Inorganic carbon chemistry

Composed of CaCO3 in the form of mineral calcite e.g. marble and chalkChalk – made of shells of marine algaeOther limestones – made from debris of animal structures.Marble – forms when limestone is subjected to high pressures or high temperatures or both acting together, to create crystals of caco3 in the rock.Direct uses of limestoneConstruction industryCalcium oxide manufactureCement manufactureIron extractionGlass makingSodium carbonate manufactureRoad buildingNeutralization of acid soilPowdered limestone is used because it is:Cheaper than limeSlow actingAn excess does not make the soil alkalineCO32- + 2H+ – CO2 + H2OManufacture of iron and steelCaCO3 used to remove earthy and sandy materials in the blast furnace.A liquid slag is then formed which is mainly calcium silicate.Manufacture of cement and concreteCaCO3 mixed with clay in a heated rotary kiln using coal or oil as fuel.This produces cement – contains Ca (Alo2)2 and CaSiO3The dry product is ground to a powderCaso4 is added to slow down the setting rate of the cement.When h2o is added, crystals of hydrated Ca (Alo2)2 and CaSiO3CaCO3 is obtainedConcrete = cement+ stone chippings and sandConcrete is poured into wooden moulds to set hardReinforced concrete – concrete allowed to set around steel rods or mesh to give it more tensile strength – required for large bridgesManufacture of Na2CO3 – the Solvay processUsed to manufacture soaps, detergents, dyes, drugsNa2CO3 = CaCO3 + NaCl + CO2 + NH3 in a Solvay tower, the centre of reactionsCO2 and NH3 are recycled continuouslyIndirect uses of limestoneLime manufactureWhen CaCO3 is heated, it breaks up reversibly to form CaO and CO2. CaCO3 – CaO + CO2 .Thus reaction can go in either direction, depending on the temp and the pressure and takes place in a lime kiln.CaO – base; used by farmers:To neutralize soil acidityTo improve drainage of water through soilsUsed as a drying agentUsed to make soda glass (= heat soda (Na2CO3) +sand + lime)Used to convert to Ca (OH)2 – slaked limeManufacture of Ca(OH)2 – slaked limeUsed to:Make bleaching powderReduce soil acidityManufacture whitewashManufacture glassPurify waterCa(OH)2 in its white powder form, is produced by adding an equal amount of water to CaO – more exothermicA weak solution of Ca(OH)2 in water is called limewaterIt is used to test for CO2as when CO2 is added to it, CaCO3 is formed.Ca(OH)2 + CO2 – CaCO3 + H2OIf CO2 is bubbled further, the CaCO3 then dissolves and a solution of calcium bicarbonate is produced.CaCO3 + CO2 + H2O -Ca(HCO3)2Ca(OH)2 +sand – mortar – when allowed to set in water – strongly bonded materialCa(OH)2 + CO2- CaCO3 + H2OCarbonates and hydrogen carbonatesCarbonates: salts of H2CO3; contain CO32- ion.Properties of carbonatesWhen metal carbonates are heated – metal oxide + CO2Group 1 metals except for lithium, do not dissociate on heatingMore reactive metals = less dissociationCO3 generally insoluble except Na, K, NH4CO32- + acid – metal salt + CO2+ H2OCaCO3 + 2HCl – CaCl2 + CO2 + H2O: used to prepare CO2 in lab because the reaction produces CO2 which causes effervescence and if bubbled through Ca(OH)2 turns it chalky white.Na2CO3 – an important industrial chemicalUsed for:Manufacturing glassTreating brinePurifying waterManufacturing detergents and textilesHydrogen carbonatesToo unstable to exist as solidsContain HCO3-NaHCO3 – most common: found in indigestion remedies because it reacts with HCl (in stomach) which causes indigestion.NaHCO3 + HCl – NaCl + H2O + CO2Added to plain flour as when the NaHCO3 is heated it releases CO2 which makes the cake riseNaHCO3 – Na2CO3 + H2O + CO2Ca(HCO3)2 and Mg(HCO3)2 – cause hardness in waterHardness in waterRainwater dissolves CO2A small fraction of CO2 reacts with H2O – H2CO3H2O + CO2 – H2CO3As this solution passes over rocks containing limestone, compounds called Mg or CaCO3 are produced