Back to: O level chemistry notes full Uganda syllabus
This is the systematic examination of a chemical compound to find out the elements or ions present in it, type of bond or nature of reactivity. It is also referred to as property analysis.
In a practical experiment, a student is normally given one or more substances and he/she is expected to carryout a number of tests on the substance(s). He/she is further expected to record the observations and make relevant deductions about the nature of substance(s) given.
Preliminary test
The purpose of this test is to give general guidelines on the nature of substance being analyzed. Preliminary analysis is based on three major aspects namely:
- Appearance
- Flame colour
- Action of heat
- Appearance
The appearance of substances can further be divided in to three categories i.e. physical state, colour and smell.
i) Physical state
Carbonates and oxides appear powdery whereas nitrates, sulphates, and chlorides appear crystalline.
ii) Smell
Some compounds especially ammonium compounds and sulphides posses a noticeable smell.
iii) Colour
Colour serves as a good guide as to what metallic ions may be present in a given substance. However it is important to note that a substance cannot be completely identified basing on its colour.
Action of heat
Usually, a little of the substance is heated in a dry, hard glass test tube gently at first and then strongly until there is no further change.
Since the action of heat on solids leads to decomposition of the substance, there is need to always observe the gases given off (if any) and the residue left in the test tube after heating. Observations made when some solids are heated are given in the table below.
Making a test solution
Procedure
Dissolve a little of the substance provided e.g. a spatula end full in about 5cm3 of water. Shake the test tube well until the substance dissolves. The solution obtained is now your test solution.
If the solid is insoluble in water, you will be required to dissolve it in cold dilute nitric acid or dilute hydrochloric acid.
The main reaction between these two reagents and the cations is the precipitation of insoluble metal hydroxides. Some of these metal hydroxides dissolve in excess of these reagents forming solutions of complex ions and some do not dissolve.
Identification using sodium hydroxide
Procedure
- To a small portion of the test solution, add a few drops of sodium hydroxide and shake the mixture. Check if precipitate is formed and note its colour.
- Add more sodium hydroxide solution until in excess and then shake the solution. Note whether the precipitate formed is soluble in excess alkali or insoluble in excess alkali.
NB. Incase no precipitate appears after adding sodium hydroxide solution, warm the mixture gently and test for ammonia.
Confirmatory tests for cations
A confirmatory test is a specific sensitive reaction undergone by a particular ion. Since sodium hydroxide solution or ammonia cannot completely identify some cations, confirmatory tests are then used.
These ions can be divided into three major groups as:
- Those that react with (major) dilute acids to give off gases
- Those that react with concentrated sulphuric acid to liberate gases
- Those that react with neither dilute acids nor concentrated sulphuric acid
Anions that react with dilute acids
Test procedure
To the test solution, add 2-3 drops of dilute nitric acid or dilute hydrochloric acid. If no reaction occurs, warm the mixture gently, if effervescence occurs, smell the gas with care and identify it.
Carbonate
All carbonates and hydrogen carbonates give off carbon dioxide when they are treated with dilute acids. This serves as a temporary test for carbonates and hydrogen carbonates.
Apart from potassium, sodium and ammonium carbonates, all other carbonates decompose upon heating to give off carbon dioxide.
Test of a soluble carbonate
Procedure
To the test solution, add
i) Barium nitrate solution/barium chloride solution followed by excess dilute nitric acid.
Observations
A white precipitate is formed which dissolves in excess dilute nitric acid with effervescence.
Equation
A part from potassium, sodium and ammonium nitrates, all other nitrates liberate brown fumes of nitrogen dioxide gas when heated strongly. The liberation of brown fumes serves as a temporary confirmatory test for a nitrate.
Brown ring test
Procedure
To the solution of a substance under test, add an equal volume of freshly prepared iron(II) sulphate solution. Then carefully add a few drops of concentrated sulphuric acid down the sides of a slanting test tube so that the acid sinks to the bottom and forms a separate layer.
