1 Fertilisers, ammonia, sulfuric acidUnit 16.1 Fertilisers, ammonia, sulfuric acid

2 Did you remember to give out the worksheet? 16.1 Learning outcomes:Describe why we need to use fertilisers. State the major components in fertilisers. Describe the displacement of ammonia and its salts.

3 What is a fertiliser? It is a substance added to the soil to replace elements taken up by plants. Artificial fertilizer containing nitrogen, phosphorus and potassium are added to the soil to replace essential plant nutrients that have been lost. These fertlisers are are called NPK fertilisers after the symbols for these three elements.

4 Did you know..? Nearly 10% of the ENTIRE chemical industry is involved in making fertilisers. Many fertilisers contain ammonium salts. The preparation of a fertiliser in the lab involves the following equipment: a measuring cylinder to measure a particular volume of an alkali solution, a 'burette' to add acid a little at a time and a filter funnel to remove the solid crystals of fertiliser from the solution.

5 Over supply of nutrients. What can go wrong?excessive richness of nutrients in a lake or other body of water, frequently due to run-off from the land, which causes a dense growth of plant life. A major problem with the use of fertilisers occurs when they're washed off the land by rainwater into rivers and lakes. The resulting increase of nitrate or phosphate in the water encourages algae growth, which forms a bloom over the water surface.

6 Making fertilisers… Most fertilisers are made by the reaction of an acid and an alkali. FERTILISER ACID ALKALI Ammonium nitrate nitric acid ammonia Ammonium phosphate phosphoric acid Ammonium sulfate sulfuric acid Potassium nitrate potassium hydroxide

7 Ammonium nitrate. HNO3 (aq) + NH3 (aq) → NH4NO (aq)Ammonium phosphate. H3PO4 (aq) + NH3 (aq) → (NH4)3PO4 (aq)

8 What the heck is slaked lime?Many fertilisers are slightly acidic. Many plants do not grow in acidic soil. The solution? Slaked lime. Lime (CaO) reacts with water to form calcium hydroxide (slake lime). This can then be added to the soil, where it reacts with the slightly acidic ammonium salts to release ammonia (An alkali)

9 The manufacture of ammoniaUnit 16.2 The manufacture of ammonia

10 Boardworks GCSE Additional Science: Chemistry Reversible ReactionsWhat is ammonia? Ammonia is an important compound in the manufacture of fertilizer and other chemicals such as cleaning fluids and floor waxes. It is made industrially by reacting nitrogen with hydrogen in the Haber process. It is a reversible reaction, so it never goes to completion. Why is this a problem for companies making ammonia? Photo credit: © 2007 Jupiterimages Corporation hydrogen nitrogen + ammonia N2 (g) 3H2 (g) 2NH3 (g)

11 Boardworks GCSE Additional Science: Chemistry Reversible ReactionsThe Haber process Teacher notes This five-stage animation illustrates how ammonia is made in the Haber process.

12 Boardworks GCSE Additional Science: Chemistry Reversible ReactionsWhat is yield? The amount of product made in a reaction is called the yield and is usually expressed as a percentage. The yield of ammonia produced by the Haber process depends on the temperature and pressure of the reaction. ammonia yield (%) pressure (atm) Teacher notes See the ‘Quantitative Chemistry’ presentations for more information on yields.

13 What is the Haber compromise?Boardworks GCSE Additional Science: Chemistry Reversible Reactions The highest yield of ammonia is theoretically produced by using a low temperature and a high pressure. In practice, though, these conditions are not used. Why? Lowering the temperature slows down the rate of reaction. This means it takes longer for ammonia to be produced. Increasing the pressure means stronger, more expensive equipment is needed. This increases the cost of producing the ammonia. Photo credit: Davis Hay Jones / Science Photo Library A compromise is reached to make an acceptable yield in a reasonable timeframe while keeping costs down.

14 Temperature, pressure and yieldBoardworks GCSE Additional Science: Chemistry Reversible Reactions Temperature, pressure and yield Teacher notes This activity could be used as a plenary or revision exercise on the effects of temperature and pressure on the Haber process.

