures of the Metalloids

Covalent bonding is the key to the crystal structures of the metalloids. In this regard, these elements resemble nonmetals in their behavior.

Elemental silicon, germanium, arsenic, antimony, and tellurium are lustrous, metallic-looking solids. Silicon and germanium crystallize with a diamond structure. Each atom within the crystal has covalent bonds to four neighboring atoms at the corners of a regular tetrahedron. Single crystals of silicon and germanium are giant, three-dimensional molecules. There are several allotropes of arsenic with the most stable being layer like and containing puckered sheets of arsenic atoms. Each arsenic atom forms covalent bonds to three other atoms within the sheet. The crystal structure of antimony is similar to that of arsenic, both shown in Figure1. The structures of arsenic and antimony are similar to the structure of graphite, covered later in this chapter. Tellurium forms crystals that contain infinite spiral chains of tellurium atoms. Each atom in the chain bonds to two other atoms.

Explore a cubic diamond crystal structure. (Note that the video has no narration. You can access the audio description using the widget below the video.)

Figure1. (a) Arsenic and (b) antimony have a layered structure similar to that of (c) graphite, except that the layers are puckered rather than planar. (d) Elemental tellurium forms spiral chains.

(Video) Lesson 7.3.3 The Structure and Properties of Metalloids

Figure2. An icosahedron is a symmetrical, solid shape with 20 faces, each of which is an equilateral triangle. The faces meet at 12 corners.

Pure crystalline boron is transparent. The crystals consist of icosahedra, as shown in Figure2, with a boron atom at each corner. In the most common form of boron, the icosahedra pack together in a manner similar to the cubic closest packing of spheres. All boron-boron bonds within each icosahedron are identical and are approximately 176 pm in length. In the different forms of boron, there are different arrangements and connections between the icosahedra.

The name silicon is derived from the Latin word for flint, silex. The metalloid silicon readily forms compounds containing Si-O-Si bonds, which are of prime importance in the mineral world. This bonding capability is in contrast to the nonmetal carbon, whose ability to form carbon-carbon bonds gives it prime importance in the plant and animal worlds.

## Occurrence, Preparation, and Compounds of Boron and Silicon

Boron constitutes less than 0.001% by weight of the earth’s crust. In nature, it only occurs in compounds with oxygen. Boron is widely distributed in volcanic regions as boric acid, B(OH)3, and in dry lake regions, including the desert areas of California, as borates and salts of boron oxyacids, such as borax, Na2B4O7•10H2O.

Elemental boron is chemically inert at room temperature, reacting with only fluorine and oxygen to form boron trifluoride, BF3, and boric oxide, B2O3, respectively. At higher temperatures, boron reacts with all nonmetals, except tellurium and the noble gases, and with nearly all metals; it oxidizes to B2O3 when heated with concentrated nitric or sulfuric acid. Boron does not react with nonoxidizing acids. Many boron compounds react readily with water to give boric acid, B(OH)3 (sometimes written as H3BO3).

Reduction of boric oxide with magnesium powder forms boron (95–98.5% pure) as a brown, amorphous powder:

${\text{B}}_{2}{\text{O}}_{3}\left(s\right)+\text{3Mg}\left(s\right)\rightarrow\text{2B}\left(s\right)+\text{3MgO}\left(s\right)$

An amorphous substance is a material that appears to be a solid, but does not have a long-range order like a true solid. Treatment with hydrochloric acid removes the magnesium oxide. Further purification of the boron begins with conversion of the impure boron into boron trichloride. The next step is to heat a mixture of boron trichloride and hydrogen:

${\text{2BCl}}_{3}\left(g\right)+{\text{3H}}_{2}\left(g\right)\stackrel{1500^{\circ}\text{C}}{\longrightarrow}\text{2B}\left(s\right)+\text{6HCl}\left(g\right)$$\qquad\Delta H^{\circ}=\text{253.7 kJ}$

Silicon makes up nearly one-fourth of the mass of the earth’s crust—second in abundance only to oxygen. The crust is composed almost entirely of minerals in which the silicon atoms are at the center of the silicon-oxygen tetrahedron, which connect in a variety of ways to produce, among other things, chains, layers, and three-dimensional frameworks. These minerals constitute the bulk of most common rocks, soil, and clays. In addition, materials such as bricks, ceramics, and glasses contain silicon compounds.

