9.4.C - Ion Analysis


A variety of quanitative tests are available to determine the concentration of common anions (negative ions) and cations (positive ions) in industry. These include techniques such as mass spectroscopy, infra-red spectroscopy and atomic absorption spectroscopy. Quantitative tests are those tests that determine how much of the concentration of an ion in a sample.

In a school laboratory these are not available and only qualitative tests can be carried out. Such tests determine whether the ion is present or not in the sample. These include flame tests, precipitation reactions and other simple tests that can be carried out safely and quickly but will only identify the presence or absence of an ion.

In this topic students:

  • Conduct a range ofdefinitive chemical tests to determine if an ion is present in a sample of water
  • Explain the role of precipitation reactions and flame tests in identifying ions
  • Use gravimetric analysis to determine the sulfate content of lawn fertiliser
  • Explain how atomic absoprtion spectroscopy is used to measure small concentrations of metal ions
  • Describe and explain the need for monitoring levels of ions in substances used by society

Identification of Common Ions

Precipitation Reactions

In the classroom, our methods of qualitative ion analysis are usually limited precipitation reactions and flame tests. The formation of precipitates is based on the solubility rules and flame tests depend on some ions having characteristic colours in a Bunsen flame.

A precipitation reaction is a reaction in which soluble ions in separate solutions are mixed together to form an insoluble compound that settles out of solution as a solid. That insoluble compound is called a precipitate. The table below shows some characteristic precipitates and their colours. The cation-anion combinations that form a precipitate are labelled white or with a colour.

Cations
Anions
phosphate sulfate carbonate chloride hydroxide nitrate
barium white white white no prec no prec no prec
calcium white white white no prec no prec no prec
lead white white white prec white no prec
copper blue no prec green no prec blue no prec
iron(II) green no prec grey no prec green no prec
iron(III) no prec no prec no prec no prec brown no prec
silver yellow No prec white white brown no prec

The solubility rules for some common anions and cations.

Solubility rules are useful summaries of information about which ionic compounds (or combinations of ions) are soluble in water and which are not. They are also important tools for making predictions about whether certain ions will react with one another to form a precipitate. In addition, they are useful for figuring out what ions might be involved when a precipitation reaction has been observed.

As you can see from the table, many precipitates are white and as such cannot be identified on this basis alone. For example, barium and calcium ions form the same white precipitates with the same anions. They can be distinguished by placing the precipitate in a colourless Bunsen burner flame (with the air hole open). The barium produces a dull green colour. In contrast, the calcium produces a dull red colour in the flame. Copper or copper salts form a bright green colour in a flame, easily distinguished from the dull green colour of barium.

Apart from precipitation reactions and flame tests, there are other reactions that confirm the presence of some ions. The presence of carbonate ions can be confirmed by the addition of hydrochloric acid to a sample. The production of gas, which when bubbled into a sample of limewater turns it milky, indicates the presence of carbonate ions (acid + carbonate → salt + water + carbon dioxide). The presence of phosphate ions can be confirmed by the addition of a solution of ammonium molybdate and nitric acid to a sample. The precipitation of a yellow solid indicates phosphate ions.

Ion Identification Tests

The table below describes the test and results for identifying some common ions in samples of water. There are many different ways of identifying ions and the lists below represent one possible test for each ion.

Cations
Lead Pb2+

A paint-like yellow precipitate of lead iodide forms when sodium iodide is added.

A yellow precipitate of lead chromate forms when sodium chromate is added.

Copper Cu2+

A blue precipitate of copper(II) hydroxide forms when sodium hydroxide os added. This precipitate darkens on standing.

It produces a green flame test.

Iron (II) Fe2+

Forms a deep blue precipitate with potassium ferricyanide.

Forms a white precipitate of iron (II) hydroxide when sodium hydroxide is added. This precipitate turns brown on standing.

Iron (III) Fe3+

Forms a rusty/brown precipitate of iron (III) hydroxide when sodium hydroxide is added.

