GCSE OCR Gateway C4:2b Carbonate and sulphate, halide and nitrate tests.

GCSE OCR Gateway C4:2b Carbonate and sulphate, halide and nitrate tests.

Learning Objectives:

C4.2b To be able to describe tests to identify various aqueous anions: carbonates using hydrochloric acid and limewater; sulfates using hydrochloric acid followed by aqueous barium chloride solution; chloride, bromide and iodide using nitric acid followed by silver nitrate solution.

Anions are negatively charged ions so named because they are attracted to and discharged from the positive anode during electrolysis. 

There are several common anions in school and college chemistry and each one comes from a simple acid.

Acid	Hydrochloric (HCl)	Sulphuric (H2SO4)	Nitric (HNO3)	Carbonic (H2CO3)
Anion	Chloride  Cl–	Sulphate  SO42–	Nitrate NO3–	Carbonate CO32–

How to test for halide ions

Halide ion refers to the anions of chlorine, bromine and iodine.  The halides: chlorides, bromides and iodides require just one simple test to identify them.

First to the solution you suspect contains halide ions add a few drops of nitric acid.  The nitric acid is there to remove any other ions particularly carbonate ions that would also form a precipitate with silver nitrate since silver carbonate (Ag₂CO₃) is a white insoluble solid.

Second, add several drops of 0.05M silver nitrate solution (AgNO₃(aq)) to the solution of halide ion.  The colour of the precipitate will tell you which halide ion is present. 

The silver ions in the silver nitrate solution form insoluble coloured precipitates with chlorides, bromides and iodides.

With Chlorides:

Silver nitrate forms a white precipitate with chlorides.

AgNO₃ (aq)    +     NaCl(aq)       ⟶      AgCl(s)    +    NaNO₃ (aq)

                                                        White ppt

The ionic equation:

Ag⁺ (aq)    +     Cl^—(aq)        ⟶         AgCl(s)

With Bromides:

Silver nitrate forms a cream precipitate with bromides.

AgNO₃ (aq)    +     NaBr(aq)    ⟶       AgBr(s)       +     NaNO₃ (aq)

                                                       Cream ppt

The ionic equation:

Ag⁺ (aq)    +     Br^—(aq)       ⟶          AgBr(s)

With Iodides:

Silver nitrate forms a yellow precipitate with iodides.

AgNO₃(aq)    +     NaI(aq)    ⟶         AgI(s)       +      NaNO₃(aq)

                                                     Yellow ppt

The ionic equation:

Ag⁺ (aq)    +     I^—(aq)     ⟶        AgI(s)

Further confirmatory tests:

If dilute ammonia solution is added to the white precipitate of silver chloride, the precipitate should dissolve.

If concentrated ammonia solution is added to the cream precipitate of silver bromide, the precipitate should dissolve.

The yellow precipitate of silver iodide is insoluble in ammonia solution.

Effect of bright sunlight

Lastly, put the three precipitates in bright light or sunlight and after a few minutes the white precipitate of silver chloride turns dark purple, the bromide turns dark green but the iodide remains yellow.


Summary

How to test for a sulphate SO₄^2–

First to the solution you suspect contains sulfate ions add a few drops of hydrochloric acid.  The hydrochloric acid is there to remove any other ions particularly carbonate ions that would also form a precipitate with barium chloride since barium carbonate (BaCO₃) is a white insoluble solid.

Second, add several drops of 0.1M barium chloride solution (BaCl₂(aq)) to the solution of sulfate ion.

The barium ions in the barium chloride solution form an insoluble white precipitate with sulphate ions.

BaCl₂ (aq)    +     Na₂SO₄(aq)       ⟶      2NaCl(aq)    +    BaSO₄(s)

                                                                                  White ppt

The ionic equation:

Ba²⁺(aq)    +     SO₄^2—(aq)        ⟶         BaSO₄(s)

How to test for a carbonate CO₃^2–

First, carbonates are usually solids, only Group 1 carbonates are soluble in water.  So to the solid (or solution) you suspect contains carbonate ions add a few drops of hydrochloric acid.  The hydrochloric acid reacts to form carbon dioxide and fizzing occurs instantly. 

Second, pass the gas produced into limewater and if it is carbon dioxide you will see cloudiness or chalkiness in the limewater.

Hydrochloric acid neutralizes the carbonate ions to from carbon dioxide.

CaCO₃ (s)    +    2HCl(aq)   ⟶      CaCl₂(aq)    +    CO₂(g)   +  H₂O(l)

The ionic equation:

CO₃^2-(aq)    +     2H⁺(aq)     ⟶      CO₂(g)   +  H₂O(l)  

And when the carbon dioxide gas enters limewater the following reaction takes place.

CO₂(g)    +   Ca(OH)₂(aq)         ⟶         CaCO₃(s)  +   H₂O(l)

                                                              chalkiness

How to test for the nitrate ion NO₃^–

Now this test isn’t always included in school and college chemistry but so what it’s different and a great test with great chemistry.

Of course the reason for the absence of nitrate tests in school and college chemistry is that they are not always reliable.

I think I can agree with Wiki when it says:

Testing for the presence of nitrate via wet chemistry is generally difficult compared with testing for other anions, as almost all nitrates are soluble in water. In contrast, many common ions give insoluble salts, e.g. halides precipitate with silver, and sulphate precipitate with barium.