It is possible to produce silicon by the high-temperature reduction of silicon dioxide with strong reducing agents, such as carbon and magnesium:

$\begin{array}{rll}{\text{SiO}}_{2}\left(s\right)+\text{2C}\left(s\right)&\stackrel{\Delta}{\longrightarrow}&\text{Si}\left(s\right)+\text{2CO}\left(g\right)\\ {\text{SiO}}_{2}\left(s\right)+\text{2Mg}\left(s\right)&\stackrel{\Delta}{\longrightarrow}&\text{Si}\left(s\right)+\text{2MgO}\left(s\right)\end{array}$

Figure3. A zone-refining apparatus used to purify silicon.

Extremely pure silicon is necessary for the manufacture of semiconductor electronic devices. This process begins with the conversion of impure silicon into silicon tetrahalides, or silane (SiH4), followed by decomposition at high temperatures. Zone refining, illustrated in Figure3, completes the purification. In this method, a rod of silicon is heated at one end by a heat source that produces a thin cross-section of molten silicon. Slowly lowering the rod through the heat source moves the molten zone from one end of the rod to other. As this thin, molten region moves, impurities in the silicon dissolve in the liquid silicon and move with the molten region. Ultimately, the impurities move to one end of the rod, which is then cut off.

This highly purified silicon, containing no more than one part impurity per million parts of silicon, is the most important element in the computer industry. Pure silicon is necessary in semiconductor electronic devices such as transistors, computer chips, and solar cells.

(Video) Metals, Nonmetals & Metalloids

Like some metals, passivation of silicon occurs due the formation of a very thin film of oxide (primarily silicon dioxide, SiO2). 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:

$\text{Si}\left(s\right)+{\text{4OH}}^{-}\left(aq\right)\rightarrow{\text{SiO}}_{4}{}^{4-}\left(aq\right)+{\text{2H}}_{2}\left(g\right)$

Silicon reacts with halogens at high temperatures, forming volatile tetrahalides, such as SiF4.

Unlike carbon, silicon does not readily form double or triple bonds. Silicon compounds of the general formula SiX4, where X is a highly electronegative group, can act as Lewis acids to form six-coordinate silicon. For example, silicon tetrafluoride, SiF4, reacts with sodium fluoride to yield Na2[SiF6], which contains the octahedral ${\left[{\text{SiF}}_{6}\right]}^{2-}$ ion in which silicon is sp3d2 hybridized:

$\text{2NaF}\left(s\right)+{\text{SiF}}_{4}\left(g\right)\rightarrow{\text{Na}}_{2}{\text{SiF}}_{6}\left(s\right)$

Antimony reacts readily with stoichiometric amounts of fluorine, chlorine, bromine, or iodine, yielding trihalides or, with excess fluorine or chlorine, forming the pentahalides SbF5 and SbCl5. Depending on the stoichiometry, it forms antimony(III) sulfide, Sb2S3, 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 15.

### Boron and Silicon Halides

Boron trihalides—BF3, BCl3, BBr3, and BI3—can be prepared by the direct reaction of the elements. These nonpolar molecules contain boron with sp2 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:

${\text{BCl}}_{3}\left(g\right)+{\text{3H}}_{2}\text{O}\left(l\right)\rightarrow\text{B}{\text{(OH)}}_{3}\left(aq\right)+\text{3HCl}\left(aq\right)$

Boron trifluoride reacts with hydrofluoric acid, to yield a solution of fluoroboric acid, HBF4:

${\text{BF}}_{3}\left(aq\right)+\text{HF}\left(aq\right)+{\text{H}}_{2}\text{O}\left(l\right)\rightarrow{\text{H}}_{3}{\text{O}}^{+}\left(aq\right)+{\text{BF}}_{4}{}^{-}\left(aq\right)$

In this reaction, the BF3 molecule acts as the Lewis acid (electron pair acceptor) and accepts a pair of electrons from a fluoride ion:

All the tetrahalides of silicon, SiX4, have been prepared. Silicon tetrachloride can be prepared by direct chlorination at elevated temperatures or by heating silicon dioxide with chlorine and carbon:

${\text{SiO}}_{2}\left(s\right)+\text{2C}\left(s\right)+{\text{2Cl}}_{2}\left(g\right)\stackrel{\Delta}{\longrightarrow}{\text{SiCl}}_{4}\left(g\right)+\text{2CO}\left(g\right)$

Silicon tetrachloride is a covalent tetrahedral molecule, which is a nonpolar, low-boiling (57 °C), colorless liquid.

