Chemistry of Metals

Slag is the fusible mass formed when a flux combines with the gangue (infusible impurities) during smelting. It is lighter than the molten metal, so it floats on top and is easily removed.

Example: in the extraction of iron, the basic flux CaO combines with the acidic gangue SiO$_2$:

$\mathrm{CaO + SiO_2 \rightarrow CaSiO_3\ (slag)}$

(i) Bleaching powder with dilute H$_2$SO$_4$ liberates chlorine (this is the chlorine that does the bleaching):

$\mathrm{CaOCl_2 + H_2SO_4 \rightarrow CaSO_4 + H_2O + Cl_2}$

(ii) Gypsum first loses water to give plaster of Paris at about 120 °C. When heated above 130 °C it loses all its water and forms anhydrous calcium sulphate (dead burnt plaster), which has no setting property:

$\mathrm{CaSO_4\cdot 2H_2O \xrightarrow{>130^\circ C} CaSO_4 + 2H_2O}$

(i) Hydrometallurgy

It is a metallurgical process in which the metal is extracted from its ore by dissolving it in a suitable aqueous reagent (leaching), and then the metal is recovered from the solution by displacement with a more electropositive metal or by electrolysis. It is used for low-grade ores of noble metals such as gold, silver and also copper.

Example (cyanide process for gold/silver): the ore is leached with dilute NaCN solution in the presence of air:

$\mathrm{4Au + 8NaCN + 2H_2O + O_2 \rightarrow 4Na[Au(CN)_2] + 4NaOH}$

Gold is then displaced by zinc:

$\mathrm{2Na[Au(CN)_2] + Zn \rightarrow Na_2[Zn(CN)_4] + 2Au}$

(ii) Zone refining

It is used to obtain metals of very high purity (e.g. Ge, Si, Ga, B, which are used in semiconductors). Principle: impurities are more soluble in the molten metal than in the solid metal.

A rod of impure metal is fitted with a circular heater at one end. The heater creates a narrow molten zone and is slowly moved along the rod. As the heater moves on, pure metal crystallises out behind it while the impurities pass into the adjacent molten zone and are swept to the other end. The process is repeated several times, and the end containing the impurities is cut off.

Principle: Sodium is a very strong reducing agent, so it cannot be extracted by chemical reduction. It is obtained by electrolysis of fused (molten) sodium chloride. A little $\mathrm{CaCl_2}$ is added to lower the melting point from $801^\circ\mathrm{C}$ to about $600^\circ\mathrm{C}$.

Process (Down's cell): The cell is a rectangular steel tank lined with firebrick. A graphite anode is placed at the centre, surrounded by a circular iron cathode. A wire-gauze diaphragm between the electrodes keeps the products apart. Above the cathode is an inverted trough that collects the molten sodium, and above the anode is a dome that collects chlorine.

        Cl2 out            Na out
          |                  |
       [ dome ]        [ trough ]
     ----|----------------|-------
     |   anode (C)  gauze  cathode (Fe) |
     |        molten NaCl + CaCl2       |
     ------------------------------------

On passing current through the molten mixture:

$\mathrm{NaCl \rightarrow Na^+ + Cl^-}$

At cathode: $\mathrm{Na^+ + e^- \rightarrow Na}$ (liquid sodium floats up and is collected)

At anode: $\mathrm{2Cl^- \rightarrow Cl_2 + 2e^-}$ (chlorine is collected as by-product)

Overall: $\mathrm{2NaCl \rightarrow 2Na + Cl_2}$

Why a sodium fire is not put out with water: Sodium reacts violently with water to form NaOH and hydrogen, $\mathrm{2Na + 2H_2O \rightarrow 2NaOH + H_2}$. The reaction is highly exothermic, so the hydrogen catches fire and the fire spreads. Dry sand or a dry powder extinguisher is used instead.

Two advantages of Down's process:

• Fused NaCl with $\mathrm{CaCl_2}$ melts at a lower temperature (about $600^\circ\mathrm{C}$), so less energy is used.
• The diaphragm keeps sodium and chlorine apart, so pure sodium is obtained along with useful chlorine as a by-product. (Sodium also does not vapourise or burn as easily at this temperature.)

Sodium to washing soda: Sodium reacts with water to give NaOH, which absorbs $\mathrm{CO_2}$ to form sodium carbonate. Crystallisation from water gives washing soda.

$\mathrm{2Na + 2H_2O \rightarrow 2NaOH + H_2}$ $\mathrm{2NaOH + CO_2 \rightarrow Na_2CO_3 + H_2O}$ $\mathrm{Na_2CO_3 + 10H_2O \rightarrow Na_2CO_3\cdot 10H_2O}$

Gangue: The earthy and rocky impurities (sand, clay, etc.) present in an ore are called gangue. Example: silica ($\mathrm{SiO_2}$) in iron ore.

