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Chemical equilibrium represents a state in which the concentrations of reactants and products no longer change with time. For reactions that reach equilibrium, we can quantify their chemical balance using the equilibrium constant, denoted as \( K_c \).
In the reaction \( 2 \mathrm{NOBr}(g) \rightleftharpoons 2 \mathrm{NO}(g) + \mathrm{Br}_2(g) \), the equilibrium constant expression is determined from the concentrations of the products and reactants at equilibrium. The formula for \( K_c \) is:

• \( K_c = \frac{[\mathrm{NO}]^2[\mathrm{Br}_2]}{[\mathrm{NOBr}]^2} \)

The brackets "[]" indicate the molar concentrations of the reactants and products.
The value of \( K_c \) indicates the position of equilibrium. A small \( K_c \) (less than 1) suggests that the concentration of the reactants is higher than that of the products at equilibrium, which was the case here with a \( K_c \) of 0.157.
Understanding these balances is crucial for predicting how changes in conditions, like temperature and pressure, can shift equilibria according to Le Chatelier's principle.

The Ideal Gas Law is a fundamental principle in chemistry and physics that relates the pressure, volume, temperature, and number of moles of a gas. It is utilized to predict the behavior of gases under various conditions. The equation for the Ideal Gas Law is:

• \( PV = nRT \)

Where:

• \( P \) stands for pressure in atmospheres,
• \( V \) is the volume in liters,
• \( n \) represents the number of moles,
• \( R \) is the ideal gas constant (0.0821 L·atm/mol·K), and
• \( T \) is the temperature in Kelvin.

In the exercise, the Ideal Gas Law was applied to find the total pressure exerted by the gases in the container. Knowing the total moles of gases, the volume, and temperature allowed us to solve for \( P \), yielding approximately 1.021 atm.
This relation is key to understanding gas mixtures and is essential for solving many chemistry problems involving gases. Always ensure that you properly convert temperature to Kelvin and use consistent units for accurate calculations.

Calculating molar masses correctly is vital when doing any quantitative chemical analysis. The molar mass of a compound is the mass of one mole of its molecules. It's determined by summing the atomic masses of all the atoms in a formula.
For example, to find the molar mass of \( \mathrm{NOBr} \):

• Nitrogen (N) = 14.0067 g/mol
• Oxygen (O) = 15.999 g/mol
• Bromine (Br) = 79.904 g/mol

Add these values to get 109.910 g/mol for \( \mathrm{NOBr} \).
These calculations help determine moles from a given mass, which was crucial in assessing the composition of the equilibrium mixture.
To find moles, you rearrange the formula:

• \( \text{Moles} = \frac{\text{Mass}}{\text{Molar Mass}} \)

By computing correct moles for each component, you can derive important equilibrium and pressure data, making molar mass assessment a foundational skill in chemistry.