Thermochemical equations
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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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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


Phase Changes

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


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


Enthalpy of Combustion


Phase Changes

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


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)


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)


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.


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


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


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


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


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


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


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


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


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


Phase Change Graph


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


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


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


Phase Diagrams

  • Different for each substance because of the different boiling/freezing points


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