Enthalpy and Fuels

1. Using and Controlling Reactions Enthalpy and Fuels

2. Energy • The capacity for doing work • Energy is inter-convertible between different forms

3. Chemical Energy • The energy from reactions comes from the breaking and forming of bonds between atoms. • When chemical bonds are formed energy is released. • When chemical bonds are broken energy is required. • This energy can be heat, light or electrical.

5. Heat Energy • Measurement of the heat change in chemical reactions is called calorimetry. • The heat energy involved in a chemical reaction can be observed from the temperature changes resulting from that reaction. • A calorimeter is used to measure these heat changes in chemical reactions.

7. Exothermic reactions • Exo: Out Thermic: Heat • Releases heat energy to the surroundings. • Temperature of the surroundings increases. • Energy released when forming bonds > Energy used to break bonds

8. Exothermic reactions Exothermic H < 0

9. Endothermic reactions • Endo: To take in • Heat energy is taken from the surroundings. • Temperature of the surroundings decreases. • Energy required to break bonds > Energy released when forming bonds.

10. Endothermic reactions Endothermic H > 0

12. Enthalpy • The quantity of heat released or absorbed when specific amounts of substances react is called the Heat of Reaction. • The difference in heat energy between products and reactants at constant pressure is called the Enthalpy Change. H = Hproducts – Hreactants

13. Enthalpy • The Molar Enthalpy Change for a reaction is the quantity of heat energy released or absorbed when 1.00 mole of a substance reacts in a chemical reaction under constant pressure. • Measured in kJmol-1 or kJ g-1 • The quantity of heat released or absorbed is directly proportional to the number of moles of reactants and products.

14. Enthalpy • The quantity of heat energy released or absorbed can be calculated by Heat = n x H • EXAMPLE: Calculate the enthalpy change for the neutralisation of 112.2g of KOH (H= – 56.0kJmol-1) • n = m/M = 112.2/56.1 = 2.00 moles • Energy released = 2 x 56.0 = 112 kJ released

15. Thermochemical Equations • An equation showing the enthalpy change is worth 4 marks. • Correct formulae • Must be BALANCED • The physical states of matter MUSTbe shown. • The sign MUST be included with the H value. 6CO2(g) + 6H2O(l) --> C6H12O6(s) + 6O2(g) ∆H = +2800 kJ

16. Thermochemical Equations • If the coefficients are doubled the H value must be doubled. • CH4(g)+2O2(g)→ CO2(g)+2H2O(l) H = -890 kJmol-1 • 2CH4(g)+4O2(g)→2CO2(g)+4H2O(l) H = -1780 kJ • If the reaction is reversed then the H value is equal but opposite in sign • N2(g) + 3H2(g) → 2NH3(g) H = - 92.4 kJmol-1 • 2NH3(g)→ N2(g) + 3H2(g) H = + 92.4 kJmol-1

17. Enthalpy of Combustion • The molar enthalpy of combustion of a substance is the quantity of heat energy released when 1.00 mole of a pure element or compound is burnt completely in oxygen under constant pressure. • The products are gaseous carbon dioxide and liquid water.

18. Enthalpy of Combustion • When a fuel is burnt in air rather than oxygen, the non reactive gases (N2) absorb some of the heat energy released. Lower flame temperature. CH4(g) + 2O2(g)→ CO2(g) + 2H2O(l)ΔH= -890kJ

19. Enthalpy of Solution • The molar enthalpy of solution is the quantity of heat energy released or absorbed when 1.00 mole of substance dissolves in sufficient solvent so that further dilution causes no further change in heat energy. NH4NO3(s) + aq  NH4+(aq)+ NO3–(aq) H = +26kJ

20. Enthalpy of Neutralisation • The molar enthalpy of neutralisation is the quantity of heat energy released when 1.00 mole of hydrated hydrogen ions from an acid is neutralised by 1.00 mole of hydrated hydroxide ions from an alkali. • Heat change per mole of water formed when dilute solutions of acids and bases are mixed. H+(aq) + OH–(aq) H2O(l) H= –57kJ

21. m = mass of water cp (specific heat capacity of water) = 4.18 J g-1oC-1 ΔT = change in temperature n = number of moles of reactant Calculating Enthalpy values • Experimental data can be used to calculate the enthalpy change for chemical reactions.