It is possible to prepare silicon tetrafluoride by the reaction of silicon dioxide with hydrofluoric acid:

(Video) What Are Allotropes of Metalloids and Metals | Properties of Matter | Chemistry | FuseSchool

${\text{SiO}}_{2}\left(s\right)+\text{4HF}\left(g\right)\rightarrow{\text{SiF}}_{4}\left(g\right)+{\text{2H}}_{2}\text{O}\left(l\right)\qquad\Delta H^{\circ}=-\text{191.2 kJ}$

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, SiCl4 reacts rapidly with hydroxide groups, replacing all four chlorine atoms to produce unstable orthosilicic acid, Si(OH)4 or H4SiO4, which slowly decomposes into SiO2.

### Boron and Silicon Oxides and Derivatives

Boron burns at 700 °C in oxygen, forming boric oxide, B2O3. Boric oxide is necessary for the production of heat-resistant borosilicate glass, like that shown in Figure4 and certain optical glasses. Boric oxide dissolves in hot water to form boric acid, B(OH)3:

${\text{B}}_{2}{\text{O}}_{3}\left(s\right)+{\text{3H}}_{2}\text{O}\left(l\right)\rightarrow\text{2B}{\text{(OH)}}_{3}\left(aq\right)$

Figure4. 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. (credit: “Tweenk”/Wikimedia Commons)

Figure5. Boric acid has a planar structure with three –OH groups spread out equally at 120° angles from each other.

The boron atom in B(OH)3 is sp2 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 Figure5, 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:

$\text{B}{\left(\text{OH}\right)}_{3}\left(aq\right)+{\text{2H}}_{2}\text{O}\left(l\right)\rightleftharpoons\text{B}{\left(\text{OH}\right)}_{4}{}^{-}\left(aq\right)+{\text{H}}_{3}{\text{O}}^{+}\left(aq\right)\qquad{K}_{\text{a}}=5.8\times {10}^{-10}$

Heating boric acid to 100 °C causes molecules of water to split out between pairs of adjacent –OH groups to form metaboric acid, HBO2. At about 150 °C, additional B-O-B linkages form, connecting the BO3 groups together with shared oxygen atoms to form tetraboric acid, H2B4O7. 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. Borate anions range from the simple trigonal planar ${\text{BO}}_{3}{}^{3-}$ ion to complex species containing chains and rings of three- and four-coordinated boron atoms. The structures of the anions found in CaB2O4, K[B5O6(OH)4]•2H2O (commonly written KB5O8•4H2O) and Na2[B4O5(OH)4]•8H2O (commonly written Na2B4O7•10H2O) are shown in Figure6. Commercially, the most important borate is borax, Na2[B4O5(OH)4]•8H2O, which is an important component of some laundry detergents. Most of the supply of borax comes directly from dry lakes, such as Searles Lake in California, or is prepared from kernite, Na2B4O7•4H2O.

(Video) Elements and their Physical Properties - Lesson Preview

Figure6. The borate anions are (a) CaB2O4, (b) KB5O8•4H2O, and (c) Na2B4O7•10H2O. The anion in CaB2O4 is an “infinite” chain.

Silicon dioxide, silica, occurs in both crystalline and amorphous forms. The usual crystalline form of silicon dioxide is quartz, a hard, brittle, clear, colorless solid. It is useful in many ways—for architectural decorations, semiprecious jewels, and frequency control in radio transmitters. Silica takes many crystalline forms, or polymorphs, in nature. Trace amounts of Fe3+ in quartz give amethyst its characteristic purple color. The term quartz is also used for articles such as tubing and lenses that are manufactured from amorphous silica. Opal is a naturally occurring form of amorphous silica.

