Thermochemical Principles

1. Thermochemical Principles

2. What happens in a chemical reaction? • When a reaction takes place, bonds are broken and new chemical bonds are made: • Breaking bonds is an endothermic process, because energy has to be taken in from the surroundings to break the bonds • Making bonds is an exothermic process, because energy is released in the formation of new chemical bonds • It is the balance between exothermic and endothermic reactions which decides the overall reaction type.

3. In or out?

4. Heat • Enthalpy is a measure of the chemical potential energy stored in the bonds of the substances involved. The amount of energy released or absorbed depends on the amount of the reagents involved. Therefore the unit for enthalpy is kJ/mol. • Heat is the sum of the kinetic energies of all the particles in a sample. • Temperature is proportional to the average kinetic energy of the sample. • The particles in a spark has high average kinetic energy (i.e. temperature) but will not burn you because the heat is low due to the few particles involved.

5. Specific heat capacity • Different substances absorb heat to a different extent, some absorb lots of heat (e.g. water) and some don’t absorb much heat (metals etc). The capability for a substance to absorb heat is called its specific heat capacity. Water has a specific heat capacity of 4.18J/g/K, which means that it takes 4.18J of energy to raise 1g of water by 1 degrees. • The absolute temperature is temperature on the Kelvin scale where 0 oC is 273K • At 0K (absolute zero) the kinetic energy of the particles is zero.

6. Boltzmann energy graph

7. States or phase change • Solid, liquid and gas are considered a phase of a substance but also referred to as allotropes (different forms of the same element) • Changing of state involves addition of energy • Particles move more rapidly when heated • Temperature is proportional to the average kinetic energy of particles • At absolute zero (0 K, -273 oC), the kinetic energy of particles is zero

8. Heating / Cooling Curves F Temperature D E B C A • On heating, any pure substance goes through patterns of constant temperature and temperature increase •  The melting point and boiling points of a substance can be used to indicate the strength of the forces between particles in the solid or liquid phase Time

9. Heating/cooling curve F Temperature D E B C A • A – B heating involves particles vibrating more quickly –kinetic energy of molecules is increasing. • B – C Phase change from solid → liquid (fusion) Heat energy is absorbed and used to weaken intermolecular bonds between molecules ( This is the enthalpy of fusion – Latent heat of fusion) • C – D Periods where kinetic energy of molecules is increasing. • E – F Molecules in solid vibrate quicker, molecules in liquid and gas move around more quickly • D – E Phase change from liquid → gas (vaporisation) Heat energy is absorbed and is used to completely break intermolecular bonds between molecules ( This is the enthalpy of vaporisation – Latent heat of Vaporisation)

10. Definitions • Enthalpy of fusion (or molar heat of fusion or latent heat of fusion): • Enthalpy of sublimation (or molar heat or latent heat of sublimation) • Enthalpy of vaporisation (or molar heat of vaporisation or latent heat of vaporisation)

11. Latent Enthalpy (Heat) of Fusion (∆fusHo) Involves the phase change solid → liquid • Provides a measure of the strength of the force holding the particles together in the solid phase • “The heat required to fuse / melt 1 mole of a substance at its standard state” • Example H2O (s) ↔ H2O (l) ∆fusHo = 6.0 kJ mol-1 • A greater ∆fusHo indicates stronger intermolecular forces between particles and results in a higher melting point

12. Latent Enthalpy (Heat) of Vaporisation (∆vapHo) Involves the phase change liquid → gas • Provides a measure of the strength of the force holding the particles together in the liquid phase • “The heat required to change 1 mole of a substance in its standard state from a liquid to a gas” • Example H2O (l) ↔ H2O (g) ∆vapHo = 40.7 kJ mol-1 • A greater ∆vapHo indicates stronger intermolecular forces between particles and results in a higher melting point

13. Latent Enthalpy (Heat) of Sublimation (∆subHo) Involves the phase change solid → gas • Provides a measure of the strength of the force holding the particles together • “The heat required to sublimate (Change form solid to gas)1 mole of a substance at its standard state”

14. Other points: • ΔfusHoand ΔvapHoare always positive forces are being broken so energy is absorbed • ΔvapHo is significantly more than ΔfusHo as many more attractive forces are being broken as particles completely separate • The greater the ∆fusHo the higher the melting point • The greater the ∆vapHo the higher the boiling point • Practical

15. Enthalpy PowerPoint PresentationINDEX Start Show Enthalpy and Standard Enthalpy definitions Hess’s Law Enthalpy of Formation Enthalpy of Combustion H reaction from H formation Enthalpy of neutralisation Enthalpy of Solution and Hydration of ions Experimental Determination of H - theory Example calculations Bomb Calorimeter Bond Energies; definitions and calculations Changes of State and Enthalpy Changes Dept Index

