Calculating pH from deltaG° and Ion Concentrations

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In summary, the conversation discusses finding the pH required for a reaction to be at equilibrium at 298 K, given the concentrations of species and the value of deltaG°. The correct answer is 9.21, which can be obtained by using K in terms of concentrations and solving for [OH-]. The initial K value provided was correct, but one of the concentrations was incorrect, leading to an incorrect final answer. The method provided by the expert was correct and helped the individual find their mistake.
  • #1
RNix25
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I've been working on this one for a few days, and I just can't seem to get the right answer.

Consider the reaction I2(s) --> I- (aq) + IO3- (aq)
Balanced= 3I2 + 6OH- --> 5I- + IO3- + 3H2O

deltaG°= -153.8 kJ/mol

The question: What pH is required for the reaction to be at equilibrium at 298 K when [I-]= 0.5 M and [IO3-]= 0.5 M

I know the answer is 9.21, but I just can't figure out how to get there. I started by finding Keq, which is 9.113 * 10^26, but I'm stumped from there. I would be extremely appreciative for a walkthrough to the final answer.
 
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  • #2
Sounds like this should go in the homework help section.

Your K value looks correct. Write out K in terms of the concentrations of species raised to the power of their stoichiometric coefficients and solve for [OH-]. pH = 14 + log[OH-], I get 9.206 using your values.
 
  • #3
I slept on it and when I woke up, I looked over it and figured out where I was going wrong. Had one of my concentrations wrong, so I kept getting the wrong value when I worked it. I already had the answer, I just couldn't figure out why my math wasn't producing it. Anyway, your method was exactly right, so thanks for the help.

Don't know protocol on these boards, but if it is proper, a mod can close/delete this thread.
 

Related to Calculating pH from deltaG° and Ion Concentrations

1. How do I calculate pH from deltaG° and ion concentrations?

To calculate pH from deltaG° and ion concentrations, you will need to use the Nernst equation. This equation relates the standard Gibbs free energy change (deltaG°) to the equilibrium constant (K) and the concentrations of ions involved in the reaction. Once you have calculated the equilibrium constant, you can use the pH formula (pH = -log[H+]) to determine the pH.

2. What is the Nernst equation and how is it used in calculating pH?

The Nernst equation is a mathematical formula that relates the standard Gibbs free energy change (deltaG°) to the equilibrium constant (K) and the concentrations of ions involved in the reaction. It is used to determine the pH of a solution by taking into account the concentration of hydrogen ions (H+) and the equilibrium constant of the reaction.

3. What is the relationship between deltaG° and pH?

DeltaG° and pH are related through the Nernst equation. This equation takes into account the concentration of hydrogen ions (H+) and the equilibrium constant of the reaction to calculate the standard Gibbs free energy change (deltaG°). From there, the pH can be determined using the pH formula (pH = -log[H+]).

4. Can pH be calculated using only deltaG° and ion concentrations?

Yes, pH can be calculated using only deltaG° and ion concentrations, as long as the reaction is at equilibrium. The Nernst equation takes into account the equilibrium constant of the reaction, which is dependent on the ion concentrations. However, it is important to note that this calculation assumes ideal conditions and may not accurately reflect the actual pH of a solution.

5. How does temperature affect the calculation of pH from deltaG° and ion concentrations?

Temperature can affect the calculation of pH from deltaG° and ion concentrations because it can change the equilibrium constant (K) of the reaction. This is because temperature affects the rate of the reaction, which in turn affects the concentrations of ions involved. Therefore, when calculating pH using the Nernst equation, it is important to take into account the temperature at which the reaction is occurring.

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