4.2: Metals, Metalloids and Nonmetals

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Figure 4-5. The unabridged version of the Periodic Table of the Elements, showing all 118 elements and emphasizing the fact that the vast majority of them are metals (shaded blue). Artificial elements are denoted by the outlined symbols. Metalloids are shaded lavender and nonmetals are rose. Notice the heavier line running diagonally from boron to astatine that separates the metals and nonmetals. This is Oliver Sacks' "Hadrians wall".

 

 

 

In the previous section, we introduced the organization of the Periodic Table along with Mendeleev's logic in organizing it the way he did. To summarize, elements in the same column often share chemical and physical properties, and each row shows a repeating, or "periodic", pattern of chemical and physical characteristics. But if we step back, as Oliver Sacks described his own viewing at the Science Museum of London when he was young, the general classification of elements according to their metallic properties is unavoidable (Figure 4-5). Of the ninety naturally occurring elements, 66 of them are metals, and almost all of the 28 artificially produced elements are metallic (or are expected to be if more than a handful of atoms of some of them could be generated). All tolled, 93 out of 118 elements are metallic. But what does it mean for an element to be metallic? 

The properties commonly associated with the term “metal” include:

• a characteristic shininess or reflective luster;
• malleability, which is the ability to be worked or rolled into thin sheets;
• ductility, which is the ability to be drawn into wires;
• exceptionally good electrical conductivity (metals typically conduct electric currents easily);
• very good thermal conductivity (they conduct heat very well).

All true – these characteristics define a particular substance as a metal, although mercury, being a liquid at room temperature, lacks a typical metal's workability (but it's a "normal" metal below its freezing point of -39°C). Several of these metallic properties are illustrated in the case of copper, probably one of the more familiar metals to you, in Figure 4-6.

 

Figure 4-6. Copper exhibits all of the classic physical characteristics of metals, including high reflectivity, malleability and ductility, as well as thermal and electrical conductivity. (Images, left: "Copper Wire Ring with Copper Bead" by bleggg is licensed under CC BY-NC 2.0; center:"Copper Wire" by Sam-Cat is licensed under CC BY-ND 2.0; and, right:  "Copperware in bazaar" by Gustible is licensed under CC BY-NC-SA 2.0)

 

The 18 elements that are non-metals have a much broader range of physical properties than do the metals: eleven are gases, one of which, helium, has a boiling point just four degrees above absolute zero, six are solids, including carbon, which has the highest melting point for any element, over 6,000°C, and one, bromine, is a liquid. Some are highly colored, including bright yellow sulfur and deep red-orange bromine, some are black, and some have no color at all. These elements are usually described, as they are named as a group, in the negative sense: they lack metallic properties and hence are non-metals. They cannot be drawn into wires, are usually poor conductors of heat and electricity, and lack any metallic luster. 

The properties listed above for metals and nonmetals are examples of physical characteristics. These describe characteristics in ways that do not involve interactions with other materials. Malleability or conductivity, to give two examples, are inherent properties of elements and do not refer to any reaction or interaction with other materials. But what about the chemical characteristics of metals and nonmetals? Chemical characteristics refer to the reactivity of materials. Flammability is probably the most familiar chemical characteristic, one that we referred to when discussing hydrocarbons and alcohols way back in Chapter 1. Are there any generalities of the reactivity of metals and nonmetals that would be useful to giving a general overview of these two broad classes of elements? Indeed there are: metals tend to lose electrons fairly easily when they react with nonmetals. Recall that all neutral atoms (and molecules) have exactly the same number of protons and electrons. Recall further that while the number of protons in a given atom is invariant, the number of electrons is not. To recap some ideas we introduced in Chapter 3, when a neutral atom gains or loses electrons, the charges will not balance and the resulting species is called an ion: those with an excess of positive charge are called cations, while anions are those that are negatively charged. Thus, one of the defining chemical properties of metallic elements is their tendency to form cations when reacting with nonmetals.

