A spontaneous reaction in the forward direction does not require any external energy. Standard conditions means that temperature and pressure are constant. When ΔG is negative in the Gibbs Free Energy equation ΔG = H - TΔS, the reaction is said to be spontaneous. A positive ΔG means the reaction is nonspontaneous in the forward direction. When ΔG = 0, the reaction is at equilibrium.
Looking at K, the equilibrium constant, we should look at the equilibrium expression
Keq = [products] / [reactants]. If the reaction is proceeding in the forward direction, the concentration of reactants decreases because reactants are being consumed, while the concentration of products is increasing because products are being formed. This means that in the equilibrium expression, the denominator becomes smaller while the numerator becomes larger, making Keq increase and therefore become greater than 1.
Therefore, the answers are ΔG is negative and K > 1. I hope this helps!
Just FYI: If K = 0, no products are formed and therefore there can't be an equilibrium. If K < 1, this means the equilibrium favors the reactants over the products..