Calculate the standard cell emf for galvanic

cells with the cell reaction
2 H2(g) + O2(g) → 2 H2O(ℓ)
in acidic solution.
Answer in units of V.

What is the standard reaction free energy?
Faraday’s constant is 96485 C/mol.
Answer in units of kJ/mol.

I got 1.23 V for the first one

I did not check you 1.25 v but if you post your work I'll check it. For the other part use dGo = -nEoF

n = 4

I got -474706.2 kJ/mol but it's saying it's wrong

I assume you are checking into a data base. A common problem with these is that you are reporting too many significant figures. Since your V is to 3 s.f. I would keep the reporting to 3 s.f. Second, the problem asks for kJ/mol. Your -4.75E5 kJ is for the reaction and that is for 2 mols. So divide by 2.

-4.747062E5/2 = -2.37031E5 kJ/mol which I would round to -2.37 kJ/mol.

4.

To calculate the standard cell emf for the given galvanic cell reaction, we need to use the standard reduction potentials of the half-reactions involved. The standard cell emf (E°cell) can be calculated using the Nernst equation:

E°cell = E°cathode - E°anode

First, let's find the standard reduction potentials for the half-reactions involved in the cell:

1. Reduction half-reaction at the cathode:
O2(g) + 4H+(aq) + 4e- → 2H2O(ℓ)

The standard reduction potential for this half-reaction is 1.23 V. This is provided in most electrochemical tables.

2. Oxidation half-reaction at the anode:
2H2(g) → 4H+(aq) + 4e-

To get the standard reduction potential for this half-reaction, we can use the formula:

E°red(anode) = -E°red(cathode)

So, the standard reduction potential for the anode half-reaction is -1.23 V.

Now, substitute these values into the Nernst equation:

E°cell = E°cathode - E°anode
E°cell = 1.23 V - (-1.23 V)
E°cell = 1.23 V + 1.23 V
E°cell = 2.46 V

Therefore, the standard cell emf for the given galvanic cell reaction is 2.46 V.

Now, to calculate the standard reaction free energy (ΔG°), we can use the equation:

ΔG° = -nE°cell

where n is the number of moles of electrons transferred by the reaction.

In this case, for the reaction: 2H2(g) + O2(g) → 2H2O(ℓ)

n = 4 (since 4 moles of electrons are transferred in the balanced equation)

Faraday's constant (F) is 96485 C/mol, which represents the charge of one mole of electrons.

To convert C to J, multiply the charge (C) by the potential (V).

1 C = 1 J/C

Therefore, the equation becomes:

ΔG° = -4(96485 C/mol)(2.46 V) = -473.58 kJ/mol

So, the standard reaction free energy (ΔG°) for the given cell reaction is -473.58 kJ/mol.