Metals and Metallurgy Notes

Metals

Minerals

Most metals are found in solid inorganic compounds known as minerals. Minerals are named by common names, not chemical names.

Important Metal Minerals

Most important metals are found in minerals as oxides, sulfides, or carbonates. Examples include:

• Aluminum:

  • Corundum: $Al<em>2O</em>3$

  • Gibbsite: $Al(OH)_3$

• Chromium:

  • Chromite: $FeCr<em>2O</em>4$

• Copper:

  • Chalcocite: $Cu_2S$

  • Chalcopyrite: $CuFeS_2$

  • Malachite: $Cu<em>2CO</em>3(OH)_2$

• Iron:

  • Hematite: $Fe<em>2O</em>3$

  • Magnetite: $Fe<em>3O</em>4$

• Lead:

  • Galena: $PbS$

• Manganese:

  • Pyrolusite: $MnO_2$

• Mercury:

  • Cinnabar: $HgS$

• Molybdenum:

  • Molybdenite: $MoS_2$

• Tin:

  • Cassiterite: $SnO_2$

• Titanium:

  • Rutile: $TiO_2$

  • Ilmenite: $FeTiO_3$

• Zinc:

  • Sphalerite: $ZnS$

Metallurgy

Metallurgy is the science and technology of extracting metals from their natural sources and preparing them for practical use. It involves:

• Mining

• Concentrating ores

• Reducing ores to obtain free metals

• Purifying metals

• Mixing metals to form alloys with desired properties

Pyrometallurgy

Pyrometallurgy is the use of high temperature to alter or reduce minerals.

Calcination

Calcination is heating an ore to bring about its decomposition and elimination of a volatile product. For example:

$PbCO<em>3(s) \longrightarrow PbO(s) + CO</em>2(g)$

Roasting

Roasting is a thermal reaction between an ore and the furnace atmosphere (often oxygen). Examples:

$2 MoS<em>2(s) + 7 O</em>2(g) \longrightarrow 2 MoO<em>3(s) + 4 SO</em>2(g)$

$HgS(s) + O<em>2(g) \longrightarrow Hg(g) + SO</em>2(g)$

Smelting

Smelting is a melting process in which materials formed during reactions separate into two or more layers.

Refining

Refining is the treatment of a crude, relatively impure metal to improve its purity and better define its composition.

Reduction of Iron

Hematite ($Fe<em>2O</em>3$), magnetite ($Fe<em>3O</em>4$), and other iron oxides are reduced in blast furnaces. Purified iron exits the furnace at the bottom.

Steel

Crude molten iron contains many impurities, including silicon, manganese, phosphorus, sulfur, and carbon. The impurities are oxidized by $O_2$ (except phosphorus, which reacts with $CaO$) to compounds easily separated from the molten iron. Purified molten steel is then poured into molds.

Hydrometallurgy

Hydrometallurgy includes techniques in which metal is extracted from ore via the use of aqueous reactions.

Leaching

Leaching is a process in which a metal-containing compound is selectively dissolved. Water can be used if the compound is water-soluble, but often acid, base, or a salt solution is required. For example:

$4 Au(s) + 8 CN^-(aq) + O<em>2(g) + 2 H</em>2O(l) \longrightarrow 4 Au(CN)_2{}^-(aq) + 4 OH^-(aq)$

$2 Au(CN)<em>2{}^-(aq) + Zn(s) \longrightarrow Zn(CN)</em>4{}^{2-}(aq) + 2 Au(s)$

Bayer Process

This is a method of purifying bauxite (aluminum ore):

$Al<em>2O</em>3 ∙ H<em>2O(s) + 2 H</em>2O(l) + 2 OH^-(aq) \longrightarrow 2 Al(OH)_4{}^-(aq)$

The soluble aluminate ion is separated from the insoluble impurities ($SiO<em>2$ and $Fe</em>3O_3$) by filtration.

Electrometallurgy

The reduction of metal ores or refining of metals by use of electricity is called electrometallurgy.

