Doc Brown Chemistry Practicals 5u6jh
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PART 1 INTRODUCTION to CHEMICAL TESTS and ANALYSING SUBSTANCES Of what use is chemical analysis and chemical testing? Well, rather a lot, as it happens! The production of many products we use in our society involved some kind of chemical analysis at some stage or other. It might analysing mineral ores for their metal content or testing the final purity of some drug/medicine from the pharmaceutical industry.
Environmental agencies monitor levels of oxygen and pollutants in rives and lakes. Hospitals carry out complex blood analysis for iron, proteins, cholesterol etc. and this is very important diagnostic information for clinicians and doctors seeking to get you well again prevent diseases developing. Blood can be analysed for specific proteins to indicate particular medical conditions e.g. cancer, and ions such as sodium, chloride and iron compounds. At water treatment plants test are done to determine the levels of metal ions, insecticides and lots of other substances to check that their levels are not high enough to be harmful to humans. Apart from a multitude of forensic tests for DNA, powder burns from guns, explosives etc. the police use breathalyser kits to test for alcohol levels in your breath and may request a blood sample for analysis too. There are two types of chemical analysis Qualitative tests and quantitative analysis. Whatever the nature of the investigation, all tests of analyses should be carried out by using 'tried and tested' standard procedures. They should be the most accurate, reliable and safest methods that can be devised. It means, whatever laboratory you work in, anywhere in the world, you should get the same results as long as the sample is collected, stored and analysed by the same standard procedures.
(a) QUALITATIVE CHEMICAL ANALYSIS Qualitative chemical analysis indicates whether a particular substance is present or not. It does not tell how much of the substance is there or its concentration. However, if a substance is potentially harmful, even toxic, its a good idea to know whether the substance is there at all. The larger the sample you have, the better. With more to work with you are more likely to able to detect minute traces of substances with qualitative tests AND have spare material to repeat tests several times if the results seem uncertain at first. At school/college level, the simple tests you learn enable to identify the cation (+ve ion e.g. metal ions) and anion (-ve ion, e.g. chloride, sulfate) in a salt, and the salt usually does contain only two ions, but some salts do have three ions e.g. iron(II) ammonium sulfate which might take a bit of sorting out. Most tests at this level are done on soluble substances so that you can dissolve the substance in water and carry out tests on the aqueous solution. Aqueous means a
solution in water. You may come across a non-aqueous solution using a non-aqueous solvent like alcohol or hexane. AND don't forget, negative test results are just as important as positive results, you may need to eliminate possibilities as well as confirm the presence of a particular ion or gas etc. UNFORTUNATELY, not all tests are unique for a particular ion, but this shouldn't be a problem in school chemistry! Associated qualitative analysis links See index at top of page Typical qualitative tests are described in Parts 2 to 5. Summary of some cation and anion tests for GCSE/IGCSE/O Level students
(b) QUANTITATIVE CHEMICAL ANALYSIS Quantitative analysis gives you (hopefully) a precise measure of how much of a substance is present or its concentration in a sample being analysed e.g. ore analysis tells you whether it is worth exploiting for a metal, the purity of drug ensure no harmful impurities in it, blood sample analysis for alcohol allows the police to decide to prosecute for 'being over the limit'. With a large sample you have spare material to repeat the quantitative analysis several times to get the most statistically valid result. In schools and colleges you can do quite accurate titrations to illustrate quantitative analysis. Associated quantitative links % purity of a product (GCSE/IGCSE/O level, introduction for Advanced Level) Volumetric titration analysis methods and calculations (GCSE/IGCSE/O level, introduction for Advanced Level) Various non-redox titration methods and questions (Advanced Level, acid-alkali, EDTA, silver nitrate etc.) Various redox titration methods and questions (Advanced Level, potassium manganate(VII), thiosulfate/iodine etc.)
PLEASE NOTE: Most of the tests describe use simple apparatus like test tubes, teat pipette, wire for flame test (nichrome, platinum best but costly) and standard chemical reagents accessible in most school or college laboratories. Where possible balanced symbol equations are given for the reactions occurring in doing the test.
Sometimes a precipitate (ppt) initially forms with a limited amount of a reagent, it may then dissolve in excess of reagent to give a clear solution. Both observations will be crucial for a positive id. There are no tests specific to identify a compound e.g. (i) there is no test for calcium chloride, but there are tests for the calcium ion and the chloride ion, i.e. using specific ion tests. (ii) Similarly, in organic tests, all you can do is identify a functional group i.e. a particular bit of the molecular structure of a member of a homologous series, rather a particular unique molecule. Not all the reactions are good definitive tests, but they may well be important reactions of cations or anions you need to know about. The first tests in the 'inorganic' section are typical of GCSE Science level, but finally these overlap and extend into those needed for GCE Advanced AS or A2 level. In the organic section, only the alkene test is in GCSE double award science, but some others might be found in a full single or coordinated triple award GCSE syllabus. If any GCSE/IGCSE/GCE/AS/A2/IB/US grade 8–12 K12 test seems missing, just let me know by email These days more emphasis is given to modern spectroscopic methods of analysis such as NMR, Infrared, Mass spectrometry, Atomic Emission etc. Quite correctly, though updating A level chemistry is intellectually challenging at times, it isn't always as much fun! The methods described give no recipe details or risk assessment, just basically what is needed, what you see and what you can or cannot deduce. Consult teacher, 'practical' text books and Hazcards before attempting any analysis. Most tests involve 'standard' chemical reactions and few tests are totally specific so observations should be viewed in context, i.e. is this a realistic deduction in that particular situation? Please each syllabus has its own 'list' of required tests – so do not 'over learn' – check out what is needed! There is a web page covering the methods some safety aspects of "Preparing and collecting gases". Use the alphabetical list to find the test you need.
METHODS OF MAKING SALTS & tests for ions Doc Brown's Chemistry GCSE/IGCSE Science–Chemistry Revision Notes The pH scale of acidity and alkalinity, acids, alkalis, salts and neutralisation 6. Methods of making salts and chemical tests for salts
Original page now split into four sections, NEW links below How do we make salts? What preparations are available to us?
Four basic methods for preparing salts are described on this page, with annotated diagrams. BEFORE preparing a salt there are two important facts to know ... (i) Is the salt is soluble or insoluble? (ii) If using a base, is it soluble (alkali)? or insoluble? ... because these facts decide which method you use!
Method (a) Making a salt by neutralising a soluble acid with a soluble base (alkali) Method (b) Making a salt by from an acid with a metal or insoluble base – oxide, hydroxide or carbonate Method (c) Preparing an insoluble salt by mixing solutions of two soluble compounds Method (d) Making a salt by directly combining its constituent elements
A summary of chemical tests to identify ions in a salt, hence the identity of a salt Apart from knowing how to make salts, you need to know how to identify salts and other compounds from their constituent ions. There is no single test for a salt, you must do at least two tests to confirm the identity of the two constituent ions. Most of the methods described below are simple precipitation tests.
Tests for METAL IONS – cations (positive ions)
Simple method for a flame test to identify metal ions: The metal salt or other compound is mixed with concentrated hydrochloric acid and a sample of the mixture is heated strongly in a bunsen flame on the end of a cleaned nichrome wire (or platinum if you can afford it!). Before doing the test the nichrome/platinum wire should be cleaned in conc. hydrochloric acid and heated in the hottest part of the flame to make sure there is no contaminating flame colours. It doesn't matter whether the salt compound is soluble or insoluble.
the lithium ion Li+ gives a red/crimson (carmine–red) colour in the flame the sodium ion Na+ gives a yellow/orange colour in the flame the potassium ion K+ gives a lilac/purple colour in the flame the calcium ion Ca2+ gives a brick red colour in the flame the copper ion Cu2+ gives a blue–green colour in the flame 6a. METHOD (a) Neutralising a soluble acid with a soluble base (alkali) to give a soluble salt Salt solubility affects the method you choose to make a salt and so 8. contains tables of informationdata on salt solubility which will help you decide on the method to prepare a salt, but a brief summary is given below below.
A solubility guide for salts - information required to decide on the method used to prepare a salt
salts
solubility?
common salts of sodium, potassium and ammonium ions
usually soluble in water
common sulfates (sulphates)
usually quite soluble except for calcium sulfate (slightly soluble), lead sulfate and barium sulfate are both insoluble
common chlorides (similar rule for bromides and iodides)
usually soluble except for insoluble lead(II) chloride and silver chloride
common nitrates
all soluble 6a. A Method of Making a Water Soluble Salt
6a. METHOD (a) Neutralising a soluble acid with a soluble base (alkali) to give a soluble salt One important point is to recognise that all the reactants are soluble here, which is why you need a titration procedure to work out how much of the acid is to be added to a given volume of alkali. e.g. the hydroxide of an alkali metal like sodium hydroxide, potassium hydroxide or ammonia solution (wrongly called )ammoium hydroxide. Steps (1) to (3) below is called a titration. Typical common soluble bases (alkalis) used for preparing soluble salts:
NaOH sodium hydroxide, KOH potassium hydroxide and some soluble carbonates Typical examples shown by the word and symbol equations below include ... sodium hydroxide + hydrochloric acid ==> sodium chloride + water NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l)
sodium hydroxide + sulphuric acid ==> sodium sulphate + water 2NaOH(aq) + H2SO4(aq) ==> Na2SO4(aq) + 2H2O(l)
potassium hydroxide + sulphuric acid ==> potassium sulphate + water 2KOH(aq) + H2SO4(aq) ==> K2SO4(aq) + 2H2O(l)
sodium hydroxide + nitric acid ==> sodium nitrate + water NaOH(aq) + HNO3(aq) ==> NaNO3(aq) + 2H2O(l)
ammonia + nitric acid ==> ammonium nitrate NH3(aq) + HNO3(aq) ==> NH4NO3(aq)
potassium hydroxide + hydrobromic acid ==> potassium bromide + water KOH(aq) + HBr(aq) ==> KBr(aq) + H2O(l)
sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide Na2CO3(s) + 2HCl(aq) ==> 2NaCl(aq) + H2O(l) + CO2(g)
More examples of neutralization equations are given in section 4.