Observation
A brown ring slowly forms where the acid meets the top layer. This confirms the presence of a nitrate.
Detection and identification of gases
Gases are normally given off when substances are heated or when reagents like acids are added to substances.
How do you know that a gas is being evolved?
i) Effervescence (rapid emission of bubbles) occurs
ii) Colour. A few gases posses noticeable colours e.g. brown fumes for nitrogen dioxide.
iii) Smell. Some gases have got distinctive smells and they are usually unpleasant. E.g. ammonia (choking smell) and hydrogen dioxide (smell of rotten eggs)
After identifying that a gas is given off, it can then be identified using a chemical test.
Common tests to identify gases
i) Support of combustion
a) Oxygen relights a glowing splint
b) For a burning splint, if the flame is blue in colour, then suspect hydrogen or carbon monoxide.
ii) Litmus paper
Damp/wet litmus paper is always used. Alkaline gases such as ammonia turn red litmus paper blue while acidic gases like carbon dioxide turn blue litmus pink/red.
iii) Lime water/calcium hydroxide
This serves as a confirmatory test for carbon dioxide gas, indicating the presence of a carbonate or hydrogen carbonate. Carbon dioxide turns lime water milky.
iv) Water vapour
This is seen as a colourless condensed liquid on the wall of the test tube when heating.
Sample questions with solutions
- You are provided with substance X that contains one cation and one anion. Carry out the following tests to identify the cation and anion in X. Identify any gases that may be evolved. Record your observations and deduction in the table below.
Ideal Chemistry – A complete approach for “O” level Page 220
- You are provided with substance Y that contains two cations and two anions. Carry out the following tests to identify any gases that may be evolved. Record your observations and deductions in the table below
NITROGEN AND ITS COMPOUNDS
Nitrogen is in period 2 and group (V) of the periodic table of elements. It has atomic number of 7 and electronic configuration 2:5.
Occurrence
Nitrogen exists freely in the air as diatomic molecules and makes up to 78% of the air by volume.
It occurs in combined states as nitrites, nitrates and most of these are distributed every where in soil as ammonium sulphate, sodium nitrate, potassium nitrate and calcium nitrate,. It is also constituent of living matter of pants and animals.
Laboratory preparation of nitrogen
a) From air
Set up
The nitrogen gas can now be collected over water.
If the nitrogen is required dry, it is then passed over fused calcium chloride to remove water vapour or it can be passed through a U tube containing beads soaked in concentrated sulphuric acid to dry the nitrogen gas.
NB. The nitrogen obtained by this method is denser than ordinary nitrogen since it contains impurities like argon and other inert gases.
b) Preparation of nitrogen by action of heat on ammonium nitrite
The ammonium nitrite is formed by the reaction between sodium nitrite and ammonium chloride.
Chemical properties
Nitrogen is generally an unreactive gas. This is because of the presence of the strong triple covalent bonds between its atoms in a molecule (N≡N). The triple covalent bonds are hard to break rendering nitrogen inert.
Some reactive metals like magnesium and calcium burn in nitrogen to form nitrides i.e. magnesium nitride and calcium nitride
Uses of nitrogen
- Nitrogen is used in the synthesis of ammonia gas.
- It is used in the manufacture of fertilizers like ammonium phosphate.
- The atmospheric nitrogen is fixed by thunder or bacteria in root nodules of leguminous plants to nitrates which can be used by plants.
- Because of its low boiling points, liquid nitrogen is used to cool materials to very low temperatures.
- It is used in the manufacture of nitric acid.