The contrast in structure and physical properties between silicon dioxide and carbon dioxide is interesting, as illustrated in Figure7. Solid carbon dioxide (dry ice) contains single CO2 molecules with each of the two oxygen atoms attached to the carbon atom by double bonds. Very weak intermolecular forces hold the molecules together in the crystal. The volatility of dry ice reflect these weak forces between molecules. In contrast, silicon dioxide is a covalent network solid. In silicon dioxide, each silicon atom links to four oxygen atoms by single bonds directed toward the corners of a regular tetrahedron, and SiO4 tetrahedra share oxygen atoms. This arrangement gives a three dimensional, continuous, silicon-oxygen network. A quartz crystal is a macromolecule of silicon dioxide. The difference between these two compounds is the ability of the group 14 elements to form strong π bonds. Second-period elements, such as carbon, form very strong π bonds, which is why carbon dioxide forms small molecules with strong double bonds. Elements below the second period, such as silicon, do not form π bonds as readily as second-period elements, and when they do form, the π bonds are weaker than those formed by second-period elements. For this reason, silicon dioxide does not contain π bonds but only σ bonds.

Figure 7. Because carbon tends to form double and triple bonds and silicon does not, (a) carbon dioxide is a discrete molecule with two C=O double bonds and (b) silicon dioxide is an infinite network of oxygen atoms bridging between silicon atoms with each silicon atom possessing four Si-O single bonds. (credit a photo: modification of work by Erica Gerdes; credit b photo: modification of work by Didier Descouens)

At 1600 °C, quartz melts to yield a viscous liquid. When the liquid cools, it does not crystallize readily but usually supercools and forms a glass, also called silica. The SiO4 tetrahedra in glassy silica have a random arrangement characteristic of supercooled liquids, and the glass has some very useful properties. Silica is highly transparent to both visible and ultraviolet light. For this reason, it is important in the manufacture of lamps that give radiation rich in ultraviolet light and in certain optical instruments that operate with ultraviolet light. The coefficient of expansion of silica glass is very low; therefore, rapid temperature changes do not cause it to fracture. CorningWare and other ceramic cookware contain amorphous silica.

Silicates are salts containing anions composed of silicon and oxygen. In nearly all silicates, sp3-hybridized silicon atoms occur at the centers of tetrahedra with oxygen at the corners. There is a variation in the silicon-to-oxygen ratio that occurs because silicon-oxygen tetrahedra may exist as discrete, independent units or may share oxygen atoms at corners in a variety of ways. In addition, the presence of a variety of cations gives rise to the large number of silicate minerals.

Many ceramics are composed of silicates. By including small amounts of other compounds, it is possible to modify the physical properties of the silicate materials to produce ceramics with useful characteristics.

### Key Concepts and Summary

The elements boron, silicon, germanium, arsenic, antimony, and tellurium separate the metals from the nonmetals in the periodic table. These elements, called metalloids or sometimes semimetals, exhibit properties characteristic of both metals and nonmetals. The structures of these elements are similar in many ways to those of nonmetals, but the elements are electrical semiconductors.