Is there a connection between the physical characteristics that metallic elements share and their tendency to lose electrons? Yes, there is. Both sets of properties are ultimately due to the relatively weak hold metal atoms have on some of their electrons. In any “bulk” metal, meaning a sample of the pure element that is large enough to make practical use of, e.g., an ingot of pure gold such as that illustrated in Figure 3-3, you can visualize the individual atoms as being arranged in an orderly, crystalline lattice. To a first approximation, you can view the atoms as packed spheres (like the cannon balls in Figure 3-6, presented again here at right); in such a model you will recognize that there must be some empty spaces, or voids, between the spheres. These interstitial pockets play an important role in how we think about the properties of metals, as explained below. Moreover, any bulk sample of metal will be neutral overall as it is a collection of a very large number of neutral atoms. Because some of the outermost electrons of each neutral metal atom are only loosely held, some tend to diffuse away. You can think of this as the electrons sort of wandering off into the interstitial spaces between the atoms, meaning they are no longer associated with the atom from which they originated but, rather, interact with several atoms simultaneously. Obviously when a neutral atom loses an electron, the remnant must be positively charged because now the number of protons is greater than the electrons. Thus, when neutral metal atoms lose electrons to the interstices of the lattice as just described, they are left as metal cations. The material as a whole remains electrically neutral because the total number of protons and electrons is not altered, but some of the electrons are free to migrate throughout the material. 

For the above reasons, bulk metals are usually thought of as crystalline arrays of metal cations held together by a completely delocalized “electron gas” that is freely mobile in the interstitial spaces between them. Think of the cations as being tightly packed marbles (or cannon balls) immersed in a negatively charged fluid, or “sea of electrons”, which is a commonly used description; the negatively charged fluid fills the interstitial spaces between the cations. In such an arrangement, the cations will tend to repel each other and hence are easily shifted with respect to one another, making the material malleable and ductile. On the other hand, metals can have extremely high tensile strength (hence their use in steel cables), as the electrostatic attraction between the “sea of electrons” and the cations is very strong and resists breaking. Finally, like a liquid that could flow between the spheres, the sea of electrons can also flow under the influence of an applied electric potential, explaining why metals are such good conductors of electricity.

If metallic properties arise from atoms “weakly holding on to their electrons" as described above, one could easily surmise that the absence of metallic properties of an element is associated with a tendency for its atoms to tightly hold on to their electrons. This would be correct. And not only do nonmetals not only hold their electrons tightly, many of them enthusiastically take electrons from other elements, especially metals. The rusting of steel (and most other forms of metal corrosion) is the result of such a reaction, wherein oxygen in the air reacts with iron, forming oxide anions, O^2-, and iron cations, Fe³⁺. This is a very common reaction pattern and can be summarized as follows: when atoms of nonmetals react with those of metals, the former take electrons away from the latter. Thus nonmetals tend to form anions. 

 

At a Glance:

Metals: lose electrons to form cations when reacting with non-metals

Non-metals: gain electrons to form anions when reacting with metals

 

We pause here to point out an important exception to the statement that nonmetals tend to form anions. The so-called inert gases (sometimes called the Noble Gases), are a family of elements that neither gains nor loses electrons easily; they comprise Group 8A on the extreme right hand side of Periodic Table. Elements in this family share with other nonmetals the tendency to tightly hold on to their own electrons, but they do not seek to gain more electrons by reacting with other elements. This reluctance underlies the dominant chemical property of these elements: their seeming inability to undergo chemical reactions of any type [8]. They form neither cations nor anions, nor do they readily form molecules by sharing electrons - they normally exist solely as isolated atoms in the gas phase. We saw this in the previous chapter: atoms of helium, the lightest inert gas, exist as isolated atoms, not molecules like the H₂ formed by hydrogen.  The other non-metallic elements, however, react quite easily.

The above focuses on how metals and nonmetals also react with each other, that is, with how metals react with nonmetals and vice versa. But nonmetals can also readily react with other nonmetals, and such reactions do not result in the formation of ions. Why? Because two different non-metals don't "complement" each other in the way that a metal and a nonmetal do, with one willing to give up electrons and they other willing to take them [9]. No, in the case of two non-metals, both want additional electrons and neither is particularly willing to give theirs away. We saw in the last chapter with the example of hydrogen and oxygen, both non-metals, reacting to form water; no ions are formed as the result of this reaction because the electrons involved in the bond formation are not completely gained or lost by the participating atoms. Rather, the electrons are shared, albeit unequally. 

