7.1.5 Identification of Species

Practical Skills Development

• Use of appropriate qualitative reagents and techniques to analyse and identify unknown samples or products including gas tests, flame tests, precipitation reactions
• Safe use of appropriate heating devices and techniques including use of a Bunsen burner and a water bath or an electric heater

Aim

• To carry out flame tests, cation tests and anion tests

Health & Safety Aspects

• Safety glasses should be worn throughout the tests
• Small volumes should be used to minimise the risk of these hazardous substances

Flame Tests

Materials

• Flame test wires
• Solids samples of the chlorides of sodium, lithium, potassium, calcium and copper(II)
• Concentrated hydrochloric acid
• Bunsen burner

Method

1. Dip the loop of an unreactive metal wire such as nichrome or platinum in dilute acid, and then hold it in the blue flame of a Bunsen burner until there is no colour change
2. This cleans the wire loop and avoids contamination
   • This is an important step as the test will only work if there is just one type of ion present
   • Two or more ions means the colours will mix, making identification erroneous
3. Dip the loop into the solid sample and place it in the edge of the blue Bunsen flame
4. Avoid letting the wire get so hot that it glows red otherwise this can be confused with a flame colour

Diagram showing the technique for carrying out a flame test

The colour of the flame and the metal ion present

Diagram showing the colours formed in the flame test for metal ions

• Metal ions produce a colour if heated strongly in a flame
• Ions from different metals produce different colours
• The flame test is thus used to identify metal ions by the colour of the flame they produce

Cation Tests

Materials

• 0.5 mol/dm³ sodium hydroxide solution
• Solutions containing the cations: Zn²⁺/Al³⁺, Mg²⁺, Ca²⁺, Cu²⁺, Fe²⁺ & Fe³⁺
• Dropping pipettes
• Test tubes

Method

1. A few drops of NaOH are added at first and very slowly to the test solution
   • If it is added too quickly and the precipitate is soluble in excess, then you run the risk of missing the formation of the initial precipitate which dissolves as quickly as it forms if excess solution is added
2. Any colour changes or precipitates formed are noted
3. Then the NaOH is added in excess and the reaction is observed again

Results

Cation Tests Results

Testing Anions - Sulfate

Materials

• 0.2 mol/dm³ barium chloride solution
• 1.0 mol/dm³ nitric acid
• Test tubes & dropping pipettes
  

Method

1. Acidify the sample with dilute hydrochloric acid and then add a few drops of aqueous barium chloride
   • The test can also be carried out with barium nitrate solution
2. If a sulfate is present then a white precipitate of barium sulfate is formed:

Ba²⁺ (aq) + SO₄^2- (aq) → BaSO₄ (s)

A white precipitate of barium sulfate is a positive result for the presence of sulfate ions

Testing Anions- Halides

Materials

• 0.05 mol dm^-3 silver nitrate solution
• 1.0 mol dm^-3 nitric acid
• Test tubes & dropping pipettes
  

Method

1. Acidify the sample with dilute nitric acid, HNO₃, followed by the addition of silver nitrate solution, AgNO₃
   • The acidification is done to remove carbonate ions that might give a false positive result
2. If a halide is presents a silver halide precipitate is formed

Ag⁺ (aq) + X^– (aq) → AgX (s)

• • Depending on the halide present, a different coloured precipitate is formed, allowing for identification of the halide ion
  • Silver chloride is white, silver bromide is cream and silver iodide is yellow

Each silver halide produces a precipitate of a different colour

Testing Anions- Carbonate

Materials

• 1.0 mol dm^-3 hydrochloric acid
• limewater solution
• Delivery tube  assembly

Method

1. Add a few cm³ of dilute hydrochloric acid  and pass the gas evolved through limewater
2. If a carbonate compound is present then effervescence should be seen and the gas produced is CO₂ which forms a white precipitate of calcium carbonate when bubbled through limewater:

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

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

Results

Limewater turns milky in the presence of CO₂ caused by the formation of insoluble calcium carbonate