15 Changing the yield of ammoniaBoardworks GCSE Additional Science: Chemistry Reversible Reactions Changing the yield of ammonia

16 HABER PROCESS N2(g) + 3H2(g) 2NH3(g) : DH = - 92 kJ mol-1Conditions Pressure kPa (200 atmospheres) Temperature °C Catalyst iron

17 HABER PROCESS N2(g) + 3H2(g) 2NH3(g) : DH = - 92 kJ mol-1Conditions Pressure kPa (200 atmospheres) Temperature °C Catalyst iron Equilibrium theory favours low temperature exothermic reaction - higher yield at lower temperature high pressure decrease in number of gaseous molecules

18 HABER PROCESS N2(g) + 3H2(g) 2NH3(g) : DH = - 92 kJ mol-1Conditions Pressure kPa (200 atmospheres) Temperature °C Catalyst iron Equilibrium theory favours low temperature exothermic reaction - higher yield at lower temperature high pressure decrease in number of gaseous molecules Kinetic theory favours high temperature greater average energy + more frequent collisions high pressure more frequent collisions for gaseous molecules catalyst lower activation energy

19 HABER PROCESS N2(g) + 3H2(g) 2NH3(g) : DH = - 92 kJ mol-1Conditions Pressure kPa (200 atmospheres) Temperature °C Catalyst iron Equilibrium theory favours low temperature exothermic reaction - higher yield at lower temperature high pressure decrease in number of gaseous molecules Kinetic theory favours high temperature greater average energy + more frequent collisions high pressure more frequent collisions for gaseous molecules catalyst lower activation energy Compromise conditions Which is better? A low yield in a shorter time or a high yield over a longer period. The conditions used are a compromise with the catalyst enabling the rate to be kept up, even at a lower temperature.

20 IMPORTANT USES OF AMMONIA AND ITS COMPOUNDSHABER PROCESS IMPORTANT USES OF AMMONIA AND ITS COMPOUNDS MAKING FERTILISERS 80% of the ammonia produced goes to make fertilisers such as ammonium nitrate (NITRAM) and ammonium sulphate NH HNO3 ——> NH4NO3 2NH H2SO4 ——> (NH4)2SO4 NITRIC ACID ammonia can be oxidised to nitric acid nitric acid is used to manufacture... fertilisers (ammonium nitrate) explosives (TNT) polyamide polymers (NYLON)

21 Boardworks GCSE Additional Science: Chemistry Reversible ReactionsThe Haber compromise To produce a high yield of ammonia, but with a fast rate of reaction and without the need for overly expensive equipment, the Haber process is carried out at 450 °C and 200 atmospheres. The most important factor in deciding what conditions to use is therefore not yield, but total cost. What costs are involved in the industrial production of ammonia? Photo credit: © 2007 Jupiterimages Corporation raw materials equipment energy wages

22 Maximizing productivityBoardworks GCSE Additional Science: Chemistry Reversible Reactions Maximizing productivity What else can be done to maximise productivity in the manufacture of ammonia? An iron catalyst is used to increase the rate of reaction. It speeds up both the forward and backward reaction, so the position of equilibrium is not affected. The ammonia is cooled, liquefied and then removed as it is produced. This causes the equilibrium to shift to the right to produce more ammonia. Unreacted nitrogen and hydrogen are recycled and given another chance to react.

23 Sulfur and sulfuric acid

24 Sulfuric acid – the contact processThe contact process, for making sulfuric acid, is a process which involves a reversible reaction. The raw materials needed to make sulfuric acid are: sulfur air water

25 Stage 1 - making sulfur dioxideIn the first stage of the contact process, sulfur is burned in air to make sulfur dioxide: sulfur + oxygen  sulfur dioxide S(l) + O2 (g)  SO2 (g) This is not a reversible reaction - (l) means liquid and (g) means gas. Sulfur dioxide must not be released as it contributes to acid rain.

26 Stage 2 - making sulfur trioxideIn the second stage, sulfur dioxide reacts with more oxygen to make sulfur trioxide: Sulfur dioxide + oxygen ↔ sulfur trioxide 2SO2(g) + O2(g) ↔ 2SO3(g) This reaction is reversible. The conditions needed for it are: a catalyst of vanadium(V) oxide, V2O5 a temperature of around 450°C a pressure of approximately 2 atmospheres

27 Stage 3 - making sulfuric acidIn the final stage, sulfur trioxide reacts with water to make sulfuric acid: H2O(l) + SO3(g)  H2SO4(aq) This is not a reversible reaction, just like the first stage - (aq) means aqueous or dissolved in water.

