Lecture11 - Chemistry 132 Lecture 11 Equilibrium Continued...

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Chemistry 132 – Lecture 11 – Equilibrium Continued Thermodynamic Equilibrium Constant We have discussed the equilibrium constants K c and K p . However, we often run into the “thermodynamic equilibrium constant”, K . K is a unit less constant found by using ratios of partial pressures to a “reference pressure” of 1.0 atm or by using ratios of concentrations to a “reference concentration” of 1.0 mol/L. e.g. Consider the reaction 3 NO(g) N 2 O(g) + NO 2 (g) Kp = [P N2O(g) ][P NO2(g) ]/[P NO(g) ] 3 which has units atm -1 K = [P N2O(g) / P ref ][P NO2(g) / P ref ]/[P NO(g) / P ref ] 3 which has no units Note : if you use 1.0 atm as the reference pressure and 1.0 mol/L as the reference concentration, the numerical values of K and K p will be identical and the numerical values of K compared to K c will be identical. 1
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Chemistry 132 – Lecture 11 – Equilibrium Continued The thermodynamic equilibrium constant ( K ) has some benefits over K c and K p . It can be calculated from G o (that is, from H o and S o ) so the extent of any equilibrium reaction can be calculated using calorimetric data alone . K is always dimensionless. Heterogeneous Equilibrium When all reactants and products in an equilibrium situation are gases, the system is an example of homogeneous equilibrium .
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