The nitrate anion is an oxidant, and many tests for the nitrate anion are based on this property. Unfortunately, other oxidants present may interfere and give erroneous results.

A test that is reliable uses an alloy of copper, aluminium and zinc called Devarda’s Alloy. 

If you suspect a solution contains a nitrate ion then add a few drops sodium hydroxide solution and a piece of Devarda’s Alloy to the solution. 

Heat the solution gently in a fume cupboard and test the gas given off with a piece of damp red litmus paper.  If the paper turns blue the gas is likely to be ammonia and the ion in the solution is a nitrate!!

3 NO3−   +    8Al    +    5OH−  +   18H2O   →    3NH3    +    8[Al(OH) 4]−

It is unusual for any alkaline gas to evolve in any test in wet chemistry other than ammonia so you can be pretty sure that the evolution of an alkaline gas confirms the presence of the nitrate anion.

In my next posts, I’m going to discuss titration technique and develop and update some earlier work in this blog.

GCSE OCR Gateway C4:2 b Sodium Hydroxide Tests for Cations

GCSE OCR Gateway C4:2 b Sodium Hydroxide Tests for Cations

Learning Objectives:

C4.2b To be able to describe tests to identify the aqueous cations: calcium, copper, iron (II), iron (III)
and zinc using sodium hydroxide solution.

One of the most powerful things in chemistry is our ability to identify unknown chemicals that are causing problems of pollution or contamination or sickness.

Of course, some chemists have to put up their hands and confess to polluting the planet.  But sadly chemists have not been quick to do this.  So pollutants have led to human misery and to the association of chemistry with names such as Minamata, Bhopal or Thalidomide.

Minamata is the name of a Bay on the coast of Japan. Back in the 1900’s it was a beautiful place where Japanese fishermen plied their trade as they had done for centuries catching fish, mainly tuna.

Then came the industry.  A chemical factory was built outside Minamata town and started production of fertiliser in 1908.  As it expanded both before and after WW2 its range of products increased.

The Minamata factory started ethanal (CH₃CHO) production in 1932, producing 210 tons that year. By 1960, production reached a peak of 45,245 tons.

Ethanal production used mercury sulphate as a catalyst. From August 1951, they changed the co-catalyst from manganese dioxide (MnO₂) to ferric sulfide (Fe₂S₃).

What they did not realise was that a side reaction then started. The process produced a small amount of the highly toxic organic mercury compound, methyl mercury (CH₃HgCl). The company released this compound into Minamata Bay from the change of the co-catalyst in 1951 until 1968, when they stopped this production method.

The identification of mercury in the water of the bay around the town was easy using typical chemical tests like those required for your school or college chemistry course.

Linking mercury to the disease symptoms emerging in the local population was more difficult.

Photos below show the effects of the mercury poisoning:

You can read more about the investigation here.

How to use Sodium Hydroxide to detect Metal Cations.

The procedure to detect metal cations is not complicated.

You place a few drops of the suspected metal cation solution on a spotting tile or in a test-tube and add a couple of drops of 1M sodium hydroxide solution. 

A cloudiness in the solution shows that a precipitate has formed.  The colour of the precipitate is characteristic of the metal cation. 

Metal Cation Mn+	Calcium Ca2+	Copper Cu2+	Iron (II) Fe2+	Zinc Zn2+	Iron (III) Fe3+
Colour of the precipitate	white	Pale blue	green	white	brown

Adding excess sodium hydroxide will tell the calcium and zinc apart because the zinc precipitate will dissolve in enough excess sodium hydroxide.

What’s happening in these precipitation reactions?

Each precipitate is a metal hydroxide (M(OH)_n) where n is the charge on the ion.

So for example:  iron (III) sulphate (Fe₂(SO₄)₃).

Iron(III)sulphate  +  sodium hydroxide  ⟶ Sodium sulphate  +  iron(III)hydroxide

Fe₂(SO₄)₃(aq)  +   6NaOH(aq)   ⟶  3Na₂SO₄(aq)   +  2Fe(OH)₃(s)

                                                                               brown ppt      

But there are in the reaction ions that do not change in the reaction these are called Spectator Ions

The spectator ions are highlighted in Red

The remaining ions form the ionic equation:

2Fe³⁺(aq)   +    6OH^—(aq)    ⟶    2Fe(OH)₃(s)

                                                  brown ppt      

Can you build the equations for the reactions between the other metal ions and sodium hydroxide where the charge on the metal ion is 2+?

The zinc precipitate dissolves because it is amphoteric. 

Amphoteric means that the precipitate reacts with both acids and bases.

Here are the equations:

Zinc sulphate  +  sodium hydroxide ⟶  sodium sulphate +  zinc hydroxide.

ZnSO₄(aq)   +    2NaOH(aq)    ⟶    Na₂SO₄(aq)    +   Zn(OH)₂(s)

                                                                              white ppt

Zn²⁺(aq)   +    2OH^—(aq)    ⟶    Zn(OH)₂(s)

Then with excess sodium hydroxide

Zinc hydroxide +  sodium hydroxide ⟶   sodium zincate

Zn(OH)₂ (s)    +   2NaOH (aq) ⟶       Na₂Zn(OH)_{4  (aq)} 

You might find that this post of mine here helps you build the precipitation equations.

In my next post I’m going to describe anion tests.