OXIDES OF NITROGEN
NITROGEN DIOXIDE (NO2)
Laboratory preparation
It is conveniently prepared in the laboratory by heating lead(II) nitrate crystals (this is because lead(II) nitrate forms crystals without water of crystallization which is not common to other metallic nitrates and would otherwise interfere with the preparation)
Nitrogen dioxide is finally obtained as brown fumes
Properties of nitrogen dioxide
Physical properties
- It is reddish brown gas at room temperature
- It is soluble in water
- It is denser than air
- It is highly poisonous and produces nitric acid in the lungs when inhaled (it causes septic pneumonia)
- It has an irritating pungent smell
- It is easily liquefied
Chemical properties
- Burning of metals. Metals e.g. magnesium and non metals like phosphorus burn even more highly in nitrogen dioxide. This suggests that nitrogen dioxide supports combustion and this will be so if the material reacting is hot enough to decompose nitrogen dioxide into oxygen (a gas that supports burning)
Uses of nitrogen dioxide
- It is used in the manufacture of plastics, explosives, nylon materials e.t.c.
- Used in the manufacture of nitric acid.
NITROGEN MONOXIDE (NO)
Laboratory preparation
Set up
DINITROGEN OXIDE, N2O (LAUGHING GAS)
Laboratory preparation
Set up
Procedure
- Grind a mixture of sodium nitrate and ammonium chloride.
- Put the mixture in a round bottom flask and assemble the apparatus as shown in the diagram above.
NB. On heating the mixture, ammonium nitrate is formed which quickly decomposes to give dinitrogen oxide and water vapour.
Physical properties of dinitrogen oxide
i) It is a colourless gas with a faint sweet smell.
ii) In it fairly soluble in cold water but insoluble in hot water.
iii) It is denser than air.
iv) It can be easily liquefied.
v) It is a neutral gas and has no effect on litmus paper.
Chemical properties of dinitrogen oxide
- It supports combustion and relights a glowing splint that is hot enough to decompose it into oxygen and nitrogen.
- Burning metals and non metals like magnesium and sulphur continue to burn in the gas giving their respective oxides.
Use of dinitrogen oxide
It is used as an anaesthetic in less complex surgical operations like dentistry. It has an effect that makes a patient laugh hysterically and it is therefore referred to as ―”laughing gas”.
AMMONIA
Laboratory preparation
Setup
Procedure
- Assemble the apparatus as shown in the diagram above
- Grind a mixture of ammonium chloride and calcium hydroxide and place it in a round bottom flask of a hard glass
NB. The neck of the flask should bend down wards and the flask should be in a slanting position because the formed water vapour will condense and if allowed to run back on the hot flask causes breakage
b) Preparation of nitrogen by action of heat on ammonium nitrite
The ammonium nitrite is formed by the reaction between sodium nitrite and ammonium chloride.
On slight warming, the solution of ammonium nitrite decomposes to give nitrogen gas. The nitrogen gas produced can then be collected over water as shown.
Industrial preparation of nitrogen
Nitrogen is obtained in the industry by fractional distillation of liquid air. Liquid air is fractionally distilled and nitrogen is obtained at a temperature of -196˚C (77 K at standard pressure). Oxygen with a higher boiling point (-183˚C ) is left behind. The separated nitrogen is liquefied and stored in specially designed container ready for use. The nitrogen may also be sold as compressed air.
Properties of nitrogen
Physical properties
- It is a tasteless, colorless and odourless gas
- It is slightly soluble in water (almost insoluble)
- It is slightly denser than air
- It is a neutral gas i.e. neither acidic nor basic
- It does not support burning though a few metals burn in it.
Chemical properties
Nitrogen is generally an unreactive gas. This is because of the presence of the strong triple covalent bonds between its atoms in a molecule (N≡N). The triple covalent bonds are hard to break rendering nitrogen inert.
Some reactive metals like magnesium and calcium burn in nitrogen to form nitrides i.e. magnesium nitride and calcium nitride
Uses of nitrogen
- Nitrogen is used in the synthesis of ammonia gas.
- It is used in the manufacture of fertilizers like ammonium phosphate.
- The atmospheric nitrogen is fixed by thunder or bacteria in root nodules of leguminous plants to nitrates which can be used by plants.
- Because of its low boiling points, liquid nitrogen is used to cool materials to very low temperatures.
