Unit 3 - Chemistry in Society

1. Unit 3 - Chemistry in Society

2. 2Al + Fe2O3 2Fe + Al2O3 Calculate the mass of iron produced from 40g of iron oxide. Example from white board

3. C9H2O + 14 O29 CO2 + 10 H2O Calculate the mass of water produced when 6.4 grams of nonane is burned. Show your working clearly. . Example from white board

4. The idea of excess When a chemical reaction involving two reactants is carried out, usually one of the reactant gets completely used up and some of the other reactant is left over. The reactant which is left over is said to be in excess.

5. n c V(l) Worked Example 1 Which reactant would be in excess if 6.5g of zinc was added to 25cm3 of dilute sulphuric acid of concentration 2mol l-1 Example from white board Nat 5 reminder n = no of mole (moles) c = concentration (mol l-1) v = volume (litres) divide cm3 by 1000

6. Worked Example 2 • Which reactant would be in excess if 0.972g of magnesium was added to 50cm3 of 0.10 mol l-1 hydrochloric acid? b) Calculate the mass of hydrogen produced Example from white board

7. Worked Example 3 • Which reactant is in excess when 0.506g magnesium carbonate is added to 100cm3 of 0.5moll-1 nitric acid. b) Calculate the mass of carbon dioxide produced Example from white board

8. Molar Volume Molar volume = formula mass density Volume = mass density When dealing with gases its more appropriate to use volume instead of it’s mass. The molar volume of a gas (measured at 0oC @ 1atm) can be worked out using the middle equation. The molar volume is the volume that one mole of the gas takes up. Molar volume allows us to relate gases to one another.

9. Density Worked example Example from white board What is the density of butane (C4H10) in gl–1 if its molar volume is 22 l mol–1?

10. uestion V n Vm normally given in Q As the temperature and pressure change, the volume a gas takes up changes. Provided the molar volume of a gas is known, the volume of a gas can be calculated from the number of moles of the gas using the following relationship. where; V = volume of gas (litre) n = no of moles (moles) Vm = molar volume (l/mol.)

11. Molar Volume Worked example 1 Example from white board What volume of carbon dioxide is produced by roasting 25g of calcium carbonate? (molar volume is 24 lmol –1)

12. Molar Volume Worked example 2 Example from white board If 300cm3 of a gas weigh 0.55g, what is the formula mass of the gas? (molar volume 24 lmol–1)

13. Gas Mixtures (common in multiple choice) Example from white board If 25cm3 of ethene are burned in 100cm3 of oxygen, what would be the composition of the resulting gas mixture?

14. Balanced Equation + Gases example Example from white board What volume of oxygen is required for the complete combustion of 4 litres of methane?

15. The Chemical Industry

16. the chemical industryis one of the largest industries in Britain. its products are vital to many aspects of modern life and many are used for the benefit of society. the chemical industry involves the investment of large sums of moneybut employs relatively few peoplemaking it a capital intensiveand not a labour intensiveindustry.

17. Top 5 categories of ‘chemicals’ made no particular order Table 11.1 on textbook page 124 Basic inorganics and fertilisers Dyestuffs, paint and pigments Petrochemicals and polymers Pharmaceuticals Specialities (e.g. explosives)

18. Stages in the manufacture of a new product

19. Raw Materials and Feedstocks • A feedstockis a chemical from which other chemicals are manufactured. • Feedstocks are made from raw materials (the basic resources that the Earth supplies to us) They are: • Water –(used in hydration of ethene to ethanol) • Air –(N2 used in Haber process) • Fossil fuels – coal, crude oil and natural gas • Metallic ores– e.g. aluminium extracted from bauxite (Al2O3) • Minerals – e.g. chlorine from sodium chloride • Organic materials– from plant and animal origin e.g. veg oils

20. Batch and Continuous Processing • There are two main types of chemical processing. Batch and continuous. • In batch processingthe chemicals are loaded into the reaction vessel and the reaction is monitored. At the end of the reaction the product is collected and the reaction vessel is cleaned out ready for the next batch. • In continuous processingthe reactants are continuously added at one end of the reaction vessel and the products are removed at the other end. • Each process has advantages and disadvantages.