### Try It

1. Give the hybridization of the metalloid and the molecular geometry for each of the following compounds or ions. You may wish to review the chapters on chemical bonding and advanced covalent bonding for relevant examples.
1. GeH4
2. SbF3
3. Te(OH)6
4. H2Te
5. GeF2
6. TeCl4
7. ${\text{SiF}}_{6}{}^{2-}$
8. SbCl5
9. TeF6
2. Write a Lewis structure for each of the following molecules or ions. You may wish to review the chapter on chemical bonding.
1. H3BPH3
2. ${\text{BF}}_{4}{}^{-}$
3. BBr3
4. B(CH3)3
5. B(OH)3
3. Describe the hybridization of boron and the molecular structure about the boron in each of the following:
1. H3BPH3
2. ${\text{BF}}_{4}{}^{-}$
3. BBr3
4. B(CH3)3
5. B(OH)3
4. Using only the periodic table, write the complete electron configuration for silicon, including any empty orbitals in the valence shell. You may wish to review the chapter on electronic structure.
5. Write a Lewis structure for each of the following molecules and ions:
1. (CH3)3SiH
2. ${\text{SiO}}_{4}{}^{4-}$
3. Si2H6
4. Si(OH)4
5. ${\text{SiF}}_{6}{}^{2-}$
6. Describe the hybridization of silicon and the molecular structure of the following molecules and ions:
1. (CH3)3SiH
2. ${\text{SiO}}_{4}{}^{4-}$
3. Si2H6
4. Si(OH)4
5. ${\text{SiF}}_{6}{}^{2-}$
7. Describe the hybridization and the bonding of a silicon atom in elemental silicon.
8. Classify each of the following molecules as polar or nonpolar. You may wish to review the chapter on chemical bonding.
1. SiH4
2. Si2H6
3. SiCl3H
4. SiF4
5. SiCl2F2
9. Silicon reacts with sulfur at elevated temperatures. If 0.0923 g of silicon reacts with sulfur to give 0.3030 g of silicon sulfide, determine the empirical formula of silicon sulfide.
10. Name each of the following compounds:
1. TeO2
2. Sb2S3
3. GeF4
4. SiH4
5. GeH4
11. Write a balanced equation for the reaction of elemental boron with each of the following (most of these reactions require high temperature):
1. F2
2. O2
3. S
4. Se
5. Br2
12. Why is boron limited to a maximum coordination number of four in its compounds?
13. Write a formula for each of the following compounds:
1. silicon dioxide
2. silicon tetraiodide
3. silane
4. silicon carbide
5. magnesium silicide
14. From the data given in Ionization Constants of Weak Bases , determine the standard enthalpy change and the standard free energy change for each of the following reactions:
1. ${\text{BF}}_{3}\left(g\right)+{\text{3H}}_{2}\text{O}\left(l\right)\rightarrow\text{B}{\left(\text{OH}\right)}_{3}\left(s\right)+\text{3HF}\left(g\right)$
2. ${\text{BCl}}_{3}\left(g\right)+{\text{3 H}}_{2}\text{O}\left(l\right)\rightarrow\text{B}{\left(\text{OH}\right)}_{3}\left(s\right)+\text{3 HCl}\left(g\right)$
3. ${\text{B}}_{2}{\text{H}}_{6}\left(g\right)+{\text{6 H}}_{2}\text{O}\left(l\right)\rightarrow\text{2B}{\left(\text{OH}\right)}_{3}\left(s\right)+{\text{6H}}_{2}\left(g\right)$
15. A hydride of silicon prepared by the reaction of Mg2Si with acid exerted a pressure of 306 torr at 26 °C in a bulb with a volume of 57.0 mL. If the mass of the hydride was 0.0861 g, what is its molecular mass? What is the molecular formula for the hydride?
16. Suppose you discovered a diamond completely encased in a silicate rock. How would you chemically free the diamond without harming it?

Show Selected Solutions

## Glossary

amorphous: solid material such as a glass that does not have a regular repeating component to its three-dimensional structure; a solid but not a crystal

borate
compound containing boron-oxygen bonds, typically with clusters or chains as a part of the chemical structure

polymorph: variation in crystalline structure that results in different physical properties for the resulting compound

(Video) Metalloids

silicate: compound containing silicon-oxygen bonds, with silicate tetrahedra connected in rings, sheets, or three-dimensional networks, depending on the other elements involved in the formation of the compounds

## FAQs

### What are the general properties of a metalloid? ›

This article will describe the six most important properties of metalloids and list some key metalloids characteristics.
• Metalloids Are Solids. ...
• Metalloids Have a Metallic Luster and Appear to be Metals. ...
• Metalloids Are Brittle and Easily Broken. ...
• Metalloids Have the Ability To Conduct Electricity, but Not As Well as Metals.

What is the structure of metalloid? ›

Structures of the Metalloids. Covalent bonding is the key to the crystal structures of the metalloids. In this regard, these elements resemble nonmetals in their behavior. Elemental silicon, germanium, arsenic, antimony, and tellurium are lustrous, metallic-looking solids.

