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
All salts are ionic compounds.
A salt is formed when a metallic ion or an
ammonium ion (NH4+) replaces one or more
hydrogen ions of an acid.
replaced by
HCl
NaCl
replaced by
HNO3
NH4NO3
Find out what is an acid salt!
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Where have you heard of salts?
Salts are essential to animal life in small
quantities, but in large excess will be very
harmful.
Are all salts salty?
Group I ions similar in size to sodium tend to
give salty taste. Which ions do you think give
a salty taste?
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How do you form CaSO4?
Acid + Base  Salt + water
CaSO4
Comes from
base




Comes from acid
One base that can be used is Calcium hydroxide.
One acid that can be used is Sulfuric acid.
Ca(OH)2 + H2SO4  CaSO4 + 2H2O
Can you use calcium carbonate or calcium?
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How do you form NaNO3?
Acid + Base  Salt + Water
NaNO3
Comes from
base



Comes from acid
One base that can be used is Sodium hydroxide.
One acid that can be used is Nitric acid.
NaOH + HNO3  NaNO3 + H2O
Negative ion
Acid used
SO42- (sulfate ion)
Sulfuric acid (H2SO4)
NO3- (nitrate ion)
Nitric acid (HNO3)
Cl- (chloride ion)
Hydrochloric acid (HCl)
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What acids do you use to get the following
salts?
Sodium nitrate
Potassium phosphate
Ammonium ethanoate
Copper(II) iodide
Sodium citrate
Aluminium sulfate
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Many salts combine with water molecules to
form crystals.
These water molecules are called water of
crystallization.
Salts that contain water of crystallization are
called hydrated salts.
Salts that do not contain water of
crystallization are called anhydrous salts.
Name of salt
copper(II) sulfate
Formula of
anhydrous salt
CuSO4
Formula of
hydrated salt
CuSO4.5H2O
magnesium sulfate MgSO4
MgSO4.7H2O
sodium carbonate Na2CO3
Na2CO3.10H2O
zinc sulfate
ZnSO4.7H2O
ZnSO4
What does the ‘dot’ mean?

Heating a hydrated salt removes water of
crystallization.
hydrated copper(II)
sulfate
CuSO4.5H2O
Heat
Heat
anhydrous copper(II) sulfate +
water
CuSO4 + 5H2O

Cobalt(II) chloride
hydrated cobalt(II)
chloride
CoCl2.6H2O
Heat
Heat
anhydrous cobalt(II) chloride + water
CoCl2 + 6H2O
Find out the use of cobalt(II) chloride paper in the lab.
Chlorides/
bromides/
iodides
All are soluble
except:
•
•
lead(II)
chloride/brom
ide/iodide
silver
chloride/brom
ide/iodide
Sulfates
Carbonates
All are soluble
except:
All are NOT
soluble except:
•
barium
sulfate
• Sodium
•
calcium
sulfate
• Potassium
•
lead(II) sulfate
carbonate
carbonate
• Ammonium
carbonate
Sodium salts
Potassium salts
Ammonium salts
All are soluble.
Nitrates
All are soluble.
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There are 3 ways of preparing salts. But to
choose which way depends on:
1. Whether the salt is soluble in water?
2. Whether the starting materials are
soluble in water?
Methods of preparing salts
Is the salt soluble?
Yes
No
Reaction with acids
• Acid + metal
• Acid + base
• Acid + carbonate
Method 3:
Precipitation
Are the starting materials soluble?
No
Yes
Method 1: Reaction of acids with insoluble substances
• Acid + metal
• Acid + base
• Acid + carbonate
1) Filter the
mixture
2) Collect filtrate
Salt solution
Salt crystals (dry with filter paper)
Method 2: Titration
• Acid + alkali
• Acid + carbonate solution
1) Concentrate
2) Crystallize
3) Filter
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Recap:
Acid + Metal oxide/hydroxide  Salt + Water
Acid + Carbonate  Salt + Water + Carbon
dioxide
Acid + Metal  Salt + Hydrogen gas
In Method 1, all the substances in red are
insoluble.
The acid is reacted with an excess of the substances
(metal, carbonate or base).
 Why?

Acid
Insoluble base

Acid + Metal  Salt + Hydrogen gas
ZnSO4
Comes from metal

Comes from acid
Zinc is insoluble in water and reacts with
sulfuric acid. We can use zinc here.
sulfuric acid
Keep adding zinc until no
more effervescence is
observed.
zinc sulfate solution +
unreacted zinc
1) Filter the
mixture
2) Collect filtrate
zinc sulfate crystals
zinc sulfate solution
1) Concentrate
2) Crystallize
3) Filter
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
Only moderately reactive metals like zinc,
aluminum, magnesium and iron can be used.
Not suitable for
1. Very reactive metals such as sodium, potassium
and calcium. Why?
2. Unreactive metals like copper and silver. Why?

