Synthesising qualitative analysis of salts

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Malaysia SPM Form 4 Chemistry, Chapter 8: Salt

Contents 1 Identification of Anion 2 Physical state 2.1 Colour 3 Heating 3.1 Heating Effect on Carbonate Salt 3.1.1 Summary 3.2 Heating Effect on Nitrate Salt 3.2.1 Summary 3.3 Heating effect on sulphate salt 3.4 The heating effect on chloride salts 3.4.1 Summary 4 Identification of Gases 4.1 Summary 5 Testing for Anions (Negative Ions) 5.1 Confirmation test for nitrate ion( IMPORTANT) 5.2 Summary 6 Testing for Cations (Positive Ions)

[edit] Identification of Anion

• There are 10 cations and 4 anions to be studied in our syllabus. :

Cation

Sodium	Na+	Iron (II)	Fe2+
Calcium	Ca2+	Iron (III)	Fe3+
Magnesium	Mg2+	Lead(II)	Pb2+
Aluminium	Al3+	Copper (II)	Cu2+
Zinc	Zn2+	Ammonium	NH4+

Anion

Chloride ion	Cl-
sulphate ion	SO42-
nitrate ion	NO3-
carbonate ion	CO32-

• Steps in qualitative analysis.

1. Examine the physical quantity of the salt ( State, solubility, colour…).
2. Heat he salt and collect the gas been released (if there is any.).
3. Prepare an aqueous solution of the salt to do the anion and cation test.
4. Confirmation test for certain ion.

[edit] Physical state

Solid	Ionic compound ( salt/metal oxide)
Liquid/gas	Covalent compound
Aqueous solution	Soluble salt

• Normally, ionic salts are in crystal form, metal oxide and metal sulphide are in powder form.

[edit] Colour

Salt or metal oxide	Solid	Aqueous solution
Salt of Sodium, Calcium, Magnesium, Aluminium, zinc, Lead, ammonium	White	colourless
Salt of Chloride, sulphate, nitrate, carbonate	White	colourless
Salt of Copper(II). Copper(II) Carbonate Copper(II) sulphate, Copper(II) nitrate, Copper(II) chloride Copper(II) oxide	Green Blue Black	Insoluble Blue Insoluble
Salt of Iron (II) Iron(II) sulphate; Iron(II) nitrate; Iron(ID chloride	Green	Green
Salt of Iron (III) Iron(III) sulphate; Iron(III) nitrate; Iron(III) chloride	Brown	Brown
Zink oxide	Yellow when it is hot and white when it is cold.	Insoluble
Lead(II) oxide-	Brown when it is hot and yellow when it is cold.	Insoluble
Magnesium oxide, Aluminium oxide	White	Insoluble
Potassium oxide, Sodium oxide, Calcium oxide	White	Colourless

[edit] Heating

[edit] Heating Effect on Carbonate Salt

[edit] Summary

Carbonate Salt	Equation of The Reaction
Potassium carbonate Sodium carbonate	Not decomposible
Calcium carbonate Magnesium carbonate Aluminium carbonate Zinc carbonate Iron (III) carbonate Lead(II) carbonate Copper(II) carbonate	CaCO3 ---> CaO + CO2 MgCO3 ---> MgO + CO2 Al2(CO33 ---> Al2O3 + 3CO2 ZnCO3 ---> ZnO + CO2 Fe2(CO3)3---> Fe2O3 + 3CO2 PbCO3 ---> PbO + CO2 CuCO3 ---> CuO + CO2
Mercury(II) carbonate Silver(I) carbonate	2HgCO3 ---> 2Hg + 2CO2 + O2 2Ag2CO3 ---> 4Ag + 2CO2 + O2
Ammonium carbonate	(NH4)2CO3 ---> NH3 + CO2 + H2O