What are 5 properties of metalloids? ›

Five Main Properties of Metalloids
• Properties intermediate between metals and nonmetals.
• Physical appearance similar to metals.
• Semi-conductors of electricity.
• Brittle.
• Chemical properties are more similar to nonmetals than to metals.
20 Jul 2022

Where are metalloids found Quizizz? ›

For metalloids on the periodic table, how do the group number and the period number relate? The lower the group number, the lower the period numbers, so the metalloids are found in a diagonal moving down from left to right.

What is metalloid short answer? ›

metalloid, in chemistry, an imprecise term used to describe a chemical element that forms a simple substance having properties intermediate between those of a typical metal and a typical nonmetal.

What are the general properties of metals? ›

Properties of metals
• high melting points.
• good conductors of electricity.
• good conductors of heat.
• high density.
• malleable.
• ductile.

Why is it called metalloid? ›

The origin and usage of the term metalloid is convoluted. Its origin lies in attempts, dating from antiquity, to describe metals and to distinguish between typical and less typical forms. It was first applied to metals that floated on water (lithium, sodium and potassium), and then more popularly to nonmetals.

Why are metalloids important? ›

The ability of the metalloids to conduct electricity and heat is far better than the nonmetals, for example, diamond , which are insulators. As such, the metalloids play a major role in the electronics industry, and modern society is built around the properties of the metalloids and their compounds.

What are the 10 example of metalloids? ›

Copper,Sulphur,Aluminium,Oxygen,Silicon,Nitrogen,Germanium,Mercury,Chlorine,Sodium.

What are metalloids give Example answer? ›

Solution : Metalloids are elements that show physical and chemical properties of metals and non metals both. Elements like boron, silicon, germanium, arsenic, antimony, tellurium are recognized as metalloids.

### Are there 10 metalloids? ›

In the modern periodic table there are six metalloids which are boron , silicon , germanium , arsenic , antimony and tellurium . Metalloids usually look like metals but they behave largely like non metals.

How is metalloids formed? ›

Alloys formed when combined with transition metals are extremely well-represented when it comes to the lighter metalloids. Boron has the ability to form intermetallic compounds. This element also has the ability to form alloys with these MnB composition metals if the value of n is greater than 2.

Where are metalloids used? ›

Metalloids are usually too brittle to have any structural uses. They and their compounds are used in alloys, biological agents, catalysts, flame retardants, glasses, optical storage and optoelectronics, pyrotechnics, semiconductors, and electronics.

What are 3 facts about metalloids? ›

Metalloids share many similar properties including:
• They appear to be metal in appearance, but are brittle.
• They can generally form alloys with metals.
• Some metalloids such as silicon and germanium become electrical conductors under special conditions. ...
• They are solids under standard conditions.

How many metalloids are there? ›

6 metalloids. Listed in order of increasing atomic number, the metalloids are: Boron (B) Silicon (Si)

What is metalloid metal? ›

The metals are to the left of the line (except for hydrogen, which is a nonmetal), the nonmetals are to the right of the line, and the elements immediately adjacent to the line are the metalloids. When elements combine to form compounds, there are two major types of bonding that can result.

What is the state of metalloids? ›

Metalloids are all solid at room temperature. They can form alloys with other metals. Some metalloids, such as silicon and germanium, can act as electrical conductors under the right conditions, thus they are called semiconductors.

Are metalloids good conductors? ›

Metalloids typically conduct heat and electricity better than nonmetals but not as well as metals.

What are the general properties of materials? ›

Physical properties of materials
• density.
• melting point.
• thermal conductivity.
• electrical conductivity (resistivity)
• thermal expansion.
• corrosion resistance.

Which properties do metalloids share with metals? ›

Metalloids share characteristics of both metals and non-metals and are also called semimetals. Metalloids are typically semi-conductors, which means that they both insulate and conduct electricity.

### What are 5 examples of metalloids? ›

The elements classified as metalloids are - boron, silicon, germanium, arsenic, antimony, and tellurium (and sometimes bismuth, polonium, and astatine).

What are metalloids 9 examples? ›

Example of metalloids:
• Only a few elements are metalloids such as Germanium , Silicon , Arsenic , Antimony , and Tellurium .
• Germanium is act as a metalloid because it shows both metallic and nonmetallic properties.

Is 13 a metalloid? ›

Therefore, in Group 13 Boron is a metalloid.