METHOD (a) procedure for making a soluble salt by neutralising a soluble base (alkali) with an acid. You need to know the exact amount of acid to just neutralise completely the alkali (soluble base). (1) A known volume of acid is pipetted into a conical flask and universal indicator added. The acid is titrated with the alkali from the burette. (2) The acid is added until the indicator turns green, pH 7 neutral. This means all the acid has been neutralised to form the salt. I've illustrated the method using universal indicator BUT it isn't that accurate an indicator for titrations. You should use a more precise indicator like phenolphthalein or methyl orange. I didn't repeat all the titration details here again, I've just kept to the basic ideas and description, but there lots of detailed examples on the page ( more examples - diagrams, descriptions of titration procedures) (3) The volume of alkali needed for neutralisation is then noted, this is called the endpoint volume. (1)-(3) are repeated with both known volumes mixed together BUT without the contaminating indicator, such as phenolphthalein or methyl orange. (4) The solution is transferred to an evaporating dish and heated to partially evaporate the water causing crystallisation or can be left to very slowly evaporate - which tends to give bigger and better crystals. (5) The residual liquid can be decanted away and the crystals can be carefully collected and dried by 'dabbing' with a filter paper OR the crystals can be collected by filtration and dried (as above). See also the preparation of ammonium salts using this method. AND also more examples - diagrams, descriptions of titration procedures Note (i) You can put the acid in the burette and the alkali in the flask. (ii) Parts (1) to (3) are known specifically as an acid-base (alkali) titration, and the general method is known as a volumetric titration by which it possible to find out exactly what volume ratios are needed for neutralisation. So knowing one concentration, you can calculate the other. (iii) Concentration calculations are on calculations pages sections 11. and 12. (iv) Apparatus used: (1) pipette and conical flask; (2)-(3) burette and conical flask; (4) evaporating (crystallising) dish, bunsen burner, tripod and gauze; (5) filter paper. (v) Other indicators e.g. phenolphthalein can be used instead (pink alkaline, colourless acid). (vi) The burette and pipette are both used for the accurate measurement of volume.
(vii)
The pH changes in this preparation are described in section 7
6b. A 2nd Method of Making a Water Soluble Salt 6b. METHOD (b) Reacting an acid with a metal or with an insoluble base to give a soluble salt e.g. an insoluble base such as a metal oxide, metal hydroxide or a metal carbonate, often of a Group 2 metal like calcium, magnesium or a Transition Metal like nickel, copper or zinc. Copper metal won't dissolve in acids, but its oxide and carbonate will. Using the same procedure you can also start with a metal that has a low reactivity towards water e.g. magnesium, zinc or iron. One important point is to recognise that all the reactants are soluble here, which is why you need a titration procedure to work out how much of the acid is to be added to a given volume of alkali. Typical common insoluble bases used for preparing soluble salts: MgO magnesium oxide, MgCO3 magnesium carbonate, CaO Calcium oxide, CaCO3 calcium carbonate, Ca(OH)2 calcium hydroxide, NiO nickel(II) oxide, ZnO zinc oxide, Zn(OH)2, zinc hydroxide, ZnCO3 zinc carbonate, CuO copper(II) oxide, CuCO3 copper(II) carbonate, PbCO3lead(II) carbonate (with nitric acid to make lead(II) nitrate), FeCO3 iron(II) carbonate (to make iron(II) salts), MnCO 3 manganese(II) carbonate Typical examples shown by the word and symbol equations below include ... copper(II) oxide + sulphuric acid ==> copper(II) sulphate + water CuO + H2SO4 ==> CuSO4 + H2O CuO(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) magnesium hydroxide + sulfuric acid ==> magnesium sulfate + water Mg(OH)2 + H2SO4 ==> MgSO4 + 2H2O
Mg(OH)2(s) + H2SO4(aq) ==> MgSO4(aq) + 2H2O(l)
magnesium hydroxide + hydrochloric acid ==> magnesium chloride + water Mg(OH)2(s) + 2HCl(aq) ==> MgCl2(aq) + 2H2O(l)
Zinc carbonate + nitric acid ==> zinc nitrate + water + carbon dioxide ZnCO3 + 2HNO3 ==> Zn(NO3)2 + H2O + CO2 ZnCO3(s) + 2HNO3(aq) ==> Zn(NO3)2(aq) + H2O(l) + CO2 (g)
zinc oxide + hydrochloric acid ==> zinc chloride + water ZnO(s) + 2HCl(aq) ==> ZnCl2(aq) + H2O(l) Similar for many other Group 2 and Transition metal oxides e.g. Mg, Ca, Ba and Co, Ni, Cu instead of Zn
zinc + sulfuric acid ==> zinc sulfate + hydrogen Zn + H2SO4 ==> ZnSO4 + H2 Zn(s) + H2SO4(aq) ==> ZnSO4(aq) + H2(g)
calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide CaCO3(s) + 2HCl(aq) ==> CaCl2(aq)+ H2O(l) + CO2(g) Similar for many other Group 2 and Transition metal carbonates e.g. Mg, Sr, Ba and Ni, Co, Zn instead of Ca
Carbonates are frequently used in this method of salt making, e.g. using copper carbonate to make copper salts
o
The equations are give with, and without sate symbols.
copper(II) carbonate + hydrochloric acid ==> Copper(II) chloride + water + carbon dioxide
o
CuCO3 + 2HCl ==> CuCl2 + H2O + CO2
o
and with sulphuric acid a blue solution of copper(II) sulphate is formed.
copper(II) carbonate + sulphuric acid ==> Copper(II) sulphate + water + carbon dioxide
o
CuCO3 + H2SO4 ==> CuSO4 + H2O + CO2
CuCO3(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) + CO2(g)
copper(II) carbonate + nitric acid ==> Copper(II) nitrate + water + carbon dioxide
o
CuCO3 + 2HNO3 ==> Cu(NO3)2 + H2O + CO2
CuCO3(s) + 2HCl(aq) ==> CuCl2(aq) + H2O(l) + CO2(g)
CuCO3(s) + 2HNO3(aq) ==> CuSO4(aq) + H2O(l) + CO2(g)
Similar equations for other carbonates to give soluble salts which can be crystallised from solution e.g.
o
calcium carbonate CaCO3, to make two salts - calcium chloride/nitrate (calcium sulfate is not very soluble)
o
iron(II) carbonate FeCO3, to make three salts - iron(II) chloride/sulfate/nitrate
o
magnesium carbonate MgCO3, to make three salts - magnesium chloride/sulfate/nitrate
o
manganese(II) carbonate MnCO3, to make three salts - manganese(II) chloride/sulfate/nitrate
o
zinc carbonate ZnCO3, to make three salts - zinc chloride/sulfate/nitrate
o
lead(II) carbonate PbCO3, only nitric acid to make lead(II) nitrate, lead(II) chloride and lead(II) sulfate are insoluble and must be prepared by method (c)
More examples of neutralization equations are given in section 4. METHOD (b) Procedure for making a soluble salt from an insoluble base, carbonate or metal
(1) The required volume of acid is measured out into the beaker with a measuring cylinder. The excess of insoluble metal, oxide, hydroxide or carbonate is weighed out and the solid added in small portions to the acid in the beaker with stirring. Doing the weighing will minimise trial and error especially if the reaction is slow, as long as you know how to do the theoretical calculation! (2) The mixture may be heated to speed up the reaction. When no more of the solid dissolves it means ALL the acid is neutralised and there should be a little excess solid. You should see a residue of the solid (oxide, hydroxide, carbonate) left at the bottom of the beaker. (3) The hot solution (with care!) is filtered to remove the excess solid metal/oxide/carbonate, into an evaporating dish. (4) You may need to carefully heat the solution to evaporate some of the water. Then hot solution is left to cool and crystallise. After crystallisation, you collect and dry the crystals with a filter paper. Note (i) Apparatus used: (1) balance, measuring cylinder, beaker and glass stirring rod. (2) beaker/rod, bunsen burner, tripod and gauze; (3)-(4) filter funnel and filter paper, evaporating (crystallising) dish. (ii) A measuring cylinder is adequate for measuring the acid volume, you do not need the accuracy of a pipette or burette required in method (a).
(iii) How to calculate amounts required and % yield is dealt with in 14
Chemical Calculations Part
Salt solubility affects the method you choose to make a salt and so section 8. contains tables of information-data on salt solubility which will help you decide on the method to prepare a salt.
Multiple choice revision quizzes and other worksheets
GCSE/IGCSE foundation-easier multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE higher-harder multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE Structured question worksheet on Acid Reaction word equations and symbol equation questions
Word equation answers and symbol equation answers)
GCSE/IGCSE word-fill worksheet on Acids, Bases, Neutralisation and Salts
GCSE/IGCSE matching pair quiz on Acids, Bases, Salts and pH
See also
Advanced Level Chemistry Students Acid-Base Revision Notes - use index
Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel Science Chemistry IGCSE Chemistry & OCR 21st Century Science Chemistry, OCR Gateway Science Chemistry, WJEC gcse science chemistry CCEA/CEA gcse science chemistry (revise science chemistry courses equal to US grade 8, grade 9 grade 10. Revision notes on acids, bases, alkalis, salts, solution pH word equations balanced symbol equations science chemistry courses revision guides How to make copper(II) sulfate (copper sulphate), how to make zinc sulfate (zinc sulphate)
ALL Website content copyright © Dr Phil Brown 2000-2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * [email protected] Describing how to make a salt by four different methods, how to do it, what laboratory equipment is needed and how to crystallise or collect the salt in the end 6c. Method (c) Preparing an Insoluble Salt Procedure for making an insoluble salt my mixing solutions of soluble compounds to form a precipitate. The two soluble compounds must each provide one of the constituent ions of the desired insoluble salt which precipitates out when the solutions are mixed. NOTE definition: A precipitation reaction is generally defined as 'the formation of an insoluble solid on mixing two solutions of soluble substances or bubbling a gas into a solution'. Note that several precipitation reactions are used as simple tests for e.g. for sulfate, chloride, bromide, iodide Salt solubility affects the method you choose to make a salt and so 8. contains tables of informationdata on salt solubility which will help you decide on the method to prepare a salt, but a brief summary is given below
A solubility guide for salts - information required to decide on the method used to prepare a salt
salts
solubility?
common salts of sodium, potassium and ammonium ions
usually soluble in water
common sulfates (sulphates)
usually quite soluble except for calcium sulfate (slightly soluble), lead sulfate and barium sulfate are both insoluble
common chlorides (similar rule for bromides and iodides)
usually soluble except for insoluble lead(II) chloride and silver chloride
common nitrates
all soluble
common carbonates
most metal carbonates are insoluble apart from sodium & potassium carbonate. Ammonium carbonate is also soluble.
