Categories
Chemistry

Calcium and Aluminium

Calcium

Calcium is too reactive to occur as a free metal in nature. It occurs abundantly in the combined state as calcium trioxocarbonate (iv) in limestone, marble, chalk, aragonite, calcite, coral, dolomite, calcium fluoride. etc In Nigeria, limestone is found at Nkalagu in Ebonyi, Ewekoro at Abeokuta and Ukpilla in Delta state.

Extraction

Since calcium are very stable, metallic calcium is commonly extracted electrolytically from fused calcium chloride a byproduct of solvay process. Some calcium chloride is usually added to the fused calcium chloride to lower the melting point from 850oC to about 650oC. The mixture is placed in a large crucible lined on the inside with graphite which serves as the anode of the cell. The cathode consists of iron rod which just touches the surface of the electrolyte. As electrolysis proceeds metallic calcium collects on the cathode which is gradually raised so that an irregular stick of calcium is formed on it. Chlorine is liberated at the cathode.

Chemistry of the Reaction

At the cathode – the calcium ions receive two electrons each to become reduced to the metal.

At the anode – two chloride ions give up an electron each to become atomic chlorine. The two atoms then combine to become liberated as a gaseous molecule.

Cl →  Cl + e

Cl + Cl →  Cl2

Overall electrolytic reaction

Ca2(l) + 2Cl(l)  →  Ca(s) + Cl2(g)

Physical Properties of Calcium

  1. Appearance – Silvery grey solid
  2. Relative density is 1.55
  3. Calcium is malleable and ductile
  4. It has relatively low tensile strength
  5. Melting point is 850oC
  6. Calcium is a good conductor of heat and electricity

Chemical Properties of Calcium

Reaction with air – Calcium is a very electropositive and reactive metal. On exposure to air, it rapidly tarnishes ad loses its metallic lustre due to the formation of white film of calcium oxide or quick lime on the surface of the metal. When calcium is heated in air, it burns with a brick red flame to form calcium oxide

2Ca(s) + O2(g) → 2CaO(s)

CaO(s) + H2O(l)→ Ca(OH)2(g)

Reaction with non-metals – on heating, calcium combines directly with nitrogen, chlorine, Sulphur and hydrogen

3Ca(s) + N2(g) → Ca3N2(s)

Ca(s) + Cl2(g) → CaCl2(s)

Reaction with water – Calcium reacts slowly with cold water and rapidly with warm water to form calcium hydroxide and hydrogen

Ca(s) + 2H2O(g) → Ca(OH)2(aq) + H2(g)

Reaction with Ammonia – if ammonia is passed over heated calcium, it reacts as follows

3Ca(s) + 2NH3(g) → Ca3N2(s) + 3H2(g)

Test for calcium ions

Flame test – calcium compounds give an orange – red colour to a non-luminous flame. Moisten the unknown compound with a few drops of concentrated hydrochloric acid. Dip the tip of a clean platinum wire into the mixture and hold it in a non-luminous Bunsen flame. If a bright brick red flame through a blue glass is produced, the unknown ions of the compound are calcium ions.

With sodium hydroxide – Add a few drops of NaOH solution to an unknown salt. The formation of white precipitate which is insoluble in excess sodium hydroxide indicate the presence of calcium ions

Uses of Calcium

  1. Calcium is used as a deoxidant in steel castings and copper alloys.
  2. It is also used in the manufacture of calcium fluoride and calcium hydride.
  3. It is used in the extraction of uranium
  4. It is needed in the diet of young children for development of strong bones and teeth
  5. Calcium metal is used as a reducing agent in preparing other metals such as thorium and uranium. It is also used as an alloying agent for aluminium, beryllium, copper, lead and magnesium alloys.
  6. Calcium compounds are widely used. There are vast deposits of limestone (calcium carbonate) used directly as a building stone and indirectly for cement.
  7. Gypsum (calcium sulfate) is used by builders as a plaster and by nurses for setting bones, as ‘plaster of Paris’.

Aluminium

Aluminium is the most common metal in the Earth’s crust, making up 7.5% by mass. Its main ore is bauxite-a clay mineral which you can think of as impure aluminium oxide. It is the most important element in group III.

Extraction of Aluminium
Aluminium is obtained largely from the ore bauxite (Al2O3.2H2O). Its production is a two-step process: the purification of bauxite and extraction by electrolysis.