Forms a blood red solution of FeSCN2+ when potassium thiocyanate (KSCN) solution is added.

Calcium Ca2+

Forms a white precipitate of calcium sulfate when sulfuric acid is added.

Produces a brick-red flame colour.

Barium Ba2+

Forms a white precipitate of barium sulfate when sulphuric acid is added.

Produces a pale green flame colour.

 

Anions
Carbonate CO32-

Bubbles of gas are produced when a few drops of hydrochloric acid are added. The gas can be identified as carbon dioxide because it turns limewater milky when bubbled through it.

A white precipitate of silver carbonate is firmed when silver nitrate is added.

Chloride Cl- A white precipitate of silver chloride is formed when silver nitrate is added. This white precipitate turns purple or brown when exposed to sunlight.
Phosphate PO43-

A yellow precipitate of silver phosphate forms when silver nitrate is added.

A fine white precipitate forms when barium nitrate is added.

Sulfate SO42-

A fine white precipitate of barium sulfate forms when barium nitrate is added. This precipitate will pass through filter paper.

A faint white precipitate of silver sulfate forms when silver nitrate is added.

 

Atomic Absorption Spectroscopy

The Atomic Absorption Spectrometer, AAS, was invented in 1953 by Alan Walsh at the CSIRO. It is used to determine the concentration of metal ions present in a sample even when the concentrations are very low.

A simplified diagram of a typical atomic absorption spectrometer.

A special cathode ray lamp is inserted into the AAS machine which emits a distinctive spectrum for that particular metal element. This spectrum passes through an air-acetylene flame into which a solution of the metal being tested for is sprayed using a nebuliser to obtain a fine mist. The metal ions being tested absorb some of the wavelengths passing through the flame. The light beam then passes through a monochromator which acts as a filter to separate light of the chosen wavelength from other light. The intensity of the light is measured by a photomultiplier which is effectively a computer.

The amount of light absorbed by the sample is proportional to the concentration of the metal ion in the solution being tested.

When known concentrations a placed into the spectrometer,
absorbance values are measured and the calibration graph
above is obtained. From this graph, the concentration of
unknown samples can be interpolated as shown.

AAS is widely used in a range of areas including:

  • To analyse rock and mineral samples in the mining industry for purity of metal ions

  • To monitor pollution levels in water when water is polluted with metal ions.

  • To monitor pollution levels of metal ions in air.

  • To monitor essential trace elements in soil by the agricultural industry.

  • To analyse blood and urine samples for the presence of an excess or deficiency of a metal ion.

Trace elements are those elements that are required by living organisms in very low concentrations (1 to 100 ppm). For example zinc, cobalt and copper ions are required by humans for enzyme function in concentrations lower than 100 ppm. Manganese, copper and zinc are required by plants for growth. When AAS was developed it enabled these trace elements to be measured in soil in concentrations that could not previously be measured. Now various symptoms of plant disease can be linked to lower than needed concentrations of trace elements.

Example

A 1.00 g sample of tuna meat was ground up in a mortar and dissolved in 10.0 mL acid. The sample was analysed in an Atomic Absorption Spectrometer to determine its lead concentration. A lead cathode ray tube was selected that emitted a distinctive wavelength of light that would only be absorbed by lead. The absorbance recorded was 0.600. The concentration of lead could then be determined by comparing the absorption of light by standards of known concentration plotted on a calibration curve shown above.

This is effective data to gather for the purpose of pollution control. If 2.4 ppm is high then there must be lead in the water. Steps can then be taken to find the cause and reduce it.

 

Monitoring Lead Levels

The levels of ions present in substances used in society must be monitored to ensure they are at suitable concentration to maintain the health of plants and animals.

