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GCSE Chemistry Notes: Index of methods of salt preparations and chemical tests

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6. Index of METHODS OF MAKING SALTS and a summary of TESTS for ions & gases

Many of the tests are associated with salt-like compounds

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Index of all my GCSE notes on acids, bases and salts

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 These revision notes on methods of making salts and chemical tests for ions (positive cations and negative anions) should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (9–1, 9-5 & 5-1) science courses.

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 soluble salt by neutralising a soluble acid with a soluble base (alkali)

Method (b) Making a soluble salt by from an acid with a metal/insoluble base – oxide, hydroxide, 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 important tests for common gases and ions (cations and anions)


Doc Brown's chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE  chemistry, O level & ~US grades 8, 9, 10 school science courses for ~14-16 year old science students for national examinations in chemistry topics including acids bases alkalis salts preparations reactions


A summary of chemical tests to identify ions in a salt, hence the identity of a salt

For chemical analysis analysts have developed a wide range of qualitative tests to detect specific chemicals which may be molecules or ions. Chemical tests are based on reactions that produce a gas with distinctive properties, or a colour change produced by adding a reagent or the production of an insoluble solid that appears as a precipitate.

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:

Before doing the test a piece of mounted nichrome/platinum wire should be cleaned in concentrated hydrochloric acid and rubbed with fine emery paper, and heated in the hottest part of the flame to check there is no contaminating flame colour.

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!).

It doesn't matter whether the salt compound is soluble or insoluble.

The method can only work if only one metal ion is present - otherwise the colours get mixed up with each other, one colour might mask another, either way you can't be sure which metal ions are present!

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 (or reddish-orange) colour in the flame

the copper ion Cu2+ gives a  blue–green colour in the flame

 

These tests are nice verification of the origin of the bright colours you see in fireworks.

There are various materials you test e.g. seawater or baking soda for sodium, dissolved eggshell for calcium.

Quiz on identifying ions, salts and other compounds

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A non–chemical test method for identifying elements – atomic emission line spectroscopy
 
FLAME EMISSION SPECTROSCOPY - an instrumental method for METALS from LINE SPECTRA

If the atoms of an element are heated to a very high temperature in a flame they emit light of a specific set of frequencies (or wavelengths) called the line spectrum. 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.

Every element atom/ion has its own unique and particular set of electron energies so each emission line spectra is unique for each element (atom/ion) because of a unique set of electron level changes. This produces a different pattern of lines i.e. a 'spectral fingerprint' by which to identify any element in the periodic table .

e.g. the diagram above on the left shows some of the visible emission line spectra for the elements hydrogen, helium, neon, sodium and mercury - all the wavelengths become reference data, either in a book or computer. A modern spectrometer will be linked to a computer system of spectral analysis and database for immediate element identification.

Each line results from a particular electronic energy level change - so each line depends on the electron arrangement of the excited particle, which may be an atom, or an ion of specific charge - the mechanism is illustrated below for the formation of the yellow lines of sodium's line spectra - the excitation can be caused by a very high temperature e.g. in a bunsen flame of the Sun!

For more on theory of light emission from atoms see Electromagnetic spectrum - including excitation of atoms gcse physics

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 yellow spectral lines.

The intensity of the line is a measure of the atom/ion's concentration (see 2nd section on emission spectroscopy below)

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 flame emission spectroscopy, is a fast, reliable, accurate and sensitive (can detect minute traces of elements) 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.

The technique has another important advantage. Because the lines can be accurately measured and each element has characteristic spectral lines, you can analyse mixtures - which I've tried to illustrate with the diagram on the left.

I've superimposed the spectra of hydrogen, helium and neon. Although some lines may overlap, you can easily pick out lines that match one element, but no other element.

From the individual intensities you can analyse a mixture of elements.

 


You can use the flame emission effect to measure the concentration of metal ions in solution.

Using a flame photometer instrument you can do quantitative analysis based on the light emitted from a solution of a metal ion. The intensity of light emission is proportional to the amount of element in the sample and therefore you can measure concentration using flame emission spectroscopy.