How can we make an insoluble salt? How do we prepare an insoluble salt from two soluble compounds?
This section describes the preparation of insoluble salts like silver chloride AgCl, lead(II) chloride (lead chloride) PbCl2, lead(II) iodide (lead iodide) PbI2, calcium carbonate CaCO3, barium sulfate (barium sulphate) BaSO4, lead(II) sulfate (lead sulphate, lead sulfate)PbSO4, and 'slightly soluble' calcium sulfate (calcium sulphate) CaSO4, which can all be made by a precipitation reaction.
o
All of the above insoluble salts are white (as in diagram), except lead(II) iodide which is yellow.
Many of the salt precipitate are WHITE, but lead iodide is pale yellow.
METHOD (c)
An insoluble salt can be made by mixing two solutions of soluble salts in a process is called precipitation.
o
The method is quite simple – illustrated above, assuming in this case the insoluble salt is colourless–white.
o
One solution contains the 1st required ion, and the other solution contains the 2nd required ion.
o
So you must prepare two solutions of soluble compounds, each of which provides one of the two ions required to combine and precipitate out as the insoluble salt.
o
Each soluble compound is weighed out into its own beaker and dissolved in a suitable volume of water until the solutions are both quite clear.
One solution is then poured into the other, order doesn't really matter.
The two solutions of SOLUBLE compounds must be thoroughly mixed together to ensure all the reactants are used up, so the maximum amount of INSOLUBLE salt precipitate is formed.
You see the two clear solutions on mixing forming a cloudy mixture as the insoluble compound is formed, known as theprecipitate.
o
The mixture is then carefully poured into a funnel holding a filter paper.
o
The precipitated salt can then be filtered off with the filter funnel and paper.
o
While still in the filter paper and funnel, the collected solid precipitate is washed with distilled/deionised water to remove any remaining soluble salt impurities and just the damp, but otherwise pure, insoluble salt is left.
o
The precipitate is then carefully removed from the filter paper into a clean dish or basin to be dried e.g. left out in a dry room or warmed in a pre–heated oven.
Examples ...
o
(i) Silver chloride is made by mixing solutions of solutions of silver nitrate and sodium chloride.
silver nitrate + sodium chloride ==> silver chloride + sodium nitrate
AgNO3(aq) + NaCl(aq) ==> AgCl(s) + NaNO3(aq)
in of ions it could be written as
Ag+NO3–(aq) + Na+Cl–(aq) ==> AgCl(s) + Na+NO3–(aq)
or: Ag+(aq) + NO3–(aq) + Na+(aq) + Cl–(aq) ==> AgCl(s) + Na+(aq) + NO3–(aq)
but the spectator ions are nitrate NO3– and sodium Na+ which do not change at all,
The 'active ions' and resulting precipitate are highlighted in yellow.
so the proper ionic equation is simply: Ag+(aq) + Cl–(aq) ==> AgCl(s)
Note (i) the use of state symbols (aq) and (s) AND
(ii) that ionic equations omit ions that do not change there chemical or physical state.
In this case the nitrate, NO3–(aq) and sodium Na+(aq) ions do not change physically or chemically and are called spectator ions,
BUT the aqueous silver ion, Ag +(aq), combines with the aqueous chloride ion, Cl–(aq), to form the insoluble salt silver chloride, AgCl (s), thereby changing their states both chemically and physically.
More Ionic equations explained with all spectator ions indicated
If you use barium chloride the word and symbol equations are ...
barium chloride + silver nitrate ==> silver chloride + barium nitrate
BaCl2(aq) + 2AgNO3(aq) ==> 2AgCl(s) + Ba(NO3)2(aq)
which can be written as
Ba2+(aq) + 2Cl–(aq) + 2Ag+(aq) + 2NO3–(aq) ==> 2AgCl(s) + Ba2+(aq) + 2NO3–(aq)
the spectator ions are Ba2+ and NO3–
so the ionic equation is: Ag+(aq) + Cl–(aq) ==> AgCl(s)
o
(ii) Lead(II) iodide, a yellow precipitate (insoluble in water!) can be made by mixing lead(II) nitrate solution with e.g. potassium iodide solution.
lead(II) nitrate + potassium iodide ==> lead(II) iodide + potassium nitrate
Pb(NO3)2(aq) + 2KI(aq) ==> PbI2(s) + 2KNO3(aq)
which can be written as
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2I–(aq) ==> PbI2(s) + 2K+(aq) + 2NO3–(aq)
the ionic equation is: Pb2+(aq) + 2I–(aq) ==> PbI2(s)
because the spectator ions are nitrate NO3– and potassium K+.
In a similar way you can make lead(II) chloride by e.g. using dilute hydrochloric acid
lead(II) nitrate + hydrochloric acid ==> lead(II) chloride + nitric acid
Pb(NO3)2(aq) + 2HCl(aq) ==> PbCl2(s) + 2HNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2H+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2H+(aq) + 2NO3–(aq)
the proper ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)
because the spectator ions are nitrate NO3– and hydrogen H+.
You can make lead(II) bromide by e.g. using sodium bromide solution
lead(II) nitrate + sodium bromide ==> lead(II) bromide + sodium nitrate
Pb(NO3)2(aq) + 2NaBr(aq) ==> PbBr2(s) + 2NaNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2Na+(aq) + 2Br–(aq) ==> PbBr2(s) + 2Na+(aq) + 2NO3–(aq)
the proper ionic equation is: Pb2+(aq) + 2Br–(aq) ==> PbBr2(s)
because the spectator ions are nitrate NO3– and sodium Na+.
and you can make lead(II) chloride by e.g. using sodium chloride solution
o
o
lead(II) nitrate + sodium chloride ==> lead(II) chloride + sodium nitrate
Pb(NO3)2(aq) + 2NaCl(aq) ==> PbBr2(s) + 2NaNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2Na+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2Na+(aq) + 2NO3–(aq)
the proper ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)
because the spectator ions are nitrate NO3– and sodium Na+.
You can also use lead nitrate and dilute hydrochloric acid to precipitate lead chloride
lead(II) nitrate + hydrochloric acid ==> lead(II) chloride + nitric acid
Pb(NO3)2(aq) + 2HCl(aq) ==> PbCl2(s) + 2HNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2H+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2H+(aq) + 2NO3–(aq)
so the ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)
because the spectator ions are nitrate NO3– and hydrogen H+.
(iii) Calcium carbonate, a white precipitate, forms on e.g. mixing calcium chloride solution and sodium carbonate solutions ...
calcium chloride + sodium carbonate ==> calcium carbonate + sodium chloride
CaCl2(aq) + Na2CO3(aq) ==> CaCO3(s) + 2NaCl(aq)
Ca2+(aq) + 2Cl–(aq) + 2Na+(aq) + CO32–(aq) ==> CaCO3(s) + 2Na+(aq) + 2Cl–(aq)
the ionic equation is: Ca2+(aq) + CO32–(aq) ==> CaCO3(s)
because the spectator ions are chloride Cl– and sodium Na+.
(iv) Barium sulphate (barium sulphate), a white precipitate, forms on mixing e.g. barium chloride solution and dilute sulphuric acid ...
barium chloride + sulphuric acid ==> barium sulphate + hydrochloric acid
o
BaCl2(aq) + H2SO4(aq) ==> BaSO4(s) + 2HCl(aq)
Ba2+(aq) + 2Cl–(aq) + 2H+(aq) + SO42–(aq) ==> BaSO4(s) + 2H+(aq) + 2Cl–(aq)
the ionic equation is: Ba2+(aq) + SO42–(aq) ==> BaSO4(s)
because the spectator ions are chloride Cl– and hydrogen H+.
Or you can use sulphate salts like sodium sulphate, so the word and symbol equations are ..
barium chloride + sodium sulfate ==> barium sulfate + sodium chloride
BaCl2(aq) + Na2SO4(aq) ==> BaSO4(s) + 2NaCl(aq)
The ionic equation is the same: Ba2+(aq) + SO42–(aq) ==> BaSO4(s)
because the spectator ions are sodium Na+ and chloride Cl–
You can do exactly the same preparation using potassium sulfate and barium nitrate or barium chloride, basically any solutions of a soluble barium salt and a soluble sulfate salt can be used to prepare barium sulfate.
e.g. barium nitrate + potassium sulfate ==> barium sulfate + potassium nitrate
or barium chloride + potassium sulfate ==> barium sulfate + potassium chloride
See at the end of the page the uses of barium sulfate.