- It is used in the manufacture of nitric acid.
OXIDES OF NITROGEN
NITROGEN DIOXIDE (NO2)
Laboratory preparation
It is conveniently prepared in the laboratory by heating lead(II) nitrate crystals (this is because lead(II) nitrate forms crystals without water of crystallization which is not common to other metallic nitrates and would otherwise interfere with the preparation)
Nitrogen dioxide is finally obtained as brown fumes
appears pale yellow when it is pure). The oxygen gas passes on and escapes as a colourless, harm less gas which can be collected if required over water as shown.
- As the white lead(II) nitrate crystals are heated, they make a crackling sound and melt. This is known as decrepitation. The residue in the boiling tube is Lead(II) oxide which is yellow when it cools.
Properties of nitrogen dioxide
Physical properties
- It is reddish brown gas at room temperature
- It is soluble in water
- It is denser than air
- It is highly poisonous and produces nitric acid in the lungs when inhaled (it causes septic pneumonia)
- It has an irritating pungent smell
- It is easily liquefied
Chemical properties
- Burning of metals. Metals e.g. magnesium and non metals like phosphorus burn even more highly in nitrogen dioxide. This suggests that nitrogen dioxide supports combustion and this will be so if the material reacting is hot enough to decompose nitrogen dioxide into oxygen (a gas that supports burning)
Uses of nitrogen dioxide
- It is used in the manufacture of plastics, explosives, nylon materials e.t.c.
- Used in the manufacture of nitric acid.
NITROGEN MONOXIDE (NO)
Laboratory preparation
Set up
Procedure
Place some copper turnings in a flask, add some water to cover it, then add moderately concentrated (50%) nitric acid (same volume as that of the water).
Observation
Vigorous effervescence occurs and the flask is filled with brown fumes. The brown fumes are nitrogen dioxide produced partly by the action of the acid upon the copper and partly by the oxidation of the main product, nitrogen monoxide by the oxygen of the air in the flask.
Alternative test
When the gas is bubble though iron(II) sulphate solution, the solution changes from pale green to dark brown or black. The dark brown of black coloration is due to formation of a compound, (FeSO4).NO. When this compound is heated, pure nitrogen monoxide is formed.
DINITROGEN OXIDE, N2O (LAUGHING GAS)
Laboratory preparation
Set up
Physical properties of dinitrogen oxide
i) It is a colourless gas with a faint sweet smell.
ii) In it fairly soluble in cold water but insoluble in hot water.
iii) It is denser than air.
iv) It can be easily liquefied.
v) It is a neutral gas and has no effect on litmus paper.
Chemical properties of dinitrogen oxide
- It supports combustion and relights a glowing splint that is hot enough to decompose it into oxygen and nitrogen.
- Burning metals and non metals like magnesium and sulphur continue to burn in the gas giving their respective oxides.
Use of dinitrogen oxide
It is used as an anaesthetic in less complex surgical operations like dentistry. It has an effect that makes a patient laugh hysterically and it is therefore referred to as ―”laughing gas”
AMMONIA
Laboratory preparation
Setup
Procedure
- Assemble the apparatus as shown in the diagram above
- Grind a mixture of ammonium chloride and calcium hydroxide and place it in a round bottom flask of a hard glass
NB. The neck of the flask should bend down wards and the flask should be in a slanting position because the formed water vapour will condense and if allowed to run back on the hot flask causes breakage - Heat the mixture in the flask to evolve ammonia gas
Since the gas is less dense than air and very soluble in water, it is collected by up ward delivery method.
Ammonia gas can be prepared in the laboratory by heating any ammonium salt with an alkali.
Properties of ammonia
Physical properties
i) Ammonia is a colourless gas with a characteristic choking smell.
ii) It is less dense than air.
iii) It has a very low boiling point (-34˚C) and liquefied under pressure.
iv) It is very soluble in water. It is in fact the most soluble gas known.
v) It turns moist red litmus paper blue. It is the only known alkaline gas.