[edit] Heating Effect on Nitrate Salt

[edit] Summary

Nitrate Salt	Equation of The Reaction
Potassium nitrate Sodium nitrate	2KNO3 ---> 2KNO2 + O2 2NaNO3 ---> 2NaNO2 + O2
Calcium nitrate Magnesium nitrate Aluminium nitrate Zink nitrate Iron (III) nitrate Lead(II) nitrate Copper(II) nitrate	2Ca(NO3)2 ---> 2CaO + 4NO2 + O2 Mg(NO3)2 ---> 2MgO + 4NO2 + O2 4Al(NO3)3 ---> 2Al2O3 + 12NO2 + 3O2 Zn(NO3)2 ---> 2ZnO + 4NO2 + O2 4Fe(NO3)3 ---> 2Fe2O3 + 12NO2 + 3O2 Pb(NO3)2 ---> 2PbO + 4NO2 + O2 Cu(NO3)2 ---> 2CuO + 4NO2 + O2
Mercury(II) nitrate Silver(I) nitrate	Hg(NO3)2 ---> Hg + 2NO2 + O2 2AgNO3 ---> 2Ag + 2NO2 + O2
Ammonium nitrate	NH4NO3 ---> N2O + 2H2O

[NOTES: Nitrogen dioxide, NO₂ is acidic gas and is brown in colour.]

[edit] Heating effect on sulphate salt

Heating effect on sulphate salt

• Most sulphate salts do not decompose by heat. For instance, sodium sulphate, potassium sulphate, and calcium sulphate are not decomposable by heat.
• Only certain sulphate salts are decomposed by heat when heated strongly.
• For instance:
  • Strong heating of green crystal iron (II) sulphate will release steam, sulphur dioxide, sulphur trioxide and leave behind a reddish solid iron (III) oxide residue. The steam released comes from the hydrated water of the crystallize salt.

2FeSO₄•7H₂O ---> Fe2O₃(p) + SO₂(g) + SO₃(g) + 14H₂O(g)

• Meanwhile, zinc sulphate, copper (II) sulphate, and iron (III) sulphate decompose when heated strongly to evolve sulphur trioxide gas and form a metal oxide.

Example
Zinc sulphate

ZnSO₄ ---> ZnO + SO₃

Copper (II) sulphate

CuSO₄ ---> CuO + SO₃

Iron (III) sulphate

Fe₂(SO₄)₃₄ ---> Fe₂O₃ + SO₃

1. When ammonium sulphate is heated strongly, this white solid sublimate and is decomposed to form ammonia gas and sulphuric acid. vapour

(NH₄)₂SO₄ ---> 2NH₃ + H₂SO₄

[edit] The heating effect on chloride salts

• All chloride salts are not decomposable by heat except ammonium chloride.

For instance:

NH4Cl ¾®

[edit] Summary

Heating effect on sulphate salt	The heating effect on chloride salts
Most sulphate salts do not decompose by heat. Only certain sulphate salts are decomposed by heat when heated strongly. Zinc sulphate, Copper (II) sulphate, Iron (III) sulphate ZnSO4 ---> ZnO + SO3 CuSO4 ---> CuO + SO3 2Fe2(SO4)3 ---> Fe2O3 + SO2 + SO3 Ammonium sulphate (NH4)2SO4 ---> 2NH3 + H2SO4	All chloride salts are not decomposable by heat except ammonium chloride. Example: NH4Cl ---> NH3 + HCl

[edit] Identification of Gases

• When testing for gases it is not sufficient to say that the gas is colourless, odourless (without smell) or insoluble in water.
• Most gases are colourless (oxygen, hydrogen, carbon dioxide, ammonia, etc.) and many gases are odourless (hydrogen, oxygen, carbon dioxide, etc.) or fairly insoluble in water (oxygen, hydrogen, etc.).
• Testing a gas with moist litmus paper is only acceptable if the result only happens with that gas.
• For example, many gases are acidic to litmus but only one common gas bleaches moist litmus (chlorine) and one common gas is alkaline to moist litmus (ammonia).
• Finally, testing gases to see if they put out a burning splint will work with nearly all gases so it is not acceptable as an identification test.
• However there is only one common gas which does the reverse and relights a glowing splint (oxygen). A flowchart for the identification of gases is given in Figure 14.6.
• A useful summary of the tests for gases is given below.