(v) Lead(II) sulphate (lead sulphate), a white precipitate, forms in a similar way e.g.
lead(II) nitrate + sodium sulphate ==> lead(II) sulphate + sodium nitrate
Pb(NO3)2 (aq) + Na2SO4(aq) ==> PbSO4(s) + 2NaNO3 (aq)
ionic equation: Pb2+(aq) + SO42–(aq) ==> PbSO4(s)
because the spectator ions are sodium Na+ and nitrate NO3–
Lead(II) sulphate can also be precipitated using dilute sulfuric acid.
o
lead(II) nitrate + sodium sulphate ==> lead(II) sulphate + nitric acid
Pb(NO3)2 (aq) + H2SO4(aq) ==> PbSO4(s) + 2HNO3 (aq)
ionic equation: Pb2+(aq) + SO42–(aq) ==> PbSO4(s)
because the spectator ions are hydrogen H+ and nitrate NO3– ions
(vi) Calcium sulphate (calcium sulphate), a white precipitate, forms on mixing e.g. calcium chloride solution and dilute sulphuric acid ...
calcium chloride + sulphuric acid ==> barium sulphate + hydrochloric acid
CaCl2(aq) + H2SO4(aq) ==> CaSO4(s) + 2HCl(aq)
Ca2+(aq) + 2Cl–(aq) + 2H+(aq) + SO42–(aq) ==> CaSO4(s) + 2H+(aq) + 2Cl–(aq)
the ionic equation is: Ca2+(aq) + SO42–(aq) ==> CaSO4(s)
because the spectator ions are chloride Cl– and hydrogen H+.
Or you can use sulphate salts like sodium sulphate, so the word and symbol equations are ..
calcium chloride + sodium sulfate ==> barium sulfate + sodium chloride
CaCl2(aq) + Na2SO4(aq) ==> CaSO4(s) + 2NaCl(aq)
The ionic equation is the same: Ba2+(aq) + SO42–(aq) ==> BaSO4(s)
because the spectator ions are sodium Na+ and chloride Cl–
You can do exactly the same preparation using potassium sulfate and calcium nitrate or calcium chloride, basically any solutions of a soluble calcium salt and a soluble sulfate salt can be used to prepare calcium sulfate.
The yield is a little lower than the theoretical because calcium sulfate is slightly soluble in water.
See at the end of the page the uses of barium sulfate.
o
NOTE reminder definition: A precipitation reaction is generally defined as 'the formation of an insoluble solid on mixing two solutions or bubbling a gas into a solution'.
General rules which describe the solubility of common types of compounds in water:
o
All common sodium, potassium and ammonium salts are soluble e.g. NaCl, K 2SO4, NH4NO3
o
All nitrate salts are soluble e.g. NaNO3, Mg(NO3)2, Al(NO3)3, NH4NO3
o
Some ethanoate salts are soluble e.g. CH 3COONa
o
Common chloride salts are soluble except those of silver and lead e.g.
o
o
soluble: KCl, CaCl2, AlCl3 or insoluble AgCl, PbCl2
Common sulfates are soluble except those of lead, barium and calcium: soluble e.g.
soluble: Na2SO4, MgSO4, Al2(SO4)3
insoluble: PbSO4, BaSO4, CaSO4 is slightly soluble.
Common oxides, hydroxides and carbonates are usually insoluble (e.g. Group 2 and Transition Metals) except those of the Group 1 Alkali Metals sodium, potassium etc. and ammonium:
soluble bases–alkalis oxides, hydroxides or carbonates: K 2O, KOH, NaOH, NH3(aq), Na2CO3, (NH4)2CO3
insoluble bases – basic oxides, hydroxides or insoluble carbonates: MgO, CuO, ZnO, Mg(OH)2, Fe(OH)2, Cu(OH)2, CuCO3, ZnCO3, CaCO3
Some uses of insoluble salts (their preparation has been described above)
o
Calcium sulfate CaSO4 (calcium sulphate) is used in plaster of Paris and plaster for domestic wall covering.
o
Barium sulfate BaSO4 (barium sulphate)
Barium sulfates quite a dense material is used in with X-rays for particular medical examinations.
Barium sulfate, like all barium salts is toxic, BUT, because it is insoluble, non of it is absorbed by the body into the bloodstream and eventually it will right through the gut system and be expelled in the normal faeces.
X-rays are used to investigate bone structure e.g. fractured or broken bones and the technique works because bone material is too dense to let weak X-rays through, so on shooting an X-ray photograph you get the bone structure as a sort of shadow effect where the X-rays have been absorbed by the bone.
Therefore, normally you can't use X-rays to examine soft tissue areas like the gut.
However, if the patient takes a 'barium meal', a thick harmless fluid containing a suspension of the insoluble barium sulfate, this can through the stomach and into the gut.
Therefore it is then possible to X-ray the intestinal gut system because the barium sulfate absorbs the X-rays just like bone.
So, because barium sulfate is opaque to X-rays, the X-ray photograph via the barium meal will highlight the structure of the tissue lining of the gut and show up any structural abnormalities and blockages.
6d. Method (d) Making a salt by direct combination of elements Sometimes it isn't appropriate to prepare a soluble salt by reacting an acid with an insoluble base or alkali, so it may be possible to prepare the salt by directly combining the metal and the non-metal elements. Two such examples are the preparation of anhydrous aluminium chlorideand anhydrous iron(III) chloride (anhydrous here means without any water of crystallisation).
Preparation of aluminium chloride AlCl3
Preparation of iron(III) chloride FeCl3
How can we make aluminium chloride? How do we prepare iron(III) chloride?
METHOD (d) both preparations illustrated above.
These compounds can be made by direct combination of the elements to form anhydrous salts e.g. if dry chlorine gas Cl2 is ed over heated iron or aluminium, the chloride is produced. These experiment preparations (shown above) should be done very carefully by the teacher in a fume cupboard.
o
o
The aluminium can burn intensely with a violet flame, white fumes of aluminium chloride sublime from the hot reacted aluminium and the white solid forms on the cold surface of the flask.
aluminium + chlorine ==> aluminium chloride
2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
The aluminium chloride is often discoloured yellow from the trace chlorides of copper or iron that may be formed from traces of these metals that might be present in the original aluminum).
The iron (e.g. as steel wool) glows red and brown fumes of iron(III) chloride stream off, the brown solid collects on the cold flask surface.
iron + chlorine ==> iron(III) chloride
2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)
o
Note (i): Both these chlorides react exothermically and hydrolyse with water to give the metal hydroxide and fumes of hydrogen chloride, and so dry conditions are needed.
o
Note (ii): Both these chlorides cannot be made in an anhydrous form from aqueous solution neutralisation. This is because on evaporation the compounds contain 'water of crystallisation'. On heating the hydrated salt hydrolyses and decomposes into water, the oxide or hydroxide and fumes of hydrogen chloride, and maybe some impure anhydrous chloride, basically it a mess in of trying to make pure AlCl 3 and FeCl3 in this way.
Multiple choice revision quizzes and other worksheets
GCSE/IGCSE foundation-easier multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE higher-harder multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE Structured question worksheet on Acid Reaction word equations and symbol equation questions
Word equation answers and symbol equation answers)
GCSE/IGCSE word-fill worksheet on Acids, Bases, Neutralisation and Salts
GCSE/IGCSE matching pair quiz on Acids, Bases, Salts and pH
See also
Advanced Level Chemistry Students Acid-Base Revision Notes - use index
Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel Science Chemistry IGCSE Chemistry & OCR 21st Century Science Chemistry, OCR Gateway Science Chemistry, WJEC gcse science chemistry CCEA/CEA gcse science chemistry (revise science chemistry courses equal to US grade 8, grade 9 grade 10. Revision notes on acids, bases, alkalis, salts, solution pH word equations balanced symbol equations science chemistry courses revision guides
ALL Website content copyright © Dr Phil Brown 2000-2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * [email protected] Describing how to make a salt by four different methods, how to do it, what laboratory equipment is needed and how to crystallise or collect the salt in the end
6d. Method (d) Making a salt by direct combination of elements Sometimes it isn't appropriate to prepare a soluble salt by reacting an acid with an insoluble base or alkali, so it may be possible to prepare the salt by directly combining the metal and the non-metal elements. Two such examples are the preparation of anhydrous aluminium chlorideand anhydrous iron(III) chloride (anhydrous here means without any water of crystallisation).
Preparation of aluminium chloride AlCl3
Preparation of iron(III) chloride FeCl3
How can we make aluminium chloride? How do we prepare iron(III) chloride?
METHOD (d) both preparations illustrated above.
These compounds can be made by direct combination of the elements to form anhydrous salts e.g. if dry chlorine gas Cl2 is ed over heated iron or aluminium, the chloride is produced. These experiment preparations (shown above) should be done very carefully by the teacher in a fume cupboard.
o
The aluminium can burn intensely with a violet flame, white fumes of aluminium chloride sublime from the hot reacted aluminium and the white solid forms on the cold surface of the flask.
aluminium + chlorine ==> aluminium chloride
2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
o
The aluminium chloride is often discoloured yellow from the trace chlorides of copper or iron that may be formed from traces of these metals that might be present in the original aluminum).
The iron (e.g. as steel wool) glows red and brown fumes of iron(III) chloride stream off, the brown solid collects on the cold flask surface.
iron + chlorine ==> iron(III) chloride
2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)
o
Note (i): Both these chlorides react exothermically and hydrolyse with water to give the metal hydroxide and fumes of hydrogen chloride, and so dry conditions are needed.
o
Note (ii): Both these chlorides cannot be made in an anhydrous form from aqueous solution neutralisation. This is because on evaporation the compounds contain 'water of crystallisation'. On heating the hydrated salt hydrolyses and decomposes into water, the oxide or hydroxide and fumes of hydrogen chloride, and maybe some impure anhydrous chloride, basically it a mess in of trying to make pure AlCl 3 and FeCl3 in this way.
