I'm stuck...that is i believe I've come to an impossible answer:

I need to calculate the equilibrium constant

4Fe(s) + 3O2(g) --> 2Fe2O3(s)

First off: Delta G G Fe(s) = 0, G O2(g) = 0, G Fe2O3(s) = -741 (x2)
-741 (2) - 0 = -1480 = delta G of reaction
at equilibrium:

0 = (delta G rxn) + RT (lnk) or lnk = -(delta G rxn) / RT

so, lnk = 1480 / 2.478 = 598 = lnk

k = e^598 = Impossible? This would be an humongous number so I think I went astray somewhere.

Any help?

To calculate the equilibrium constant (K) for the given reaction, you've correctly started by calculating the change in Gibbs free energy (ΔG) of the reaction. However, there seems to be a mistake in your calculations.

The correct equation for calculating ΔG at equilibrium is:
ΔG = -RT ln K

Where:
- ΔG is the change in Gibbs free energy
- R is the gas constant (8.314 J/(mol·K))
- T is the temperature in Kelvin
- K is the equilibrium constant

Given that ΔG of the reaction is -1480 J, you need to rearrange the formula to solve for K. The correct equation is:
K = e^(-ΔG / (RT))

Now let's calculate K:
Given:
ΔG = -1480 J
R = 8.314 J/(mol·K)
T = (assumed temperature in Kelvin)

Step 1: Convert ΔG to its proper units
ΔG = -1480 J / (1000 J/kJ) ≈ -1.48 kJ/mol

Step 2: Plug the values into the equation
K = e^(-(-1.48 kJ/mol) / (8.314 J/(mol·K) × T))

Note: To correctly calculate K, you need to specify the temperature (T). Please provide the temperature in Kelvin so we can proceed with the calculation.