Thermochemistry

1. Thermochemistry

2. Relative Space Occupied

3. Liquid Properties • Liquids are fluids. (So are gases.) • Liquids have higher densities than their corresponding gases. • Liquids can diffuse readily. • Liquids are relatively incompressible. • Liquids have a definite volume but no definite shape.

4. Liquids Solids IMF Strength Stronger than very strong in gases Fluid yes no Density high high Compressibility No No Diffusion Slower than in No gases

5. water mercury B. Liquid Properties • Capillary Action • movement of a fluid against the pull of gravity due to adhesive/cohesive forces and pressure differences

6. B. Liquid Properties • Surface Tension • cohesive forces between the molecules at the surface of a liquid, draws surface molecules inward

7. Solid Properties • Even more incompressible than liquids • Cannot diffuse well at all into neighboring media • Have a definite shape and volume • Have molecules that are primarily rotating or vibrating in place • Are usually more dense than their corresponding liquids

8. Types of Solids • Ionic crystals – formed by atoms that are ionically bound together (metal to nonmetal) – hard, brittle, high melting points – examples: NaCl, MgF2 • Covalent molecular crystals – formed by covalent molecules (nonmetal to nonmetal) – generally have low melting points, relatively soft, easily vaporized – example: dry ice (CO2)

9. Phase Changes

10. Phase Changes • Exothermic phase changes -release energy (heat). Examples: Condensation, Deposition, Freezing • Endothermic phase changes -require the input of energy (heat). Examples: Sublimation, Vaporization, Melting

11. Phase Changes • Vaporization = phase change of a liquid to a gas • Two types: • Evaporation = vaporization of molecules at the surface of a liquid – molecules at the surface gain enough energy to overcome intermolecular forces (IMF’s) • Boiling = vaporization of molecules throughout the entire volume of the liquid • Condensation = phase change from gas to liquid

12. Equilibrium • Dynamic condition in which two opposing changes occur at equal rates in a closed system.

13. Equilibrium • Equilibrium vapor pressure = pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature • Remember, pressure = number of collisions of the gas particles

14. Vaporization • Volatile liquids evaporate readily due to weak intermolecular forces. Nonvolatile liquids evaporate slowly. Examples: alcohol and acetone • Boiling occurs when equilibrium vapor pressure equals atmospheric pressure. What is atmospheric pressure? What happens to the boiling point when you increase altitude?

15. ▶ What temperature does water boil at on Everest - Live Experiments (Ep 31) - Head Squeeze - YouTube.mht

16. v.p. IMF v.p. temp A. Phase Changes • Vapor Pressure pressure of vapor above a liquid at equilibrium p.478 v.p. • depends on temp. & IMF’s • directly related to volatility related to volatility temp

17. IMF Patm b.p. A. Phase Changes • Boiling Point • temp at which the v.p. of liquid equals external pressure • depends on Patm & IMF’s • atmospheric pressure at sea level = 1atm b.p.

18. IMF m.p. A. Phase Changes • Melting Point • equal to freezing point • Which has a higher m.p.? • covalent or ionic? ionic

19. Heating Curve

20. Phase Diagrams

21. Phase Diagrams • Graph of pressure versus temperature that shows the conditions under which the phases of a substance exist. • Triple point = temperature and pressure conditions at which solid, liquid, and vapor of a substance can coexist at equilibrium • Critical temperature = temperature above which the substance cannot exist as a liquid • Critical pressure = lowest pressure at which the substance can exist as a liquid at the critical temperature

22. What can stress a reaction? • The addition or removal of reactants or products • The addition or removal of heat. • The increase or decrease of pressure (if gases are involved in the reaction).

23. Thermochemistry Thermochemistry is the study of heat change in chemical reactions. • Exothermic reactions -give off heat as a product of the reaction – ex. Burning wood 2H2(g) + O2(g) 2H2O (l) + energy • Endothermic reactions -must absorb heat from the surroundings as a reactant –ex. cold packs for sports injuries energy + 2HgO (s) 2Hg (l) + O2(g)

24. Thermochemistry • Heat travels from hot to cold surfaces. With arrows, show the transfer of heat for a hot cup of coffee and an ice pack and label what is endothermic and exothermic.