and hey slowly dissolve in themCaCO3 + H2CO3 – Ca(HCO3)2Some rocks may contain:Gypsum(CaSO4.2 H2O)Anhydrite(CaSO4) sparingly soluble in water and their presence makes the water hardKieserite(mgso4.h2o)Hardness :temporary-removed by boiling- Caused by Ca or Mg HCO3s:permanent – more difficult to removed- Caused by Ca or Mg SO4sWhen water containing ca or Mg SO4 is evaporated, a white solid deposit is left behind called ca or mg SO4s or CaCO3 – causes furring in kettles – removed by dilute acid2H+ + CO32- – CO2 + H2OAlso blockages of hot water pipes are caused by a process like furring.Stalactites and stalagmites – found in underground caverns in limestone areas v- formed by Ca/Mg HCO3s decomposition Ca(HCO3)2 – CaCO3 + CO2 + H2OEffect of hard water on soapHard water areas – no lather – H2O becomes cloudy because of the precipitate formed by the reaction of the dissolved substances in water with soap (NaSt). This precipitate – scum: 2 NaSt + Ca(HCO3)2 – Ca(St)2 + NaHCO3To overcome scum – soap less detergents have been made – don’t react with the dissolved substancesRemoval of hardnessTemp – boiling because when heated the cahco32 decomposesPerm – when heated Ca(HCO3)2 does not decomposeSo the temp and perm hardness can be removedBy adding washing soda crystals (Na2CO3.10 H2O)Ca2+ + CO32- – CaCO3Ion exchange: H2O is passed through a container filled with a suitable resin containing sodium ions. The Ca or Mg ions are exchanged for the Na ions in the resin.Ca2+ +Na- – Ca2+-r +2Na+The Na ions can be regenerated by pouring a solution of suitable Na salt through it.Distillation: very expensiveDisadvantagesAdvantagesWastes soapHas a nice tasteCauses kettles to furCalcium ions in hard water are required by the body for bones and teethCan cause hot water pipes to blockCoats lead pipes with a thin layer of lead sulphate or lead carbonate and cuts down the possibility of lead poisoning.Can spoil the finish of some fabricsSome industries favor hard water for e.g. – leather industryCarbon dioxideTwo oxides: CO and CO2CO2 is:Produced by burning fossil fuelsProduced by all living organisms through aerobic respiration : C6H12O6 + 6O2 – 6 CO2+ 6 H2OTaken by plants and used with h2o to synthesize sugars” photosynthesis : 6CO2+ 6 H2O – C6H12O6 + 6O2More CO2 = more temperature of the earth = greenhouse effectSome energy from the sun is absorbed by the earth due to the greenhouse gases in the atmosphere. This energy is heat and helps heat up the earth as the gases don’t let the heat escape.More temperature = melting of ice caps , floods in low lying areas.Uses of CO2CO2is bubbled into the liquid under pressure, which increases its solubilityFire extinguishers – CO2 is denser than air and forms a layer around the burning material. It covers the fire and starves it of O2. CO2 does not burn and so the fire is put out.Refrigerants – used because it is colder than ice and it sublimes.Special effects – dry ice is placed in boiling water and it forms thick clouds of white smoke. It stays close to the floor since CO2 is denserHeat transfer agents – in nuclear power stationsLab preparation of CO2 gasBy reacting dil. HCl on CaCO3 (marble chip)CaCO3 +HCl – CaCl2 + H2O+ CO2Properties of CO2Physical propertiesColorless gasSparingly soluble in H2ODenser than airWhen bubbled in water, some dissolves slightly and some forms H2CO3Supports the combustion of strongly burning substances as they burning substances as they break CO2 to C and O2 and get O2 to support the combustion.When co2 is bubbled through limewater, a white precipitate called caco3 is formed.CO2 + Ca(OH)2 – CaCO3 + H2OIf CO2 is bubbled through this solution, the solution becomes clear because of the formation of Ca(HCO3)2CaCO3 + CO2 – Ca(HCO3)2CO2 reacts with strong alkalis to form carbonates. If excess CO2 is bubbled through a solution of an alkali then a white precipitate may be obtained.