Slag: The fusible compound formed when flux reacts with gangue during smelting is called slag. Example: calcium silicate, formed in the blast furnace:

$\mathrm{CaO + SiO_2 \rightarrow CaSiO_3}$

(i) On heating at $120^\circ\mathrm{C}$, gypsum loses part of its water of crystallisation and forms plaster of Paris:

$\mathrm{2CaSO_4\cdot 2H_2O \xrightarrow{120^\circ C} (CaSO_4)_2\cdot H_2O + 3H_2O}$

(ii) Chlorine reacts with slaked lime to form bleaching powder:

$\mathrm{Ca(OH)_2 + Cl_2 \rightarrow CaOCl_2 + H_2O}$

(i) Carbon reduction process: Oxides of less reactive metals are reduced by heating with carbon (coke), which acts as the reducing agent. The metal is obtained and carbon monoxide or dioxide is released. The oxide is obtained first if the ore is a carbonate (by calcination) or a sulphide (by roasting).

$\mathrm{ZnO + C \xrightarrow{\Delta} Zn + CO}$ $\mathrm{Fe_2O_3 + 3CO \xrightarrow{\Delta} 2Fe + 3CO_2}$

It is used for metals such as Fe, Zn, Sn and Pb, whose oxides are easily reduced by carbon at feasible temperatures.

(ii) Froth floatation process: This is used to concentrate sulphide ores (e.g. galena PbS, zinc blende ZnS, copper pyrites). The finely powdered ore is mixed with water, pine oil (a frothing agent) and a collector such as sodium ethyl xanthate in a tank, and air is blown in. The oil wets the sulphide particles, which stick to the air bubbles and rise with the froth. The gangue is wetted by water and settles. The froth is skimmed off and the concentrated ore is obtained. Depressants such as NaCN can be used to separate ores of different metals (e.g. ZnS from PbS).

Four important ores of Nepal:

1. Haematite ($\mathrm{Fe_2O_3}$), iron ore, found at Phulchoki (Lalitpur).
2. Copper pyrites ($\mathrm{CuFeS_2}$), found in Kavrepalanchok and Dhading.
3. Galena ($\mathrm{PbS}$), lead ore, found in the Ganesh Himal area.
4. Magnesite ($\mathrm{MgCO_3}$), found at Kharidhunga, Dolakha.

(Zinc blende, ZnS, and limestone are also found.)

(i) Sodium is very reactive. In air it reacts with oxygen, moisture and CO₂ to form a dull layer of Na₂O, NaOH and Na₂CO₃ on its surface, so the shiny metal tarnishes. (It is therefore kept under kerosene.)

(ii) Alkali metals are very strong reducing agents, and their oxides are much more stable than the oxides of carbon, so carbon cannot reduce them (and at high temperature carbides form). They are extracted by electrolysis of their fused chlorides.

Epsom salt is hydrated magnesium sulphate, $\mathrm{MgSO_4\cdot 7H_2O}$. It is a colourless, crystalline, bitter-tasting solid, very soluble in water.

Preparation. It is obtained by dissolving magnesite or magnesium oxide in dilute sulphuric acid, filtering, and crystallising below $48^\circ\mathrm{C}$ (the heptahydrate separates):

$\mathrm{MgCO_3 + H_2SO_4 \rightarrow MgSO_4 + H_2O + CO_2}$

Industrially it is also obtained from the mineral kieserite, $\mathrm{MgSO_4\cdot H_2O}$, by dissolving it in water and recrystallising.

Action of heat. On heating it loses water in stages:

$\mathrm{MgSO_4\cdot 7H_2O \xrightarrow{150^\circ C} MgSO_4\cdot H_2O \xrightarrow{>200^\circ C} MgSO_4}$

At very high temperature the anhydrous salt decomposes to $\mathrm{MgO}$ and $\mathrm{SO_3}$.

Chemical reactions.

• With $\mathrm{BaCl_2}$ it gives a white precipitate of $\mathrm{BaSO_4}$ (test for sulphate): $\mathrm{MgSO_4 + BaCl_2 \rightarrow BaSO_4\downarrow + MgCl_2}$
• With $\mathrm{NaOH}$: $\mathrm{MgSO_4 + 2NaOH \rightarrow Mg(OH)_2\downarrow + Na_2SO_4}$
• With $\mathrm{Na_2HPO_4}$ in presence of $\mathrm{NH_4Cl}$ and $\mathrm{NH_4OH}$ it gives a white precipitate of magnesium ammonium phosphate, $\mathrm{MgNH_4PO_4}$.
• With sodium carbonate it gives a white precipitate of basic magnesium carbonate.

Uses.

• As a purgative (laxative) in medicine and for relieving muscle pain in baths.
• In dyeing and calico printing, and for making cotton and silk fabrics fire-proof and heavy.
• As a source of magnesium in fertilisers and in tanning leather.

An alloy is a homogeneous mixture of a metal with one or more other metals or non-metals, made by melting them together. Example: brass ($\mathrm{Cu + Zn}$).

An amalgam is a special alloy in which one of the metals is mercury. Example: sodium amalgam, $\mathrm{Na-Hg}$ (or the silver-tin amalgam used in dental filling).

So every amalgam is an alloy, but an alloy is an amalgam only if it contains mercury.

