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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

(c) doc b

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

 

A non–chemical test method for identifying elements – atomic emission line spectroscopy
 
An instrumental method for METALS from LINE SPECTRA

If the atoms of an element are heated to a very high temperature they emit light of a specific set of frequencies (or wavelengths). These are all due to electronic changes in the atoms, the electrons are excited and then lose energy by emitting energy as photons of light. These emitted frequencies can be recorded on a photographic plate, or these days a digital camera.

Each emission line spectra is unique for each element and so offers a different pattern of lines i.e. a 'spectral fingerprint' by which to identify any element in the periodic table .e.g. the diagram on the left shows some of the visible emission line spectra for the elements hydrogen, helium, neon, sodium and mercury.

Note the double yellow line for sodium, hence the dominance of yellow in its flame test colour. In fact the simple flame test colour observations for certain metal ions relies entirely on the observed amalgamation of these spectral lines.

This is an example of an instrumental chemical analysis called spectroscopy and is performed using an instrument called an optical spectrometer (simple ones are called spectroscopes). This method, called spectroscopy, is a fast and reliable method of chemical analysis. This type of optical spectroscopy has enabled scientists to discover new elements in the past and today identify elements in distant stars and galaxies. The alkali metals caesium (cesium) and rubidium were discovered by observation of their line spectrum and helium identified from spectral observation of our Sun.

 

Some metal ions (cations) can be identified by the formation of coloured precipitates with sodium hydroxide solution

The non–metallic ion, the ammonium ion, can be detected with the same reagent, sodium hydroxide, because ammonia gas is released, especially if the mixture is gently warmed.

A few drops of sodium hydroxide solution are added to a solution of the salt under investigation to see if any precipitate (insoluble solid) is formed, and, from the observations e.g.

Metal ion detected colour of precipitate with NaOH ionic equation for the reactions
calcium, Ca2+

colourless

white precipitate Ca2+(aq) + 2OH(aq) ==> Ca(OH)2(s)
copper(II), Cu2+

blue

blue precipitate (3 in diagram below) Cu2+(aq) + 2OH(aq) ==> Cu(OH)2(s)
iron(II), Fe2+

pale green

dark green precipitate (1 in diagram below) Fe2+(aq) + 2OH(aq) ==> Fe(OH)2(s)
iron(III), Fe3+

orange

orange–brown precipitate (2 in diagram below) Fe3+(aq) + 3OH(aq) ==> Fe(OH)3(s)
zinc, Zn2+

colourless

white precipitate (4a in diagram below), which dissolves in excess to give a clear colourless solution  (4b in diagram below) Zn2+(aq) + 2OH(aq) ==> Zn(OH)2(s)

Zn(OH)2(s) + 2OH(aq) ==> Zn(OH)4]2–(aq)

aluminium, Al3+

colourless

white precipitate, which dissolves in excess, to give a clear colourless solution  (same as zinc ion, 4a + 4b in diagram below)

Al3+(aq) + 3OH(aq) ==> Al(OH)3(s)

Al(OH)3(s) + OH(aq) ==> [Al(OH)4](aq)

ammonium, NH4+

colourless

no precipitate formed, but ammonia gas released which you can smell, the gas turns damp red litmus paper blue NH4+(aq) + OH(aq) ==> H2O(l) + NH3(g)
**************************** ***************************************************************************** **********************************************************

The above reactions are illustrated in the diagram below

 

For more on ionic equations see How to write equations and Making salts by precipitation


Tests for NON–METAL IONS – anions (negative ions)

 

Tests to detect and identify halide ions X,  the negative ions (anions) formed from the halogens, chloride, bromide and iodide.

To the suspected halide ion solution add a little dil. nitric acid and a few drops of silver nitrate solution.

Depending on the halide ion you get a different coloured silver halide precipitate, summarised below.

halide ion Colour of precipitate with silver nitrate Ionic equation to show precipitate formation
chloride Cl white precipitate of AgCl silver chloride (slowly darkens when exposed to light) Ag+(aq) + Cl(aq) ==> AgCl(s)
bromide Br cream precipitate of AgBr silver bromide Ag+(aq) + Br(aq) ==> AgBr(s)
Iodide I yellow precipitate of  AgI silver iodide  Ag+(aq) + I(aq) ==> AgI(s)

The silver nitrate tests for halide ions is illustrated in the diagram below.

You can only use this silver nitrate test on soluble chlorides.

 

test for CO2Test for the carbonate ion CO32–

Addition of dil. hydrochloric acid to any carbonate or hydrogen carbonate results in fizzing! The effervescence is due to the evolution of carbon dioxide gas. If a sample of the evolved gas is carefully collected and bubbled into limewater a white precipitate is formed. The formation of the carbon dioxide confirms the original compound was a carbonate.  It doesn't matter whether the compound is soluble or insoluble.

carbonate + acid ==> salt + water + carbon dioxide

The ionic equation is

CO32–(s) + 2H+(aq) ==> H2O(l) + CO2(g)

 

 

Test for the sulfate ion SO42–

The suspected sulfate is dissolved in water. A little dilute hydrochloric acid added followed by a few drops of barium chloride solution. If a sulfate is present a white precipitate of barium sulfate is formed.

barium ion + sulfate ion ==> barium sulfate

Ba2+(aq) + SO42–(aq) ==> BaSO4(s)

This can only be done on soluble sulfate compounds.

 


Quiz on identifying salts and other compounds

ALL chemical tests for GCSE/IGCSE/A Level etc.

GCSE/IGCSE Acids & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids

See also Advanced Level Chemistry Students Acid–Base Revision Notes – use index

6. METHODS of MAKING SALTS – salt preparation procedures, now on four separate pages


6a. METHOD (a) Neutralising a soluble acid with a soluble base (alkali) to give a soluble salt


6b. METHOD (b) Reacting an acid with a metal or with an insoluble base to give a soluble salt


6c. Method (c) Preparing an Insoluble Salt


6d. Method (d) Making a salt by direct combination of elements


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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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