What are metalloids called? ›

The elements boron, silicon, germanium, arsenic, antimony, and tellurium separate the metals from the nonmetals in the periodic table. These elements, called metalloids or sometimes semimetals, exhibit properties characteristic of both metals and nonmetals.

What is the most useful property of a metalloid? ›

The most useful property of metalloids is their varying ability to conduct electricity. Whether or not a metalloid conducts electricity can depend on the temperature or the exposure to light. For this reason metalloids, such as silicon or germanium, are used to make semiconductors.

Are there 7 or 8 metalloids? ›

The number of metalloids in the modern periodic table are 6. The names are : boron, silicon, germanium, polonium, tellurium, arsenic.

What are the six metalloids? ›

The six commonly recognised metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. Five elements are less frequently so classified: carbon, aluminium, selenium, polonium, and astatine.

What are metalloids in one sentence? ›

Iodine is most commonly given as an example of a metalloid because of its metallic appearance. Some sources refer to polonium as a metalloid, although it has some metallic properties. Astatine is usually classified as either a nonmetal or a metalloid. It is a brittle and easily pulverized metalloid.

Which one is an example of metalloid? ›

Arsenic and antimony are the examples of metalloids.

Is water a metalloid? ›

Water is made up of the two nonmetals oxygen and hydrogen. Even counting the halogens and the noble gases there are only 18 elements in the periodic table that are considered nonmetals. Many nonmetals can gain metallic properties under very high pressures.

### Which group contains metalloids? ›

Groups 13–16 of the periodic table (orange in the Figure below) are the only groups that contain elements classified as metalloids. Unlike other groups of the periodic table, which contain elements in just one class, groups 13–16 contain elements in at least two different classes.

Which of the first 20 elements are metalloids? ›

Metalloids : Boron and Silicon.

When was metalloid found? ›

The first metalloid used dates back to ancient Egypt, where antimony was used as makeup and as a colorant for paint and stains. However, it was only classified as a metalloid in the 1500s. Arsenic was also widely used in the middle ages. It was likely first isolated by a German physician, Albertus Magnus, around 1250.

What type of bond is a metalloid? ›

Metalloids: Boron, silicon, germanium, arsenic, antimony, and tellurium. Metalloids generally form covalent bonds with non-metals. But sometimes, they form ionic compounds with other elements. SiO₂, silicon dioxide, is a covalent compound.

What color are metalloids? ›

Metalloids can also be called semimetals. On the periodic table, the elements colored yellow, which generally border the stair-step line, are considered to be metalloids.

What are the general properties of a metalloid quizlet? ›

Match
• Property(1) Properties of both metals and nonmetals.
• Property(2) Metalloids are more brittle than metals, less brittle than most nonmetallic solids.
• Property(3) Metalloids are semiconductors of electricity.
• Property(4) Some metalloids possess metallic luster.
• Examples. Silicon, arsenic, boron, anitomy.

What are the general properties of metals nonmetals and metalloids? ›

Here are a few properties of metals, non-metals, and metalloids:
• Metals are generally shiny, malleable, and hard. Metals are also good conductors of electricity. ...
• Non-metals do not conduct heat or electricity very well. ...
• Metalloids share characteristics of both metals and non-metals and are also called semimetals.

What is the importance of metalloid? ›

In plants, metalloids play a role in diverse physiological processes. Some of the metalloids like boron, selenium, and silicon are beneficial or essential for healthy plant growth, whereas others, like arsenic and germanium, are highly toxic.

What is the most important property of metalloids? ›

They fall between metals and nonmetals in their ability to conduct heat, and if they can conduct electricity, they usually can do so only at higher temperatures. Metalloids that can conduct electricity at higher temperatures are called semiconductors.

What are the properties of metal short answer? ›

Metals are malleable and ductile. Metals are good conductors of heat and electricity. Metals are lustrous (shiny) and can be polished. Metals are solids at room temperature (except mercury, which is liquid).

### What are the physical and chemical properties of the metalloid group? ›

Metalloids can conduct electricity, but not as well as metals. Chemically, they act more like nonmetals, easily forming anions, having multiple oxidation states, and forming covalent bonds. Their ionization energies and electronegativities are in between the values of metals and nonmetals.

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