Could I have made Cu(NO2)2?
The direct synthesis of copper(II) nitrite, Cu(NO2)2, as a stable, isolable solid is unlikely under typical laboratory conditions. Copper(II) nitrite is generally known only in solution form and tends to decompose upon drying into copper(II) nitrate, copper oxide, and nitrogen oxides. Therefore, while a copper-nitrite complex might form transiently in solution, producing solid Cu(NO2)2 is highly improbable using the described method.
Analysis of the Experimental Procedure
The experiment begins with copper metal reacting with a mixture of hydrogen peroxide (3% H2O2) and sulfuric acid (37% H2SO4) to produce a copper sulfate (CuSO4) solution. Following this, the solution was neutralized carefully using ammonia and acetic acid to avoid generating nitrogen oxides (NO, NO2) gases during later steps.
Upon adjusting the pH to near neutral (~7), sodium nitrite (NaNO2) was added. This produced a dark green complex in solution, which aligns with observations of copper and nitrite interactions in aqueous media.
Boiling the solution concentrated it, causing precipitation of a solid assumed to be poorly soluble in water. Addition of acid reversed the color change to blue, suggesting a loss or transformation of the nitrite complex into copper sulfate or nitrate species.
The precipitate was filtered and dried. Residual solid on the filter was treated with hydrochloric acid (HCl), causing noticeable emission of NO gas, indicating nitrate or nitrite residues. This also hints at incomplete washing and the presence of nitrite trapped in the solid.
Identity of the Solid Product
The exact nature of the precipitate remains uncertain. The following hypotheses arise:
- Copper nitrate: Copper nitrates (Cu(NO3)2) are stable and blue-colored, commonly formed by copper reaction with nitric acid. There is limited data on solid copper nitrite.
- Basic copper salts: The solid might be copper hydroxide or basic copper carbonate due to partial neutralization and atmospheric CO2 absorption.
- Copper nitrite: Though the green complex suggests copper nitrite presence, isolated solid Cu(NO2)2 is reportedly unstable.
The acid addition causing a blue solution indicates excess copper sulfate or nitrate formation rather than a stable copper nitrite solid.
Known Characteristics of Copper(II) Nitrite
Scientific literature and practical reports show copper(II) nitrite exists stably only in solution. One preparation involves mixing solutions of copper(II) sulfate and lead(II) nitrite, precipitating lead(II) sulfate, and leaving behind copper(II) nitrite in solution.
When attempts are made to isolate copper(II) nitrite as a solid, it readily decomposes. Typical decomposition products include copper(II) nitrate, copper oxide (CuO), and nitrogen oxides (NO, NO2), hence drying solid copper nitrite is challenging.
Implications for Your Experiment
The dark green complex formed after adding sodium nitrite probably corresponds to copper coordinated with nitrite ions in solution. The color change upon acidification, precipitation, and gas evolution during washing strongly support the presence of mixed copper-nitrogen-oxygen species rather than a stable copper nitrite salt.
Thus, the initial hypothesis that the isolated precipitate is copper nitrite is doubtful. Instead, it likely contains basic copper salts, copper nitrate, or decomposed species due to the instability of copper nitrite.
Recommended Tests to Confirm Compound Identity
To clarify the composition of the solid, several tests can help:
- Redissolve the precipitate in water and test for nitrite ions (NO2-) using standard colorimetric tests such as the Griess reagent.
- Test for nitrate ions (NO3-) using brown ring test or diphenylamine test after reduction.
- Analyze the solid using techniques such as infrared spectroscopy (IR), X-ray diffraction (XRD), or elemental analysis to determine its composition.
Comments on Experimental Approach
Producing copper(II) sulfate by dissolving copper tubing in H2O2 and sulfuric acid is practical but introduces variability. Commercial copper sulfate pentahydrate is inexpensive and highly pure, and its use is recommended for reproducibility and accuracy.
Neutralization with ammonia requires careful pH control. Overshooting leads to ammine complexes or basic copper salts, which complicate product identity.
Boiling the solution to concentrate it may promote decomposition of sensitive copper-nitrite complexes, causing unwanted side reactions.
Handling of precipitates and washing steps must be thorough to avoid contamination with unreacted nitrites or nitrates that can confound analysis.
Summary Table: Key Points on Copper Nitrite Synthesis
| Aspect | Observation | Interpretation | Recommendation |
|---|---|---|---|
| Copper source | Copper tubing dissolved in H2O2 + H2SO4 | Variable purity and concentration | Use commercial CuSO4 for consistency |
| pH control | Attempted neutralization, overshoot with ammonia + acetic acid | Potential formation of complex species and basic salts | Careful pH monitoring with titration |
| Complex formation | Dark green complex appears with NaNO2 | Likely copper-nitrite complex in solution | Verify ion presence, avoid prolonged boiling |
| Precipitate identity | Solid forms, green/dark green color, brown precipitate on filter | Likely basic copper salt or decomposed products | Perform ion tests and spectral analysis |
| Gas evolution upon acid wash | NO or NO2 smell noticed | Possible nitrite/nitrate decomposition | Ensure thorough washing to remove residual ions |
Conclusion: Can Cu(NO2)2 be made this way?
Based on chemical knowledge and experimental results, copper(II) nitrite as a stable solid is not formed by mixing copper sulfate solutions with sodium nitrite followed by concentration and drying. The green solution complex is consistent with copper-nitrite interaction in solution. However, the isolated solid likely comprises other copper species due to instability of copper nitrite upon drying.
Future work to confirm compound identity should emphasize careful ion analysis, avoidance of strong heating, and possibly in situ spectroscopic characterization. Purchasing high purity starting materials can enhance reproducibility.
Key Takeaways
- Copper(II) nitrite exists stably only in solution; attempts to isolate solid lead to decomposition.
- Dark green complexes formed upon adding NaNO2 to copper(II) solution represent transient copper-nitrite species in aqueous media.
- The precipitate obtained is unlikely Cu(NO2)2 but rather basic copper salts or copper nitrate.
- Proper pH control and use of commercial copper sulfate improve experimental consistency.
- Testing for nitrate and nitrite ions, along with spectral analysis, helps confirm compound identity.
Could I have made Cu(NO2)2 in the experiment described?
Most likely not. Copper(II) nitrite is known only in solution and decomposes upon drying. The solid you obtained is probably a mixture like copper nitrate or copper hydroxide.
Why did adding acid change the solution color from green to blue?
The blue color indicates not enough nitrite present to maintain the green complex. Acid causes the complex to break down, shifting color and composition.
Can copper nitrite be isolated as a solid easily?
No. Copper nitrite tends to decompose when dried, forming copper nitrate, copper oxide, and nitrogen oxides. It is stable only in solution.
What tests can confirm the identity of the compound formed?
Dissolve the solid and test for nitrite (NO2-) and nitrate (NO3-) ions. This helps distinguish copper nitrite from copper nitrate or other compounds.
Is making copper sulfate from copper tubing recommended?
It’s better to buy pure copper sulfate. Homemade solutions vary and can complicate reactions. Commercial copper sulfate is cheap and consistent.
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