25. Thermochemistry Temperature – a measure of the average kinetic energy of the particles in a sample of matter Heat – the energy transferred between samples of matter because of a difference in their temperatures

26. Thermochemistry • q = mcDt • q = heat absorbed or released – in Joules (J) • c = the specific heat of a substance - the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius. • m = mass of sample (g) • Dt = tfinal – tinitial

27. Gold has a specific heat of 0.129 J/(g×C). How many joules of heat energy are required to raise the temperature of 15 grams of gold from 220C to 850C? • An unknown substance with a mass of 100 grams absorbs 1000 J while undergoing a temperature increase of 150C. What is the specific heat of the substance?

28. Heat of Fusion (Hfus) • energy required to melt 1 gram of a substance at its m.p. • Heat of Vaporization (Hvap) • energy required to boil 1 gram of a substance at its b.p.

29. What is the heat in Joules to convert 16g of ice at -20 C to steam at 140 C? heat of fusion of water = 334 J/gheat of vaporization of water = 2257 J/gspecific heat of ice = 2.09 J/g·°Cspecific heat of water = 4.18 J/g·°Cspecific heat of steam = 2.09 J/g·°C

30. Homework 5/9/2017 What is the heat in Joules required to convert 25 grams of -10 ̊̊C ice into 150 ̊̊C steam?

31. Thermochemistry • Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. • DH = H (products) – H (reactants) • DH = heat given off or absorbed during a reaction at constant pressure

32. Hproducts < Hreactants Hproducts > Hreactants DH > 0 DH < 0

33. H2O (s) H2O (l) DH = 6.01 kJ Thermochemical Equations Is DH negative or positive? System absorbs heat Endothermic DH > 0 and is positive 6.01 kJ are absorbed for every 1 mole of ice that melts at 00C and 1 atm. 6.4

34. DH = -890.4 kJ CH4(g) + 2O2(g) CO2(g) + 2H2O (l) Thermochemical Equations Is DH negative or positive? System gives off heat Exothermic DH < 0 and is negative 890.4 kJ are released for every 1 mole of methane that is combusted at 250C and 1 atm. 6.4

35. Reaction Pathways

36. Hess’s Law: When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

37. Hess’s Law Calculate H for the reaction: C2H4(g) + H2(g) C2H6(g) Given: C2H4(g) + 3O2(g) 2CO2(g) + 2H2O(l) H = -1411 kJ 2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(l) H = -1560 kJ 2H2(g) + O2(g) 2H2O(l) H = -285.8 kJ

38. Hess’s Law • Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values: N2H4(l)  +  CH4O(l)  →  CH2O(g)  +  N2(g)  +  3H2 (g) • 2NH3(g)  →  N2H4(l)  +  H2(g)                 ΔH = 22.5 kJ 2NH3(g)  →  N2(g)  +  3H 2(g)                 ΔH = 57.5 kJ CH2O(g) +  H2(g)  →  CH 4O(l)               ΔH = 81.2 kJ

39. Calculating Enthalpy using Bond Energies • Bond Energy is the amount of energy that would be required to break the bond. • Similarly, it is also the amount of energy that is released when the bond is formed • Every chem rxn involves breaking and forming bonds

40. The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy

41. Bond Energies Steps: • Draw Lewis diagrams to determine the number and type of bonds in the molecules • Look up the bond energies for each bond. • Add the reactants (don’t forget to multiple by molar coefficients). • Add the products (don’t forget to multiple by molar coefficients) • ΔH = Σenergy of bonds broken – Σenergy of bonds formed

43. Consider the reaction: H2 + F2 2HF Determine the enthalpy change for the reaction.

44. Consider the following equation for the reaction of methane with chlorine. CH4(g) + 3Cl2(g) → CHCl3(g) + 3HCl(g) Use bond energies to estimate the enthalpy change for the reaction

45. Homework 5/16 Use bond energies to estimate the enthalpy change for the reactions below. • N2(g) + 3H2(g)  2NH3(g) • CH4 (g)+  2 O2 (g)   CO2 (g)+ 2 H2O (l)