Acid + Metal oxide/ hydroxide  Salt + Water
CuSO4
Comes from metal oxide Comes from acid

Copper(II) oxide is insoluble in water and
reacts with sulfuric acid. We can use Copper(II)
oxide here. Why can’t we use copper metal?
sulfuric acid
Keep adding copper(II)
oxide until no more
effervescence is
observed.
Copper(II) sulfate
solution + unreacted
Copper(II) oxide
1) Filter the
mixture
2) Collect filtrate
Copper(II) sulfate
crystals
Copper(II) sulfate solution
1) Concentrate
2) Crystallize
3) Filter

Acid + Carbonate  Salt + Carbon dioxide + Water
MgCl2
Comes from carbonate Comes from acid


Magnesium carbonate is insoluble in water and
reacts with hydrochloric acid.
NOTE: ALL carbonates are insoluble except
potassium, sodium and ammonium carbonate!
hydrochloric acid
Keep adding magnesium
carbonate until no more
effervescence is
observed.
Magnesium chloride
solution + unreacted
magnesium carbonate
1) Filter the
mixture
2) Collect filtrate
Magnesium chloride
crystals
Magnesium chloride solution
1) Concentrate
2) Crystallize
3) Filter
Add the metal/metal carbonate/ base slowly with stirring to hot acid (what acid
depends on what salts you want) until no more dissolves. (This means all the acid is
used up)
Remove the excess metal/metal carbonate/ base by filtering and collect the filtrate.
The filtrate contains the solution of the salt you want.)
Pour the filtrate into an evaporating dish and heat to evaporate most of the water.
This produces a hot saturated solution of the salt. Let the solution cool.
Filter the crystals and dry them by squeezing them between sheets of filter paper.
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Notice that all sodium, potassium and
ammonium salts are SOLUBLE in water.
So you cannot use Method 1 for any of such
salts! Why?
To tell when all the acid has been completely used
up, we have to use titration, by using an indicator.
 What is an indicator?

Indicator
methyl orange
screened methyl
orange
litmus
bromothymol blue
phenolphthalein
Colour in acidic
solution
pH range at which
indicator changes
colour
Colour in alkaline
solution
red
3–5
yellow
violet
3–5
green
red
5–8
blue
yellow
6–8
blue
colourless
8 – 10
pink
V1 cm3
V2 cm3
Fill up a burette with dilute nitric
acid and note down the initial
burette reading (V1 cm3).
Record the final burette reading
3 or two drops of methyl
Add
one
(V
2 cm ). Hence, the volume of
orange
to thefor
NaOH
solution.
acid
required
complete
The solution turns
neutralization
= (V2yellow.
– V1) cm3.
Pipette 25.0 cm3 of dilute sodium
hydroxide into a conical flask.
Add dilute HNO3 from the burette
slowly until the solution turns
orange permanently. This is the
end-point. The acid is all used up.
1.
2.
3.
4.
5.
Pipette 25.0cm3 of NaOH into a beaker. Then
add (V2 – V1) cm3 of dilute nitric acid from
the burette. This time do not add indicator.
Why?
Heat to evaporate the water till it is
saturated.
Cool the saturated solution so that the salt
crystallizes.
Filter to collect the crystals.
Dry the crystals between a few sheets of
filter paper.

Simulation
Fill up a burette with dilute acid (depending on what salt you want) and note down
the initial burette reading (V1 cm3).
Pipette 25.0 cm3 of dilute alkali (depending on what salt you want) into a conical
flask.
Add one or two drops of methyl orange to the alkali solution. The solution turns
yellow. (Note, if you have a strong acid and weak base, you use methyl orange, if
you have a strong base and weak acid, use phenolphthalein. If both are strong, you
can use either indicator.
Add dilute HNO3 from the burette slowly until the solution turns orange
permanently. This is the end-point.
Record the final burette reading (V2 cm3). Hence, the volume of acid required for
complete neutralization = (V2 – V1) cm3.
Pipette 25.0cm3 of NaOH into a beaker. Then add (V2 – V1) cm3 of dilute nitric acid
from the burette.
Heat to evaporate the water till it is saturated.
Cool the saturated solution so that the salt crystallizes.
Filter to collect the crystals.
Dry the crystals between a few sheets of filter paper.
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Easiest to prepare
Just need to use precipitation
Mix a solution containing the positive ions of
the salt with another solution containing the
negative ions of the salt to be prepared.
What salts are insoluble?

1.
2.
3.
4.
Using lead(II) nitrate and dilute sulfuric acid
First, pour 50 cm3 of lead(II) nitrate solution
into a small beaker. Add sulfuric acid (in
excess) and stir until no more precipitate
forms.
Filter and collect precipitate.
Wash the precipitate with a small amount of
distilled water to remove impurities.
Allow the precipitate to dry on filter paper.
First, pour 50 cm3 of one reagent (depending on what salt you want) into a small
beaker.
Add another reagent (again depending on what salt you want) and stir until no
more precipitate forms.
Filter and collect precipitate.
Wash the precipitate with a small amount of distilled water to remove impurities.
Allow the precipitate to dry on filter paper.

1.
2.
3.
4.
5.
How do you get the following salts:
Magnesium sulfate
Lead(II) chloride
Potassium nitrate
Sodium sulfate
Zinc nitrate

Are these salts soluble?
1.
2.
3.
4.
5.
6.
7.
8.
9.
iron(III) nitrate
potassium carbonate
sodium ethanoate
silver chloride
lead(II) nitrate
copper(II) carbonate
ammonium iodide
titanium(IV) chloride
barium sulfate
Yes
Yes
Yes
No
Yes
No
Yes
Yes
No
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1.
2.
3.
4.
5.
6.
7.
8.
Which method will you use to get the following
salts:
Magnesium sulfate
Lead(II) chloride
Potassium nitrate
Sodium sulfate
Copper(II) chloride
Lead(II) carbonate
Silver chloride
Zinc chloride
When an acid Z is added to a solution of
lead(II) nitrate, an insoluble precipitate is
formed.
 When Z is added to a solution of silver
nitrate, an insoluble precipitate is formed too.
 What acid could Z be?
A) hydrochloric acid B) sulfuric acid
C) nitric acid


A metal oxide A dissolves in sulfuric acid,
hydrochloric acid and nitric acid and does
NOT give any precipitate with any of the
acids. Which of the following could be A?
A) Barium oxide
C) Silver oxide
B) Calcium oxide
D) Sodium oxide