To what pH should you adjust a standard hydrogen electrode to get an electrode potential of -0.122 V? (Assume that the partial pressure of hydrogen gas remains at 1 atm.)

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Hey everyone in this example we need to calculate the ph needed for the standard hydrogen electrode to have an electoral potential of negative 0.153 volts. So we should recall that our charge of our battery is equal to our standard Self potential value subtracted from our new Ernst equation constant where we have a value of 0.0592V divided by our electrons transferred at a reduction n. Which is then multiplied by the log of our equilibrium constant. Q. So according to the prompt, our Standard cell potential or rather our charge of our battery is equal to a value of negative 0.153V. And our standard cell potential value for hydrogen, Which we would call exists as a diatonic molecule is equal to a value of 0V. And we would find this from our table in our textbooks for standard electrode potential. And then lastly we should recall that the equilibrium constant Q. Is equal to our concentration of products divided by our concentration of reactant. However we don't have any products mentions. So we would just go ahead and make our numerator one. Where in the denominator we would have our concentration of our reactant which is hydrogen. And this is going to then be placed to an expletive too Because again we recall that hydrogen exists as a diatonic molecule and so N. As well as that exponents would be equal to one another which which makes it clear that we have the transfer of two electrons. So right now we want to go ahead and plug in what we know into our formula to solve for Q. So what we should have is that for our Charge of our battery again we said we have a value of negative 0.153V. We're going to set this equal to our standard cell potential value for hydrogen, equal to zero volts subtracted from our first equation constant 0.592 volts divided by our electrons transferred which we agreed is two electrons for hydrogen. And then this is going to be multiplied by the log of Q. And so simplifying our equation out, we would simplify this so that now we have the log of Q Is equal to 5.1689. And so to get rid of the log term, we're going to take 10 and make our right hand side and exponents As well as doing so on the left hand side for the long term. So that we can cancel out the log term. And what we would get is that our equilibrium constant Q is equal to a value of 1.475 times 10 to the fifth power. So now we can actually calculate our value for Q because we recalled above that Q is equal to one over our concentration of H plus squared where now we would take one over our concentration of H plus which we calculated above as 1.475 times 10 to the fifth power. And we're going to make this under a square root to get rid of the square term. And this is going to give us a value and correction. So this to be clear, is solving for our concentration of H plus and we would get a value equal to 2.60 times 10 to the negative third power. So now that we have our concentration of H plus, we should recall that we can calculate ph by taking the negative log of our concentration of H plus. And this means that we would say our ph is equal to the negative log of our concentration of H plus, which above we stated as 2.60 times 10 to the negative third power, which will give us a ph equal to 2.584. And this would be our final answer for our ph of our hydrogen electrode. So I hope that everything I reviewed was clear. If you have any questions, please leave them down below. Otherwise, I will see everyone in the next practice video

To what pH should you adjust a standard hydrogen electrode to get an electrode potential of -0.122 V? (Assume that the partial pressure of hydrogen gas remains at 1 atm.)

Verified Solution

Key Concepts

Nernst Equation

Standard Hydrogen Electrode (SHE)

pH and Hydrogen Ion Concentration