Multiple choice revision quizzes and other worksheets
GCSE/IGCSE foundation-easier multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE higher-harder multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE Structured question worksheet on Acid Reaction word equations and symbol equation questions
Word equation answers and symbol equation answers)
GCSE/IGCSE word-fill worksheet on Acids, Bases, Neutralisation and Salts
GCSE/IGCSE matching pair quiz on Acids, Bases, Salts and pH
See also
Advanced Level Chemistry Students Acid-Base Revision Notes - use index
Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel Science Chemistry IGCSE Chemistry & OCR 21st Century Science Chemistry, OCR Gateway Science Chemistry, WJEC gcse science chemistry CCEA/CEA gcse science chemistry (revise science chemistry courses equal to US grade 8, grade 9 grade 10. Revision notes on acids, bases, alkalis, salts, solution pH word equations balanced symbol equations science chemistry courses revision guides
ALL Website content copyright © Dr Phil Brown 2000-2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * [email protected] Describing how to make a salt by four different methods, how to do it, what laboratory equipment is needed and how to crystallise or collect the salt in the end
Environmental agencies monitor levels of oxygen and pollutants in rives and lakes. Hospitals carry out complex blood analysis for iron, proteins, cholesterol etc. and this is very important diagnostic information for clinicians and doctors seeking to get you well again prevent diseases developing. Blood can be analysed for specific proteins to indicate particular medical conditions e.g. cancer, and ions such as sodium, chloride and iron compounds. At water treatment plants test are done to determine the levels of metal ions, insecticides and lots of other substances to check that their levels are not high enough to be harmful to humans. Apart from a multitude of forensic tests for DNA, powder burns from guns, explosives etc. the police use breathalyser kits to test for alcohol levels in your breath and may request a blood sample for analysis too. There are two types of chemical analysis Qualitative tests and quantitative analysis. Whatever the nature of the investigation, all tests of analyses should be carried out by using 'tried and tested' standard procedures. They should be the most accurate, reliable and safest methods that can be devised. It means, whatever laboratory you work in, anywhere in the world, you should get the same results as long as the sample is collected, stored and analysed by the same standard procedures.
(a) QUALITATIVE CHEMICAL ANALYSIS Qualitative chemical analysis indicates whether a particular substance is present or not. It does not tell how much of the substance is there or its concentration. However, if a substance is potentially harmful, even toxic, its a good idea to know whether the substance is there at all. The larger the sample you have, the better. With more to work with you are more likely to able to detect minute traces of substances with qualitative tests AND have spare material to repeat tests several times if the results seem uncertain at first. At school/college level, the simple tests you learn enable to identify the cation (+ve ion e.g. metal ions) and anion (-ve ion, e.g. chloride, sulfate) in a salt, and the salt usually does contain only two ions, but some salts do have three ions e.g. iron(II) ammonium sulfate which might take a bit of sorting out. Most tests at this level are done on soluble substances so that you can dissolve the substance in water and carry out tests on the aqueous solution. Aqueous means a
solution in water. You may come across a non-aqueous solution using a non-aqueous solvent like alcohol or hexane. AND don't forget, negative test results are just as important as positive results, you may need to eliminate possibilities as well as confirm the presence of a particular ion or gas etc. UNFORTUNATELY, not all tests are unique for a particular ion, but this shouldn't be a problem in school chemistry! Associated qualitative analysis links See index at top of page Typical qualitative tests are described in Parts 2 to 5. Summary of some cation and anion tests for GCSE/IGCSE/O Level students
(b) QUANTITATIVE CHEMICAL ANALYSIS Quantitative analysis gives you (hopefully) a precise measure of how much of a substance is present or its concentration in a sample being analysed e.g. ore analysis tells you whether it is worth exploiting for a metal, the purity of drug ensure no harmful impurities in it, blood sample analysis for alcohol allows the police to decide to prosecute for 'being over the limit'. With a large sample you have spare material to repeat the quantitative analysis several times to get the most statistically valid result. In schools and colleges you can do quite accurate titrations to illustrate quantitative analysis. Associated quantitative links % purity of a product (GCSE/IGCSE/O level, introduction for Advanced Level) Volumetric titration analysis methods and calculations (GCSE/IGCSE/O level, introduction for Advanced Level) Various non-redox titration methods and questions (Advanced Level, acid-alkali, EDTA, silver nitrate etc.) Various redox titration methods and questions (Advanced Level, potassium manganate(VII), thiosulfate/iodine etc.)
PLEASE NOTE: Most of the tests describe use simple apparatus like test tubes, teat pipette, wire for flame test (nichrome, platinum best but costly) and standard chemical reagents accessible in most school or college laboratories. Where possible balanced symbol equations are given for the reactions occurring in doing the test.
Sometimes a precipitate (ppt) initially forms with a limited amount of a reagent, it may then dissolve in excess of reagent to give a clear solution. Both observations will be crucial for a positive id. There are no tests specific to identify a compound e.g. (i) there is no test for calcium chloride, but there are tests for the calcium ion and the chloride ion, i.e. using specific ion tests. (ii) Similarly, in organic tests, all you can do is identify a functional group i.e. a particular bit of the molecular structure of a member of a homologous series, rather a particular unique molecule. Not all the reactions are good definitive tests, but they may well be important reactions of cations or anions you need to know about. The first tests in the 'inorganic' section are typical of GCSE Science level, but finally these overlap and extend into those needed for GCE Advanced AS or A2 level. In the organic section, only the alkene test is in GCSE double award science, but some others might be found in a full single or coordinated triple award GCSE syllabus. If any GCSE/IGCSE/GCE/AS/A2/IB/US grade 8–12 K12 test seems missing, just let me know by email These days more emphasis is given to modern spectroscopic methods of analysis such as NMR, Infrared, Mass spectrometry, Atomic Emission etc. Quite correctly, though updating A level chemistry is intellectually challenging at times, it isn't always as much fun! The methods described give no recipe details or risk assessment, just basically what is needed, what you see and what you can or cannot deduce. Consult teacher, 'practical' text books and Hazcards before attempting any analysis. Most tests involve 'standard' chemical reactions and few tests are totally specific so observations should be viewed in context, i.e. is this a realistic deduction in that particular situation? Please each syllabus has its own 'list' of required tests – so do not 'over learn' – check out what is needed! There is a web page covering the methods some safety aspects of "Preparing and collecting gases". Use the alphabetical list to find the test you need.
METHODS OF MAKING SALTS & tests for ions Doc Brown's Chemistry GCSE/IGCSE Science–Chemistry Revision Notes The pH scale of acidity and alkalinity, acids, alkalis, salts and neutralisation 6. Methods of making salts and chemical tests for salts
Original page now split into four sections, NEW links below How do we make salts? What preparations are available to us?
Four basic methods for preparing salts are described on this page, with annotated diagrams. BEFORE preparing a salt there are two important facts to know ... (i) Is the salt is soluble or insoluble? (ii) If using a base, is it soluble (alkali)? or insoluble? ... because these facts decide which method you use!
Method (a) Making a salt by neutralising a soluble acid with a soluble base (alkali) Method (b) Making a salt by from an acid with a metal or insoluble base – oxide, hydroxide or carbonate Method (c) Preparing an insoluble salt by mixing solutions of two soluble compounds Method (d) Making a salt by directly combining its constituent elements
A summary of chemical tests to identify ions in a salt, hence the identity of a salt Apart from knowing how to make salts, you need to know how to identify salts and other compounds from their constituent ions. There is no single test for a salt, you must do at least two tests to confirm the identity of the two constituent ions. Most of the methods described below are simple precipitation tests.
Tests for METAL IONS – cations (positive ions)
Simple method for a flame test to identify metal ions: The metal salt or other compound is mixed with concentrated hydrochloric acid and a sample of the mixture is heated strongly in a bunsen flame on the end of a cleaned nichrome wire (or platinum if you can afford it!). Before doing the test the nichrome/platinum wire should be cleaned in conc. hydrochloric acid and heated in the hottest part of the flame to make sure there is no contaminating flame colours. It doesn't matter whether the salt compound is soluble or insoluble.
the lithium ion Li+ gives a red/crimson (carmine–red) colour in the flame the sodium ion Na+ gives a yellow/orange colour in the flame the potassium ion K+ gives a lilac/purple colour in the flame the calcium ion Ca2+ gives a brick red colour in the flame the copper ion Cu2+ gives a blue–green colour in the flame 6a. METHOD (a) Neutralising a soluble acid with a soluble base (alkali) to give a soluble salt Salt solubility affects the method you choose to make a salt and so 8. contains tables of informationdata on salt solubility which will help you decide on the method to prepare a salt, but a brief summary is given below below.
A solubility guide for salts - information required to decide on the method used to prepare a salt
salts
solubility?
common salts of sodium, potassium and ammonium ions
usually soluble in water
common sulfates (sulphates)
usually quite soluble except for calcium sulfate (slightly soluble), lead sulfate and barium sulfate are both insoluble
common chlorides (similar rule for bromides and iodides)
usually soluble except for insoluble lead(II) chloride and silver chloride
common nitrates
all soluble 6a. A Method of Making a Water Soluble Salt
6a. METHOD (a) Neutralising a soluble acid with a soluble base (alkali) to give a soluble salt One important point is to recognise that all the reactants are soluble here, which is why you need a titration procedure to work out how much of the acid is to be added to a given volume of alkali. e.g. the hydroxide of an alkali metal like sodium hydroxide, potassium hydroxide or ammonia solution (wrongly called )ammoium hydroxide. Steps (1) to (3) below is called a titration. Typical common soluble bases (alkalis) used for preparing soluble salts:
NaOH sodium hydroxide, KOH potassium hydroxide and some soluble carbonates Typical examples shown by the word and symbol equations below include ... sodium hydroxide + hydrochloric acid ==> sodium chloride + water NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l)
sodium hydroxide + sulphuric acid ==> sodium sulphate + water 2NaOH(aq) + H2SO4(aq) ==> Na2SO4(aq) + 2H2O(l)
potassium hydroxide + sulphuric acid ==> potassium sulphate + water 2KOH(aq) + H2SO4(aq) ==> K2SO4(aq) + 2H2O(l)
sodium hydroxide + nitric acid ==> sodium nitrate + water NaOH(aq) + HNO3(aq) ==> NaNO3(aq) + 2H2O(l)
ammonia + nitric acid ==> ammonium nitrate NH3(aq) + HNO3(aq) ==> NH4NO3(aq)
potassium hydroxide + hydrobromic acid ==> potassium bromide + water KOH(aq) + HBr(aq) ==> KBr(aq) + H2O(l)
sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide Na2CO3(s) + 2HCl(aq) ==> 2NaCl(aq) + H2O(l) + CO2(g)
More examples of neutralization equations are given in section 4.
