CHEMISTRY COMMUNITY

Can someone explain the difference between ∆G and ΔG° conceptually. Why don't we always need to calculate for ΔG°?

It's just a difference of values, the concept of are still the same. I think that it just depends of the conditions of the reaction. I'm not entirely sure though, that's a good question.

ΔG° is in standard condition, so you would have to calculate ΔH° and ΔS°, just like you would in the regular Gibbs Free Energy equation.

∆G° is considered the standard free energy change of a reaction, while ∆G is considered the free energy change of a reaction. Remember, standard free energy change implies that the reactants and products are in their standard states at 1 atm and generally 25 celsius.

∆G° is the standard free energy of the reaction. This means that the reactants and products are at their standard state: 1 atm for gases and 1 M of aqueous solutions, typically at 25°C. ∆G is just the free energy of the reaction that doesn't have to be under the standard state conditions.

is the standard free energy change in a reaction at 1 bar for gases or 1M for solutions. Because 1 bar is roughly equivalent to 1 atm, we can use the values of the standard free energy for most reactions