Purifying the bauxite (aluminium oxide) – the Bayer Process

Crushed bauxite is treated with moderately concentrated sodium hydroxide solution. The concentration, temperature and pressure used depend on the source of the bauxite and exactly what form of aluminium oxide it contains. Temperatures are typically from 140°C to 240°C; pressures can be up to about 35 atmospheres.

High pressures are necessary to keep the water in the sodium hydroxide solution liquid at temperatures above 100°C. The higher the temperature, the higher the pressure needed.

With hot concentrated sodium hydroxide solution, aluminium oxide reacts to give a solution of sodium aluminate (III) (NaAl(OH)4).

Extraction of Aluminium by Electrolysis

After purification, aluminium oxide is mixed with cryolite (sodium aluminium fluoride) Na3AlF6 to lower the melting point from 2000º to 1000º, which saves money. This mixture is heated and the molten liquid used as the electrolyte. Both electrodes are made of graphite (carbon).  The anode (+ve) is graphite and the cathode (-ve) is a graphite lining to a steel case.

The anode disintegrates. The hot oxygen produced here reacts with the hot carbon anode to give carbon dioxide. Hence it must be replaced regularly.

Aluminium ions are attracted to the cathode (the negative electrode) and are reduced to aluminium by gaining electrons.

Al3+ (l) + 3e →  Al (l)

The molten aluminium produced sinks to the bottom of the cell.

The oxide ions are attracted to the anode and lose electrons to form oxygen gas.

2O2- (l) → O2 (g) + 4e

Note: The extraction of aluminium is an expensive process because the large amount of electricity needed to keep the electrolytes molten is expensive. Hence using cryolite saves energy and money, as it acts as a solvent for the aluminium oxide and melts at a much lower temperature.

Physical Properties of Aluminium

  1. Aluminium is a silvery white metal which is comparatively soft
  2. It is a strong, malleable metal element.
  3. It has a low density.
  4. It is resistant to corrosion
  5. It is a good conductor of heat and electricity.
  6. It can be polished to give a highly reflective surface.

Chemical Properties of Aluminium

  • Action with air: Aluminium burns in air at high temperature to form the oxide and the nitride

4Al (s) + 3O2 (g) ——> 2Al2O3 (s)

  • Reaction with non-metals: On heating, aluminium combines directly with non-metals like the halogens, sulphur, nitrogen, phosphorus and carbon with evolution of heat

2Al (s) + 3Cl2 ——> 2AlCl3 (s)

  • Action with acids: Aluminium reacts more rapidly with the concentrated hydrochloric acid to displace hydrogen but more slowly with dilute one

2Al (s) + 6HCl (aq) ——> 2AlCl3 (aq) + 3H2 (g)

  • Reaction with alkalis: Aluminium reacts with both sodium and potassium hydrogen solutions giving hydrogen gas and soluble tetrahydroxoaluminate (III)

2Al (s) + 2NaOH (aq) + 6H2O (l) ——> 2NaAl(OH)4 (aq) + 3H2 (g)

  • Reaction with Iron (III) oxide: Aluminium reduces iron (III) oxide to molten iron. The reaction is used in thermit process and it gives out a great deal of energy

2Al (s) + Fe2O3 (s) ——> Al2O(s) + 2Fe (s)

Uses

  1. Low density and strength make aluminium ideal for construction of aircraft, lightweight vehicles, and ladders.
  2. An alloy of aluminium called duralumin is often used instead of pure aluminium because of its improved properties.
  3. Easy shaping and corrosion resistance make aluminium a good material for drink cans and roofing materials.
  4. Corrosion resistance and low density leads to its use for greenhouses and window frames.
  5. Good conduction of heat leads to its use for boilers, cookers and cookware.
  6. Good conduction of electricity leads to its use for overhead power cables hung from pylons (low density gives it an advantage over copper).
  7. High reflectivity makes aluminium ideal for mirrors, reflectors and heat resistant clothing for fire fighting.

ASSESSMENT (POST ANSWERS BELOW USING THE BOX)

Explain the method of extraction of Aluminium in few lines and bullet points.

Mention 3 uses of Calcium

Join Discussion Forum and do your assignment: Find questions at the end of each lesson, Click here to discuss your answers in the forum

Ad: Get a FREE Bible: Find true peace. Click here to learn how you can get a FREE Bible.

For advert placement/partnership, write [email protected]

Download our free Android Mobile application: Save your data when you use our free app. Click picture to download. No subscription. stoplearn

We are interested in promoting FREE learning. Tell your friends about Stoplearn.com. Click the share button below!