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Contents
Enthalpy
- The heat produced by a chemical reaction can be calculated from the First Law of Thermodynamics and the concept of enthalpy
First Law of Thermodynamics
- First Law
- The change in a system’s internal energy equals the heat added to the system from its surroundings minus the work done by the system on its surroundings.
- That is, change in internal energy ΔU = heat added Q – work done W
- ΔU = Q – W
- Definitions
- The internal energy of a system is the total kinetic and potential energies of its constituent particles.
- Heat is the energy transferred from one body to another as the result of a difference in temperature.
- Work is the energy transferred as the result of a force moving an object.
- View First Law
Enthalpy
- Enthalpy = EN-THOWL-pea
- The enthalpy of a system equals its internal energy + the pressure of the system times its volume:
- H = U + PV
- PV is in units of energy:
- P = force / length2 and V = length3. So PV = force x length = work.
- Thus enthalpy is the internal energy of a system plus the work it does in the form of pressure times volume.
- The standard enthalpy of formation of a compound is the change of enthalpy during the formation of one mole of the substance from its constituent elements under standard conditions:
- ΔH = ΔU + PΔV
- Standard Conditions:
- Temperature of 293.15 K = 68 degrees Fahrenheit
- Pressure of 101325 pascals = 1 atm = pressure at sea level
- Substances are in their standard state: liquid, gas, solid
- Some standard enthalpies (from wikipedia.org/wiki/Standard_enthalpy_of_formation)
- CO2 (gas) = −393.509
- O2 (gas) = 0
- All elements have zero enthalpies
- H2O(liquid) = −285.8
- O3 (gas) = +143 (ozone)
Heat produced by combustion of methane
Argument that the combustion of methane produces 890.345 kJ/mol of heat
- Premise 1: If the only work done by a chemical reaction is a change of volume at constant pressure, the heat added to the reaction (Q) equals the change in enthalpy (ΔH).
- The proof of Q=ΔH principle is below.
- Premise 2: Change in enthalpy (ΔH) = -890.345 kJ/mol
- Reactant Side
- Standard Enthalpy of Formation for CH4 = -74.848 kJ/mol
- Standard Enthalpy of Formation for O2 = 0 kJ/mol
- Total (-74.848 + 2 (0)) = -74.848 kJ/mol
- Product Side
- Standard Enthalpy of Formation for CO2 = -393.513 kJ/mol
- Standard Enthalpy of Formation for H2O = -285.84 kJ/mol
- Total (-393.513 + 2 (-285.84)) = -965.193 kJ/mol
- Change in Enthalpy = -965.193 – (-74.848) = -890.345 kJ/mol
- Reactant Side
- Premise 3: The only work done by the combustion of methane is a change of volume at constant pressure
- Experimental conditions
- Therefore, the heat added to the reaction is -890.345 kJ/mol, meaning the reaction produced 890.345 kJ/mol of heat.
Proof of the Q=ΔH Principle
- Principle
- If the only work done by a chemical reaction is a change of volume at constant pressure, the heat added to the reaction (Q) equals the change in enthalpy (ΔH).
- Proof
- #1 ΔU = Q − W
- First Law of Thermodynamics
- Change in internal energy = heat added − work done
- #2 W = PΔV
- Assumption that the only work done is a change of volume at constant pressure
- #3 Therefore, ΔU = Q − PΔV
- From #1 and #2
- #4 Q = ΔU + PΔV
- Rearranging #3
- #5 ΔH = ΔU + PΔV
- From the definition of enthalpy (H = U + PV) for a system at constant pressure
- Therefore, Q = ΔH
- Frome #4 and #5
- That is, the heat transferred to the system = change in enthalpy
- #1 ΔU = Q − W
Gibbs Free Energy
- In the 1870s Josiah Willard Gibbs developed the ingenious notion of Gibbs free energy, combing the internal energy of the First Law of Thermodynamics with the entropy of the Second.
- The change in Gibbs free energy of a chemical reaction equals the change in enthalpy minus the temperature of the reaction times the change in entropy.
- ΔG = ΔH -T ΔS
- A chemical process at constant temperature and pressure is
- exergonic and spontaneous if ΔG < 0
- e.g. combustion of methane
- endergonic and non-spontaneous if ΔG > 0
- e.g. photosynthesis
- in chemical equilibrium if ΔG = 0
- e.g. carbonation
- exergonic and spontaneous if ΔG < 0
- Definitions
- A spontaneous process is one that, once started, requires no external energy.
- A non-spontaneous process, by contrast, requires constant external energy to keep it going.
- A chemical reaction is exoergic (exergonic) if it releases energy.
- Exoergic = ex-oh-ERR-jick
- Exergonic = ex-er-GONE-ick (used in biology)
- A chemical reaction is endoergic (endergonic) if it absorbs energy.
- Endoergic = end-oh-ERR-jick
- Exergonic = ex-er-GONE-ick (used in biology)
- A chemical reaction is in chemical equilibrium if the forward and reverse reactions proceed at the same rate, so that the concentrations of reactants and products remain the same.
- The change ΔG in the Gibbs free energy of a chemical reaction equals the Gibbs free energy of formation of the products minus the Gibbs free energy of formation of the reactants
- ΔG = Gf of products− Gf of reactants
- The Gibbs free energy of formation of a compound is the change of Gibbs free energy in the formation of the compound from its elements.