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| Thermal energy transferred between two objects as the result of a temperature difference. |
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| the measure of the kinetic energy of molecular motion |
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| The First Law of Thermodynamics |
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| Energy can neither be created nor destroyed |
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| Neither mass nor energy can be exchanged with the surroundings |
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| Both mass and energy can be exchanged with the surroundings |
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| Energy, but not mass can be exchanged with the surroundings |
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the force that produces the movement of an object times the distance moved
F * d |
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| Most common type of work in chemical systems |
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Definition
Expansion work aka Pressure-Volume Work or PV work
W = -P * A * d = -P * DeltaV |
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| Total Energy Change of a System |
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Definition
q + W =
q - P * DeltaV
q is (+) if the system gains heat, (-) if the system loses heat |
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| Transferred Heat at constant volume |
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Definition
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| Transferred Heat at constant pressure |
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Definition
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Definition
The "heat of the reaction"
equivalent to transferred heat at constant pressure (this is the most common type of chemical reaction in practice)
DeltaH = DeltaE + P * DeltaV
! A state function -> only enthalpy change is important (same with internal energy)
! values assume rxn is going in the direction as it is written; switch sign for reverse rxn
! it is useful to remember the equation P * DeltaV = change in mol * R * T ..so..
DeltaH = DeltaE + (Change in mol)*R*T |
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| Thermodynamic Standard State |
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Definition
| Most stable form of a substance at 1 atm pressure and at a specified temperature usually 25 degrees celsius, 1 M concentration for all substances in solution |
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| the direct conversion of a solid to a vapor without going through a liquid. Equals the sum of the heat of fusion and heat of vaporization. |
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| Products have more enthalpy than the reactants, DeltaH is positive; heat flows into the system. |
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| Products have less enthalpy than the reactants, DeltaH is negative; heat flows out of the system. |
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| A stirrer, thermometer and loose-fitting lid to keep the contents at atmospheric pressure; thermometer measures change in temperature to calculate enthalpy. |
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Definition
Measures enthalpy for a combustion rxn.
Used to calculate heat capacity:
C = q/DeltaT
!Note that capital C is for heat capacity |
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Definition
The amount of heat required to raise the temperature of an object or substance a given amount.
An extensive property so its value depends on both the size of an object and its composition. |
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Definition
the amount of heat necessary to raise the temperature of 1 g of a substance by 1 degree celsius; the amount of heat necessary to raise the temperature of a given object, then is the specific heat times the mass of the object times the rise in temperature:
q = (specific heat) * (Mass of substance) * DeltaT = C * m * DeltaT |
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Definition
| The overall enthalpy change for a reaction is equal to the sum of the enthalpy changes for the individual steps in the reaction. |
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| Standard Heats of Formation |
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Definition
| The enthalpy change for the formation of 1 mol of a substance in its standard state from its constituent elements in their standard states. |
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Definition
A Process that, once started, proceeds on its own without a continuous external influence.
Favored by a decrease in H (-) and increase in S (+) |
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Definition
By contrast takes place only in the presence of a continuous external influence.
Favored by an increase in H (+) and decrease in S (-) |
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Definition
| (S) the amount of molecular randomness in a system |
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| Bond Dissociation Energies |
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Definition
(D) a standard enthalpy change for the corresponding bond-breaking reaction
! always positive because energy must always be put into bonds to break them
We can calculate an approximate enthalpy change for any reaction by subtracting the sum of the bond dissociation energies in the products from the sum of the bond dissociation energies in the reactions. |
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Definition
| Standard enthalpy values, but per gram instead of per mole, to compare efficiency |
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Definition
DeltaG = DeltaH - T * DeltaS
If negative => process is spontaneous
If 0 => process is at equilibrium
If positive => process is nonspontaneous |
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Definition
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| The Third Law of Thermodynamics |
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Definition
| The entropy of a perfectly ordered crystalline substance at 0 K is zero |
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| The Second Law of Thermodynamics |
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Definition
| In any spontaneous process, the total entropy of a system and its surroundings always increases |
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| The standard entropy of reaction |
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Definition
For the rxn aA + bB -> cC + dD,
DeltaS = [(c * S value of C) + (d * S value of D)] - [(a * S value of A) + (b * S value of B)] |
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Definition
| DeltaS = DeltaS for system + DeltaS for surroundings |
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| Total entropy change for surroundings |
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Definition
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| Standard State Conditions |
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Definition
Solids, liquids, and gases in pure form at 1 atm pressure
Solutes at 1 m concentration
A specified temperature, usually 25 degrees celsius |
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| Free-energy using REaction Quotient |
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Definition
DeltaG + R * T lnQ
R is the gas constant
T is the absolute temperature in Kelvins and Q is the reaction quotient |
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Term
| Free Energy for a reaction moving towards equilibrium |
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Definition
At equilibrium DeltaG equals 0, so...
Standard DeltaG = -R * T * lnK
lnK > 0 K>1 mainly products
lnK < 0 K<1 mainly reactants
lnK = 0 K=1 comparably amounts of reactants and products |
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