METHOD (a) procedure for making a soluble salt by neutralising a soluble base (alkali) with an acid. You need to know the exact amount of acid to just neutralise completely the alkali (soluble base). (1) A known volume of acid is pipetted into a conical flask and universal indicator added. The acid is titrated with the alkali from the burette. (2) The acid is added until the indicator turns green, pH 7 neutral. This means all the acid has been neutralised to form the salt. I've illustrated the method using universal indicator BUT it isn't that accurate an indicator for titrations. You should use a more precise indicator like phenolphthalein or methyl orange. I didn't repeat all the titration details here again, I've just kept to the basic ideas and description, but there lots of detailed examples on the page ( more examples - diagrams, descriptions of titration procedures) (3) The volume of alkali needed for neutralisation is then noted, this is called the endpoint volume. (1)-(3) are repeated with both known volumes mixed together BUT without the contaminating indicator, such as phenolphthalein or methyl orange. (4) The solution is transferred to an evaporating dish and heated to partially evaporate the water causing crystallisation or can be left to very slowly evaporate - which tends to give bigger and better crystals. (5) The residual liquid can be decanted away and the crystals can be carefully collected and dried by 'dabbing' with a filter paper OR the crystals can be collected by filtration and dried (as above). See also the preparation of ammonium salts using this method. AND also more examples - diagrams, descriptions of titration procedures Note (i) You can put the acid in the burette and the alkali in the flask. (ii) Parts (1) to (3) are known specifically as an acid-base (alkali) titration, and the general method is known as a volumetric titration by which it possible to find out exactly what volume ratios are needed for neutralisation. So knowing one concentration, you can calculate the other. (iii) Concentration calculations are on calculations pages sections 11. and 12. (iv) Apparatus used: (1) pipette and conical flask; (2)-(3) burette and conical flask; (4) evaporating (crystallising) dish, bunsen burner, tripod and gauze; (5) filter paper. (v) Other indicators e.g. phenolphthalein can be used instead (pink alkaline, colourless acid). (vi) The burette and pipette are both used for the accurate measurement of volume.
(vii)
The pH changes in this preparation are described in section 7
6b. A 2nd Method of Making a Water Soluble Salt 6b. METHOD (b) Reacting an acid with a metal or with an insoluble base to give a soluble salt e.g. an insoluble base such as a metal oxide, metal hydroxide or a metal carbonate, often of a Group 2 metal like calcium, magnesium or a Transition Metal like nickel, copper or zinc. Copper metal won't dissolve in acids, but its oxide and carbonate will. Using the same procedure you can also start with a metal that has a low reactivity towards water e.g. magnesium, zinc or iron. One important point is to recognise that all the reactants are soluble here, which is why you need a titration procedure to work out how much of the acid is to be added to a given volume of alkali. Typical common insoluble bases used for preparing soluble salts: MgO magnesium oxide, MgCO3 magnesium carbonate, CaO Calcium oxide, CaCO3 calcium carbonate, Ca(OH)2 calcium hydroxide, NiO nickel(II) oxide, ZnO zinc oxide, Zn(OH)2, zinc hydroxide, ZnCO3 zinc carbonate, CuO copper(II) oxide, CuCO3 copper(II) carbonate, PbCO3lead(II) carbonate (with nitric acid to make lead(II) nitrate), FeCO3 iron(II) carbonate (to make iron(II) salts), MnCO 3 manganese(II) carbonate Typical examples shown by the word and symbol equations below include ... copper(II) oxide + sulphuric acid ==> copper(II) sulphate + water CuO + H2SO4 ==> CuSO4 + H2O CuO(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) magnesium hydroxide + sulfuric acid ==> magnesium sulfate + water Mg(OH)2 + H2SO4 ==> MgSO4 + 2H2O
Mg(OH)2(s) + H2SO4(aq) ==> MgSO4(aq) + 2H2O(l)
magnesium hydroxide + hydrochloric acid ==> magnesium chloride + water Mg(OH)2(s) + 2HCl(aq) ==> MgCl2(aq) + 2H2O(l)
Zinc carbonate + nitric acid ==> zinc nitrate + water + carbon dioxide ZnCO3 + 2HNO3 ==> Zn(NO3)2 + H2O + CO2 ZnCO3(s) + 2HNO3(aq) ==> Zn(NO3)2(aq) + H2O(l) + CO2 (g)
zinc oxide + hydrochloric acid ==> zinc chloride + water ZnO(s) + 2HCl(aq) ==> ZnCl2(aq) + H2O(l) Similar for many other Group 2 and Transition metal oxides e.g. Mg, Ca, Ba and Co, Ni, Cu instead of Zn
zinc + sulfuric acid ==> zinc sulfate + hydrogen Zn + H2SO4 ==> ZnSO4 + H2 Zn(s) + H2SO4(aq) ==> ZnSO4(aq) + H2(g)
calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide CaCO3(s) + 2HCl(aq) ==> CaCl2(aq)+ H2O(l) + CO2(g) Similar for many other Group 2 and Transition metal carbonates e.g. Mg, Sr, Ba and Ni, Co, Zn instead of Ca
Carbonates are frequently used in this method of salt making, e.g. using copper carbonate to make copper salts
o
The equations are give with, and without sate symbols.
copper(II) carbonate + hydrochloric acid ==> Copper(II) chloride + water + carbon dioxide
o
CuCO3 + 2HCl ==> CuCl2 + H2O + CO2
o
and with sulphuric acid a blue solution of copper(II) sulphate is formed.
copper(II) carbonate + sulphuric acid ==> Copper(II) sulphate + water + carbon dioxide
o
CuCO3 + H2SO4 ==> CuSO4 + H2O + CO2
CuCO3(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) + CO2(g)
copper(II) carbonate + nitric acid ==> Copper(II) nitrate + water + carbon dioxide
o
CuCO3 + 2HNO3 ==> Cu(NO3)2 + H2O + CO2
CuCO3(s) + 2HCl(aq) ==> CuCl2(aq) + H2O(l) + CO2(g)
CuCO3(s) + 2HNO3(aq) ==> CuSO4(aq) + H2O(l) + CO2(g)
Similar equations for other carbonates to give soluble salts which can be crystallised from solution e.g.
o
calcium carbonate CaCO3, to make two salts - calcium chloride/nitrate (calcium sulfate is not very soluble)
o
iron(II) carbonate FeCO3, to make three salts - iron(II) chloride/sulfate/nitrate
o
magnesium carbonate MgCO3, to make three salts - magnesium chloride/sulfate/nitrate
o
manganese(II) carbonate MnCO3, to make three salts - manganese(II) chloride/sulfate/nitrate
o
zinc carbonate ZnCO3, to make three salts - zinc chloride/sulfate/nitrate
o
lead(II) carbonate PbCO3, only nitric acid to make lead(II) nitrate, lead(II) chloride and lead(II) sulfate are insoluble and must be prepared by method (c)
More examples of neutralization equations are given in section 4. METHOD (b) Procedure for making a soluble salt from an insoluble base, carbonate or metal
(1) The required volume of acid is measured out into the beaker with a measuring cylinder. The excess of insoluble metal, oxide, hydroxide or carbonate is weighed out and the solid added in small portions to the acid in the beaker with stirring. Doing the weighing will minimise trial and error especially if the reaction is slow, as long as you know how to do the theoretical calculation! (2) The mixture may be heated to speed up the reaction. When no more of the solid dissolves it means ALL the acid is neutralised and there should be a little excess solid. You should see a residue of the solid (oxide, hydroxide, carbonate) left at the bottom of the beaker. (3) The hot solution (with care!) is filtered to remove the excess solid metal/oxide/carbonate, into an evaporating dish. (4) You may need to carefully heat the solution to evaporate some of the water. Then hot solution is left to cool and crystallise. After crystallisation, you collect and dry the crystals with a filter paper. Note (i) Apparatus used: (1) balance, measuring cylinder, beaker and glass stirring rod. (2) beaker/rod, bunsen burner, tripod and gauze; (3)-(4) filter funnel and filter paper, evaporating (crystallising) dish. (ii) A measuring cylinder is adequate for measuring the acid volume, you do not need the accuracy of a pipette or burette required in method (a).
(iii) How to calculate amounts required and % yield is dealt with in 14
Chemical Calculations Part
Salt solubility affects the method you choose to make a salt and so section 8. contains tables of information-data on salt solubility which will help you decide on the method to prepare a salt.
Multiple choice revision quizzes and other worksheets
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Advanced Level Chemistry Students Acid-Base Revision Notes - use index
Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel Science Chemistry IGCSE Chemistry & OCR 21st Century Science Chemistry, OCR Gateway Science Chemistry, WJEC gcse science chemistry CCEA/CEA gcse science chemistry (revise science chemistry courses equal to US grade 8, grade 9 grade 10. Revision notes on acids, bases, alkalis, salts, solution pH word equations balanced symbol equations science chemistry courses revision guides How to make copper(II) sulfate (copper sulphate), how to make zinc sulfate (zinc sulphate)
ALL Website content copyright © Dr Phil Brown 2000-2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * [email protected] Describing how to make a salt by four different methods, how to do it, what laboratory equipment is needed and how to crystallise or collect the salt in the end 6c. Method (c) Preparing an Insoluble Salt Procedure for making an insoluble salt my mixing solutions of soluble compounds to form a precipitate. The two soluble compounds must each provide one of the constituent ions of the desired insoluble salt which precipitates out when the solutions are mixed. NOTE definition: A precipitation reaction is generally defined as 'the formation of an insoluble solid on mixing two solutions of soluble substances or bubbling a gas into a solution'. Note that several precipitation reactions are used as simple tests for e.g. for sulfate, chloride, bromide, iodide Salt solubility affects the method you choose to make a salt and so 8. contains tables of informationdata on salt solubility which will help you decide on the method to prepare a salt, but a brief summary is given below
A solubility guide for salts - information required to decide on the method used to prepare a salt
salts
solubility?
common salts of sodium, potassium and ammonium ions
usually soluble in water
common sulfates (sulphates)
usually quite soluble except for calcium sulfate (slightly soluble), lead sulfate and barium sulfate are both insoluble
common chlorides (similar rule for bromides and iodides)
usually soluble except for insoluble lead(II) chloride and silver chloride
common nitrates
all soluble
common carbonates
most metal carbonates are insoluble apart from sodium & potassium carbonate. Ammonium carbonate is also soluble.
