|
GCSE level chemistry exam revision notes on
acids, alkalis, salts etc.
6. Index of METHODS OF MAKING SALTS and a summary of TESTS
for ions and gases[Author
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pH-reactions-acids-bases-neutralisation-salt preps,
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GCSE level easier m/c QUIZ on pH, Indicators, Acids,
Bases, Neutralisation and Salts
GCSE level harder m/c QUIZ on pH, Indicators, Acids,
Bases, Neutralisation and Salts
Index of all GCSE level (~US grades 8-10) notes on acids, bases
and salts
Ion/gas test section added because many of the tests are associated with salt-like compounds
Sub-index for this page on methods of making salts and tests for ions
(1)
What methods can we use for making
salts? and is the salt soluble or insoluble?
(a)
Making a soluble salt by neutralising a soluble acid with a soluble base (alkali)
(b)
Making a soluble salt by from an acid with a metal/insoluble base
(c) Preparing an
insoluble salt by mixing solutions of two soluble compounds
(d) Making a
salt by directly combining its constituent elements
(2)
A solubility guide
for salts and other compounds (this affects choice of method)
(3)
Notes on important tests for common gases and ions (cations & anions)
Test details for positive metal ions (cations),
flame tests
Test details for positive metal ions
(cations), hydroxide precipitate formation
Test details for negative 'non-metal' ions
(anions)
Tests for common gases
Quick revision summary of tests
Index of all GCSE level (~US grades 8-10) notes on acids, bases
and salts
What next?
Associated Pages
(1)
What methods can we use for making salts? and is the salt soluble or insoluble?
How do we make salts? What
preparations are available to us?
Four basic methods for preparing
salts are described on 4 separate pages, with annotated diagrams.
BEFORE preparing a salt there
are two important facts to know ...
(i)
Is the salt 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
BUT, check out the solubility of
the salt you want to prepare
|
(2) A solubility guide
for salts and other compounds
Information required to decide on the method used to
prepare a salt |
|
salts and other compounds |
solubility? |
|
common salts of sodium, potassium and
ammonium ions |
usually soluble in water |
|
common sulfates (sulfates) |
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. |
|
common hydroxides |
most metal hydroxides are insoluble apart from
sodium, potassium and ammonium hydroxide |
General rules
which describe the solubility of common types of compounds in water:
-
All common sodium,
potassium and ammonium salts are soluble e.g. NaCl, K2SO4,
NH4NO3
-
All nitrate salts are
soluble e.g. NaNO3, Mg(NO3)2, Al(NO3)3,
NH4NO3
-
Some ethanoate salts are
soluble e.g. CH3COONa
-
Common chloride salts are
soluble except those of silver and lead e.g.
-
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:
-
K2O, 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
Four methods are
described in detail on four separate pages
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
(3) 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.
Quiz on
identifying ions, salts and other compounds
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
|
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 physicsNote
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.
Quiz on
identifying ions, salts and other compounds
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 giving characteristic hydroxide
precipitates
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(aq) |
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 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
|
 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. |
Revision summary of Chemical tests for
cations, anions, and gases
which are core
practical and theoretical content across all major GCSE and IGCSE
chemistry exam boards.
Students must
memorise test procedures, expected observations, and correct ionic
names.
These tests are examinable
across AQA, Edexcel, OCR Gateway, OCR 21st Century, WJEC, CCEA,
and CIE IGCSE syllabi.
Cation Tests (Positive Ions)
Flame Tests
(Group 1 and examples from Group 2 and Transition metals)
|
Ion |
Flame
Colour |
|
Lithium (Li⁺) |
Crimson red |
|
Sodium (Na⁺) |
Yellow/orange |
|
Potassium (K⁺) |
Lilac |
|
Calcium (Ca²⁺) |
Brick red |
|
Copper (Cu²⁺) |
Blue-green |
Tip:
Use a clean platinum/nichrome wire and hydrochloric acid to avoid
contamination.
Sodium Hydroxide Test
(Transition and ammonium ions)
|
Ion |
Observation with NaOH |
|
Ammonium (NH4⁺) |
No precipitate;
pungent gas turns red litmus blue (ammonia) |
|
Copper(II) (Cu²⁺) |
Blue precipitate
(copper(II) hydroxide) |
|
Iron(II) (Fe²⁺) |
Green precipitate
(iron(II) hydroxide) |
|
Iron(III) (Fe³⁺) |
Brown precipitate
(iron(III) hydroxide) |
|
Calcium (Ca²⁺) |
White precipitate
(calcium hydroxide) |
|
Aluminium (Al³⁺) |
White precipitate,
dissolves in excess NaOH |
|
Magnesium (Mg²⁺) |
White precipitate,
insoluble in excess NaOH |
Anion Tests (Negative Ions)
| Ion |
Test |
Observation |
| Carbonate (CO3²⁻) |
Add dilute acid (e.g. HCl) |
Effervescence; gas turns limewater
milky (CO2) |
| Sulfate (SO4²⁻) |
Add dilute HCl, then barium
chloride solution |
White precipitate (barium sulfate) |
| Chloride (Cl⁻) |
Add dilute nitric acid, then
silver nitrate |
White precipitate (silver
chloride) |
| Bromide (Br⁻) |
Same as above |
Cream precipitate (silver bromide) |
| Iodide (I⁻) |
Same as above |
Yellow precipitate (silver iodide) |
Gas Tests
|
Gas |
Test |
Observation |
|
Hydrogen |
Lit splint |
Squeaky pop |
|
Oxygen |
Glowing splint |
Relights |
|
Carbon dioxide |
Bubble through
limewater |
Turns milky/cloudy |
|
Ammonia |
Damp red litmus
paper |
Turns blue
(alkaline gas) |
|
Chlorine |
Damp blue litmus
paper |
Turns red then
bleaches white |
Student Exam Tips
-
Use acidified
reagents: Always add
dilute nitric or hydrochloric acid before silver nitrate or barium
chloride to remove interfering ions.
-
Memorise colours:
Flashcards or colour-coded tables help with flame and precipitate
colours.
-
Practise writing
ionic names: e.g.,
“iron(III)” not “iron 3”.
-
Link to practicals:
Know how to carry out and interpret these tests in lab settings.
Common Misconceptions
-
Mixing up flame
colours: Sodium’s intense
yellow can mask others—test separately.
-
Forgetting gas
tests: Especially ammonia
and chlorine—often overlooked.
-
Confusing white
precipitates: Aluminium,
calcium, and magnesium all form white precipitates—test solubility in
excess NaOH.
-
Skipping acid step:
Omitting acid before silver nitrate or barium chloride leads to
incorrect results.
|
What next?
Associated Pages
GCSE/IGCSE level revision notes on acids, bases,
neutralisation and salts
INDEX of ALL pH, Acids, Alkalis, Neutralisation and 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.
Methods
of making salts index and chemical tests for ions in salts
& tests for common gases
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
advanced Acid–Base Theory and Weak and Strong Acids
See also
Advanced Level Chemistry Students Acid–Base Revision
Notes – use index
Multiple choice revision quizzes and other worksheets
Quiz on identifying ions, salts and other compounds
GCSE level
QUIZ on identifying ions,
salts, gases and other compounds
GCSE level easier multiple choice quiz on pH, Indicators, Acids,
Bases, Neutralisation and Salts
GCSE level harder multiple choice quiz on pH, Indicators, Acids,
Bases, Neutralisation and Salts
GCSE level 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
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