Iron and Calcium: Properties, Reactions, and Effects

Posted on Dec 11, 2024 in Chemistry

The Iron Cation

Pure iron is a silvery-white, tenacious, and ductile metal. Commercial iron is rarely pure and often contains small amounts of carbides, silicides, phosphides, and sulfides of iron, and a bit of graphite. Iron dissolves in diluted hydrochloric acid and concentrated sulfuric acid, diluted with hydrogen evolution, and the formation of ferrous salt. With hot, concentrated sulfuric acid, it produces sulfur dioxide and ferric sulfate. Under experimental conditions with concentrated nitric acid, cold iron is made passive. In this state, it does not react with dilute nitric acid and does not displace copper from a copper salt solution.

Reactions of Ferrous Ion Fe++

  1. Sodium Hydroxide Solution: Forms a white precipitate of ferrous hydroxide, Fe(OH)2, insoluble in excess reagent but soluble in acids. Upon exposure to air, ferrous hydroxide is rapidly oxidized, finally obtaining reddish ferric hydroxide. Under normal conditions, it appears as a brownish-green precipitate.
  2. Ammonium Hydroxide Solution: This produces a precipitate of ferrous hydroxide in ways similar to the reaction. In the presence of excess ammonium chloride solution, the common ion effect of ammonium ions decreases the concentration of hydroxyl ions to an extent that the solubility product of ferrous hydroxide precipitate is not reached.

Environmental Effects of Iron

Iron(III)-O-arsenite, pentahydrate can be hazardous to the environment. Special attention should be paid to plants, air, and water. It is strongly recommended not to let the product enter the environment because it persists.

  • Ferric Ferrous Oxide (Fe3O4): This form of oxide tends to occur when iron rusts underwater and is therefore frequently found in tanks or under the waterline.
  • Ferric Oxide (Fe2O3): This is commonly seen in iron oxide and oxidized steel structures. It attacks from bridges to car bodies and is extremely destructive.

Effects of Iron on Health

May cause conjunctivitis, chorioretinitis, and retinitis if it contacts the tissues and remains in them. Chronic inhalation of excessive concentrations of vapors or dust of iron oxide can result in the development of a benign pneumoconiosis, called siderosis, which is observable as a change in the X-ray. No physical damage to lung function has been associated with siderosis. Inhalation of excessive concentrations of iron oxide may increase the risk of developing lung cancer in workers exposed to lung carcinogens.

Calcium

Calcium is the fifth most abundant element in Earth’s crust (3.6% by weight) but not in the native state. It forms compounds with high industrial interest, such as carbonate and dolomite and sulfate (gypsum, alabaster), from which quicklime, plaster, cement, etc., are obtained.

Features

Calcium is a soft, malleable, and ductile alkaline earth metal that burns with a red flame, forming calcium oxide and nitrure. The recent surfaces are silvery-white but soon become slightly pale yellow when exposed to air and ultimately gray or white. It reacts violently with water to form hydroxide Ca(OH)2, releasing hydrogen.

Reactions

  • Ammonium Oxalate Precipitation: Calcium produces a white precipitate of calcium oxalate, CaC2O4.H2O, slowly formed in concentrated solutions and slowly in diluted ones.
  • Flame Test: When the calcium mineral appears in a state such that it can be volatilized by heat, it will give a characteristic orange flame.
  • Alkaline Reaction: Calcium is an alkaline earth metal. Therefore, if a mineral has calcium in combination with a volatile acid, once burned, its residue will give an alkaline reaction on damp turmeric paper.
  • Precipitation as Calcium Sulfate: Calcium is precipitated from a hydrochloric acid solution as sulfate by adding a little dilute sulfuric acid. The precipitate is soluble in water and therefore does not form a diluted solution.

Calcium Carbonate

Calcium carbonate is also known as calcite, limestone, chalk, lithographic stones, aragonite, etc. It is made (referring to the pure substance): CaCO3. Calcium carbonate is obtained by combining CO2 and Ca(OH)2, but if used in excess, CO2 originates bicarbonate.

Calcium Oxide

It is called quicklime and has the formula CaO. Pure calcium oxide is white, and lime is usually grayish or yellowish due to certain impurities. Lime is obtained by burning calcium in the air, although on an industrial scale, it is produced from calcium carbonate. The reaction is reversible: The limestone is burned in kilns called calderas.

Environmental Effects of Calcium

Calcium and magnesium are associated with water hardness. In accordance with the analytical determination, the total hardness is expressed as the sum of the concentrations of calcium and magnesium, both expressed as calcium carbonate. Calcium phosphate is very toxic to aquatic life. It eliminates CO2 and SO2 by calcium.

New research suggests that technology using calcium oxide has the potential to cheaply capture the CO2 from power plants. In this method, called the carbonation-calcination cycle, CO2 is removed from the smoke through a chemical reaction with calcium oxide (CaO). This produces CaCO3, where it is heated to create a reaction that separates added CO2 from CaO. This produces CO2 that can be stored and heated. This technique has the added advantage that it also captures sulfur dioxide (SO2), a sulfate contaminant found in coal and responsible for acid rain.