How to synthesize anhydrous zinc(II) nitrate?

I'm interested in making anhydrous zinc nitrate. I know it decomposes fairly rapidly if you heat it. It is also soluble in ethanol.


Would another desiccant or a molecular sieve do the trick? Or maybe it can be made directly somehow? There is surprisingly little information on ranking desiccate materials.




Reaction with dinitrogen tetroxide\* $\ce{N2O4}$ is probably the most prominent and convenient way of obtaining anhydrous zinc(II) nitrate, as well as alkali metal and mercury(II) nitrates according to the scheme (the reaction with zinc goes to completion and is therefore studied in more detail [[1](https://doi.org/10.1039/jr9510002829)]):


$$\ce{M + x N2O4 -> M(NO3)\_x + x NO} \quad \ce{M} = \ce{Li, Na, K, Rb, Zn, Hg}$$


A piece of zinc (5 g) with a freshly cleaned surface is immersed in liquid nitrogen tetroxide (10 ml).
Within 10 minutes the metal is covered with a thin layer of white fine-crystalline substance, nitrogen oxide is released, and the liquid in direct contact with the metal turns green due to release of nitrogen trioxide $\ce{N2O3}$.


As the product layer becomes thicker, $\ce{Zn(NO3)2·2N2O4}$ solvate crystals begin to form, which are easily separated from the metal surface with gentle stirring.
In 24 h, about 1–2 g of the product is formed, but its formation in the future is increasingly difficult to control.
After removal of excess metal, nitrogen tetroxide is distilled off by evaporation.
Here, a free-flowing fine-crystalline powder is formed.
Nitrogen tetroxide is weakly bound to zinc nitrate and is released when the powder is heated.
When about $10\,\%$ of the attached tetroxide is removed, the solid turns into a liquid.
The change in the phase state is caused by a decrease in the melting temperature of the initial solvate as a result of the formation of zinc nitrate.
Since the reaction


$$\ce{Zn(NO3)2·2N2O4 <=> Zn(NO3)2 + 2 N2O4}$$


is reversible, a complete removal of the tetroxide occurs by heating for 6 h at 100 °C under reduced pressure.
The resulting anhydrous zinc nitrate is a free-flowing white powder.


Nitrogen trioxide formed in solution significantly affects on the reaction rate. [[2](https://doi.org/10.1039/JR9510002838)]
If the piece of metal is stationary, the reaction rate gradually increases, but if the solution is stirred by the metal itself, the rate *decreases*.
This behavior distinguishes this system from the usual heterogeneous processes of solid-liquid interaction when a concentration gradient of active forms is established near the surface, and stirring *increases* the reaction rate.


This phenomenon is explained by the presence of nitrogen trioxide at the metal surface.
The molecules of this substance are more polar than the molecules of tetroxide, and therefore they increase the dielectric constant of the medium and the concentration of the active form of $\ce{NO+}.$


The presence of the trioxide at the surface facilitates the reaction, and with stirring, the substance is distributed in solution.
The trioxide does not directly participate in the reaction; otherwise, the formation of metal nitrite would inevitably occur as a result of the ionic dissociation


$$\ce{N2O3 <=> NO+ + NO2-}.$$


The rate of dissolution of zinc is proportional to $\sqrt{[\ce{N2O3}]}$.
At 0 °C the reaction rate in a solution containing $40\,\%$ nitrogen trioxide is eight times greater than in pure nitrogen tetroxide, but the product does *not* contain traces of nitrites.




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For an extensive overview I suggest reading chapter *Anhydrous Metal Nitrates*, section II *Preparative Methods* in [3, pp. 73–136].
Below are some relevant points and considerations taken from that review.


* Thermal dehydration as well as dehydration of the solid hydrate over phosphoric oxide in a vacuum desiccator at room temperature yields insoluble *basic* salt.
* Desiccants and molecular sieves won't be effective due to the fact that water molecules in crystallohydrates usually belong to inner coordination sphere and are not merely held by hydrogen or van der Waals bonding or host-guest interactions.
* Displacement in nonaqueous solvents is not a viable option due to poor solubility of nitrates (silver(I) nitrate is typically used) in organic solvents and highly problematic or impossible solvent adduct removal from obtained solvated metal nitrate.
* Anhydrous nitric acid, as straightforward as it seems, usually doesn't react completely with the dried salt, and only a partial replacement occurs, often yielding in double salts.
Additionally, there is always some water present due to equilibrium
$$\ce{2HNO3 <=> N2O5 + H2O}.$$
* Addition of dinitrogen pentoxide to nitric acid assists in binding water and provides an additional source of water-free nitrate anion due to ionic dissociation in $\ce{HNO3(conc)}$:
$$\ce{N2O5 <=> NO2+ + NO3-},$$
although conversion of zinc(II) salts into anhydrous nitrate using either this method or chlorine/acetyl nitrates $(\ce{ClNO3},$ $\ce{CH3C(O)ONO2})$ hasn't been reported at the time.
Pure $\ce{N2O5}$ reacts directly with many metals, albeit with low selectivity giving mixed products.




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\* In a lab pure [dinitrogen tetroxide](https://en.wikipedia.org/wiki/Dinitrogen_tetroxide) is easily obtained by thermal decomposition of lead nitrate.
Under normal conditions, $\ce{N2O4}$ is a liquid substance (m.p. −11.2 °C, b.p. 21.15 °C), and the reactions are almost always carried out at room temperature.
Traces of moisture in $\ce{N2O4}$ can be detected simply by freezing the liquid, and this is one of the most sensitive methods of testing.
Pure $\ce{N2O4}$ freezes to form a colorless glassy substance, but in the presence of traces of moisture nitrogen trioxide $\ce{N2O3}$ is formed, which turns the frozen tetroxide green.
Usually it is recommended to cool the liquid to about 10 °C prior to synthesis, otherwise the heat released during the reaction may cause vigorous boiling and the release of toxic fumes.


### References


1. Addison, C. C.; Lewis, J.; Thompson, R. 628. The Liquid Dinitrogen Tetroxide Solvent System. Part VII. Products of Reaction of Zinc with Liquid Dinitrogen Tetroxide. *J. Chem. Soc.* **1951**, 2829. DOI: [10.1039/jr9510002829](https://doi.org/10.1039/jr9510002829).
2. Addison, C. C.; Lewis, J.; Thompson, R. 630. The Liquid Dinitrogen Tetroxide Solvent System. Part IX. Products, Rates, and Possible Mechanisms of the Reaction of Zinc with Liquid Dinitrogen Trioxide–Tetroxide Mixtures. *J. Chem. Soc.* **1951**, 2838. DOI: [10.1039/JR9510002838](https://doi.org/10.1039/JR9510002838).
3. *Advances in Inorganic Chemistry and Radiochemistry*; Emeléus, H. J., Sharpe, A. G., Eds.; Academic Press: New York, **1964**; Vol. 6.


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