The Chemistry of Nitrites and nitrates
Note: IUPAC names of nitrite and nitrate are nitrate(III) and nitrate(V), respectively.
Introduction
This series of experiments is designed to survey the chemistry of nitrites in relation to nitrates
and other simple nitrogen compounds. In order to interpret the behaviour of the reagents used, it
will be necessary to ascertain whether a reaction has occurred and, if so, what products have been
formed. Thus, very careful observations must be made , and additional test may be required. The
half-equations given below are listed in order of decreasing standard reduction potential and,
when you have established the identity of the products of a reaction, these tables will be of help
in interpreting the reaction and writing an overall equation for it.
Table 1
Complete the following half-equations by inserting the number of electrons involved in each case.
Oxidizing agent
Reducing agent
(1) MnO4- + 8H+ + 5e<==> Mn2+ + 4H2O
Strength of
Strength of
(2) NO2-+ H+ + e- <==> N2 + 4H2O
(3) Br2 + e
<==> 2 Br
Oxidizing
reducing
+
(4) HNO2 + H + e
<==> NO + H2O
agent
agent
(5) NO3- + H+ + e<==> HNO2 + H2O
decreases
decreases
(6) Fe3+ + e<==> Fe2+
(7) I2 + e
<==> 2I
(8) N2 + H+ + e<==> 2NH4+
(9) NO3- + H2O + e- <==> NH3 + 9OH(10)NO2- + H2O + e- <==> NH3 + 7OH(11) [Al(OH4)- +
e<==> Al + 4OHApplication of the table
Under standard conditions, no oxidizing agent can react with a reducing agent above it in the
above table. Conversely, an oxidizing agent will react with a reducing agent below it in the table,
the reaction may be very fast or so slow as to be of no practical significance.
Example
If permanganate ions react with iron(II) ions, in acid solution, it can be shown that manganese(II)
ions and iron(III) ions are formed. Thus half-equations (1) and (6) are involved. Combining these
half-equations the overall equations can be constructed.
MnO4- + 8H+ + 5Fe2+ <===> Mn2+ + 4H2O + 5Fe3+
Questions for discussion
1. What are the 'standard conditions' for the half-equations?
2. Predict whether iron(III) ions would oxidize (a) bromide ions and (b) iodide ions under standard
conditions.
3. Assign each of the nitrogen-containing species to a position in table 2 showing the oxidation
number of the nitrogen in the species.
Table 2
Oxidation no.
of nitrogen
+V +IV +III +II +I
0
-I
-II -III
---+-----+------+------+-----+------+-----+-----+-----+---
Requirements:
In this experiment you will investigate the reaction between oxidizing agent and reducing agent.
You should observe any change carefully. Write down the observations immediately IN INK. Try
to explain the change by the use of the half-equations and finally work out the overall equation
for each of the following experiments.
Procedure
If you suspect that gas evolved may be oxygen, dinitrogen monoxide or carbon dioxide, test the
gas.
A. Test for nitrite and nitrate ions.
(1) Add about five drops of a fresh solution of iron(II) sulphate, acidified with dilute sulphuric
acid, to approximately 2 cm3 of dilute sodium nitrite solution.
(2) Add about five drops f a fresh solution of iron(II) sulphate, acidified with dilute sulphuric
acid, to approximately 2 cm3 of dilute potassium nitrate solution. Carefully run
concentrated sulphuric acid acid down the side of the tube to form a separate layer.
B. Reaction with acid.
(1) Add dilute hydrochloric acid to about 2 cm3 of aqueous sodium nitrite solution in a
test-tube.
(2) Warm the solution until the reaction is complete. Test the remaining solution for the
presence of (i) nitrite and (ii) nitrate ions.
(3) Repeat experiments (1) and (2) substituting potassium nitrate solution for sodium nitrite
solution.
C. Reaction with other common oxidizing agents.
(1) Acidify 2 cm3 of bromine water with dilute sulphuric acid. Add sodium nitrite solution
dropwise to the bromine water, until no further colour change occurs.
(2) Add sodium nitrite solution dropwise to about 2 cm3 of potassium permanganate solution
acidified with dilute sulphuric acid.
(3) Repeat experiments (a) and (b) substituting potassium nitrate solution for sodium nitrite
solution.
D. Reaction with some reducing agents
(1) Add sodium nitrite solution dropwise to about 2 cm3 of potassium iodide solution, acidified
with dilute sulphuric acid.
(2) Repeat the experiment substituting potassium nitrate solution for sodium nitrite solution.
(3) Make about 2 cm3 of sodium nitrite solution alkaline by the addition of sodium hydroxide
solution, add some aluminium turnings (they should be rinsed by dilute hydrochloric acid,
followed by water). The mixture is warmed.
(4) Repeat experiment (b) substituting potassium nitrate solution for sodium nitrite solution.
Question for discussion
Explain why the reaction is carried out in alkaline conditions, and why aluminium is used. What
other metal do you think would perform this reaction with nitrites?
E. Reaction with ammonium chloride
(1) Boil about 2 cm3 of sodium nitrite solution with excess ammonium chloride until no gas,
other than steam, is evolved. (This may be achieved by boiling the mixture until gas is no
longer evolved, and then adding more solid aluminium chloride and repeating until the
addition of further ammonium chloride does not produce more gas.) Attempt to establish
the identify of the gas evolved, and test the remaining solution for the presence of nitrite
ions.
(2) Repeat the above experiment substituting potassium nitrate solution for sodium nitrite
solution.
Results:
Tabulate your results for the above sections in the following table
Section
Nitrite
Nitrate
A. Test for nitrite and
nitrate
B. With acids
C With Br2
MnO4/H+
D. With KI(aq)
Al/OHE. With ammonium
chloride
Questions for discussion
1. Give immediate conclusion for each section.
2. What procedure would you adopt to test for nitrate ions in the presence of nitrite ions?
Conclusion
Quoting experimental evidence, compare the efficiency as an oxidizing agent of nitrous acid or the
nitrite ion with that of nitric acid or the nitrate ion.
End