[edit] Summary

Gases	Characteristics
Oxygen	Rekindle glowing splinter.
Hydrogen	Explode with a ‘pop’ sound when brought close to a lighted splinter.
Carbon Dioxide	Turns lime water chalky.
Chlorine	Bleach moist litmus paper.
Ammonia	Pungent smell. Turn moist red litmus paper to blue. Produces white fume when reacts with concentrated hydrochloric Acid.
Sulphur Dioxide	Pungent smell. Bleach the purple colour of potassium manganate(VII). Turn moist blue litmus paper to red.
Nitrogen Dioxide	Pungent smell. Brown in colour. Turn moist blue litmus paper to red.

[edit] Testing for Anions (Negative Ions)

[edit] Confirmation test for nitrate ion( IMPORTANT)

• About 2cm³ of dilute sulphuric acid is added into the solution that wants to be tested and then followed by 2cm3 iron (II) sulphate solution.
• A few drops of concentrated sulphuric acid are carefully drop through the inclined side of the test tube without shaking the test tube.

Observation:

• Explanation: Iron (II) sulphate reduce nitric acid (from the reaction between nitrate ion and concentrated sulphuric acid) to nitrogen monoxide. Afterwards, nitrogen monoxide combines with iron (II) sulphate to form the compound FeSO₄.NO which is brown in colour (brown ring).

[edit] Summary

	Diluted HCl or diluted HNO3 or diluted H2SO4	BaCl (aq) or Ba(NO3)2 (aq) follow by diluted HCl/HNO3	AgNO3 follow by diluted HNO3.	Brown Ring Test ( + FeSO4 (aq ) + concentrated H2SO4
CO32-	Carbon Dioxide is released.	White precipitate is formed. It is soluble in diluted HCl/HNO3	White precipitate is formed. It is soluble in diluted HNO3
SO42-		White precipitate is formed. It is NOT soluble in diluted HCl/HNO3
Cl-			White precipitate is formed. It is NOT soluble in diluted HNO3
NO3-				Formation of Brown Ring

[edit] Testing for Cations (Positive Ions)

• Cations can be identified by their reaction with aqueous sodium hydroxide and aqueous ammonia. Sodium hydroxide and aqueous ammonia produce hydroxide ion which will react with most anion to form precipitate.

NaOH + H₂O ---> Na⁺ + 2OH^- + H⁺

NH₃ + H₂O ---> NH₄⁺ + OH^-

• Different cations like aluminium Al³⁺,calcium Ca²⁺, copper(II) Cu²⁺, iron(II) Fe²⁺, iron(III) Fe³⁺, lead(II) Pb²⁺, zinc Zn²⁺

1. produce different coloured precipitates,
2. which may or may not dissolve in excess alkali.

• A precipitate is an insoluble solid.
• When testing for cations, these precipitates only form when a metal ion in solution joins with hydroxide ions in solution to form an insoluble metal hydroxide. For example:

• Zn(OH)₂, Al(OH)₃ and Pb(OH)₂ dissolve in excess NaOH solution, this is because Zn(OH)₂, Al(OH)₃ and Pb(OH)₂ are amphoteric, they can react with NaOH to form salt and water.

Zn(OH)₂ + NaOH ---> Na₂ZnO₂ + 2H₂O

Al₂(OH)₃ + NaOH ---> Na₃Al₂O₃ + 2H₂O

Pb(OH)₂ + NaOH ---> Na₂PbO₂ + 2H₂O

• The ammonium cation NH₄+ produces ammonia gas with aqueous sodium hydroxide.
• This is the only common alkaline gas. It is therefore, easy to identify as it turns moist red litmus paper blue, for example:

• Some tests to distinguish Fe²⁺ ion from Fe³⁺ ion.

Reagent	Observation	Ion presents
Solution of potassium hecxacianoferate(II)	Light blue precipitate	Fe2+
	Dark Blue precipitate	Fe3+
Solution of hecxacianoferate (III)	Dark blue precipitate	Fe2+
	Greenish brown solution	Fe3+
Solution of potassium thiocyianate	Pinkish solution	Fe2+
	Blood red solution	Fe3+

• The following table summarize the importance tests for cations