16. Enthalpy Change of a reaction Hr is the heat energy exchanged with the environment (the surroundings) at constant pressure. A B Hr = - 100 kJ In this exothermic reaction A has more heat energy than B, so Hris negative. A 100 Energy kJ B Hr= Heat energy in PRODUCTS - Heat energy in reactants Exothermic reactions: the chemicals give out heat to the environment; Hris negative - ve Endothermic reactions: the chemicals take in heat from the environment; Hris positive + ve

17. Standard Enthalpy Change Hr The heat exchanged with the environment (the surroundings) at a pressure of 100 kPa, a temperature of 298 k, and usually for 1 mole of substance

18. Standard Enthalpy ChangeHr The heat exchanged with the environment (the surroundings) at a pressure of 100 kPa, and a temperature of 298 k and usually for 1 mole of substance A B Hr= -100 kJ mol-1 In an exothermic reaction, the chemical A loses 100 kJ of heat forming B (more stable) A Energy100 kJoules B A starts at 298 k, Heat is given out to the surrounding, B cools down to 298 k; Then Hr= -100 kJ, if 1 mole is involved at 100 kPa pressure

19. Hess’s Law The enthalpy change for a chemical reaction is the same which ever route is taken from reactants to products. Enthalpy by route 1 = -200 kJ C Route 2 +100 By route 2( A-C-D-B) = +100 -120 - 180 - 120 Energy kJ A = -200 kJ D Route 2 -200 Enthalpy Change by route 1 = Enthalpy Change by route 2 Route 1 -180 B

20. The enthalpy change for the reverse reaction has the same magnitude, but opposite sign to the forward reaction A B Hr= + 200 kJ C D Hr= - 275 kJ C B +275 kJ Energy + 200 kJ - 200 kJ -275 kJ A D

21. Photosynthesis is an endothermic process, but respiration is exothermic; energy coming from sunlight. 6CO2(g) + 6H2O(l) = C6H12O6(s) + 6O2(g) Hr= + 5347 kJ mol-1 C6H12O6(s) + 6 O2(g) Photosynthesis Respiration Energy + 5347 kJ - 5347 kJ 6CO2(g) + 6H2O(l) Photosynthesis is arguably the most important chemical reaction for life; it produces oxygen and glucose from which animals derive much of their energy.

22. Kinetic theory • AIM – to investigate the energy needed for a change of state

23. Complete investigation 4.1

24. Melting and freezing • If energy is supplied by heating a solid, the heat energy causes stronger vibrations until the particles eventually have enough energy to break away from the solid arrangement to form a liquid. • The heat energy required to convert 1 mole of solid into a liquid at its melting point is called the molar heat of fusion or enthalpy of fusion and is represented by the symbol ΔHfus.

25. You may see it described as latent heat as the heat is hidden in the bonding and not represented in a change of temperature. • When a liquid is cooled, at some temperature, the motion of the particles is slow enough for the forces of attraction to be able to hold the particles as a solid. As the new bonds are formed, heat energy is evolved.

26. Boiling and condensing • If more heat energy is supplied, the particles eventually move fast enough to break all the attractions between them, and the liquid boils. • The heat energy required to convert 1 mole of liquid into a gas at its boiling point is called the molar heat of vaporisation or enthalpy of vaporisation and is represented by the symbol ΔHvap.

27. Again you may see this referred to as latent heat. • If the gas is cooled, at some temperature the gas particles will slow down enough for the attractions to become effective enough to condense it back into a liquid. As forces are re-established, heat energy is released.

28. The evaporation of a liquid • The average energy of the particles in a liquid is governed by the temperature. The higher the temperature, the higher the average energy. But within that average, some particles have energies higher than the average, and others have energies lower than the average.

29. Some of the more energetic particles on the surface of the liquid can be moving fast enough to escape from the attractive forces holding the liquid together. They evaporate.

30. Evaporation only takes place on the surface of the liquid. That's quite different from boiling which happens when there is enough energy to disrupt the attractive forces throughout the liquid. That's why, if you look at boiling water, you see bubbles of gas being formed all the way through the liquid.

31. If you look at water which is just evaporating in the sun, you don't see any bubbles. Water molecules are simply breaking away from the surface layer. • Eventually, the water will all evaporate in this way. The energy which is lost as the particles evaporate is replaced from the surroundings. As the molecules in the water jostle with each other, new molecules will gain enough energy to escape from the surface.

32. Sublimation • Solids can also lose particles from their surface to form a vapour. Sublimation is the direct change from solid to vapour (or vice versa) without going through the liquid stage. • The forces of attraction in most solids are too high to allow much loss of particles from the surface. However, there are some which do easily form vapours. • Examples include naphthalene (moth balls) and solid carbon dioxide - "dry ice".