How can we make an insoluble salt? How do we prepare an insoluble salt from two soluble compounds?
This section describes the preparation of insoluble salts like silver chloride AgCl, lead(II) chloride (lead chloride) PbCl2, lead(II) iodide (lead iodide) PbI2, calcium carbonate CaCO3, barium sulfate (barium sulphate) BaSO4, lead(II) sulfate (lead sulphate, lead sulfate)PbSO4, and 'slightly soluble' calcium sulfate (calcium sulphate) CaSO4, which can all be made by a precipitation reaction.
o
All of the above insoluble salts are white (as in diagram), except lead(II) iodide which is yellow.
Many of the salt precipitate are WHITE, but lead iodide is pale yellow.
METHOD (c)
An insoluble salt can be made by mixing two solutions of soluble salts in a process is called precipitation.
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The method is quite simple – illustrated above, assuming in this case the insoluble salt is colourless–white.
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One solution contains the 1st required ion, and the other solution contains the 2nd required ion.
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So you must prepare two solutions of soluble compounds, each of which provides one of the two ions required to combine and precipitate out as the insoluble salt.
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Each soluble compound is weighed out into its own beaker and dissolved in a suitable volume of water until the solutions are both quite clear.
One solution is then poured into the other, order doesn't really matter.
The two solutions of SOLUBLE compounds must be thoroughly mixed together to ensure all the reactants are used up, so the maximum amount of INSOLUBLE salt precipitate is formed.
You see the two clear solutions on mixing forming a cloudy mixture as the insoluble compound is formed, known as theprecipitate.
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The mixture is then carefully poured into a funnel holding a filter paper.
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The precipitated salt can then be filtered off with the filter funnel and paper.
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While still in the filter paper and funnel, the collected solid precipitate is washed with distilled/deionised water to remove any remaining soluble salt impurities and just the damp, but otherwise pure, insoluble salt is left.
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The precipitate is then carefully removed from the filter paper into a clean dish or basin to be dried e.g. left out in a dry room or warmed in a pre–heated oven.
Examples ...
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(i) Silver chloride is made by mixing solutions of solutions of silver nitrate and sodium chloride.
silver nitrate + sodium chloride ==> silver chloride + sodium nitrate
AgNO3(aq) + NaCl(aq) ==> AgCl(s) + NaNO3(aq)
in of ions it could be written as
Ag+NO3–(aq) + Na+Cl–(aq) ==> AgCl(s) + Na+NO3–(aq)
or: Ag+(aq) + NO3–(aq) + Na+(aq) + Cl–(aq) ==> AgCl(s) + Na+(aq) + NO3–(aq)
but the spectator ions are nitrate NO3– and sodium Na+ which do not change at all,
The 'active ions' and resulting precipitate are highlighted in yellow.
so the proper ionic equation is simply: Ag+(aq) + Cl–(aq) ==> AgCl(s)
Note (i) the use of state symbols (aq) and (s) AND
(ii) that ionic equations omit ions that do not change there chemical or physical state.
In this case the nitrate, NO3–(aq) and sodium Na+(aq) ions do not change physically or chemically and are called spectator ions,
BUT the aqueous silver ion, Ag +(aq), combines with the aqueous chloride ion, Cl–(aq), to form the insoluble salt silver chloride, AgCl (s), thereby changing their states both chemically and physically.
More Ionic equations explained with all spectator ions indicated
If you use barium chloride the word and symbol equations are ...
barium chloride + silver nitrate ==> silver chloride + barium nitrate
BaCl2(aq) + 2AgNO3(aq) ==> 2AgCl(s) + Ba(NO3)2(aq)
which can be written as
Ba2+(aq) + 2Cl–(aq) + 2Ag+(aq) + 2NO3–(aq) ==> 2AgCl(s) + Ba2+(aq) + 2NO3–(aq)
the spectator ions are Ba2+ and NO3–
so the ionic equation is: Ag+(aq) + Cl–(aq) ==> AgCl(s)
o
(ii) Lead(II) iodide, a yellow precipitate (insoluble in water!) can be made by mixing lead(II) nitrate solution with e.g. potassium iodide solution.
lead(II) nitrate + potassium iodide ==> lead(II) iodide + potassium nitrate
Pb(NO3)2(aq) + 2KI(aq) ==> PbI2(s) + 2KNO3(aq)
which can be written as
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2I–(aq) ==> PbI2(s) + 2K+(aq) + 2NO3–(aq)
the ionic equation is: Pb2+(aq) + 2I–(aq) ==> PbI2(s)
because the spectator ions are nitrate NO3– and potassium K+.
In a similar way you can make lead(II) chloride by e.g. using dilute hydrochloric acid
lead(II) nitrate + hydrochloric acid ==> lead(II) chloride + nitric acid
Pb(NO3)2(aq) + 2HCl(aq) ==> PbCl2(s) + 2HNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2H+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2H+(aq) + 2NO3–(aq)
the proper ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)
because the spectator ions are nitrate NO3– and hydrogen H+.
You can make lead(II) bromide by e.g. using sodium bromide solution
lead(II) nitrate + sodium bromide ==> lead(II) bromide + sodium nitrate
Pb(NO3)2(aq) + 2NaBr(aq) ==> PbBr2(s) + 2NaNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2Na+(aq) + 2Br–(aq) ==> PbBr2(s) + 2Na+(aq) + 2NO3–(aq)
the proper ionic equation is: Pb2+(aq) + 2Br–(aq) ==> PbBr2(s)
because the spectator ions are nitrate NO3– and sodium Na+.
and you can make lead(II) chloride by e.g. using sodium chloride solution
o
o
lead(II) nitrate + sodium chloride ==> lead(II) chloride + sodium nitrate
Pb(NO3)2(aq) + 2NaCl(aq) ==> PbBr2(s) + 2NaNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2Na+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2Na+(aq) + 2NO3–(aq)
the proper ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)
because the spectator ions are nitrate NO3– and sodium Na+.
You can also use lead nitrate and dilute hydrochloric acid to precipitate lead chloride
lead(II) nitrate + hydrochloric acid ==> lead(II) chloride + nitric acid
Pb(NO3)2(aq) + 2HCl(aq) ==> PbCl2(s) + 2HNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2H+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2H+(aq) + 2NO3–(aq)
so the ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)
because the spectator ions are nitrate NO3– and hydrogen H+.
(iii) Calcium carbonate, a white precipitate, forms on e.g. mixing calcium chloride solution and sodium carbonate solutions ...
calcium chloride + sodium carbonate ==> calcium carbonate + sodium chloride
CaCl2(aq) + Na2CO3(aq) ==> CaCO3(s) + 2NaCl(aq)
Ca2+(aq) + 2Cl–(aq) + 2Na+(aq) + CO32–(aq) ==> CaCO3(s) + 2Na+(aq) + 2Cl–(aq)
the ionic equation is: Ca2+(aq) + CO32–(aq) ==> CaCO3(s)
because the spectator ions are chloride Cl– and sodium Na+.
(iv) Barium sulphate (barium sulphate), a white precipitate, forms on mixing e.g. barium chloride solution and dilute sulphuric acid ...
barium chloride + sulphuric acid ==> barium sulphate + hydrochloric acid
o
BaCl2(aq) + H2SO4(aq) ==> BaSO4(s) + 2HCl(aq)
Ba2+(aq) + 2Cl–(aq) + 2H+(aq) + SO42–(aq) ==> BaSO4(s) + 2H+(aq) + 2Cl–(aq)
the ionic equation is: Ba2+(aq) + SO42–(aq) ==> BaSO4(s)
because the spectator ions are chloride Cl– and hydrogen H+.
Or you can use sulphate salts like sodium sulphate, so the word and symbol equations are ..
barium chloride + sodium sulfate ==> barium sulfate + sodium chloride
BaCl2(aq) + Na2SO4(aq) ==> BaSO4(s) + 2NaCl(aq)
The ionic equation is the same: Ba2+(aq) + SO42–(aq) ==> BaSO4(s)
because the spectator ions are sodium Na+ and chloride Cl–
You can do exactly the same preparation using potassium sulfate and barium nitrate or barium chloride, basically any solutions of a soluble barium salt and a soluble sulfate salt can be used to prepare barium sulfate.
e.g. barium nitrate + potassium sulfate ==> barium sulfate + potassium nitrate
or barium chloride + potassium sulfate ==> barium sulfate + potassium chloride
See at the end of the page the uses of barium sulfate.
(v) Lead(II) sulphate (lead sulphate), a white precipitate, forms in a similar way e.g.
lead(II) nitrate + sodium sulphate ==> lead(II) sulphate + sodium nitrate
Pb(NO3)2 (aq) + Na2SO4(aq) ==> PbSO4(s) + 2NaNO3 (aq)
ionic equation: Pb2+(aq) + SO42–(aq) ==> PbSO4(s)
because the spectator ions are sodium Na+ and nitrate NO3–
Lead(II) sulphate can also be precipitated using dilute sulfuric acid.
o
lead(II) nitrate + sodium sulphate ==> lead(II) sulphate + nitric acid
Pb(NO3)2 (aq) + H2SO4(aq) ==> PbSO4(s) + 2HNO3 (aq)
ionic equation: Pb2+(aq) + SO42–(aq) ==> PbSO4(s)
because the spectator ions are hydrogen H+ and nitrate NO3– ions
(vi) Calcium sulphate (calcium sulphate), a white precipitate, forms on mixing e.g. calcium chloride solution and dilute sulphuric acid ...
calcium chloride + sulphuric acid ==> barium sulphate + hydrochloric acid
CaCl2(aq) + H2SO4(aq) ==> CaSO4(s) + 2HCl(aq)
Ca2+(aq) + 2Cl–(aq) + 2H+(aq) + SO42–(aq) ==> CaSO4(s) + 2H+(aq) + 2Cl–(aq)
the ionic equation is: Ca2+(aq) + SO42–(aq) ==> CaSO4(s)
because the spectator ions are chloride Cl– and hydrogen H+.