28 Uses and hazards of sulfuric acidThe majority of sulfuric acid that is produced is used to make fertilisers. This is often by reacting the sulfuric acid with ammonia, or ammonium hydroxide solution, to make ammonium sulfate, (NH4)2SO4. ammonia + sulfuric acid → ammonium sulfate ammonium hydroxide + sulfuric acid → ammonium sulfate + water Sulfuric acid is also used in the production of detergents and paints. Sulfur is rarely found pure in the ground. It is an impurity which is produced when fossil fuels are burned and it causes acid rain.

29 The limestone industryUnit 16.5 Pgs

30 LO: Describe the manufacture of lime in terms of thermal decomposition. Name some uses of lime and slaked lime Name the uses of calcium carbonate.

31 LIMESTONE Limestone is made of calcium carbonate CaCO3Origin Formed in the sea millions of years ago from the remains of shells Extraction Quarried in large amounts e.g. 150 million tonnes each year in the UK IT IS AN IMPORTANT RAW MATERIAL Use building material and road chippings neutralising excess acid in lakes and soil making cement and concrete added to bread added to toothpaste

32 LIMESTONE PRODUCTS ARE VERY USEFUL … especially in building housesCEMENT / MORTAR GLASS CONCRETE

33 CARBONATES Na2CO3 Formulae All carbonates contain the CO3 unitThe compounds are ionic; the formula is found by balancing the charges of the ions Carbonate ions have a 2- charge CO32- CALCIUM CARBONATE CaCO3 SODIUM CARBONATE Na2CO3

34 CARBONATES Formulae All carbonates contain the CO3 unitChemical formula sodium carbonate Na2CO3 calcium carbonate CaCO3 magnesium carbonate MgCO3 copper carbonate CuCO3

35 calcium carbonate —> calcium oxide + carbon dioxidePROCESSING LIMESTONE Roasting Heating limestone very strongly makes it break up The process is known as THERMAL DECOMPOSITION calcium carbonate —> calcium oxide carbon dioxide CaCO —> CaO CO2 Calcium oxide is also known as Lime or QUICKLIME

36 ACTION OF HEAT ON METAL CARBONATESCO2 METHOD Place a small amount of one of the solid carbonates in a dry test tube. Place about 2cm3 of lime water in another test tube. Heat the solid, gently at first, then more strongly. Test any gas with the limewater. lime water metal carbonate Appearance CO2 produced Residue Conclusion Calcium carbonate CaCO3 Copper carbonate CuCO3 Magnesium carbonate MgCO3 Zinc carbonate ZnCO3

37     CaCO3 CuCO3 MgCO3 ZnCO3 ACTION OF HEAT ON METAL CARBONATESAppearance CO2 produced Residue Conclusion Calcium carbonate CaCO3 WHITE SOLID WHITE SOLID Calcium oxide formed (needs very strong heating)  Copper carbonate CuCO3 BLACK SOLID Copper oxide formed GREEN SOLID  Magnesium carbonate MgCO3 WHITE SOLID Magnesium oxide formed WHITE SOLID  Zinc carbonate ZnCO3 YELLOW SOLID WHICH TURNS WHITE WHEN COOL Zinc oxide formed WHITE SOLID  Sodium carbonate (Na2CO3) also decomposes on heating but it requires more heat than an ordinary bunsen burner can supply.

38 THERMAL DECOMPOSITION OF CARBONATESAll metal carbonates decompose when heated to form carbon dioxide and a metal oxide. The process is known as THERMAL DECOMPOSITION calcium carbonate —> calcium oxide carbon dioxide copper carbonate —> copper oxide carbon dioxide magnesium carbonate —> magnesium oxide carbon dioxide sodium carbonate * —> sodium oxide carbon dioxide zinc carbonate —> zinc oxide carbon dioxide * THE THERMAL DECOMPOSITION OF GROUP I METAL CARBONATES (SUCH AS SODIUM CARBONATE) REQUIRES A TREMENDOUS AMOUNT OF ENERGY – A BUNSEN BURNER IS NOT HOT ENOUGH +

39 THERMAL DECOMPOSITION OF CARBONATESAll metal carbonates decompose when heated to form carbon dioxide and a metal oxide. The process is known as THERMAL DECOMPOSITION calcium carbonate —> calcium oxide carbon dioxide copper carbonate —> copper oxide carbon dioxide magnesium carbonate —> magnesium oxide carbon dioxide sodium carbonate * —> sodium oxide carbon dioxide zinc carbonate —> zinc oxide carbon dioxide * THE THERMAL DECOMPOSITION OF GROUP I METAL CARBONATES (SUCH AS SODIUM CARBONATE) REQUIRES A TREMENDOUS AMOUNT OF ENERGY – A BUNSEN BURNER IS NOT HOT ENOUGH