21. Batch Pros (advantages) • suited to smaller scale production (up to 100 tons per year) • more versatile than continuous as they can be used for more than one reaction • more suited for multi step reactions or when reaction time is long Cons (disadvantages) • possibility of contamination from one batch to the next • filling and emptying takes time during which no product, and hence no money, is being made • safety – relatively large amounts of reactants may not be controllable in the event of an exothermic reaction going wrong. An example of a batch process is used in the making of pharmaceuticals

22. Continuous Pros (advantages) • suited to large scale production (>1000 tons per year) • suitable for fast single step processes • more easily automated using computer control • smaller workforce operates round the clock 365 days per year • tend to operate with relatively low volumes of reactants allowing easy removal of excess heat energy Cons (disadvantages) • very much higher capital cost before any production can occur • not versatile, can make only one product • not cost effective when run below full capacity Some examples of a continuous process is the making of sulphuric acid, iron and poly(ethene) and ammonia.

23. Percentage Yield + Atom Economy

24. What is percentage yield and why do we use it? In chemical reactions we rarely, if ever, get the amount/quantity of products we calculate from a (balanced) chemical equation. The reasons for this can be: • at the end of the reaction there may be reactant left unconverted to product (see excess) • some reactant may be converted into a by-product • the isolation of the product may be difficult Therefore chemists like to calculate the percentage yieldof a reaction (on a small scale) to see if it makes ‘economical sense’. (good ‘money’ example on textbook pages 146/147)

25. Actual quantity obtained Percentage= Yield X 100 Theoretical quantity from equation

26. Worked Example 1 Example from white board When 5g of methanol reacts with excess ethanoic acid 9.6g of methyl ethanoate is produced. What is the percentage yield in this reaction?

27. Worked Example 2 Excess ethyne was reacted with 0.1 moles of hydrogen chloride and 4.1g of the product - 1,1 dichloroethane – were obtained. Calculate the percentage yield. Example from white board

28. Worked Example 3 10kg of nitrogen reacts with excess hydrogen producing 1kg of ammonia. Calculate the percentage yield. N2 + 3H22NH3 Example from white board

29. Atom Economy Although % yield can be used to calculate the overall efficiency of a chemical reaction, it does not take into account how much of the reactants are changed into (unwanted) by-products. Atom economy allows chemists to examine the proportion of reactants that are converted into the desired product (i.e. the chemical they want.) Atom Economy = Mass of desired product(s) x100 Total mass of reactants

30. Worked Example Calculate the atom economy for the production of ethyl propanoate, assuming that all reactants are converted into products. C2H5OH + C2H5COOH C2H5OOCC2H5 + H2O Example from white board

31. Chemical Equilibrium

32. Reversible Reactions and Equilibrium In one way reactions (example 1) the reactants change completely into products. These products do not change back into the reactants. Example 1 However, there are many reactions (example 2) in which the products can react to reform the reactants. These are called reversible reactions. Example 2 Reversible reactions give rise to equilibrium.

33. In general terms… At the start with A and B, the rate of the forward reaction is high because the concentrations of A and B are high. The rate of the back reaction is initially 0 because there are no products (C and D) yet. As the reaction proceeds the concentrations of A and B decrease while the concentrations of C and D increase. This continues until the two rates become equal. At this point the concentrationof A, B, C and D are constant and the (closed) system is at chemical equilibrium. Equilibrium is only possible in a closed system– i.e. no substances are added or removed.

34. At equilibrium, the forward and backward reactions are continuing and their rates are equal. Therefore the concentrations of A, B, C + D remain constant. This is known as dynamic equilibrium Equilibrium does not mean there are 50% reactants 50% products (i.e. equal amounts)

35. Position of Equilibrium The position of equilibrium varies from reaction to reaction. Sometimes it occurs when the forward reaction is almost complete whilst other times it occurs when the forward reaction has barely started. We use the terms; ‘equilibrium lies to the left’ when the conc. of reactants is greater than the conc. of products and ‘equilibrium lies to the right’ when the conc. of products is greater than the conc. of reactants

36. Equilibrium lies to the _____ Equilibrium lies to the _____ left right

37. Factors that affect the Position of Equilibrium As many reactions are at equilibrium, it is important we understand how to alter its position as this has a bearing on the yield of reactants to products. Factors include; • Use of a catalyst • Concentration • Pressure (of gases) • Temperature

38. 1. Use of a catalyst A catalyst lowers the activation energy between reactants and products by providing an alternative reaction path. The activation energy is lowered by the same amount for both the forward reaction and the back reaction. As a catalyst speeds up both the forward and back reactions, equilibrium is reached more quickly. However a catalysthas no effect on the position of equilibrium.

