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A chemical equilibrium concerns chemical reactions. There should be at least a forward- and backward reaction between two species but more complex systems with multiple individual reactions may occur. The important observation is that there is no macroscopic change to the chemical constituents of the system, i.e. the concentrations of all reaction partners ...


12

Chemical equilibrium is a type of dynamic equilibrium, but not every dynamic equilibrium is a chemical equilibrium. In a chemical equilibrium there is no change on the macroscopic scale. That means that if you look at the system it seems like nothing is happening, but at molecular scale there are reactions going on and the rate of forward reaction = rate of ...


6

Current definition implies that $\mathrm{pH}$ is a function of relative activity. Originally, the amount concentration of $\ce{H+}$ in $\pu{mol L-1}$ was proposed, which is also often used these days as an approximation [1]. $\mathrm{pH}$ was originally defined by Sørensen in 1909 … in terms of the concentration of hydrogen ions (in modern nomenclature) ...


4

You got the solubility part reversed. The solubility of $\ce{AgCl}$ is lower than the solubility of $\ce{Ag2CrO4}:$ $$s(\ce{AgCl}) = \sqrt{K_\mathrm{sp}(\ce{AgCl})} = \sqrt{\pu{1.8E-10 mol2 L-2}} = \pu{1.34E-5 mol L-1}$$ $$s(\ce{Ag2CrO4}) = \sqrt[3]{\frac{K_\mathrm{sp}(\ce{Ag2CrO4})}{4}} = \sqrt[3]{\frac{\pu{1.1E-12 mol3 L-3}}{4}} = \pu{6.50E-5 mol L-1}$$ ...


3

Why ? Because nature of matter at molecular and atomic level is dynamic, not static. In classical mechanics, objects can be in long term mutual rest. Not in quantum mechanics and quantum chemistry. Molecules have zero knowledge about the system being in equilibrium or not. If the process is supported by the thermodynamic and kinetic aspects, nothing is ...


3

The reaction you wrote down is wrong on two counts. The reactant is not a hypothetical tetraaqua complex and the product is not a hypothetical tetraammin complex. The correct reaction is as shown below: $$\ce{[Cu(H2O)6]^2+ (aq) + 4 NH3 (aq) <=> [Cu(NH3)4(H2O)2]^2+ (aq) + 4 H2O (l)}\tag{1}$$ Note that I have used an equilibrium arrow here: the ...


3

Your textbook's derivation is done under the assumption of constant $T$, which means $T_{sys} = T_{surr} =T$. However, this does not mean $dG_{sys}$ is always zero. Let's start with the following: $$dS_{univ}=dS_{sys}+dS_{surr}= \frac{\text{đ}q_{rev, sys}}{T_{sys}}+\frac{\text{đ}q_{rev, surr}}{T_{surr}}$$ Since heat flow always affects the surroundings ...


3

Proper definitions of chemical equilibrium will not involve reaction rates whatsoever. Thermodynamics does not care about time. Chemical potential is the work required to form a molecule in solution, irregardless of the time it takes. Statistical Mechanics says chemical potential is the work required to move a molecule from infinitely far away (ideal gas ...


2

I expect it may be due to the basic definition of chemical equilibrium simply being inadequate This is the answer (sort of). In essence, when you are below the solubility limit, the chemical potential of the solid lies above that of the solubilized salt, and there is no equilibrium with the solid (because no solid can form). This is illustrated in the ...


2

You have to look at two things in terms of equillibrium. The pot with the solution, once all AgCl is dissolved and you have stirred it a bit more, is in equillibrium. Obviously. There is no chemical potential gradient, and you have only one phase. By itself, nothing will ever happen again in it. The reaction (dissolution of AgCl in water) isn`t: If you add ...


2

For a microscopic step at constant $T$ and $p$ $$\mathrm dG=0\tag{constant $T$ and $p$}$$ implies: reversibility (equilibrium) $\mathrm dS_\mathrm{univ} = 0$ $\mathrm dH_\mathrm{sys} = T\,\mathrm dS_\mathrm{sys}$ since $\mathrm dG = \mathrm dH_\mathrm{sys} - T\,\mathrm dS_\mathrm{sys} \tag{constant $T$ and $p$}$ The derivation you suggest seems strange. ...


2

The value $\pu{1.3653 mol L^{−1} atm^{−1}}$ is the solubility constant (or Henry's law solubility constant), not the solubility. The solubility is defined as the maximum possible concentration (the saturation concentration) of a solute under given solution conditions (e.g. temperature and pressure), whereas the solubility constant $H^{cp}$ defines how solute ...


2

Both chloride and bromide ions are present in 10-fold excess over silver ions. That means that the chloride and bromide concentration in solution will not drop by much (they will remain major species). The solubility product of AgCl is 200-times higher than that of AgBr. If both AgCl and AgBr precipitate, the chloride solution would be 200-times higher than ...


2

van 't Hoff equation $$\frac{\mathrm d}{\mathrm dT} \ln K_\mathrm{eq} = \frac{\Delta H^⦵}{RT^2}$$ = Le Chatelier's principle in a particular context of temperature and chemical reaction equilibrium. What the Le Chatelier's principle says qualitatively as a general principle, the van 't Hoff equation says quantitatively in context of temperature ...


1

pH is not defined versus normality or molarity of an acid: Whatever the normality or the molarity, the pH is defined from the activity (or the concentration) of the ions H+. A given value of the normality or of the molarity does not give you the activity (or the concentration) of H+. So it does not allow you to calculate the pH.


1

The problem could be solved with simultaneous equations, but the following are the wrong equations. $$K_{\mathrm{sp,}\ \ce{AgCl}} = [\ce{Ag+}][\ce{Cl-}]$$ $$K_{\mathrm{sp,}\ \ce{AgBr}} = [\ce{Ag+}][\ce{Br-}]$$ The right equations to use would be: $$K_{\mathrm{sp,}\ \ce{AgCl}} \ge [\ce{Ag+}][\ce{Cl-}]$$ $$K_{\mathrm{sp,}\ \ce{AgBr}} \ge [\ce{Ag+}][\ce{...


1

Assume surface molecules are restrained by a QM harmonic potential, with energy levels described as $$E_n = \hbar \omega(n + \frac12)$$ where $$\omega=\sqrt{\frac{k}{\mu}}$$ is the frequency of the harmonic oscillator and $$\mu=\frac{m_1m_2}{m_1 + m_2}$$ is its reduced mass. The actual shape of the potential is determined by $k$, the force constant, ie ...


1

Shouldn't it clearly be the opposite, as increasing temperature favors the reverse reaction? Yes, they made a mistake. For all the other scenarios, they paired a figure of the rate changes with a figure of the matching concentration changes. For this scenario (increase in temperature) they matched it with a correct figure of concentration changes when the ...


1

My answer starts with the premise that a static equilibrium exists (although that is debatable). How could we be sure that one certain reaction in equilibrium is dynamic? If your reaction is at dynamic equilibrium, changing the temperature or the concentration of one of the reactant or product species (by adding some solute, for example) should disturb ...


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