C2H5NO2 + 2H2O → 2CO2 + NH3 + 3H2

You begin with 1.25 × 10-2 moles of glycine and an excess amount of water that will allow the reaction to proceed completely to products. This is all in a vessel with a volume of 850 mls at 27°C. Assume that the vessel begins with air at atmospheric pressure. Since glycine is a solid and water is a liquid, neglect any volume that they occupy as reactants. For all questions, assume that temperature remains constant unless told otherwise.

a. How many moles of gas are present at the start of the process (before any glycine has been consumed)?

b. What is the number of moles of gas produced if all of the glycine is consumed?

c. What is the final pressure of the system assuming that the volume remains constant?

d. If the vessel were allowed to expand keeping the pressure constant, what would the resulting volume be?

e. If the vessel can only expand to a volume of 1,250 ml, what is the resulting pressure?

C2H5NO2 + 2H2O → 2CO2 + NH3 + 3H2

a. That will be zero unless you want to count the vapor pressure of water at 27C. I assume that is to be neglected.

b. You have 0.0125 mols glycine. You get 0.0125 mol NH3, 3*0.0125 mols H2 and 2*0.0125 mols CO2.

c. PV = nRT

d. PV = nRT

To solve this problem, we need to use stoichiometry and the ideal gas law. Here's how we can approach each question:

a. To determine the moles of gas present at the start of the process, we need to consider the gases present in the reaction equation. According to the balanced equation, there are 2 moles of CO2, 1 mole of NH3, and 3 moles of H2 produced per mole of C2H5NO2 consumed. Since we have an excess of water, we can neglect its contribution to the gas moles.

Therefore, the total moles of gas present at the start of the process is:

2 moles CO2 + 1 mole NH3 + 3 moles H2 = 6 moles of gas

b. If all the glycine is consumed, the reaction will produce the same number of moles of gases mentioned in part a. Therefore, the number of moles of gas produced would still be 6 moles.

c. To determine the final pressure of the system assuming the volume remains constant, we need to use the ideal gas law:

PV = nRT

Since the volume remains constant, we can rewrite the equation as:

P = nRT/V

The number of moles of gas produced is 6 moles, the temperature is 27°C (which needs to be converted to Kelvin), the gas constant R is 0.0821 L·atm/mol·K, and the volume is 850 mL (which needs to be converted to liters).

Plug in these values into the equation to find the final pressure:

P = (6 moles) * (0.0821 L·atm/mol·K) * (27°C + 273.15) / (0.850 L)

Solve for P to find the final pressure.

d. If the vessel were allowed to expand while keeping the pressure constant, we can use the Boyle's Law, which states that the product of the initial pressure and volume is equal to the product of the final pressure and volume.

P1V1 = P2V2

We can rearrange this equation to solve for the final volume:

V2 = (P1 * V1) / P2

Since we are given the initial pressure, P1, and the initial volume, V1, we need to use the final pressure determined in part c, P2, to calculate the final volume.

e. If the vessel can only expand to a volume of 1,250 mL (which needs to be converted to liters), we can use the Boyle's Law again to determine the resulting pressure, which is unknown in this case.

P1V1 = P2V2

Rearrange the equation to solve for the final pressure:

P2 = (P1 * V1) / V2

We know the initial pressure, P1, the initial volume, V1, and the final volume, V2, so we can substitute these values into the equation to find the resulting pressure.

By following these steps, you should be able to find the answers to all the questions using the given information and the appropriate equations.