Partial Pressures of ethanal and water

[nydream12]

Partial Pressures of ethanal and water!

Hello, I need help with a physical chemistry fundamentals problems:

1.CH3CH2OH (s)+O2-->CH3CHO+H2O (l)

a) the equilibrium constant for the reaction...K=e^(-deltaG/RT)...is this is the right equation for solving K?
b) what will be the partial pressures of ethanal and water in equilibrium with 1.00 bar each of ethanol and oxygen??...Pj=xjP where p is 100kPa; I can't seem to figure out a way to begin solving (b)

2. A buffer solution of Tris (pKa=8.3 @ 20C) was prepared by combining 100.oo mL of 0.500 M Tris with 150.00 mL of 0.500 M Tris+ in a 500.00 mL volumetric flask and diluting to the line with water

a) pH of this buffer? ...Would it be pH 8.3 (same as pH)
b) What will be the pH after addition of 5.5 mmol of HCI to 100.00 mL of this buffer...would it be pH=pKa-log (acid/base)

Please help someone, anyone...THANKS!

 

[nate808]

Hello there,

I would be happy to assist you with your physical chemistry problems.

1. For the first problem, you are correct in using the equilibrium constant equation K=e^(-deltaG/RT). However, in order to solve for K, you will need to know the values of deltaG, R, and T. DeltaG is the change in Gibbs free energy, which can be calculated using the standard free energy of formation for each species involved in the reaction. R is the gas constant (8.314 J/mol*K), and T is the temperature in Kelvin. Once you have all these values, you can plug them into the equation to solve for K. As for part b, you can use the equation Pj=xjP to calculate the partial pressures of ethanal and water. Xj represents the mole fraction of each species, which can be calculated using the ideal gas law (PV=nRT).

2. For the second problem, the pH of the buffer solution can be calculated using the Henderson-Hasselbalch equation, which is pH=pKa+log([base]/[acid]). In this case, the base is Tris and the acid is Tris+. Once you have the pH of the buffer solution, you can use the same equation to calculate the pH after the addition of 5.5 mmol of HCl. Just remember to adjust the concentrations accordingly.

I hope this helps! Let me know if you have any further questions. Good luck with your studies!

 

[Blueshift5]

1. a) Yes, the equation K=e^(-deltaG/RT) is the correct equation for solving the equilibrium constant (K) for this reaction. DeltaG is the change in Gibbs free energy, R is the gas constant, and T is the temperature in Kelvin.

b) To solve for the partial pressures of ethanal and water in equilibrium, you will need to use the ideal gas law. The equation for the ideal gas law is PV=nRT, where P is pressure, V is volume, n is the number of moles, R is the gas constant, and T is the temperature in Kelvin. In this case, you know the pressure (1.00 bar), volume (unknown), number of moles (unknown), and temperature (unknown). You can set up two equations using the ideal gas law, one for ethanal and one for water. Since the pressure and temperature are the same for both gases, you can set the equations equal to each other and solve for the unknown variables. This will give you the partial pressures of ethanal and water in equilibrium.

2. a) The pH of the buffer solution will be equal to the pKa of Tris, which is 8.3. This is because the buffer solution is made up of equal amounts of Tris and Tris+, so the concentration of the acid and its conjugate base will be equal, resulting in a pH equal to the pKa.

b) To calculate the pH after adding 5.5 mmol of HCl, you will need to use the Henderson-Hasselbalch equation. This equation is pH=pKa+log([A-]/[HA]), where pKa is the acid dissociation constant, [A-] is the concentration of the conjugate base, and [HA] is the concentration of the acid. In this case, the acid is HCl and the conjugate base is Cl-. You will need to calculate the new concentrations of HCl and Cl- in the buffer solution after adding 5.5 mmol of HCl. Once you have these values, you can plug them into the Henderson-Hasselbalch equation to calculate the new pH of the buffer solution.