Sodium

$NaCl$ is electrolyzed in a Downs cell. Gaseous $Cl_2$ is allowed to disperse, and molten $Na$ is siphoned off.

Aluminum

In the Hall process, $Al<em>2O</em>3$ is dissolved in molten cryolite ($Na<em>2AlF</em>6$), and $Al^{3+}$ is reduced to molten $Al$.

Copper

Active metal impurities are oxidized at the anode, but don’t plate out at the cathode, as $Cu^{2+}$ is more easily reduced. Less active metals deposit as sludge below the anode.

Physical Properties of Metals

Metals:

• Conduct heat and electricity.

• Are malleable (can be pressed or hammered into sheets).

• Are ductile (can be drawn into wire).

• Are packed so atoms can slip past each other, making them less brittle than other solids.

Electron-Sea Model

Metals can be thought of as cations suspended in a “sea” of valence electrons. Attractions hold electrons near cations, but not so tightly as to impede their flow. This explains many of the properties of metals, like conductivity of heat and electricity, and deformation.

Molecular Orbital Model

The electron-sea model does not explain observed trends in melting point, boiling point, heat of fusion, etc., because the model suggests these properties should increase with increasing number of valence electrons. These trends can be explained by energy bands created by large number of molecular orbitals formed as metal atoms bond with each other. As with nonmetals, bond order apexes in center of row, then decreases. Thus, attractions (and melting point, etc.) apex in center of transition metals (Group 6B).

Alloys

Alloys are mixtures of elements that have properties characteristic of metals. Many of the ordinary uses of metals involve alloys.

Solution Alloys

Components of alloys are dispersed uniformly.

• In substitutional alloys, solute particles take the place of solvent metal atoms. The particles in these alloys are quite close in size.

• In interstitial alloys, solute particles find their way into the holes between solvent metal atoms. In this type of alloy, the solute particles are smaller than the solvent particles.

Intermetallic Compounds

Intermetallic compounds are homogeneous alloys with definite properties and compositions. An example is $Co_5Sm$, which is used for permanent magnets in headsets and speakers.

Transition Metals

Many important metals are included in this group. The transition metals are the elements in the d block of the periodic table.

Physical Properties of Transition Metals

Some of their properties (such as ionization energy, atomic radius, etc.) are suggestive of isolated atoms, while others (such as density, melting point, etc.) suggest bulk solid metal.

Atomic Radii

Trends in atomic radii are similar across all three rows of transition metals. While $Z_{eff}$ increases across a row, so does the number of nonbonding electrons, which repel each other and increase the radius.

Electron Configurations and Oxidation States

Transition metals often have more than one common oxidation state. Most have +2 state due to loss of s electrons. Oxidation numbers greater than 2 are due to the loss of d electrons as well as s. Many form compounds that have colors and significant magnetic properties.

• In diamagnetic elements, all electron spins are paired, so there is no net magnetic moment.

• In paramagnetic atoms and ions, there are unpaired spins. The magnetic fields are randomly arranged, though, unless placed in an external magnetic field.

• In ferromagnetic substances, the orientations of magnetic fields from unpaired electrons are affected by spins from electrons around them. When an external field is applied and then removed, the substance maintains the magnetic moment and becomes a permanent magnet.

Chromium

Chromium is oxidized by $HCl$ or $H<em>2SO</em>4$ to form blue $Cr^{2+}$ ion. $Cr^{2+}$ is oxidized by $O<em>2$ in air to form green $Cr^{3+}$. $Cr$ is also found in the +6 state, as in $CrO</em>4{}^{2-}$ and the strong oxidizer $Cr<em>2O</em>7{}^{2-}$.

Iron

Iron exists in solution in +2 or +3 state. Elemental iron reacts with non-oxidizing acids to form $Fe^{2+}$, which oxidizes in air to $Fe^{3+}$. Brown water running from a faucet is caused by insoluble $Fe<em>2O</em>3$. $Fe^{3+}$ is soluble in acidic solution, but forms a hydrated oxide as a red-brown gel in basic solution.

Copper

In solution, copper exists in the +1 or the +2 state. +1 salts are generally white and insoluble. +2 salts