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Thermochemical Equations

Thermochemical Equations. Thermochemical equations are balanced chemical equations that include the physical states of all reactants and products and the energy change. 2H 2 O(l)  2H 2 (g) + O 2 (g) ΔH = 572 kJ CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g) ΔH = -802 kJ. Phase Changes.

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Thermochemical Equations

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  1. Thermochemical Equations • Thermochemical equations are balanced chemical equations that include the physical states of all reactants and products and the energy change. • 2H2O(l)  2H2(g) + O2(g) ΔH = 572 kJ • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) ΔH = -802 kJ

  2. Phase Changes • Occurs when energy is added or removed from a system and the substance can go from one physical phase to another

  3. Enthalpy of Combustion • Enthalpy heat of combustion (ΔHcomb)is the enthalpy change for the complete burning of one mole of the substance. -- carried out under standard conditions which are one atmospheric pressure (1 atm) and 298K (250C) - C6H1206(s) + 6O2(g)6CO2(g)+ 6H2O(l) ΔHocomb= -2808kJ

  4. Enthalpy of Combustion

  5. Phase Changes • Occurs when energy is added or removed from a system and the substance can go from one physical phase to another

  6. Changes of State • Molar enthalpy (heat) of vaporization (ΔHvap)is the heat required to vaporize one mole of liquid. -- think of water vaporizing from your skin after you take a hot shower. Your skin provides the heat needed to vaporize the water and as the water absorbs the heat you feel cool (shiver) ΔHvap = -ΔHcond (condensation)

  7. Changes of State • Molar enthalpy (heat) of fusion (ΔHfus) is the heat required to melt one mole of a solid substance. --think of ice in a drink. The drink cools as it provides the heat for the ice to melt ΔHfus= - ΔHsolid (solidification—freezing)

  8. Changes of State ??What do you notice about the magnitude of the molar enthalpy of vaporization versus the molar enthalpy of fusion? The molar enthalpy of vaporization for a substance is much larger than the molar enthalpy of fusion for the same substance. It takes much more energy to change a substance from a liquid to a gas than it does to change a solid to a liquid.

  9. Endothermic Phase Changes Melting • The energy absorbed to melt a solid is not used to raise the temperature of that solid • The energy instead disrupts the bonds holding the solid’s molecules together and cause the molecules to move into the liquid phase

  10. Endothermic Phase Changes • The amount of energy required to melt one mole of a solid depends on the strength of the forces that hold the solid together • The melting point of a crystalline solid is the temperature at which the forces holding its crystal lattice together are broken and it becomes a liquid

  11. Endothermic Phase Changes Vaporization • Particle that escape from the liquid enter the gas phase and those liquids at room temperature the gas phase is called vapor • Vaporizationis the process by which a liquid changes into a gas or vapor • Once the solid becomes a liquid then and only then does the temperature of the substance begin to increase

  12. Endothermic Phase Changes • When vaporization takes place only at the surface of the liquid it is called evaporation • Evaporation is the method by which the human body maintains and controls its temperature

  13. Endothermic Phase Changes Sublimation • Is the process by which a solid changes directly to a gas without first becoming a liquid • Dry ice (CO2) and snow are the most common examples

  14. Endothermic Phase Changes • If ice cubes are left in the freezer for extended periods of time, they will eventually sublime and become smaller • This process is also helpful in freeze drying foods for hikers and astronauts

  15. Exothermic Phase Changes Condensation • When a vapor molecule loses energy its velocity is reduced therefore colliding more with other molecules to form a liquid • Condensation is the process by which a gas or vapor becomes a liquid and it is the reverse action of vaporization

  16. Exothermic Phase Changes Deposition • Is the process by which a substance changes from a gas or vapor to a solid without first becoming a liquid • It is the reverse action of sublimation • The formation of snow crystals high up in the atmosphere is an example

  17. Exothermic Phase Changes Freezing Point • Is the temperature at which a liquid is converted into a crystalline solid • The same temperature as the melting point of a given substance

  18. Phase Change Graph

  19. Phase Change Graph • Graph shows the energy required to go from one phase to the other • Where the graph inclines, potential energy is at its greatest and temperature is increasing • Where the graph plateaus (flatregion) kinetic energy is at its greatest but the temperature remains constant

  20. Phase Diagrams • A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure

  21. Phase Diagrams • The triple point is the point on a phase diagram that represents the temperature and pressure at which three phases of a substance can coexist • The critical point is the point that indicates critical pressure and temperature above which water cannot exist as a liquid

  22. Phase Diagrams • Different for each substance because of the different boiling/freezing points

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