1. The problem statement, all variables and given/known data Calcium carbonate primarily occurs as two crystalline forms, calcite and aragonite. The value of∆!° for the transition CaCO3(calcite) ⇌ CaCO3(aragonite) is +1.04 kJ·mol-1 at 25°C. At that temperature the density for calcite is 2.710 g·cm-3, and that of aragonite is 2.930 g·cm-3. At what pressure will the two crystalline phases be at equilibrium at 25°C? 2. Relevant equations ln(ai)= ((molar V)/RT) * (P-1) ΔG = -RT * ln(Kp Kp = aaragonite/acalcite 3. The attempt at a solution So I started by finding the molar volume of each by dividing the MW by the individual density and got Vara = 34.157 cm^3/mol and Vcal = 36.930 cm^3/ mol From here I used the activity equation to get that ai = e^(((molar V)/RT) * (P-1))) And since Kp = aaragonite/acalcite and ΔG = -RT * ln(Kp I can write that ΔG = -RT * ln(aaragonite/acalcite) This leads to ΔG = -RT * (((molar Vara/RT) * (P-1) - ((molar Vcal/RT)*(P-1)) Or ΔG = - molar Vara*(P-1)+(molar Vcal*(P-1)) Continuing 1040 J/mol =(-34.157 cm^3/mol ) * P + 34.157 cm^3/mol +(36.930 cm^3/mol) * P - 39.630 cm^3/mol Solving for P I get 376 J/cm^3. I am given that the answer is 3850.9 but no units and I wasn't able to convert my units to get that number... I'm pretty sure my math is correct but I think I messed up on units somewhere. Any help would be greatly appreciated!! Thanks!