22. Assumptions • All of the heat energy produced or absorbed by the reaction is transferred to or from the known mass of water or aqueous solution in the calorimeter. No exchange of heat occurs with the calorimeter or the surrounding air. • The reaction occurs quickly enough for the max or min temperature of the liquid in the calorimeter to be reached before the liquid begins to return to room temperature.

24. Approximations • The specific heat capacity of any aqueous solution is taken to be the same as that for water.(i.e. 4.18 J g-1oC-1) • The density of any aqueous solution in the calorimeter is taken to be the same as the density of water. 1mL=1g • Always read the question carefully in any test or exam and use the specific heat capacity given.

25. Fuels • Useable fuel source provides maximum energy from minimum amount of material • Chemical fuels release energy from combustion. • Carbon based Fuels can be divided into: • Fossil Fuels e.g. Coal, Oil, Natural Gas • Biofuels e.g. Ethanol, Biodiesel, Biogas

26. Advantages of Burning Fossil Fuels • High enthalpy of combustion (High energy density) • Can be burnt at point of use. i.e. Internal combustion engine, home heating etc (80% of fossil fuels are used this way) • Versatile and inexpensive. • Provide governments with tax income.

27. Disadvantages of Burning Fossil Fuels • Finite reserves of coal, oil and gas. Oil and gas expected to run out before end of this century. • Burning releases CO2 which contributes to the enhanced greenhouse effect. • Other pollutants are released due to contaminants in the fuel or when incomplete combustion occurs e.g. soot, CO, SO2, NOx

28. Disadvantages of Burning Fossil Fuels • Fossil fuels are used as feedstock (raw material) for the chemical industry • Used to make synthetic fibres, solvents, paints, detergents, dyes and plastics • Burning fossil fuels for energy, reduces raw material required for these industries

29. Advantages of Burning Biofuels • Renewable, approximately Carbon neutral • Less harmful air pollutants released due to less fuel contaminants • Biodegradable • Can be domestically produced

30. Disadvantages of Burning Biofuels • Non renewable fossil fuels can be used in the production of ethanol • Slightly decreases fuel economy • May solidify at low temperatures • Not widely available • Many developing countries are growing plants to produce biofuels rather than food crops leading to increased famine

32. Complete Combustion • Occurs with unlimited oxygen • Products are Carbon dioxide and water • Maximum quantity of heat energy produced when a fuel undergoes complete combustion

33. Incomplete Combustion • Occurs if the oxygen supply is limited. • Additional products formed include carbon (Soot) and carbon monoxide. • Less heat energy is produced than complete combustion.

34. Carbon Monoxide CO • Persists in the environment for long periods of time. • Strongly bonds to haemoglobin in blood preventing it from carrying oxygen to cells. • Symptoms of CO poisoning range from visual disturbance, headaches, fatigue to unconsciousness and death dependant on concentration of Carbon monoxide.

35. Soot C • Particulate pollution. Settles on objects. • Clogs air filters/ inlets lowering efficiency of engines and burners. • Coats plant leaves reducing photosynthesis. • Damages the respiratory system and aggravates asthma, bronchitis and other lung diseases.

36. Comparing Fuels • Fuel comparisons can be made per mole, per gram or per litre of fuel. • Molar enthalpy of combustion: Energy released per mole of fuel (kJmol-1) • Energy density: Energy released per gram (kJg-1) or per litre (kJL-1) • Per gram: Energy density = H/M • Per Litre: Energy density(liquid fuel) = H/M x  (where  = the density of the fuel in gL-1)