EjerciciosFyQ

Calculation of equilibrium constants and degree of dissociation of N2O4 (8413)

F_y_Q — 2025-03-14T05:49:47Z

At and 1 atm, dissociates into by according to the following equilibrium:

2NO2(g)}" title="\ce{N2O4(g) <=> 2NO2(g)}" />

Calculate:

a) The values of the equilibrium constants and at this temperature.

b) The percentage of dissociation at and a total pressure of 0.1 atm.

Data: .

The first step is to determine the number of moles of each substance at equilibrium, based on the initial moles and the degree of dissociation ():

\underset{2n_0\alpha}{\ce{2NO2(g)}}}" title="\ce{\underset{n_0(1-\alpha)}{\ce{N2O4(g)}}} \ce{<=> \underset{2n_0\alpha}{\ce{2NO2(g)}}}" />

Since the degree of dissociation is , the moles can be written as:

\underset{0.4n_0}{\ce{2NO2(g)}}}" title="\ce{\underset{0.8n_0}{\ce{N2O4(g)}}} \ce{<=> \underset{0.4n_0}{\ce{2NO2(g)}}}" />

The total number of moles at equilibrium is the sum of all species:

Next, calculate the mole fraction for each substance:

a) The equilibrium constant in terms of partial pressures is:

Substitute the values into this equation to determine the constant:

The equilibrium constant in terms of concentrations is calculated as:

The change in the number of moles of gas is one, so substitute the values to find the constant:

b) When the total pressure of the system changes to 0.1 amt, the equilibrium shifts to favor greater dissociation of . The same expression for holds, and it is used to calculate the new degree of dissociation under these conditions. Be cautious to express the mole fractions in terms of the initial moles:

For simplicity, work with the denominator and substitute to make solving easier:

Application of limiting reagent, purity of reagents and reaction yield (8281)

F_y_Q — 2024-08-09T03:25:18Z

You have 87 g of silver nitrate, with purity, reacting with 50 mL of a hydrochloric acid solution, by mass concentration and density 1.07 g/mL, producing silver chloride and nitric acid, with a reaction yield of .

a) Write the chemical reaction and balance it if necessary.

b) Calculate the amount of silver chloride and nitric acid produced in the reaction.

c) Determine the amount of the excess reagent that does not react.

Atomic masses: H = 1; O = 16; N = 14; Cl = 35.5; Ag = 108.

a) First, it is necessary to write the chemical reaction and check its stoichiometry:

AgCl + HNO3}}}}" title="\fbox{\color[RGB]{192,0,0}{\textbf{\ce{AgNO3 + HCl -> AgCl + HNO3}}}}" />

The stoichiometry of the reaction is 1:1, which will greatly simplify the calculations.

b) You need to determine the limiting reagent in the reaction, and for this, you must know the moles of each reagent you have at the beginning. Be very careful with the data treatment because you must consider that the reagents are not pure.

Moles of pure silver nitrate:

Moles of pure hydrochloric acid:

It is easy to conclude that the limiting reagent is because it is the one with the fewest initial moles, and the stoichiometry is 1:1.

Since the reaction does not go to completion, you will calculate how many moles react from the initial 0.45 moles.

This means that only 0.40 moles of each reactant react, and the same moles of each product are obtained. You calculate the mass of each product:

c) The excess reagent is HCl, of which 0.40 moles react, leaving unreacted: (0.542 - 0.40) mol = 0.142 mol. You convert this to mass:

Mass of reactants in a chemical equation (3590)

F_y_Q — 2016-05-29T07:18:18Z

How many grams of are needed to react with 120 g of to obtain ?

The chemical equation is:

2NO2}}}" title="\color[RGB]{2,112,20}{\textbf{\ce{2NO + O2 -> 2NO2}}}" />

The molecular mass of is:

Oxygen has a molecular mass of:

According to the chemical equation 60 g of will require 32 g of to complete the reaction. Let's perform a mathematical proportion to determine the required mass of oxygen: