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The state of our environment is an important issue for society.
Pollution of air, land and water in urban, rural and wilderness areas is a phenomenon that affects the health and survival of all organisms, including humans. An understanding of the chemical processes involved in interactions in the full range of global environments, including atmosphere and hydrosphere, is indispensable to an understanding of how environments behave and change. It is also vital in understanding how technologies, which in part are the result of chemical research, have affected environments. This module encourages discussion of how chemists can assist in reversing or minimising the environmental problems caused by technology and the human demand for products and services. Many chemical occupations will require chemists to monitor and manage the products of reactions in order to ensure safety and efficiency, and even high yields in production processes.
Some modern technologies can facilitate gathering information about chemicals - those occurring in natural environments and those released as a result of human technological activity. Such technologies include systems that have been developed to quantify and compare amounts of substances.
This module increases understanding of the nature, practice, applications and uses of chemistry and the implications of chemistry for society and the environment.
IQ1 How are the ions present in the environment identified and measured?
IQ1 How are the ions present in the environment identified and measured?
8.1 Analysis of inorganic substances
8.1.1 Monitoring the environment.
8.1.2 Investigating aqueous ions.
8.1.3 Gravimetric analysis and precipitation titrations.
8.1.4 Colorimetry, UV-visible spectrophotometry, AAS.
The two main branches of qualitative analysis are:
1. Inorganic analysis: looks at the elemental and ionic composition of a sample, usually by examination of ions in aqueous solution. Example is flame test.
2. Organic analysis: tends to look at types of molecules, functional groups, and chemical bonds. Example is iodine test.
8.1.1 analyse the need for monitoring the environment
All reactions require some degree of monitoring, for safety if no other reason.
However there are certain reactions that will form different products under different conditions. One simple example is the combustion reaction. We know the type of combustion depends on the quantity of oxygen, so in low oxygen concentrations, fuels combust to form carbon or soot, while in high oxygen concentrations, carbon dioxide is formed.
Either way, combustion reactions release unwanted products into the atmosphere.
For reasons such as this we need to monitor our air and our water.
View video:
a) Combustion reactions
- In combustion, fuel-to-oxgyen ratios must be monitored to maximise energy output and efficiency and minimise harmful pollutants:
A plentiful supply of oxygen should be available for complete combustion to occur – the more complete the combustion, the more CO2 is produced, and so more high-energy C=O bonds are formed to release more energy.
eg complete combustion of methane:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
If there is insufficient oxygen, incomplete combustion takes place, forming undesirable products such as carbon monoxide (poisonous) and soot (lung-irritant and carcinogen), which are harmful to human health.
eg incomplete combustion of methane:
CH4(g) + 3/2O2(g) → CO(g) + 2H2O(g) AND CH4(g) + O2(g) → C (s) + 2H2O(g)
CO2 emission must be monitored, so that knowledge about effect on the enhanced greenhouse effect can be determined and data used for control strategies.
b) Potential pollutants of the air
The atmosphere is the gaseous mixture surrounding the surface of the Earth. Excluding water, the atmosphere is composed of the following gases:
Nitrogen 78.08%
Oxygen 20.95%
Argon 0.93%
Carbon Dioxide 0.036%
Other 0.004%
However, the percentage of carbon dioxide in the atmosphere has been increasing over the last 150 years, largely due to the burning of fossil fuels, and is contributing to global warming in what has become known as the ‘Greenhouse Effect’.
The percentage of water vapour in the atmosphere varies considerably, particularly in the lower atmospheric levels, eg the troposphere, where it ranges from 0.5% to 5%. A pollutant is a substance that has a negative effect on the natural environment.
There are many pollutants in the troposphere, most of which are derived from:
- the combustion of fuels in power stations and vehicles
- the smelting and purification of metals
The table below identifies the main pollutants in the lower atmosphere (troposphere) and their source.
c) Potential pollutants of the waterways
There are many factors that influence the concentration of various ions in a natural body of water:
Nature of the rain forming them:
- Minerals present in the pathways from rain to the water body.
- Solubilities of minerals present in these pathways.
- Acidity of rain, as acid rain increases concentrations of certain ions (such as calcium, magnesium and iron) in the water body because it is better able to leach these from its pathway.
- Rate at which rainwater enters the water body (high rate dilutes water).
- Rate of evaporation from the water body.
Human activities:
- Extent of agricultural practices, such as the use of fertilisers, in areas near the water body.
- Extent of discharge of waste from industry into the water body.
- Extent of discharge of effluent into the water body.
- Extent of leaching from rubbish dumps into the water body.
Due to their smaller volumes, rivers and dams are more susceptible than oceans to changes in ion concentrations.
However, ion concentration in coastal ocean water can be significantly affected, particularly by:
- discharge of sewerage
- run-off from rivers that flow through agricultural land.
There are a number of chemically important sources of water pollution and hence reasons for constantly monitoring waterways.
c) Heavy metals
A metal with a relative density of 5.0 or higher is called a heavy metal. Heavy metals have many routes of entry into the body: ingestion, inhalation, absorption by the skin.
Some heavy metals are essential micronutrients (e.g. iron, cobalt and zinc) – but can become toxic in high concentrations. Other heavy metals, such as mercury and lead, are highly poisonous, even in very small concentrations, and can cause various health problems:
- affect the central nervous system (primarily the brain) to cause intellectual impairment in young children (eg decreased concentration span), neurological disorders, and brain damage if unchecked.
- cause anaemia (by inhibiting haemoglobin formation and thus the availability of oxygen to body tissue).
- form a cumulative poison (bio-accumulates) because it builds up in bones and it is difficult to excrete.
Sources of lead are widespread. Examples include:
- Released into the atmosphere from leaded petrol. Monitoring of soils near major roads is required to ensure that people around these areas are not exposed to excessive concentrations.
- From lead compounds in paints. Hazards occur when paint layers from older homes are stripped away.
- Mining and refining, e.g. Mt Isa – risk to children. Car battery manufacture and disposal.
Because lead is dangerous and widespread, its monitoring is essential in areas of heavy traffic, areas providing drinking water and food, and areas producing and using lead.
Testing for heavy metal pollution:
The heavy metal pollution of water can be tested using:
- Chemical testing: sodium sulfide solution (Na2S) added to a highly concentrated acidified or basified water sample (heavy metal ions react with sulfide ions to form sulfide precipitates):
- If a precipitate is formed when the sample is acidified, then one or more of the following is present: lead, silver, mercury, copper, cadmium, arsenic.
- If a precipitate is formed when the sample is basified, then one or more of the following is present: chromium, zinc, iron (III), nickel, cobalt, manganese, aluminium.
2. Atomic absorption spectroscopy: using a light source specific to the heavy metal being tested for.
Eutrophication: The abundant growth of aquatic plants due to nutrient-enriched conditions, in particular, nitrate and phosphate enriched conditions.
The aquatic plants that grow abundantly in eutrophication eventually use up all of the available nutrients that they require and die.
The plants decompose, and in doing so, use up all dissolved oxygen.
After using all oxygen, they decay anaerobically, resulting in chemicals that kill all remaining life.
The decay causes sediment at the bottom of the water body.
The mains sources of nutrients that cause eutrophication are:
Sewerage. Fertiliser.
Nitrate and phosphate are monitored in waterways vulnerable to eutrophication.
The nitrogen-phosphorus ratio of waterways is often monitored, with the EPA recommending a ratio of less than 10:1.
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TASK 8.1.1
PAST HSC QS WHERE NO MARK IS INDICATED, Q IS MULTIPLE CHOICE (CURRENT EXAM HAS 20 MC QS, EARLY YEARS HAD 15).
2016 Q2
2014 Q23B 2
2013 Q23B 2
2010 Q31B 2
2009 Q25C 2
2005 Q25C 2
8.1.2 conduct qualitative investigations using
- flame tests
- precipitation reactions
- complexation reactions
to test for the presence in aqueous solution of:
a) cations:
barium (Ba2+), calcium (Ca2+), magnesium (Mg2+), lead (II)(Pb2+), silver (Ag+), copper(II) (Cu2+), iron(II) (Fe2+), iron(III) (Fe3+)
b) anions:
chloride (Cl−), bromide (Br−), iodide (I−), hydroxide (OH−), acetate (CH3COO−), carbonate (CO32−), sulfate (SO42−), phosphate (PO43-)
Conduct Qualitative Investigations
Conduct Qualitative Investigations
Qualitative analysis is a method used for identification of which ions or compounds are in a sample. Examples of qualitative tests include flame tests, and ion-precipitation reactions (solubility tests), where the identification of ions is achieved by making use of their solubility properties.
To deduce the ions present in a sample from the results of cation and anion analysis tests:
Some points to remember
- When identifying one cation present in a solution, use a new sample for each test.
- When analysing mixtures with several cations, an excess reagent is added until no further precipitate forms.
- The precipitate is separated from the filtrate using filtration or a centrifuge (quicker) before performing the next test. (A centrifuge spins test tubes at high speeds and flings the precipitate to the bottom.)
- Sufficient nitric acid ensures all carbonate is destroyed and does not interfere with other tests (all carbonates react with dilute acids to form CO2 gas)
using flame tests
using flame tests
Flame Tests
Flame tests are a simple test for qualitative analysis.
How: Strong heating from the Bunsen burner may provide sufficient energy for electrons to be promoted from their normal (called ground state) state into higher orbitals (called the excited state). As the electrons return to the ground state (either in one go, or in several jumps), energy is released, sometimes as visible light. Each jump involves a specific amount of energy being released as electromagnetic radiation so each corresponds to a particular wavelength (or frequency). The more electrons in an atom or ion, the more potential jumps and hence a spectrum of multiple lines of different wavelength will be produced. Some of these may correspond to the visible part of the electromagnetic spectrum. The observed colour is a combination of all the individual wavelengths.
Usually in flame tests, we need atoms to produce wavelengths in the visible spectra. Ions will often emit in the ultraviolet region of the spectrum. In the example shown in Figure 4, “a sodium atom in an unexcited state has the structure
1s2 2s2 2p6 3s1,
but within the flame there will be all sorts of excited states of the electrons. Sodium's familiar bright orange-yellow flame colour results from promoted electrons falling back from the 3p1 level to their normal 3s1 level”. (LibreTexts, 2019)
What: As we saw in Module 1, flame tests can be used to identify the presence of a small and distinct number of cations in solution. Flame tests are most useful to distinguish between Group 1 metals, as well as between barium and calcium in Group 2. Some other metals may also be identified or indicated using flame tests, but we can confirm these tests with a series of precipitation reactions.
Pros and cons
- Flame tests can be used to identify the presence of a particular atom/ion but do not provide any information about the concentration of the ion.
- Drawbacks include the closeness (subjectivity to distinguish) of some colours, and the frequent presence of sodium contamination giving false colour
Method
- Dissolve chloride salts in de-ionised water and spray into a Bunsen flame using an atomiser
- Repeat for each metal: barium (Ba2+), calcium (Ca2+), magnesium (Mg2+), lead (II)(Pb2+), silver (Ag+), copper(II) (Cu2+), iron(II) (Fe2+), iron(III) (Fe3+).
You need to know the colours
Why don't we test for lead with a flame test?
using precipitation reactions
using precipitation reactions
Precipitation reactions
Sometimes it is not appropriate to use flame tests to identify cations. Some metals do not produce a distinctive flame colour; others, such as lead, are extremely dangerous if heated in an open flame. In such cases, or even to confirm an ion we have identified, we can use precipitation reactions.
Recall precipitates in Module 3: A precipitate forms when the attraction between two oppositely charged ions in solution is greater than the attraction between the individual ion and the water molecules.
In this case, the oppositely charged ions will bond ionically to each other and precipitate out of the solution as an insoluble solid.
In Module 5, we examined the solubility equilibria established by ions in solution and the solid precipitate.
Recall that: some ionic salts are more easily dissolved in water than others. Some ionic salts can exist in very high concentrations in water while others cannot.
The maximum amount of solute which can dissolve in a given volume of solvent is termed the solubility of the solute.
Temperature affects solubility.
Recalling the solubility rules relating to NAAGSG, LMS and CaStroBar, can help us use our knowledge of precipitation reactions to identify ions in solution.
Depending upon how we carry out precipitation reactions, we can use these either qualitatively (to identify the presence of a particular ion) or quantitatively (by gravimetric analysis).
using complexation reactions
using complexation reactions
Complexation reactions
To understand complexation reactions, we need to understand two new terms: complex and ligand. To simplify matters, a complexation reaction occurs when a complex ion is formed during a chemical reaction.
A complex is a combination of a metal (usually a transition metal) atom or ion and a Lewis base (an electron donor, beyond our course, don't worry about understanding) which has been bonded to it. The complex often has a charge. Usually a complex contains a central atom (eg metal) which is surrounded by ions or molecules. Many complexes also produce coloured ions or precipitates, which makes them useful for identifying ions in solution.
A ligand is an ion or molecule which can link to a metal atom or ion to form a complex. This often occurs through the Lewis base donation of an electron pair to a metal atom or ion. The central metal could be surrounded with 2, 4 or 6 ligands. The most common ligands are CN-, OH- and NH3.
When ammonia is added to silver chloride, the silver ion interacts with the ammonia molecules to produce a complex. In this case, the nitrogen atom from the ammonia molecule donates an electron pair to the silver ion. This results in a complex with a +1 charge.
One complex we have already encountered is iron thiocyanate [FeSCN]2+. This was formed through the reaction of ferric ions (Fe3+) and the thiocyanate ion (SCN-). Once again the metal ion accepts an electron pair from the sulfur atom.
The stability of the complex ion can be determined through the equilibrium constant K for the formation of the complex ion. Kf for Ag(NH3)2]+ is 1.7 x 10 to the power of 7.
Complexation reactions are often useful for quantitative determination of the presence of various ions. A chart is shown below. You may have to explain the use, or the simple chemistry behind a ligand/complex, or to be able to work out the oxidation numbers, but you will not need to know the table.
View videos:
to identify cations
to identify cations
to test for the presence in aqueous solution of:
a) cations:
barium (Ba2+), calcium (Ca2+), magnesium (Mg2+), lead (II)(Pb2+), silver (Ag+), copper(II) (Cu2+), iron(II) (Fe2+), iron(III) (Fe3+)
We can identify the presence of cations through any of the three methods previously mentioned: flame tests, precipitation reactions or complexation reactions. However, there are some tests which are dangerous, eg identifying lead ions using a flame test, and some tests which may be inconclusive, eg calcium and barium both form white precipitates with sulfate ions.
Whichever tests we choose, we need to consider whether the tests are merely to indicate the presence of a particular cation (qualitative) or to measure the mass or concentration of a particular cation (quantitative). This may help us determine which test would be the most appropriate.
It is helpful to organise the information you need to remember about identifying ions in a graphical form. A flowchart and summary table are shown below.
Flowcharts are an excellent tool to use when developing a set of tests to eliminate, or positively identify, a particular cation. Which of the flowcharts shown below would be the most useful one for you to study? Could you develop something completely different?
Read this flowchart by using solubility rules eg
- why must the sample be lead (Pb2+) if it precipitates with HCl? (ie which chlorides are insoluble?)
- why is it Ca2+ or Ba2+ if it reacts with sulfuric acid? (ie which sulfates are insoluble)?
Experiment 8.1.2: Testing for Cations
Introduction
Qualitative analysis allows the chemist to determine what substances are present in a sample. In this experiment you will perform precipitation and complex ion formation to confirm the presence of a single ion in an aqueous solution.
This series of tests will identify the presence of a cation by the formation of a coloured precipitate or coloured complex ion. Spot tests may only useful if the sample contains one ‘target’ ion. Samples containing several ions may cause interference, making the test invalid.
Method
1. Spot test for barium ion: Add 1 mL of 0.1 mol L–1 sodium sulfate solution to 1 mL of a solution of 0.1 M barium nitrate in a well. Record your observations.
2. Spot test for copper(II) ion: Add 1 mL of 0.1 mol L–1 sodium hydroxide solution to 1 mL of a solution of 0.1 M copper(II) sulfate in a well. Record your observations.
3. Spot test for lead(II) ion: Add 1 mL of 0.1 mol L–1 sodium hydroxide solution to 1 mL of 0.1 mol L–1 lead nitrate in a well. Observe and record your observations. Continue adding sodium hydroxide until a change occurs. Record your observations.
4. Spot test for iron(II) ion: Add 1 mL of 0.1 mol L–1 sodium hydroxide solution to 1 mL of 0.1 mol L–1 iron(II) sulfate solution in a well. Observe and record your observations.
5. Spot test for iron(III) ion: Add 1 mL of 0.1 mol L–1 sodium hydroxide solution to 1 mL of 0.1 mol L–1 iron(III) chloride solution in a well. Observe and record your observations.
6. Spot test for calcium ion: Add 1 mL of 0.1 mol L–1 sodium sulfate solution to 1 mL of 0.1 mol L–1 calcium chloride solution in a well. Observe and record your observations.
7. Spot test for magnesium ion: Add 1 mL of 0.1 mol L–1 sodium sulfate solution to 1 mL of 0.1 mol L–1 magnesium chloride solution in a well. Observe and record your observations.
8. Spot test for silver(I) ion: Add 1 mL of 0.1 mol L–1 sodium hydroxide solution to 1 mL of 0.1 mol L–1 silver(I) nitrate solution in a well. Observe and record your observations.
VIEW Videos:
- ACI#3 https://www.youtube.com/watch?v=MytcGBLrmno&list=PLeFSFSJ9WqSC36UhgaA0FsiqNtiSo8I6H&index=3 [14.59]
- Testing for cations https://www.youtube.com/watch?v=r8-0I2DgEWA&index=20&t=0s&list=PLE4082392972BB7ED [16.05]
- Testing for positive ions summary https://www.youtube.com/watch?v=D-yDQSxI0_M [3.33]
Complete the summary table below to assist your study:
to identify anions
to identify anions
to test for the presence in aqueous solution of:
b) anions:
chloride (Cl−), bromide (Br−), iodide (I−), hydroxide (OH−), acetate (CH3COO−), carbonate (CO32−), sulfate (SO42−), phosphate (PO43-)
TESTING FOR ANIONS
VIEW Videos:
TASK 8.1.2A
Answer the questions below
TASK 8.1.2B
1. Which of the following is an example of a qualitative analysis? (Chan, et al., 2019, p. 423) a) Finding the amount of salt in a jar of peanut butter b) Detecting the presence of arsenic in a water supply c) Measuring the fat content in milk d) Finding the NaOH concentration in oven cleaner
2. Explain how the following data from a flame test can be used to prove that it is the metal ion that is responsible for the flame colour a salt produces. (Chan, et al., 2019, p. 423) (2 marks) Barium sulfate (green), potassium sulfate (lilac), barium chloride (green).
3. Cobalt(II) ions (Co2+) can form a complex ion with four chloride ions (Cl-). Predict the chemical formula and charge of this complex ion. (Chan, et al., 2019, p. 423) (2 marks)
4. Explain why all atoms of sodium will emit the same set of wavelengths of light when heated. (Davis, Disney, & Smith, 2018, p. 419) (2 marks)
TASK 1.2.2C
3.1 PAST HSC QS WHERE NO MARK IS INDICATED, Q IS MULTIPLE CHOICE (CURRENT EXAM HAS 20 MC QS, EARLY YEARS HAD 15).
2017 Q2, Q8
2016 Q13
2015 Q3
2014 Q24 5
2013 Q22C 2
2012 Q10
2011 Q10; Q28 4
2010 Q10; Q31A11 2
2009 Q4; Q8; 25A 2; 26D 1
2007 Q17A 1; 17B 2; 17C 1
2006 Q26A 1; 26B 3
2005 Q25B 2
2004 Q3; Q21C 3
1.3 conduct investigations and/or process data involving
a) gravimetric analysis.
Gravimetric analysis is a precise analytical technique which is used to determine the proportion by mass of a particular chemical substance in a compound or mixture.
Australian soils are very poor in terms of their nitrogen content and one of the most common fertilisers used to increase the nitrogen content is sulfate of ammonia (ammonium sulfate). To determine the amount of ammonium ions in a bag of fertiliser is difficult in the laboratory as ammonium ions are very soluble. However, we can use gravimetric analysis to determine the proportion of sulfate ions in a sample of fertiliser by precipitating it out as barium sulfate. From the mass of the isolated precipitate, it is possible to calculate the % composition of the fertiliser in terms of both sulfate and ammonium ions.
- Gravimetric Analysis: A method of quantitative analysis that involves the precipitation of a highly insoluble compound followed by weighing of the dried precipitate formed.
- Gravimetric analysis can be used to determine the sulfate content of a lawn fertiliser. Barium sulfate is a highly insoluble sulfate compound, and can be formed by adding barium chloride to sulfate in solution:
- The sulfate content of a lawn fertiliser can be determined using the following procedure (need to know this):
- Grind a sample of fertiliser into a fine powder using a mortar and pestle.
- Place a certain mass (around 1 g) of the ground fertiliser into a certain volume (around 100 mL) of a known concentration (around 0.1 M) of dilute hydrochloric acid.
- Stir to dissolve as much solid a possible.
- Filter off insoluble material.
- Slowly add excess barium chloride solution to the filtrate to form barium sulfate precipitate.
- Weigh a dry filter paper, and then filter off the precipitate.
- Dry the filter paper and then weigh again to determine the mass of the precipitate.
- Determine the mass of sulfate:
- Determine the percentage of sulfate in the fertiliser:
PRAC 8.1.3a Sulfate in fertiliser
Planning
1 List the glassware required for this experiment.
2 How big should the sample of solid ammonium sulfate be?
► Consider experimental error, balance sensitivity and filtration duration.
3 How much water should the ammonium sulfate be dissolved in?
► Consider filtration duration.
4 List the precision weighings you will need to make during this analysis.
Method
1 Warning: Wear safety glasses!
2 Warning: 1 mol L–1 HCl is corrosive. Avoid contact. Clean up any spills.
3 Warning: 0.2 mol L–1 BaCl2 is toxic. Avoid contact. Clean up any spills.
4 Weigh out about 0.5 g of ammonium sulfate. Record the exact mass.
5 Add 10 mL of water to the (NH4)2SO4 sample and dissolve. Add 2 mL of 1 mol L–1 HCl and 100 mL distilled water. Heat on a hot plate to near-boiling.
6 Add 25 mL of 0.20 mol L–1 BaCl2 solution dropwise with stirring. Heat (‘digest’) to near-boiling for 10 minutes. Add dropwise, with stirring, 10 mL of 0.1% agar solution.
7 Filter the mixture using a pre-weighed filter paper. Test the filtrate for complete precipitation with BaCl2.
► How can complete precipitation be judged visually?
8 Wash the precipitate three times using a wash bottle. Dry the filter paper and residue to constant weight. Record the final weight.
9 Set up a table to record all relevant experimental data.
Calculations and discussion
10 Calculate the exact moles of ammonium sulfate in your sample.
11 Calculate the exact volume of 0.20 mol L–1 BaCl2 required for complete precipitation of BaSO4 for your sample.
► Compare it with the volume used in procedure Step 6.
12 Calculate the mass of BaSO4 recovered and the mass of SO4 in the precipitate.
13 Calculate the experimental SO4% in (NH4)2SO4:
= mass of SO4 (ppt) ÷ mass of (NH4)2SO4 (sample) 100
14 Calculate the theoretical SO4% in (NH4)2SO4.
15 Calculate the experimental percentage error and comment on it.
16 Why is a slight excess of BaCl2 used for precipitation?
17 Why must the barium sulfate precipitate be rinsed/washed before drying?
18 Describe how potential filtration problems were overcome.
19 Comment on the validity of your results.
20 How could the reliability of the results be increased?
PRAC 8.1.3b Sulfate in fertiliser: Alternate
1. Grind up a sample of fertiliser into a fine powder using a mortar and pestle.
2. Weigh out 1g of the powder using an electronic balance.
3. Dissolve the powder in 100 mL of distilled water in a 250 mL beaker using a stirring rod.
4. Filter the solution in a clean 250 mL beaker using filter paper.
5. Add 25 mL of 1M HCl to remove carbonates (acid plus carbonate forms CO2 and H2O).
6. Heat the mixture just below boiling using a hotplate.
7. Slowly add 20 mL or an excess amount of 20% BaCl2 solution with constant stirring.
8. Digest and stir the mixture for 30 minutes.
9. Allow the beaker to cool and then rest in ice.
10.Filter the solution through a sintered glass funnel.
11.Wash the precipitate using a wash bottle, once with deionised water and again with acetone (acetone is a drying agent, it will evaporate away and not add to the mass).
12.Dry the precipitate on the sintered glass funnel in an oven.
13.Cool it in a desiccator (drying container).
14.Weigh and record the funnel and precipitate.
15.Repeat Step 12 – 14 until a constant mass is reached.
You do not need to memorise this process in detail as they would never ask you to recite one, rather it is important to understand the purpose of each step.
Main Problems Encountered and How it was Solved
- Ensuring no carbonate is present. carbonate may precipitate with Ba2+ so HCl was added to acidify the solution and remove any carbonate ions that could be present, prevent the precipitation of barium salts with carbonate.
- Ensuring complete precipitation of sulfate. The complete precipitation of sulfate was ensured by adding excess BaCl2.
- Filtering the very fine BaSO4 precipitate. This was overcome by use of gravimetric grade filter paper
- Ensuring the BaSO4 is free of water, Cl ions etc. This was prevented by undertaking many rinsings when transferring the solution the filter funnel (quantitative transfer), washed precipitate with distilled water to remove chloride ions, placing the precipitate in an oven and drying to constant mass.
Further Solutions to Problems Encountered
- Use excess precipitating reagent to ensure all sulfate is precipitated.
- Use a sintered glass filter (sintering glass crucible under vacuum) instead of fine filter paper, which is of bigger pore size. Its very small pore size prevent very fine particles passing through as part of the filtrate. This would mean that agar-agar solution would not be needed, thus increasing validity.
- Use a greater mass of solid, so that weight of collected is greater, reducing the % error.
- Use a more accurate electronic balance that measures up to at least 5 significant figures.
VIEW videos:
TASK 8.1.3A
b) It was found that 4.25 g had a sulfate content of 35%.
What is the mass of the dried precipitate at Step 4? Include a chemical equation in your answer. (3 marks)
2. (08, Q15, 1 mark)
A 2.45 g sample of lawn fertiliser was analysed for its sulfate content. After filtration and drying, 2.18 g of barium sulfate was recovered. What is the %w/w of sulfate in the lawn fertiliser?
3. (03, Q27, 5 marks)
A student carried out an investigation to analyse the sulfate content of lawn fertiliser.
The student weighed out 1.0 g of fertiliser and dissolved it in water. 50 mL of 0.25 mol.L−1 barium chloride solution was then added. A white precipitate of barium sulfate formed, which weighed 1.8 g.
(a) Calculate the percentage by mass of sulfate in the fertiliser. (2 marks)
(b) Evaluate the reliability of the experimental procedure used. (3 marks)
4. Four students analysed a sample of fertiliser to determine its percentage of sulfate.
Each student:
• weighed an amount of fertiliser;
• dissolved this amount in 100 mL of water;
• added aqueous barium nitrate;
• filtered, dried and weighed the barium sulfate precipitate.
Their results and calculations are shown in the table.
The percentage of sulfate calculated by Student C was significantly higher than that of the other students. Which is the most likely reason for this?
A) Student C did not dry the sample for long enough.
B) Student C added more Ba(NO3)2 solution than the other students.
C) Student C used a balance capable of measuring weight to more decimal places.
D) Student C waited longer than the other students for the Ba(NO3)2 to react completely with the sulfate.
1.3 conduct investigations and/or process data involving
b) precipitation titrations.
Volumetric analysis using the technique of titration is one with which we are already familiar. The key to any titration is to identify the end point of the titration so you know when to stop adding the titrant.
The difficulty associated with identifying ion concentration using precipitation is the often invalid process of adding excess reagent and then obtaining an inaccurate mass of the dried precipitate through filtration. It can be hard to isolate the particular ion in the precipitate or to collect all of the ions in solution through precipitation.
An alternative quantitative measure to help resolve the challenges in these processes is the formation of a precipitate during a titration. As with previous acid-base titrations, recording the volume of solution needed to reach an end point can provide data for the calculation of the concentration of particular ions in a solution.
The challenge associated with precipitation titrations is determining the end point. This cannot be determined by knowing when there is no more precipitate forming so we need an extra reaction. Usually, when conducting a precipitation titration, we select a reagent which will precipitate the desired ion, but which will also form a coloured complex with another ion once the target ion has been removed from the solution.
If we wanted to identify the chloride ion concentration in a water sample, we could add silver nitrate. This would precipitate silver chloride.
Ag+(aq) + Cl-(aq) → AgCl(s)
However this is a white precipitate and we would not know when all the chloride was out of the solution. If we add potassium chromate to our sample (the same way we added an indicator for our acid-base titrations) then once the silver ions have precipitated all the chloride ions they will start to form a precipitate with the chromate ions.
Ag+(aq) + CrO42--(aq) → Ag2CrO4(s)
Silver chromate has a brick-red colour so it is easy to determine when it starts to form. We could use an alternative indicator which formed a coloured complex once the target ion was removed. For any of these techniques to work, we need to ensure that our target ion is less soluble than our indicator ion. If silver chromate were less soluble than silver chloride, it would precipitate first making our procedure invalid.
PRAC 8.1.3B Precipitation Titration
Introduction
In Module 1, we investigated the salt content of a salt/sand/water mix using an inquiry method based on gravimetric analysis. If we consider the primary anion in beach water to be chloride (and this is a big assumption), we can use the technique of precipitation titration to determine the chloride ion concentration and hence the sodium chloride concentration in a water sample.
Due to the difficulty in determining an endpoint from precipitate formation we will use potassium chromate solution as an indicator of the end point. Once the chloride ions have precipitated out of the solution, a red silver chromate precipitate will form. This indicates the end point of the titration.
The method below was adapted from Pearson Chemistry 12 Skills and Assessment (Commons, 2018, p. 195).
Method
1. Collect a sample of beach water, or other water sample.
2. Evaporate the water using distillation apparatus to ensure only the salt remains.
3. Cool, dry and weigh this residue.
4. Separate the residue into 3 equal, weighed samples (if there is sufficient mass) and add them to separate clean and rinsed conical flasks.
5. Add sufficient distilled water to dissolve the samples.
6. NB It may be necessary to add some sodium bicarbonate to each flask to remove any acid which may interfere with the titration. This can be tested with a small quantity to determine if there is any effervescence.
7. Add 2 mL of 0.01M K2CrO4 solution to each flask.
8. Add 0.1 M AgNO3 solution to the burette and titrate the silver nitrate against each sample in the flasks until the end point is reached.
9. Record the volume of each titrant.
Discuss validity and reliability of this experiment.
TASK 8.1.3B
All natural waters contain some chloride ions. If there is a high concentration of these ions, the water may be brackish or saline and many freshwater species of plants and animals will die. High saline levels in irrigation water are causing serious problems in many parts of Australia.
A farmer thought a competitor was poisoning his crops by adding salt to the soil. He took a sample of soil to a forensic chemist and requested the soil be tested to determine whether a higher than normal concentration of chloride ions was present. The chemist dried and weighed the soil. Then 10 g of dry soil was weighed out and 100mL distilled water was added to the dry soil. The solution was filtered and the solid component discarded. The sample was then tested for the presence of chloride ions.
Confirming these were present, the concentration of chloride ions was found by titrating a sample with a silver nitrate solution of known concentration. The end point of the titration is shown by another silver compound – silver chromate (Ag2CrO4). This compound is red and appears when there is an excess of silver ions.
When 10.0mL of a water sample was titrated against a 0.100molL−1 silver nitrate solution, the volume of silver nitrate solution needed to reach the end point was 7.70mL.
a) Calculate the concentration of chloride ions in the soil sample in molL−1 and Cl− in g /g soil. (3 marks)
b) Which precipitation method was used in this titration? Explain why you identified this method. (2 marks)
c) Consider the method you identified and describe any changes that could be made to improve the accuracy of the result obtained. (2 marks)
2. A group of students was given an assignment in which they had to determine the concentration of Ba2+ ions in an unknown sample of Ba(OH)2. They decided to use two different analysis processes and compare the results to see if one process was more accurate than the other. They chose to use gravimetric analysis as one process and a precipitation titration as the other.
Following extensive research to determine a suitable end point for the titration, they decided to use conductivity to determine the endpoint. In both analytical processes, they used the same reactants − 0.100molL−1 sulfuric acid (H2SO4) and the unknown Ba(OH)2 solution.
The data table below gives the change in conductivity as sulfuric acid is added to the solution.
a) Plot points for conductivity vs volume of H2SO4 solution. (3 marks)
b) Draw two lines of best fit. (2 marks)
c) Record the volume of H2SO4 where the two lines meet. This is the equivalence point. (1 mark)
d) Calculate the moles of H2SO4. (3 marks)
e) Write the balanced equation for the reaction. (1 mark)
f) Calculate the moles and concentration of Ba(OH)2. (3 marks)
g) After submitting their results, the students were told the concentration of the Ba(OH)2 solution was 0.100mol L−1. Evaluate the accuracy of the conductance method for precipitation titrations from these results. (3 marks)
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ANSWERS AT BOTTOM OF PAGE
1.4 conduct investigations and/or process data to determine the concentration of coloured species and/or metal ions in aqueous solution, including but not limited to the use of:
a) colourimetry.
Colourimetry is an analytical technique used to quantitatively measure the amount of light absorbed by coloured solutions containing certain cations.
Colourimetry requires a light source, a filter which is set to a particular wavelength of light, a sample cell (usually a small rectangular prism) and a light detector. The intensity of light of the given wavelength which reaches the detector is calibrated using a cell which contains pure water and then with the cell containing the solution which will be analysed.
Colorimeters calculate a quantity called the absorbance which is a comparison between the intensities both with and without the sample in the path. The greater the absorbance, the greater the concentration of ions in the coloured solution. (Smith, 2004)
Figure 10 The colour intensity is related to the concentration of ions. (Chan, et al., 2019, p. 436)
To accurately determine the concentration of an unknown solution, we use a calibration curve. This is a graphical representation of the relationship between absorbance and concentration which usually resolves as a straight line of best fit. From this line we can calculate the concentration of an unknown solution by interpolating its absorbance to find its concentration.
PRAC 1.4.1 Determining the strength of copper sulfate solution using colourimetry
https://www.youtube.com/watch?v=mbUSTw1oP0E
Method
1. Warning: Wear safety glasses!
2. Warning: 1 mol L–1 CuSO4 is harmful. Avoid contact. Harmful if swallowed. Causes serious eye irritation. Causes skin irritation. Very toxic to aquatic life, with long-lasting effects.
3. Collect a 30mL sample of 1M copper sulfate solution.
4. Fill a sample tube (cuvette) with this solution.
5. Use a colourimeter to record the absorbance of this 1M solution.
6. Pour 7.5 mL of the copper sulfate solution into a 10 mL measuring cylinder and add 2.5 mL of distilled water.
7. Fill a sample tube with this solution.
8. Use a colourimeter to record the absorbance of this 0.75M solution.
9. Pour 5 mL of the copper sulfate solution into a 10 mL measuring cylinder and add 5 mL of distilled water.
10. Fill a sample tube with this solution.
11. Use a colourimeter to record the absorbance of this 0.5M solution.
12. Pour 2.5 mL of the copper sulfate solution into a 10 mL measuring cylinder and add 7.5 mL of distilled water.
13. Fill a sample tube with this solution.
14. Use a colourimeter to record the absorbance of this 0.25M solution.
15. Use the colourimeter to record the absorbance of the copper sulfate solution of unknown concentration.
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TASK 8.1.4A
1. A solution of copper(II) tetraamine, [Cu(NH3)4(H2O)2]2+ was analysed using colourimetry to determine its concentration. It was diluted by a factor of 100 and found to have an absorbance reading of 0.6.
a) Use the figure below to determine the concentration of the
diluted sample in ppm. (1 mark)
b) What is the concentration of the original undiluted sample in ppm? (2 marks)
c) Suggest a reason for diluting the solution. (1 mark)
1.4 conduct investigations and/or process data to determine the concentration of coloured species and/or metal ions in aqueous solution, including but not limited to the use of:
b) ultraviolet-visible spectrophotometry.
Spectroscopy is the study of the interaction between electromagnetic radiation (emr) and matter.
When vaporised, different elements absorb light of specific frequencies.
The type of emr absorbed by matter affects substances in different ways. When electrons absorb energy, the higher electron state can take the form of an increased rotational, vibrational or electronic excitation. By studying this change in energy state, scientists are able to learn more about the physical and chemical properties of the molecules.
Radio waves can cause nuclei in some atoms to change magnetic orientation and this forms the basis of a technique called nuclear magnetic resonance (NMR) spectroscopy.
Molecular rotations are excited by microwaves.
Vibrations of bonds are excited by infrared radiation.
Electrons are promoted to higher orbitals by ultraviolet or visible light.
(Royal Society of Chemistry, 2017, p. 2)
Absorption of ultraviolet or visible radiation, as we have seen from flame tests, is associated with the movement of electrons from a ground state to an excited state. The radiation needs to have a precise amount of energy to cause a particular transition.
Absorption bands are characterised by their:
- wavelength of maximum absorption
- intensity of absorption compared to a baseline
Flame tests are effective in ion identification because some elements have electrons which transition within the visible part of the electromagnetic spectrum. However many ions are colourless and do not show any results with flame tests. Other substances may be inconclusive or absorb and emit wavelengths beyond the visible spectra. The UV part of the em spectrum may be useful in identifying some of these ions. UV-Vis is also useful for identifying carbon compounds.
The key to these processes is to find a limited number of wavelengths which act as a fingerprint, linked to a particular element so it can be readily identified. It is often helpful to scan a sample first to help in identifying the best wavelength to choose. Some wavelengths may show much higher absorption than other wavelengths, so it is important that these become the focus of the analysis.
In UV-Visible spectroscopy, the best wavelength for analysis of a sample is where maximum absorbance occurs without interference from other components in the sample. This is not always at the wavelength at which the sample absorbs most strongly. (Chan, et al., 2019, p. 441)
UV-Vis spectrometry can also be useful in identifying transitional metal complexes. For example, in the octahedral copper complex, [Cu(H2O)6]2+, yellow light has sufficient energy to promote the d electron in the lower energy level to the higher one.
The UV-Vis spectrometer (simplified image above)
The UV-Vis spectrometer (simplified image above)
This needs a light source which gives the entire visible spectrum plus the near ultra-violet to cover the range from about 200 nm to about 800 nm. (This extends slightly into the near infra-red as well.)
Note: "Near UV" and "near IR" simply means the parts of the UV and IR spectra which are close to the visible spectrum.
You can't get this range of wavelengths from a single lamp, and so a combination of two is used - a deuterium lamp for the UV part of the spectrum, and a tungsten / halogen lamp for the visible part. A deuterium lamp contains deuterium gas [(D, or 2H), also called heavy hydrogen, isotope of hydrogen with a nucleus consisting of one proton and one neutron, which is double the mass of the nucleus of ordinary hydrogen (one proton)] under low pressure subjected to a high voltage. It produces a continuous spectrum in the part of the UV spectrum we are interested in.
The combined output of these two bulbs is focussed on to a diffraction grating.
The diffraction grating and the slit
The diffraction grating splits light into its component colours.
The various wavelengths of the light are sent off in different directions. The slit only allows light of a very narrow range of wavelengths through into the rest of the spectrometer.
By gradually rotating the diffraction grating, you can allow light from the whole spectrum (a tiny part of the range at a time) through into the rest of the instrument.
The sample and reference cells
These are small rectangular glass or quartz containers. They are often designed so that the light beam travels a distance of 1 cm through the contents.
The sample cell contains a solution of the substance you are testing - usually very dilute.
The reference cell contains the pure solvent, chosen so that it doesn't absorb any significant amount of light in the wavelength range 200 - 800 nm.
The detector and computer
The detector converts the incoming light into a current. The higher the current, the greater the intensity of the light.
For each wavelength of light passing through the spectrometer, the intensity of the light passing through the reference cell is measured. This is usually referred to as Io - that's I for Intensity.
The intensity of the light passing through the sample cell is also measured for that wavelength - given the symbol, I.
If I is less than Io, then obviously the sample has absorbed some of the light. A simple bit of maths is then done in the computer to convert this into something called the absorbance of the sample - given the symbol, A.
[Maths following: On most of the diagrams, the absorbance ranges from 0 to 1, but it can go higher than that.
An absorbance of 0 at some wavelength means that no light of that particular wavelength has been absorbed. The intensities of the sample and reference beam are both the same, so the ratio Io/I is 1. Log10 of 1 is zero.
An absorbance of 1 happens when 90% of the light at that wavelength has been absorbed - which means that the intensity is 10% of what it would otherwise be.
In that case, Io/I is 100/I0 (=10) and log10 of 10 is 1.
If you don't feel comfortable with logarithms, don't worry about it. Just accept that an absorbance scale often runs from zero to 1, but could go higher than that in extreme cases (in other words where more than 90% of a wavelength of light is absorbed).]
The chart recorder
Chart recorders usually plot absorbance against wavelength. The output might look like the one below.
This particular substance has what are known as absorbance peaks at 255 and 395 nm.
UV-Vis spectroscopy can be used to make quantitative determinations of concentration by use of a calibration curve (below right).
To make a calibration curve, at least three concentrations of the compound will be needed, but five concentrations would be better for a more accurate curve.
The concentrations should start at just above the estimated concentration of the unknown sample and should go down to about an order of magnitude lower than the highest concentration.
The calibration solution concentrations should be spaced relatively equally apart, and they should be made as accurately as possible using digital pipettes and volumetric flasks instead of graduated cylinders and beakers.
An example of absorbance spectra of calibration solutions of Rose Bengal (4,5,6,7-tetrachloro-2',4',5',7'-tetraiodofluorescein) can be seen below (left). To make a calibration curve, the value for the absorbances of each of the spectral curves at the highest absorbing wavelength, is plotted in a graph (below right).
The concentration of the unknown is then determined by applying its absorbance data from the spectrometer to the graph. In the above example, if the absorbance is 1.5, the concentration will be 17 micromoles/Litre.
