# Chemistry - Can aqua regia work with sources of chloride and nitrate other than hydrochloric acid and nitric acid?

## Solution 1:

Not even zinc would react with neutral nitrate + chloride, why should gold ?

Aqua regia must be strongly acidic for nitrates to have oxidative properties for oxidation of chlorides to chlorine and to oxidize gold to trace amounts of $$\ce{Au^3+}$$(what redox potential allows)

Mixing of acids leads to these reactions:

$$\ce{HNO3 + 3 HCl -> NOCl + Cl2 + 2 H2O}$$ $$\ce{2 NOCl -> 2 NO + Cl2}$$ $$\ce{2 NO + O2 -> NO2 }$$

That also means aqua regia dissolution ability decreases with time.

The major oxidative reagent is free chlorine. Hydrochloric acid saturated by chlorine dissolves gold as well. Chlorides additionally help gold dissolution by reaction $$\ce{Au^3+ + 4 Cl- <=>> AuCl4^-}$$, this decreases the redox potential $$\ce{Au/Au^3+}$$, making gold dissolution by chlorine and even by nitric acid easier.

$$\ce{2 Au + 3 Cl2 + 2 Cl- -> 2 AuCl4-}$$ $$\ce{Au + HNO3 + 4 HCl -> HAuCl4 + NO + 2 H2O}$$ $$\ce{Au + 3 HNO3 + 4 HCl -> HAuCl4 + 3 NO2 + 3 H2O}$$

## Solution 2:

Gold dissolves in aqua regia because there is a redox reaction: The nitrate ion oxidizes the gold in the presence of $$\ce{Cl-}$$ ion into $$\ce{Au^3+}$$, which combine with the chloride ions to form $$\ce{Au(Cl)4-}$$ ion. In this reaction, the presence of $$\ce{Cl-}$$ ion is important because it reduces the $$\ce{Au/Au^3+}$$ couple's potential as follows:

$$\ce{Au <=> Au^3+ + 3e-} \qquad E^\circ = \pu{-1.498 V} \tag{1}$$

However, in the presence of $$\ce{Cl-}$$ ion, this half-reaction becomes:

$$\ce{Au + 4Cl- <=> AuCl4^- + 3e-} \qquad E^\circ = \pu{-1.002 V} \tag{2}$$

The nitrate reduction half-reaction is:

$$\ce{NO3- + 4 H+ + 3e- <=> NO + 2H2O } \qquad E^\circ = \pu{0.957 V} \tag{3}$$

Therefore, the complete redox reaction is:

$$\ce{Au + 4Cl- + NO3- + 4 H+ <=> AuCl4^- + NO + 2H2O} \qquad E_\mathrm{rxn}^\circ = \pu{-0.045 V} \tag{4}$$

Although, $$E_\mathrm{rxn}^\circ$$ for this reaction is $$\pu{-0.045 V}$$ under standard conditions ($$\pu{1 M}$$ and $$\pu{1 atm}$$), at working conditions (e.g., concentrated $$\ce{HCl}$$ is at least $$\pu{12 M}$$), the corresponding $$E_\mathrm{rxn}$$ becomes positive and forward reaction proceeds (Recall Nernst equation).

According to the above half-reactions, having strong acidic conditions is essential for this reaction. Therefore, OP's suggestion of having just nitrates and chlorides (e.g., $$\ce{NaNO3}$$ and $$\ce{NaCl}$$) will not work as it in aqua regia. For example, reduction half reaction for $$\ce{NO3-}$$ is as follows:

$$\ce{NO3- + H2O + 2e- <=> NO2- + 2OH- } \qquad E^\circ = \pu{0.01 V} \tag{5}$$

Sum of $$(2) \times 2$$ and $$(5) \times 3$$ gives:

$$\ce{2Au + 8Cl- + 3NO3- + 3 H2O <=> 2AuCl4^- + 3NO2- + 6OH-} \qquad E_\mathrm{rxn}^\circ = \pu{-1.001 V} \tag{6}$$

This $$E_\mathrm{rxn}^\circ$$ value is too much to overcome.