One metal ion that must be monitored carefully is lead. Lead is a toxic, heavy metal. It can damage all organs of the body, especially the brain, kidneys and reproductive system, by disrupting enzyme systems. Lead ions are similar in size to calcium ions, so lead can be taken up in bones and teeth. It can also be absorbed by fatty tissues. Lead inhibits the formation of haemoglobin in red blood cells, causing anaemia and reducing the ability of blood to carry oxygen. Lead is rapidly absorbed by the brain, where it causes neurological damage. In children, exposure to lead is correlated with learning disabilities, impaired hearing and behavioural disorders such as hyperactivity, aggressiveness and attention deficit. Children and adults can also show non-specific symptoms such as tiredness, headache, abdominal pain and irritability.

Lead is released into the environment as elemental lead, inorganic lead ions, and organic lead. It can enter the body by ingestion, inhalation and absorption through the skin. Lead accumulates in the body and is difficult to excrete. Since 1998, the Australian national goal for blood lead levels has been 100 mg/L. The accumulation of heavy metals such as lead in the body is called bioaccumulation. Once lead enters the food chain, it becomes present in increasing amounts in the tissues of organisms along the chain. This is called bioconcentration.

Following are some sources of lead in the environment.

  • Leaded petrol produces most of the lead in the atmosphere of cities such as Sydney.

  • Mining and refining of lead produces fumes, dust, tailings and slag wastes contaminated with lead.

  • Paints used before 1950 contained a very high percentage of lead and until the 1980’s all NSW State government buildings were coated with such paints. Lead in paints is released by deterioration with age, sanding, burning or demolition. Today the maximum allowable concentration of lead in paint is 0.25%.

  • Lead glazing of pottery, corrosion of lead plumbing materials, making stained glass and lead-lighting.

  • Manufacturing (e.g. batteries).

It is necessary to monitor the lead content of our atmosphere, water, food and soil. This is especially necessary in areas providing drinking water and food, in areas of heavy traffic, and near industries using or producing lead. The World Health Organisation recommends a range of 0.5 to 1.0 x 10-3 mg/L as a goal for average annual lead concentration in air. The maximum recommended lead concentration in water is 50mg/L.

 

Measuring the Sulfate Content of Fertiliser

There are a variety of ways that the percentage mass of sulfate in lawn fertiliser can be determined. Some methods are more valid than others. The method outlined below has some problems but these can later be discussed and solutions put forward.

Step Description Purpose Compromise to Validity Improvement
1 A sample of fertiliser is weighed and ground using a mortar and pestle Increases the surface area of the fertiliser to assist dissolving Not all of the sample may be transferred from the mortar to the flask Rinse the mortar several times with water
2 Dissolve the crushed fertiliser in 250 mL of warm water Enhances solubility as solids have a higher solubility in warm water Some of the sulfate ions may not dissolve Use a larger volume of hot water
3 Filtration to discard the undissolved fertiliser residue Ensures the filtrate has no unndissolved solids which would affect the validity at Steps 5 and 6 Undissolved sulfate ions from the previous step may be discarded Use a larger volume of hot water to enhance solubility
4 Add excess barium chloride solution to precipitate sulfate ions Precipitates the sulfate ions Some of the sulfate ions may remain unprecipitated and ions other than sulfate may form precipitates with the barium Add concentrated, excess barium chloride
5 Filtration to collect precipitate Separates precipitate from the filtrate Barium sulfate is fine grained and passes through filter paper Use a finer analytical grade filter and a scintered glass funnel
6 Dry and weigh the barium sulfate precipitate Drives off any water that would affect the mass of precipitate Incomplete drying would alter the final mass Ensure drying is complete

 

Sample Procedure and Calculations

  1. Dissolve 1.00 g of ammonium sulfate fertiliser in 250 mL of warm water in a volumetric flask.

  2. Transfer 100mL of this solution to a flask using a pipette.

  3. Acidify with a few drops of 1 mol L-1 HCl and bring to the boil.

  4. Add 1 mol L-1 BaCl2 solution from a burette until no more white precipitate forms.

  5. Filter through weighed filter paper and dry.

  6. Calculate the mass of precipitate and from this the moles and mass of sulfate present in fertiliser (see above).