39. Le Chatelier’s Principle The effect of changes in concentration, pressure and temperature on an equilibrium can be predicted using Le Chatelier’s Principle; “If a system at equilibrium is subjected to a change, the system will adjust to oppose the effect of the change.”

40. 2. Concentration Increasing the concentration of A or B will speed up the forward reaction. The equilibrium position moves to the rightto counteract this change i.e. more C and D are produced. Decreasing the concentration of C or D will slow down the back reaction This means the equilibrium will move to the right in order to counteract the changes. i.e. more C and D are produced. Increasing the concentration of C and D or decreasing the concentration of A and B moves the equilibrium to the left.

41. right red left Increasing the conc. of either Fe3+(aq) or CNS-(aq) will result in the equilibrium position moving to the ______, using up some of the additional reactants and producing more FeCNS2+(aq). The solution will become more ____. The reverse is true i.e. decreasing the concentration of either reactant will result in less red. Increasing the concentration of FeCNS2+(aq) will result in the equilibrium position moving to the ______ to use up some of the additional product by making more Fe3+(aq) and CNS-(aq). The solution will become less red. The reverse is true i.e. decreasing the concentration of FeSCN2+(aq) will result in more red.

42. 3. Pressure (of gases) If the pressure on an equilibrium system is increased, then the equilibrium position shifts to reduce the pressure. In other words, increasing the pressure on an equilibrium system will result in the equilibrium shifting to reduce the pressure i.e. moving to the side that has the smallest number of gas particles.

43. Pressure Example 1 (g) (g) left There is 1 mole of gas on the left hand side of the reaction and 2 moles of gas on the right hand side of the reaction. Increasing the pressure on this system would result in the equilibrium position moving to the ______ i.e. consuming NO2(g) and producing more N2O4(g). The system will become a lighter in colour. Decreasing the pressure on this equilibrium system will result in the equilibrium position moving to the right i.e. the side that has the most gas particles, in order to increase the pressure. The brown colour of the system becomes darker.

44. Pressure Example 2 3 1 right left There are ____ moles of gas on the left hand side of the reaction and ______ mole(s) of gas on the right hand side of the reaction. Increasing the pressure on this system results in the equilibrium position moving to the _________. Decreasing the pressure on this equilibrium system will result in the equilibrium position moving to the _________.

45. Pressure Example 3 C(s) + H2O(g) CO(g) + H2(g) 1 2 left right There are ____ mole(s) of gas on the left hand side of the reaction and ______ mole(s) of gas on the right hand side of the reaction. Increasing the pressure on this system results in the equilibrium position moving to the _________. Reducing the pressure on this equilibrium system will result in the equilibrium position moving to the _________.

46. 4. Effect of Temperature on the Position of Equilibrium In a system at equilibrium, if the forward reaction is exothermic the back reaction must be endothermic, and vice versa. In an endothermic reaction, energy can be considered as a reactant of the reaction. If the temperature of an endothermicequilibrium system is increased, the equilibrium position shifts to use up the heat by producing more products. (movesto right) In an exothermic reaction, energy can be considered as a product of the reaction. If the temperature of an exothermic equilibrium system is increased, the equilibrium position shifts to use up the heat by producing more reactants. (moves to left)

47. Temperature Example 1 (g) (g) right left If the temperature is increased the position of equilibrium moves to the _________. If the temperature is decreased the position of equilibrium moves to the ________.

48. decreases An increase in temperature favours the back reaction and therefore _______ the concentration of methanol. This suggests that to get a high yield of methanol we should carry out the reaction at low temperature. However, low temperature means a low rate and a long time to establish equilibrium. In industry a compromise is reached at a moderately high temperature (200 to 300°C) which gives a worthwhile rate but a reduced yield of methanol.

49. Enthalpy

50. Reminder - Potential Energy Diagrams ΔH is always positive (remember you must write in the sign!) ΔH is always negative (your calculator won’t always do this for you, look at the temp. change!) ExothermicEndothermic