two flasks of equal vol. connected by a narrow tube of negligible volume are at 27c and contain Cl2 at 1 atm pressure & Cl2O at 1.5 atm pressure.d 2 gases were allowed to mix & heated to react - Cl2O- Cl2 + 1/2 O2. after d reatn 1 of d vessels was kept in a water bath maintained at 27c while other was kept in another water bath at 52c. calculate d pressure in both d vessels at this temp.

To solve this problem, we need to analyze the reaction between chlorine gas (Cl2) and dichlorine monoxide (Cl2O) and how the pressures change when the gases are allowed to mix and react at different temperatures. The reaction is given as:

Cl2O → Cl2 + 1/2 O2

Initially, we have two flasks with chlorine gas at 1 atm and dichlorine monoxide at 1.5 atm. When these gases mix and react, we need to consider the stoichiometry of the reaction and how temperature affects the pressure of the gases in each flask.

Initial Conditions

Let's summarize the initial conditions:

• Flask 1: Cl2 at 1 atm
• Flask 2: Cl2O at 1.5 atm

Reaction Stoichiometry

From the balanced equation, we see that 1 mole of Cl2O produces 1 mole of Cl2 and 0.5 moles of O2. This means that for every mole of Cl2O that reacts, we will produce an equal amount of Cl2 and half that amount of O2.

Calculating Moles of Gases

To find the total pressure after the reaction, we first need to determine how much Cl2O will react. Assuming complete reaction of Cl2O:

• Let’s denote the initial moles of Cl2O as n.
• After the reaction, we will have:
• Cl2 remaining = Initial moles of Cl2 + Moles produced from Cl2O = 1 + n
• O2 produced = 0.5n

Using Ideal Gas Law

Using the ideal gas law, we can relate pressure, volume, and temperature. The pressure of a gas is directly proportional to the number of moles and temperature:

P = nRT/V

Since the volume is constant and we are dealing with equal volumes of gas, we can simplify our calculations by focusing on the ratio of moles and the effect of temperature.

Temperature Effects

Now, we need to consider the two different temperatures:

• Flask 1 at 27°C (300 K)
• Flask 2 at 52°C (325 K)

Using the ideal gas law, we can calculate the pressure in each flask after the reaction has occurred. The total number of moles after the reaction will be:

Total moles = (1 + n) + 0.5n = 1 + 1.5n

Calculating Pressure in Each Flask

For Flask 1 at 27°C:

P1 = (1 + 1.5n)RT1/V

For Flask 2 at 52°C:

P2 = (1 + 1.5n)RT2/V

Now, substituting the values of R (0.0821 L·atm/(K·mol)) and the respective temperatures:

For Flask 1:

P1 = (1 + 1.5n)(0.0821)(300)/V

For Flask 2:

P2 = (1 + 1.5n)(0.0821)(325)/V

Final Pressure Calculation

To find the final pressures, we need to know the initial moles of Cl2O (n). If we assume that the initial moles of Cl2O were such that it completely reacted, we can set n to a specific value based on the initial conditions. However, without a specific value for n, we can express the pressures in terms of n.

Ultimately, the pressures in the two flasks will depend on the initial moles of Cl2O and the volume of the flasks. If we assume equal volumes and complete reaction, we can calculate the pressures numerically once we have a specific value for n.

In summary, the pressures in the two flasks after the reaction and at the specified temperatures can be calculated using the ideal gas law, taking into account the stoichiometry of the reaction and the initial conditions of the gases involved.