1. Thermochemistry

2. Energy Energy is the ability to do work or transfer heat.
Energy used to cause an object that has mass to move is called work.
Energy used to cause the temperature of an object to rise is called heat.
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3. Heat Energy can also be transferred as heat.
Heat flows from warmer objects to cooler objects.
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4. First Law of Thermodynamics
Energy is neither created nor destroyed.
In other words, the total energy of the universe is a constant; if the system loses energy, it must be gained by the surroundings, and vice versa.
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5. E, q, w, and Their Signs
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6. Exchange of Heat between System and Surroundings
When heat is absorbed by the system from the surroundings, the process is endothermic.
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7. Exchange of Heat between System and Surroundings
When heat is absorbed by the system from the surroundings, the process is endothermic.
When heat is released by the system into the surroundings, the process is exothermic.
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8. Enthalpy
If a process takes place at constant pressure (as the majority of processes we study do) and the only work done is this pressure-volume work, we can account for heat flow during the process by measuring the enthalpy of the system.
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9. Endothermicity and Exothermicity
A process is endothermic when H is positive.
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10. Endothermicity and Exothermicity
A process is endothermic when H is positive.
A process is exothermic when H is negative.
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11. Enthalpy of Reaction
This quantity, H, is called the enthalpy of reaction, or the heat of reaction.
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12. Enthalpies of Formation
An enthalpy of formation, Hf, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms.
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13. Standard Enthalpies of Formation
Standard enthalpies of formation, Hf°, are measured under standard conditions (25 °C and 1.00 atm pressure).
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14. Calorimetry
Since we cannot know the exact enthalpy of the reactants and products, we measure H through calorimetry, the measurement of heat flow.
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15. Heat Capacity and Specific Heat
The amount of energy required to raise the temperature of a substance by 1 K (1C) is its heat capacity.
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16. Heat Capacity and Specific Heat
We define specific heat capacity (or simply specific heat = cp) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.
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17. Heat Capacity and Specific Heat
Specific heat, then, is
Specific heat =
heat transferred
mass  temperature change
cp = Q
m  T
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18. Phase Changes
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19. Energy Changes Associated with Changes of State
The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.
The temperature of the substance does not rise during a phase change.
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20. Phase Diagrams
Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.
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21. Phase Diagrams The circled line is the liquid-vapor interface.
It starts at the triple point (T), the point at which all three states are in equilibrium.
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22. Phase Diagrams
It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.
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23. Phase Diagrams
Each point along this line is the boiling point of the substance at that pressure.
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24. Phase Diagrams
The circled line in the diagram below is the interface between liquid and solid.
The melting point at each pressure can be found along this line.
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25. Phase Diagrams
Below the triple point the substance cannot exist in the liquid state.
Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.
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26. Phase Diagram of Water
Note the high critical temperature and critical pressure.
These are due to the strong van der Waals forces between water molecules.
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27. Phase Diagram of Water The slope of the solid-liquid line is negative.
This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.
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28. Phase Diagram of Carbon Dioxide
Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.
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