# Why do these reactions not occur?

The following reactions do not occur: $$\ce{HCl(aq)}\ +\ \ce{Pb(NO3)4(aq)} \to$$ $$\ce{HCl(aq)}\ +\ \ce{Cu(NO3)2(aq)} \to$$

Is this because in both cases both reactants are entirely dissolved as their components ion? $\ce{NO3}$ is a conjugate base of a strong acid so it is a very weak base so it doesn't participate in reactions, right?

• Check the subscripts on the $\ce{Pb}$ compound, $\ce{Pb^{2+}}$ seems more likely. It depends on concentrations, but $\ce{PbCl2}$ should precipitate. – TAR86 Feb 7 '18 at 19:07
• @TAR86 I checked and it is 4 – Lorenzo B. Feb 7 '18 at 20:24

When put in water a mixture such as $\ce{HCl}$ and $\ce{Pb(NO3)4}$ goes to complete ionic dissociation. Lead nitrate is soluble in water and hydrochloric acid is a very strong acid. Nitric acid is also strong, which means those dissolved nitrate ions are not reversing back to the molecular form.
Their ions will be fairly stable in solution and will only combine in case a more stable compound is formed. You could come up with reactions in which some compounds will go out of solution (like $\ce{Cl2}$), such as $\ce{ 2HCl + HNO3 \to HNO2 + Cl2 + H2O}$. Lead could also act as an oxidizing agent. But I would argue that if you consider the change in free energy or the net redox potential of those reactions, you'd find that they are quite unfavourable. If you want to calculate it itself, you could start from a reduction potential table like this one.
• $\ce{Pb}(\rm IV)$ can definitely oxidize $\ce{Cl-}$ to $\ce{Cl2}$. – Weijun Zhou Feb 7 '18 at 19:47
• I agree, but there are more to consider besides redox. For example, as written in the comment, $\ce{PbCl2}$ precipitates. – Weijun Zhou Feb 7 '18 at 20:03
• Isn't the lead in $\ce{PbCl2}$ reduced with respect to $\ce{Pb(NO3)4}$? – Vinícius Godim Feb 7 '18 at 20:06