Limitations of UV-vis Spectroscopy
UV-vis spectroscopy works well on liquids and solutions, but if the sample is more of a suspension of solid particles in liquid, the sample will scatter the light more than absorb the light and the data will be very skewed. Most UV-vis instruments can analyse solid samples or suspensions with a diffraction apparatus, but this is not common. UV-vis instruments generally analyse liquids and solutions most efficiently.
While the Beer-Lambert Law is not explicitly mentioned anywhere in the syllabus, it is on the data sheet, so it may be wise to look at it. It provides some background to the relationship between wavelength, absorbance and concentration for different types of spectrometry.
1. For a particular frequency or wavelength of light passing through a sample, we record a particular intensity I. We do likewise for a reference cell, Io. If I < Io then some of the light has been absorbed by the sample.
2. Assumption 1: absorbance is directly proportional to the concentration of the solution in the sample (A α C).
3. Assumption 2: absorbance is directly proportional to the length of the path of light (l) which is the width of the sample tube or cuvette (A α l)
4. Since both assumptions involve absorbance, we can combine them into a single equation (A α lC)
5. We can remove the proportion symbol by replacing it with a constant of proportionality; in this case ε, molar absorptivity or molar extinction constant, so (A = εlC) This is known as the Beer-Lambert Law
6. We can also express absorbance in terms of intensities, where A = log10 (𝐼o / 𝐼) so the Beer-Lambert Law can also be expressed as
A = εCl = log10 (𝐼o / 𝐼). (LibreTexts, 2019)
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TASK 8.1.4B
1. UV-visible spectroscopy was used to measure the spectra of two solutions, A and B.
Solution A was a pink colour, while Solution B was a green colour.
The analyst recorded the absorbance of each solution over a range of wavelengths on the same axes. The resultant absorbance spectrum is shown below.
Explain why Solution A is pink and Solution B is green. (2 marks)
b) From the curve, determine the concentration of cadmium in the paint sample. (1 mark)
1.4 conduct investigations and/or process data to determine the concentration of coloured species and/or metal ions in aqueous solution, including but not limited to the use of:
c) atomic absorption spectroscopy.
THIS SECTION IS VERY IMPORTANT.
LOOK AR 2019 HSC EXAM and make sure you can answer the AAS question there.
Another important application of spectroscopy is atomic absorption spectroscopy, AAS. Atomic absorption spectroscopy (AAS) is a technique used to identify the presence and concentration of substances by analysing the spectrum produced when the substance is vaporised and atomised and absorbs certain frequencies of light.
Effectiveness of AAS in Pollution control
- Atomic absorption spectroscopy (AAS) is an important technique in measuring the concentrations of metal ions in very minute quantities. A liquid sample containing the lead ion is aspirated through a tube into a flame hot enough to vaporise the sample into atoms. A cathode lamp transmits light with the desired frequency through the atomised sample. A detector measures the amount passing through the flame and gives out the absorbance (amount absorbed) reading.
- The basis of AAS is the result of the electron structure of the analysed atom. Electrons move to higher energy levels by absorbing electromagnetic radiation of a specific frequency. Since lead has a unique set of electron energy levels, it absorbs the radiation with a specific frequency. The greater the concentration of the lead ion, the more radiation is absorbed and the less reaches the detector. The amount of light absorbed is proportional to the amount of lead present (according to the Beer-Lambert Law). A calibration (standard) graph using solutions of known concentrations allows the concentration of the unknown to be determined.
- Even very low blood levels of lead are a cause for concern. (10 mg/dL or above-however, lead can impair normal development even below 10 mg/dL).
- AAS is sensitive enough to be able to detect levels at these orders of magnitude (of ppm/ppb).
- AAS analysis has many other advantages:
- lts ease of operation and reliability/precision
- It is quick, as sample is simply aspirated into the atomiser and analysed to standards.
- It is relatively cheap to operate.
- It is a more accurate method than the use of gravimetric/volumetric analyses & colorimetry
- Its high degree of freedom from interference & its specificity. No two elements absorb at the same wavelength-the presence of another element does not directly interfere with the absorption of radiation by lead.
- The ability of the atom to absorb is independent of atomiser temperature so there is no interference from temperature.
- Many lead samples can be analysed quickly (to assist with reliability).
- Minimal contact with toxic lead (compare with gravimetric analysis).
- There are some disadvantages.
- Daily calibration with the lead ion is required for the machine to measure accurately.
- At high concentrations Beer-Lambert’s law (absorption is directly proportional to concentration) fails and precision can be compromised.
- Another problem in atomic absorption is that it is very useful for analysing liquid samples, but has not been successfully used for the direct analysis of solids or gases (the difficulty arises in the atomizer stage).
- Even though the actual AAS measurements are accurate and precise, error can occur at the dilution stage when standards are prepared.
- AAS is an easy and reliable method for determining lead concentrations once standards are prepared carefully. Many samples and effective dilutions can ensure that the linear portion of the Absorbance versus Concentration standard curve is used. Solid and air samples can still be analysed via acid digestion/making solutions so AAS sensitivity, specificity and safe ease of operation are ideal factors for analysing lead which at even low levels can be dangerous.
VIEW Videos:
- AAS How it works: https://www.youtube.com/watch?v=X-vJM62SSvc
- CMM 10 AAS: https://www.youtube.com/watch?v=xh2r7hwSxfI&list=PLeFSFSJ9WqSDgGKuJGcsoZrC359zCmh2o&t=480s&index=3 14.34
- AAS https://www.youtube.com/watch?v=j68xgp64L40&index=24&t=0s&list=PLE4082392972BB7ED 20.16
- Flame AAS https://www.youtube.com/watch?v=_KZjb9G3hB8&index=25&t=0s&list=PLE4082392972BB7ED 1.39
- Effect of trace minerals https://www.youtube.com/watch?v=D0xr6rXZSJg&index=26&t=0s&list=PLE4082392972BB7ED 10.35
- Monitoring lead pollution https://www.youtube.com/watch?v=yWsPngLExwk&index=27&t=0s&list=PLE4082392972BB7ED 11.01
- AAS data interpretation https://www.youtube.com/watch?v=b4H0QPwWeDI&index=28&t=0s&list=PLE4082392972BB7ED 11.17
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