Another way atoms of nonmetallic elements can share electrons is with other atoms of the same element. We've seen this already too. Two atoms of hydrogen react to form the H₂ molecule because that decreases their potential energy. Oxygen is another example: oxygen in the atmosphere usually exists as O₂. Just as with hydrogen, when individual oxygen atoms collide they “pair up”, forming diatomic molecules, each of which has two oxygen atoms as described by the following two (equivalent, [10]) balanced chemical equations.

\[ \ce{O (g) + O (g)  -> O2 (g)} \]

\[ \ce{2 O (g)  -> O2 (g)} \]

The tendency to form such diatomic species is common to several non-metallic elements, including hydrogen, nitrogen, fluorine, and chlorine, which make H₂, N₂, F₂, and Cl₂, respectively. In all of these cases, electrons are shared between pairs of atoms such that the negative charge of the electron cloud is, on average, closer to the positive charge of the atomic nuclei, decreasing their potential energy. The above examples are all gases at room temperature, but not all diatomic species are gaseous. Bromine and iodine also exist as diatomic molecules, but these molecules are large enough (in terms of their masses) that they have higher melting and boiling points, hence Br₂ is a liquid, as previously mentioned, while I₂ a solid at room temperature.

Thus far we have introduced two forms that nonmetallic elements assume when pure: isolated atoms, in the case of the inert gases, and the diatomic molecules listed above. Other structural types exist for some nonmetals. Sulfur, for example, is composed predominantly of S₈ molecules, with eight sulfur atoms arranged in a closed ring (Figure 4-7). One form of phosphorus exists as P₄ molecules, with each phosphorus bonding to three neighbors simultaneously, assuming a geometry in which each atom is at the vertex of a tetrahedron. 

Figure 4-7. (left) Pure sulfur is a bright yellow. The small mountain of it shown above sits at the port of Vancouver Bay waiting to be exported from Canada, which produces nearly 5 million metric tons of the element annually. (Photo: "Sulfur Piles Awaiting Export, Vancouver Bay, British Columbia, Canada" by euthman is licensed under CC BY 2.0); (right) On the molecular level, sulfur atoms coalesce to form molecules that are structured as puckered eight-member rings that resemble a crown.

Finally, to cite a particularly important example of bonding between atoms of the same element, the carbon atoms in diamond share electrons with four neighbors in an extended three-dimensional lattice. The attractive forces between adjacent atoms in this structure are exceptionally strong, so strong that it is the hardest material known (meaning it can scratch every other material but can't itself be scratched) and has a higher melting point than any other element, higher even than the temperature at the surface of the Sun. Carbon can also exist in the form of graphite, a soft grayish-black solid that is sometimes used as an industrial lubricant. Diamond and graphite are examples of allotropes, which are different physical forms of the same pure element (Figure 4-8). Allotropes are particularly prevalent among the nonmetallic elements, as multiple “strategies” exist by which the atoms can share electrons to reduce their potential energy. The above-mentioned P₄ molecules of phosphorus refer to an allotrope called white phosphorus (used for military applications, among other things), but other forms, such as red and black phosphorus also exist. Another example is ozone, O₃, an allotrope of “normal” atmospheric oxygen, O₂; ozone forms naturally in the upper atmosphere and serves the very important function of filtering harmful ultraviolet rays from the sun before they reach the earth’s surface. It can also be formed when industrial pollutants interact in sunlight and is quite toxic, causing tissue damage and respiratory problems. 

