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2

This is an approach to Question 2. The reaction equilibrium condition for the system is: $$\frac{V}{RT}\frac{n_C}{n_A n_B}=K_P\tag{1}$$Let $n_{A0}$, $n_{B0}$, and $n_{C0}$ be the number of moles at equilibrium for each species when the volume is equal to $V_0$. Then we have $$\frac{V_0}{RT}\frac{n_{C0}}{n_{A0} n_{B0}}=K_P\tag{2}$$. For a small change in ...


1

Assuming $\ce{A, B,}$ and $\ce{C}$ are gases and behave like ideal gases we can conclude following by applying $P_xV=n_xRT$ where $x$ is either $\ce{A, B,}$ or $\ce{C}$, and $V$ and $T$ are constants: $$[\ce{A}] = \frac{n_A}{V}=\frac{P_A}{RT}; \quad [\ce{B}] = \frac{n_B}{V}=\frac{P_B}{RT}; \quad \text{and } \quad [\ce{C}] = \frac{n_C}{V}=\frac{P_C}{RT} $$ ...


2

The root of your confusion, as I understand, is that you are not much familiar with other kinds of work done. Although, $p\,\mathrm{d}V$ is everyone's favourite but it's not the only one in the league. Always remember, the $p\,\mathrm{d}V$ is work done during an expansion/compression. You can change the pressure or concentration to get electric work.


0

Boyle, Gay-Lussac, Charles laws can all be deduced from the more general ideal gas law, namely : $\ce{PV = nRT}$, provided you know that P is the pressure of the gas, V is the volume of the container, n is the amount of gas, in moles, R is a constant ( $\ce{R = 8.314 J K^{-1} mol^{-1}}$), and T is the absolute temperature in Kelvin. Let's show that Boyle's ...


1

There is no contradiction. Each law is valid under a certain conditions. This is because each law of these assumes the constancy of one of the three following variables (the pressure , the volume , the absolute temperature) and then studies the relation between the other two. Isn’t this a contradiction? This is how scientist design experiments. Suppose a ...


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You could mechanically rotate a generator (i.e., do work) to make an electrical current which flows into a resistor within the gaseous volume and heats it. So you could say you did work on the gaseous system (indirectly), and some of it is recoverable as work by releasing the pressure.


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Yes, I agree with your teacher's comment that upon increasing the total pressure of a system, there are select reactions "more dependent on collisions", and such reactions have been a matter of study. For example, per this 2008 work reported in 'The Journal of Physical Chemistry', The Temperature and Pressure Dependence of the Reactions H + O2 (+M)...


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