The sample is evaporated at high temperature in a flame and the light emitted is measured with a special detector.

You can determine the precise concentration of a metal ion in dilute solution by using a calibration curve (right).

Solutions of known concentration are tested and a measure of the emitted light (flame photometer signal intensity) can be plotted against the concentration to produce a linear calibration curve with an x,y origin of 0,0

Then, a solution of unknown concentration can be tested with the same set-up, and from the emitted light value you can obtain the unknown concentration from the calibration curve.

You can use special light filters to exclude the colour produced by other ions that may be present so improving the accuracy of a specific metal ion measurement.

Many instrumental methods of analysis are available and that these can improve sensitivity, accuracy and speed of tests.  More on instrumental methods of analysis

Quiz on identifying ions, salts and other compounds


Tests for cations - positive metal ions

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

The non–metallic cation, 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.

The above reactions are illustrated in the diagram below

4a also applies to Zn2+, Ca2+, Mg2+, Al3+ , BUT 4b only applies to Zn2+ and Al3+, so watch out!

 

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

colourless

white precipitate (picture 4a) Ca2+(aq) + 2OH(aq) ==> Ca(OH)2(s)
magnesium, Mg2+

colourless

white precipitate (picture 4a)

 

Mg2+(aq) + 2OH(aq) ==> Mg(OH)2(s)

 

copper(II), Cu2+

blue

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

pale green

dark green precipitate (1 in diagram above) 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 above), which dissolves in excess to give a clear colourless solution  (4b in diagram below) (4a) Zn2+(aq) + 2OH(aq) ==> Zn(OH)2(s)

(4b) 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 above

(4a) Al3+(aq) + 3OH(aq) ==> Al(OH)3(s)

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

one non-metal cation - 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)
************************** ************************************************************** ********************************************************

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

Quiz on identifying ions, salts and other compounds


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.

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

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

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

Quiz on identifying ions, salts and other compounds

test for CO2


Test for the carbonate ion CO32–

Addition of dilute 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/hydrogencarbonate + acid ==> salt + water + carbon dioxide

e.g.

sodium carbonate + hydrochloric acid ===> sodium chloride + water + carbon dioxide

Na2CO3  +  2HCl  ===>  2NaCl  +  H2O  +  CO2

The ionic equation is

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

Eggshells and baking soda will give the same reaction!

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 ions, salts and other compounds

Qualitative TESTS for common gases

CHEMICAL TEST FOR TEST METHOD OBSERVATIONS Chemical equation and comments
hydrogen gas H2 a colourless and odourless gas Apply a lit splint or spill. A squeaky pop! (might see condensation on test tube) 2H2(g) + O2(g) ==> 2H2O(l) + energy!
Chemical test for carbon dioxide gas CO2 a colourless and odourless gas

This test is also mentioned above as part of the test for a carbonate or hydrogencarbonate

Bubble the gas into limewater (aqueous calcium hydroxide solution). It turns cloudy – fine milky white precipitate of calcium carbonate. Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l

test for CO2

Chemical test for oxygen gas O2  a colourless and odourless gas Apply a glowing splint or spill. It re–ignites to a flame. C(in wood) + O2(g)  ==> CO2(g)

The relighted splint is mainly combustible carbon.

Chemical test for ammonia gas NH3 colourless gas with a strong pungent odour Damp red litmus.

 

(i) Litmus turns blue.

(ii) Gives white clouds with HCl fumes.

(i) Ammonia is the only common alkaline gas.

(ii) It forms fine ammonium chloride crystals with HCl

Chemical test for chlorine gas Cl2 pungent green gas - irritating and potentially toxic if breathed in - take care! Apply damp blue litmus. (Can use red litmus and just see bleaching effect.) (i) litmus turns red and then is bleached white. The damp litmus initially goes red because chlorine water is acidic in water. Non–metallic chlorine is acid in aqueous solution and a powerful oxidising agent which is why you get the double colour change and its the bleaching effect that distinguishes chlorine for some other acidic gases.

See also Methods of preparing gases - apparatus and reagent chemicals needed


Quiz on identifying ions, 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



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