METALLIC BONDING 1­ Metallic Bond: Metals are good electric and heat conductors. Their characteristics include: lustre (they shine), malleability (they can be hammered into thin sheets), ductility (they can be stretched into very thin wires), hardness, high density, high melting points. On what type of bonding can these properties be explained? In their solid state, metals form crystals in which each metallic atom is surrounded by 8 to 12 similar atoms. Since these atoms have only one, two or three valence electrons, they do not form normal covalent bonds with their neighbours. Also, all atoms of the same element exert an identical attraction on their electrons. Therefore, they do not have a tendency to form ions since there is no atom that can lose or gain an electron from its neighbour: they are all identical! Any theory related to metallic bonding is valid if it can explain the main physical properties of the metals, i.e. malleability, ductility, and thermal and electrical conductivity. Although most of the pure metals are relatively malleable, they are very hard, and their melting point is high. These properties reveal that in most metals, the bond is strong and non­oriented. In other words, although it is difficult to separate the neighbouring shells of the atoms of a metal, it is relatively easy to displace them, as long as they remain in contact with each other. The simplest theory which explains these metallic properties is the electron­sea theory, in which a network of metallic cations swim in a «sea» of valence electrons. These mobile electrons can conduct heat and electricity since they can easily go from one metal atom with low electronegativity to another metal atom. These electrons are referred to as delocalized electrons, since they are not held in one particular position or locality; they do not belong to any ion or covalent bond. When subjected to an electric field, these delocalized electrons travel through the metal and produce an electric current. When they react with light, these electrons absorb and re­emit light over a wide wavelength range, producing the lustre of the metal. When a metal is hammered, there is rearrangement of the atoms. The cations slide over each other pretty much like marbles, which causes a change in the form of the metal without breakage. The electrons get into new positions in the orbitals. The metallic bond is formed by delocalized electrons which hold the atoms together. Representation of an alkali metal (group 1) with one valence electron. The number of electrons that participate in the metallic bond determines the properties of metals. Elements of group 1 have only one valence electron per atom. They are rather soft in comparison with other metals. Those of group 2 have two valence electrons. They are harder than group 1 metals. As for the transition elements, the orbital d electrons which are partially filled can also participate in the metallic bond. Many of these transition metals are hard. It is the strength of the metallic bond that makes a metal hard and solid, such as iron, chromium, and nickel. In general, transition metals are the hardest and strongest metals. As a rule: the more delocalized electrons, the harder the metal. 2­ Alloys: It is possible to strengthen certain metal elements with a small number of delocalized electrons by combining them to other metals. The properties of an alloy differ from those of its component metals. An alloy is defined as the physical mixing of two metals to form a solution. It is usually prepared by melting the metals together and cooling the mixture.
There are two types of alloys: substitutional alloys, and interstitial alloys. In a substitutional alloy, some atoms of the base metal are replaced by other metallic atoms of identical or similar size. In brass, for example, about one third of the copper atoms are replaced by zinc atoms. In an interstitial alloy, smaller atoms fit into the interstices, the spaces between the larger metal atoms. In steel, for example, carbon atoms occupy the spaces between the iron atoms. The presence of these interstitial atoms change the properties of the base metal and make it harder, stronger, and less ductile than pure iron. Certain metals dissolve in each other in any type of proportion. When they form an alloy, they produce solid solutions. Other pairs of metals do not entirely dissolve and make up heterogeneous mixtures. The degree of solubility of a metal in another metal depends on the relative size of their atoms. Metals with similar sizes tend to dissolve in each other by substitution. Metals with very different sizes dissolve in each other by insertion. In all alloys, the relative electronegativity of the atoms play an important role. Scientists know that elements can form intermetallic compounds because many of these have been prepared, such as CuMg2, Au2Bi, etc. but they still know very little about this branch of chemistry. Read p. 171­172 of your textbook. METALLIC BOND: AN OVERVIEW Metallic bonds consist of the attraction of the free­floating valence electrons for the positively charged metal ions. 1­ Theory on Metallic Bonding: Characteristics of metals:
· · · · · · · lustre (sheen)
malleability (can be hammered into thin sheets)
ductility (can be stretched into very thin wires)
hardness
high density
high melting point
electrical and thermal conductivity The electron­sea theory:
· · · a network of metallic cations swim in a sea of valence electrons
the valence electrons are delocalized, i.e. they are not held in a fixed position and can go from one metal atom to another
the number of delocalized electrons that participate in the metallic bond determine the physical properties of these metals Explanation of the properties of metals according to the electron­sea theory:
· · · · subjected to an exterior electrical field, delocalized electrons travel from one metal atom of low electronegativity to another to produce an electric current
submitted to light, delocalized electrons absorb and re­emit light over a wide wavelength range, producing the lustre of the metal
when pressure is applied to a metal by hammering it, cations slide on each other just like marbles which causes a change in the form of the metal without breaking it
elements of groups 1 and 2 have only one or two delocalized valence electrons per atom, and therefore, they are relatively soft; transition metals, on the other hand, have electrons on their d orbitals that also participate in the metallic bond, and that makes them very hard and strong 2­ Alloys:
® An alloy is a solid solution of metals. Alloys have properties that differ from those of their component metals. Alloys are generally stronger and harder than pure metals. Types of alloys:
· · substitutional alloy: some atoms from the base metal are replaced by other metallic atoms of similar size; ex: brass (1/3 of copper atoms are replaced by zinc atoms)
interstitial alloy: smaller atoms come fill the holes (spaces) between the metal atoms; ex: steel (carbon atoms fill the spaces between iron atoms) METALLIC BOND 1. What are three characteristics of metals? 2. Explain how the number of delocalized electrons affect the properties of a metal. 3. Using the electron­sea theory, explain how the metallic bond determines the electrical conductivity of metals. 4. Why are alkali metals and alkaline earth metals relatively soft? Suggest a way to make them harder. 5. To make steel, we add carbon to iron metal. Explain why adding a certain quantity of carbon will make the metal more resistant. 6. Many metals can withstand multiple reshaping and hammering without breaking. We can bend them, fold them, flatten them into sheets, stretch them, etc. to give them different shapes. How does the electron structure of the metals explain these characteristics? 7. Metals such as silver, cadmium, and gold have different uses. However, they are relatively soft in when they are pure, especially if we compare them to the metals that are located above them in the periodic table. Explain why. Suggest a method of making them stronger so they have more uses. 8. Many pure metals are good electrical conductors. However, in their crude or unrefined state, they have a lesser electrical conductivity. Why? 9. Explain briefly the two types of alloys. Give an example of each type.