Alkali metals have low ionisation energies and a single loosely held valence electron. In a flame, this electron absorbs a small amount of energy and jumps to a higher energy level. The excited electron is unstable and returns to the ground state, releasing the absorbed energy as radiation. For alkali metals this energy falls in the visible region, so the flame is coloured. Each metal has its own energy-level spacing, so the colour is characteristic: Li crimson red, Na golden yellow, K lilac (violet), Rb red-violet, Cs blue.

Plaster of Paris (POP)

Chemical name: calcium sulphate hemihydrate; formula $\mathrm{CaSO_4\cdot\tfrac{1}{2}H_2O}$ (or $\mathrm{2CaSO_4\cdot H_2O}$).

Preparation

It is prepared by heating gypsum, $\mathrm{CaSO_4\cdot 2H_2O}$, at about $120$–$130^\circ\mathrm{C}$ (the temperature must be controlled):

$\mathrm{CaSO_4\cdot 2H_2O \xrightarrow{120-130^\circ C} CaSO_4\cdot\tfrac{1}{2}H_2O + 1\tfrac{1}{2}H_2O}$

Properties

1. It is a white powder, slightly soluble in water.
2. Setting: when mixed with water to a paste, it takes up water and sets into a hard, porous solid mass of gypsum, with a slight expansion in volume:

$\mathrm{CaSO_4\cdot\tfrac{1}{2}H_2O + 1\tfrac{1}{2}H_2O \rightarrow CaSO_4\cdot 2H_2O}$

3. Overheating: at about $200^\circ\mathrm{C}$ it loses all its water to form anhydrous $\mathrm{CaSO_4}$ (dead burnt plaster), which does not set with water.

$\mathrm{CaSO_4\cdot\tfrac{1}{2}H_2O \xrightarrow{200^\circ C} CaSO_4 + \tfrac{1}{2}H_2O}$

Because of this, POP must be stored in a moisture-proof container.

Uses

• In hospitals, to make plaster casts for setting fractured bones.
• To make statues, toys, models and decorative articles.
• For plastering walls and ceilings, and false ceilings.
• In dentistry, to make moulds and casts.
• In making chalk, and as a fire-proofing material.

(c) BeCO$_3$ > MgCO$_3$ > CaCO$_3$ > BaCO$_3$

Thermal stability of Group II carbonates increases down the group, so the ease of giving off CO$_2$ decreases down the group. The small Be$^{2+}$ ion has high polarising power and distorts the carbonate ion most, weakening its C–O bond, so BeCO$_3$ decomposes most easily; the large Ba$^{2+}$ ion polarises least, so BaCO$_3$ is the most stable and needs the highest temperature. (Option (d) gives the reverse order, i.e. order of thermal stability.)

An electrolytic cell is an electrochemical cell in which electrical energy from an external source is used to bring about a non-spontaneous chemical change, e.g. the decomposition of molten NaCl into sodium metal at the cathode and chlorine gas at the anode.

Calcium is more electropositive than sodium (standard reduction potential $-2.87$ V for Ca$^{2+}$/Ca against $-2.71$ V for Na$^+$/Na), so Ca$^{2+}$ ions are harder to reduce than Na$^+$ ions. CaCl$_2$ is added only to lower the melting point of NaCl, and at the controlled voltage used, Na$^+$ is preferentially discharged at the cathode:

$\text{Na}^+ + e^- \rightarrow \text{Na}$

So sodium, not calcium, is deposited.

In aqueous NaCl, water is also present. Water is reduced at the cathode more easily than Na$^+$ (Na is much more electropositive than hydrogen), so hydrogen gas is liberated instead of sodium:

$2\text{H}_2\text{O} + 2e^- \rightarrow \text{H}_2 + 2\text{OH}^-$

Any sodium formed would also react with water. Metallic sodium is therefore obtained only by electrolysis of molten NaCl, so aqueous NaCl cannot replace it for the same product.

Roasting is the process of heating a sulphide ore strongly in the presence of excess air (oxygen) to convert it into its oxide. Sulphur is removed as sulphur dioxide gas.

$2ZnS + 3O_2 \xrightarrow{\text{heat}} 2ZnO + 2SO_2$

Calcination is the process of heating a carbonate or hydrated oxide ore strongly in the absence (limited supply) of air to convert it into its oxide. Carbon dioxide or water vapour is removed.

$ZnCO_3 \xrightarrow{\text{heat}} ZnO + CO_2$

So the key difference is the atmosphere: roasting needs air and is used for sulphide ores, while calcination is done in the absence of air and is used for carbonate ores.

Roasting is carried out in a reverberatory furnace, where the ore is heated strongly in a plentiful supply of air, below its melting point, to convert sulphide ores into oxides.

Two impurities removed during roasting:

1. Sulphur — present as sulphide ores (e.g. ZnS, PbS, Cu₂S); it is oxidised to sulphur dioxide gas and escapes.
2. Arsenic (and antimony) — present as arsenides; they are oxidised to volatile oxides (As₂O₃) which sublime off.

Roasting is the heating of an ore in a plentiful supply of air, so the impurities removed are those that form volatile products — mainly sulphur (as SO₂) and arsenic/antimony (as their volatile oxides).