40 THERMAL DECOMPOSITION OF CARBONATES Write out the chemical equations+ CaCO3 —> CaO CO2 CuCO3 —> CuO CO2 MgCO3 —> MgO CO2 Na2CO3 —> Na2O CO2 ZnCO3 —> ZnO CO2 + + + +

41 THERMAL DECOMPOSITION OF CARBONATES+ CaCO3 —> CaO CO2 CuCO3 —> CuO CO2 MgCO3 —> MgO CO2 Na2CO3 —> Na2O CO2 ZnCO3 —> ZnO CO2 + + + +

42 QUICKLIME Calcium oxide CaO Reacts with water to form SLAKED LIME (CALCIUM HYDROXIDE) calcium oxide + water ——> calcium hydroxide HEAT CaO H2O ——> Ca(OH)2 Use added to soil to make it less acidic added to lakes which have been polluted by acid rain

43 SLAKED LIME & LIME WATERLime water is a solution of calcium hydroxide in water it is an alkali it is used to test for the gas carbon dioxide (limewater goes ‘cloudy’ if CO2 is present) calcium hydroxide carbon dioxide ——> calcium carbonate + water Ca(OH) CO ——> CaCO H2O

44 AQUEOUS CALCIUM HYDROXIDELIMESTONE CYCLE CALCIUM CARBONATE (limestone) HEAT AQUEOUS CALCIUM HYDROXIDE (lime water) CALCIUM OXIDE (quicklime) HEAT SOLID CALCIUM HYDROXIDE (slaked lime)

45 All three rely on limestone for their manufactureCEMENT, CONCRETE & GLASS All three rely on limestone for their manufacture CEMENT powdered clay powdered limestone mix and roast them in a rotary kiln

46 All three rely on limestone for their manufactureCEMENT, CONCRETE & GLASS All three rely on limestone for their manufacture MORTAR cement sand water mortar is a ‘thin’ form of concrete used for bricklaying

47 REACTION OF ACIDS WITH CARBONATESmagnesium nitric —> magnesium carbon water carbonate acid nitrate dioxide zinc hydrochloric —> zinc carbon water carbonate acid chloride dioxide sodium nitric —> sodium carbon water carbonate acid nitrate dioxide zinc sulphuric —> zinc carbon water carbonate acid sulphate dioxide copper sulphuric —> copper carbon water carbonate acid sulphate dioxide potassium hydrochloric —> potassium carbon water

48 REACTION OF ACIDS WITH CARBONATESmagnesium nitric —> magnesium carbon water carbonate acid nitrate dioxide zinc hydrochloric —> zinc carbon water carbonate acid chloride dioxide sodium nitric —> sodium carbon water carbonate acid nitrate dioxide zinc sulphuric —> zinc carbon water carbonate acid sulphate dioxide copper sulphuric —> copper carbon water carbonate acid sulphate dioxide potassium hydrochloric —> potassium carbon water

49 AQA C1.2.1 Calcium carbonate - summaryLimestone, mainly composed of the compound calcium carbonate (CaCO3), is quarried and can be used as a building material. b) Calcium carbonate can be decomposed by heating (thermal decomposition) to make calcium oxide and carbon dioxide. c) The carbonates of magnesium, copper, zinc, calcium and sodium decompose on heating in a similar way to give carbon dioxide and the metal oxide. Not all carbonates of metals in Group 1 of the periodic table decompose at the temperatures reached by a Bunsen burner. d) Calcium oxide reacts with water to produce calcium hydroxide, which is an alkali that can be used in the neutralisation of acids. e) A solution of calcium hydroxide in water (limewater) reacts with carbon dioxide to produce calcium carbonate. Limewater is used as a test for carbon dioxide. Carbon dioxide turns limewater cloudy. f) Carbonates (eg Mg, Cu, Zn, Ca, Na) react with acids to produce carbon dioxide, a salt and water. Limestone is damaged by acid rain. g) Limestone is heated with clay to make cement. Cement is mixed with sand to make mortar and with sand and aggregate to make concrete.