Figure 4-8. Two forms of pure carbon (allotropes): (top,left) diamond, which exists as a 3-dimensional lattice of carbon atoms wherein every carbon atom is bonded to four other carbons atoms (bottom left), is the hardest known substance in Nature and has long been prized a jewel; (top, right) graphite, the very soft material used in pencil "lead" , exists as flat "sheets" of carbon atoms in which every carbon atom bonds to three other carbon atoms via two single bonds and one double bond (bottom, right) [11]; the bonds between the carbon atoms are very strong, but the interactions between adjacent "sheets" of graphite are quite weak, making it a good solid lubricant because the sheets easily glide over each other. (Image credits: "diamonds" by Judy ** is licensed under CC BY-NC-ND 2.0; "Graphite pencil tip" by Matthew Fells is licensed under CC BY-NC 2.0; Structure illustrations from Chemistry, 2nd ed.,produced by OpenStax College and licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/85abf193-2bd...a7ac8df6@9.110)  "

The above paragraphs focused on the two extreme ends of the metal/nonmetal continuum. If you return to Periodic Table-as-landscape metaphor, mentioned earlier, and imagine taking a walk this space, moving from the realm of metals on the left to that of nonmetals on the right, you would need to cross what Oliver Sacks likened to "Hadrian's Wall". This boundary is better viewed, however, as something more akin to a change in scenery or climate than any sort of fortification. All of the pure metallic elements have their characteristic sheen, and most have a silver color. The nonmetals have a range of colors, states, and other properties. As you go from the reflective, silvery elements to the colorful, nonconductive ones, the transition wouldn't be abrupt. There would be a narrow range where the metallic character of the elements on the left hand side of the table begin to diminish but does not disappear entirely. Elements in this transitional zone are called metalloids (the purple-shaded elements in Figure 4-2). The most familiar metalloid is probably silicon [12], the main component of semiconductor computer chips (Figure 4-9). As the image below shows, silicon has a metallic reflectivity so it resembles a metal. But its electrical conductivity is much lower than metals (it has greater resistance), but it is not an insulating material (such as diamond and nonmetals). This intermediate conductivity is why materials like silicon are called semiconductors. Beyond their physical characteristics, metalloids also have chemical properties that are intermediate between metals and nonmetals [13]. We won't be devoting a large amount of time on metalloids in this text but we will mention some important aspects of them when relevant to other topics.

 

Figure 4-9. Pure silicon (left) in the form that silicon wafers (right) are made from. This sample is located outside The Arithmeum, a museum of mathematics and computer technology located in Bonn, Germany. Note the high reflectivity of the silicon surface, a characteristic it shares with metals. It is not, however, malleable or ductile, and has much lower electrical conductivity, hence its use as a semiconductor. (Image credits: left,  "pure silicon crystal" by Dave Messina is licensed under CC BY 2.0; right "Silicon Wafer Disc" by Business_Durham is licensed under CC BY-NC-ND 2.0.

 

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Notes and References

[8] This lack of reactivity is the basis for the name of the most common inert gas, argon. Making up about 1% of Earth's atmosphere, the name comes from the Greek word for "lazy". It is so lazy, in fact, that it escaped detection for a long time (things usually get detected because they do something that can be observed!), being one of the last naturally occurring elements to be discovered.

[9] Please forgive the anthropomorphization here - chemists often speak of atoms or molecules "wanting" or "liking" things, but we know full well that they do want or like anything - they are, indeed, inanimate objects, but we can describe what they do a little more easily by ascribing our own human motivations to them. In the end, this is a shorthand for describing energetic changes: atoms "want" to decrease the potential energy and "like" it when that happens!

[10] If they are equivalent, why give two equations? Depending on what you want to emphasize, you can write an equation in various ways. The first equation emphasizes the fact that two separate oxygen atoms come together and make a single molecule. The second equation is more efficient in that it combines like terms, but that can obscure certain interactions. Either is fine though. 

[11] We will see later that a better way to think about the bonding in graphite is to view each carbon-carbon interaction to not be a single bond or a double bond, but actually to be 1¹/₃ bonds.

[12] Important note: silicon is different from silicone. Silicon is a pure element, but silicone is a type of polymer (extremely large molecules made of repeating units of simple building blocks). Silicones have many industrial uses, such as in water-repellant coatings, caulks, and prosthetics.

[13] For example, the acid/base properties of the oxides of metals and nonmetals differ: oxides of metals are basic, while those of nonmetals are acidic. Oxides of metalloids have intermediate character. More on acid/base chemistry in coming chapters.