Or you can use sulphate salts like sodium sulphate, so the word and symbol equations are ..
calcium chloride + sodium sulfate ==> barium sulfate + sodium chloride
CaCl2(aq) + Na2SO4(aq) ==> CaSO4(s) + 2NaCl(aq)
The ionic equation is the same: Ba2+(aq) + SO42–(aq) ==> BaSO4(s)
because the spectator ions are sodium Na+ and chloride Cl–
You can do exactly the same preparation using potassium sulfate and calcium nitrate or calcium chloride, basically any solutions of a soluble calcium salt and a soluble sulfate salt can be used to prepare calcium sulfate.
The yield is a little lower than the theoretical because calcium sulfate is slightly soluble in water.
See at the end of the page the uses of barium sulfate.
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NOTE reminder definition: A precipitation reaction is generally defined as 'the formation of an insoluble solid on mixing two solutions or bubbling a gas into a solution'.
General rules which describe the solubility of common types of compounds in water:
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All common sodium, potassium and ammonium salts are soluble e.g. NaCl, K 2SO4, NH4NO3
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All nitrate salts are soluble e.g. NaNO3, Mg(NO3)2, Al(NO3)3, NH4NO3
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Some ethanoate salts are soluble e.g. CH 3COONa
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Common chloride salts are soluble except those of silver and lead e.g.
o
o
soluble: KCl, CaCl2, AlCl3 or insoluble AgCl, PbCl2
Common sulfates are soluble except those of lead, barium and calcium: soluble e.g.
soluble: Na2SO4, MgSO4, Al2(SO4)3
insoluble: PbSO4, BaSO4, CaSO4 is slightly soluble.
Common oxides, hydroxides and carbonates are usually insoluble (e.g. Group 2 and Transition Metals) except those of the Group 1 Alkali Metals sodium, potassium etc. and ammonium:
soluble bases–alkalis oxides, hydroxides or carbonates: K 2O, KOH, NaOH, NH3(aq), Na2CO3, (NH4)2CO3
insoluble bases – basic oxides, hydroxides or insoluble carbonates: MgO, CuO, ZnO, Mg(OH)2, Fe(OH)2, Cu(OH)2, CuCO3, ZnCO3, CaCO3
Some uses of insoluble salts (their preparation has been described above)
o
Calcium sulfate CaSO4 (calcium sulphate) is used in plaster of Paris and plaster for domestic wall covering.
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Barium sulfate BaSO4 (barium sulphate)
Barium sulfates quite a dense material is used in with X-rays for particular medical examinations.
Barium sulfate, like all barium salts is toxic, BUT, because it is insoluble, non of it is absorbed by the body into the bloodstream and eventually it will right through the gut system and be expelled in the normal faeces.
X-rays are used to investigate bone structure e.g. fractured or broken bones and the technique works because bone material is too dense to let weak X-rays through, so on shooting an X-ray photograph you get the bone structure as a sort of shadow effect where the X-rays have been absorbed by the bone.
Therefore, normally you can't use X-rays to examine soft tissue areas like the gut.
However, if the patient takes a 'barium meal', a thick harmless fluid containing a suspension of the insoluble barium sulfate, this can through the stomach and into the gut.
Therefore it is then possible to X-ray the intestinal gut system because the barium sulfate absorbs the X-rays just like bone.
So, because barium sulfate is opaque to X-rays, the X-ray photograph via the barium meal will highlight the structure of the tissue lining of the gut and show up any structural abnormalities and blockages.
6d. Method (d) Making a salt by direct combination of elements Sometimes it isn't appropriate to prepare a soluble salt by reacting an acid with an insoluble base or alkali, so it may be possible to prepare the salt by directly combining the metal and the non-metal elements. Two such examples are the preparation of anhydrous aluminium chlorideand anhydrous iron(III) chloride (anhydrous here means without any water of crystallisation).
Preparation of aluminium chloride AlCl3
Preparation of iron(III) chloride FeCl3
How can we make aluminium chloride? How do we prepare iron(III) chloride?
METHOD (d) both preparations illustrated above.
These compounds can be made by direct combination of the elements to form anhydrous salts e.g. if dry chlorine gas Cl2 is ed over heated iron or aluminium, the chloride is produced. These experiment preparations (shown above) should be done very carefully by the teacher in a fume cupboard.
o
o
The aluminium can burn intensely with a violet flame, white fumes of aluminium chloride sublime from the hot reacted aluminium and the white solid forms on the cold surface of the flask.
aluminium + chlorine ==> aluminium chloride
2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
The aluminium chloride is often discoloured yellow from the trace chlorides of copper or iron that may be formed from traces of these metals that might be present in the original aluminum).
The iron (e.g. as steel wool) glows red and brown fumes of iron(III) chloride stream off, the brown solid collects on the cold flask surface.
iron + chlorine ==> iron(III) chloride
2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)
o
Note (i): Both these chlorides react exothermically and hydrolyse with water to give the metal hydroxide and fumes of hydrogen chloride, and so dry conditions are needed.
o
Note (ii): Both these chlorides cannot be made in an anhydrous form from aqueous solution neutralisation. This is because on evaporation the compounds contain 'water of crystallisation'. On heating the hydrated salt hydrolyses and decomposes into water, the oxide or hydroxide and fumes of hydrogen chloride, and maybe some impure anhydrous chloride, basically it a mess in of trying to make pure AlCl 3 and FeCl3 in this way.
Multiple choice revision quizzes and other worksheets
GCSE/IGCSE foundation-easier multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE higher-harder multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE Structured question worksheet on Acid Reaction word equations and symbol equation questions
Word equation answers and symbol equation answers)
GCSE/IGCSE word-fill worksheet on Acids, Bases, Neutralisation and Salts
GCSE/IGCSE matching pair quiz on Acids, Bases, Salts and pH
See also
Advanced Level Chemistry Students Acid-Base Revision Notes - use index
Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel Science Chemistry IGCSE Chemistry & OCR 21st Century Science Chemistry, OCR Gateway Science Chemistry, WJEC gcse science chemistry CCEA/CEA gcse science chemistry (revise science chemistry courses equal to US grade 8, grade 9 grade 10. Revision notes on acids, bases, alkalis, salts, solution pH word equations balanced symbol equations science chemistry courses revision guides
ALL Website content copyright © Dr Phil Brown 2000-2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * [email protected] Describing how to make a salt by four different methods, how to do it, what laboratory equipment is needed and how to crystallise or collect the salt in the end
6d. Method (d) Making a salt by direct combination of elements Sometimes it isn't appropriate to prepare a soluble salt by reacting an acid with an insoluble base or alkali, so it may be possible to prepare the salt by directly combining the metal and the non-metal elements. Two such examples are the preparation of anhydrous aluminium chlorideand anhydrous iron(III) chloride (anhydrous here means without any water of crystallisation).
Preparation of aluminium chloride AlCl3
Preparation of iron(III) chloride FeCl3
How can we make aluminium chloride? How do we prepare iron(III) chloride?
METHOD (d) both preparations illustrated above.
These compounds can be made by direct combination of the elements to form anhydrous salts e.g. if dry chlorine gas Cl2 is ed over heated iron or aluminium, the chloride is produced. These experiment preparations (shown above) should be done very carefully by the teacher in a fume cupboard.
o
The aluminium can burn intensely with a violet flame, white fumes of aluminium chloride sublime from the hot reacted aluminium and the white solid forms on the cold surface of the flask.
aluminium + chlorine ==> aluminium chloride
2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
o
The aluminium chloride is often discoloured yellow from the trace chlorides of copper or iron that may be formed from traces of these metals that might be present in the original aluminum).
The iron (e.g. as steel wool) glows red and brown fumes of iron(III) chloride stream off, the brown solid collects on the cold flask surface.
iron + chlorine ==> iron(III) chloride
2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)
o
Note (i): Both these chlorides react exothermically and hydrolyse with water to give the metal hydroxide and fumes of hydrogen chloride, and so dry conditions are needed.
o
Note (ii): Both these chlorides cannot be made in an anhydrous form from aqueous solution neutralisation. This is because on evaporation the compounds contain 'water of crystallisation'. On heating the hydrated salt hydrolyses and decomposes into water, the oxide or hydroxide and fumes of hydrogen chloride, and maybe some impure anhydrous chloride, basically it a mess in of trying to make pure AlCl 3 and FeCl3 in this way.
Multiple choice revision quizzes and other worksheets
GCSE/IGCSE foundation-easier multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE higher-harder multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts
GCSE/IGCSE Structured question worksheet on Acid Reaction word equations and symbol equation questions
Word equation answers and symbol equation answers)
GCSE/IGCSE word-fill worksheet on Acids, Bases, Neutralisation and Salts
GCSE/IGCSE matching pair quiz on Acids, Bases, Salts and pH
See also
Advanced Level Chemistry Students Acid-Base Revision Notes - use index
Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel Science Chemistry IGCSE Chemistry & OCR 21st Century Science Chemistry, OCR Gateway Science Chemistry, WJEC gcse science chemistry CCEA/CEA gcse science chemistry (revise science chemistry courses equal to US grade 8, grade 9 grade 10. Revision notes on acids, bases, alkalis, salts, solution pH word equations balanced symbol equations science chemistry courses revision guides
ALL Website content copyright © Dr Phil Brown 2000-2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * [email protected] Describing how to make a salt by four different methods, how to do it, what laboratory equipment is needed and how to crystallise or collect the salt in the end
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