Solubility of ammonia in water
The fountain experiment
The great solubility of ammonia is illustrated by the fountain experiment.
Set up
Procedure
- Fill a dry thick walled flask with ammonia gas and close it with a cork carrying tubes and clips as shown.
- Clamp the flask upside down and immerse the tubes with clips inside water which has been coloured with red litmus solution.
- Open clip B for a moment and close it. This allows few drops of water to enter in the flask. The water is made to run to the round end of the flask.
The ammonia gas in the flask dissolves in the water (forming a blue solution since it is an alkaline gas). This greatly reduces the gas pressure inside the flask. - Open clip A.
Observation
A fountain at once occurs as illustrated in the diagram. This will continue until the flask is full of water as it was with ammonia. When clip A was opened, water was forced into the flask because the atmospheric pressure from the outside was far much greater than the pressure inside the flask The fountain appears blue due to the alkaline nature of ammonia gas.
Chemical properties of ammonia
Reaction with air
When a lit splint is placed in a gas jar full of ammonia, it is extinguished showing that ammonia does not support combustion. On its own, ammonia does not burn but in plenty of oxygen, ammonia burns to produce nitrogen and water vapour.
In the presence of a catalyst e.g. platinum foil, ammonia it oxidized to nitrogen monoxide when it reacts with oxygen/air. In this case, the platinum foil continues to glow in the mixture of air and ammonia since the reaction is exothermic. Brown fumes of nitrogen dioxide are later seen as the nitrogen monoxide formed is oxidized in the presence of oxygen.
Reaction with hydrogen chloride gas
Ammonia reacts with hydrogen chloride to form dense white fumes of ammonium chloride which settle as white solids.
Setup for the preparation of ammonium solution
A mixture of ammonium chloride and calcium hydroxide is heated to produce ammonia gas which then dissolves in water to form the alkaline solution of ammonium hydroxide.
Precaution: The rim of the inverted funnel should just touch the surface of water. This prevents the water from being sucked into the flask.
Reaction with acids
Ammonia solution (aqueous ammonia) has many properties of typical alkaline solutions. E.g. it reacts with acids to form ammonium salts.
These salts can be crystallized out and are similar to ordinary metallic salts.
Reaction of aqueous ammonia with metallic ions
When aqueous ammonia is added to a solution of metallic salt, it forms precipitate of the insoluble metal hydroxide.
For example when a solution containing copper (II) ions e.g. copper (II) sulphate solution is put in a test tube and aqueous solution of ammonia added to it a little, pale blue precipitates of copper (II) hydroxide are formed.
When the ammonia solution is added until in excess, the pale blue precipitates dissolve giving a deep blue solution. The deep blue solution is due to the formation of a complex salt containing tetra amine copper (II) ion.
However, sodium hydroxide and ammonium hydroxide are different in that:
- The pale blue precipitates of copper(II) hydroxide do not dissolve in excess sodium hydroxide but dissolves in excess ammonia solution forming a deep blue solution.
- The amphoteric aluminium and lead ions dissolve in excess sodium hydroxide but not in excess ammonia solution.
- Zinc dissolves in both excess sodium hydroxide and ammonia solution but the reactions are different.
Process
Dry nitrogen and hydrogen in the ration of 1:3 by volume are mixed in the presence of a finely divided iron catalyst and the mixture heated.
Condition
The heating is carried out at a very high pressure of about 200-400 atmospheres and at a temperature of 400-500˚C in the presence of a catalyst (finely divided iron impregnated with alumina). The reaction is exothermic and reversible.
Note
- The two gases (hydrogen and nitrogen) must first be purified. This is because the impurities may poison the catalyst. The impurities that must be removed include: water vapour; dust; carbon dioxide and sulphur dioxide.
- The raw materials for the process are obtained from various sources e.g. the nitrogen is obtained from fractional distillation of liquid air, hydrogen from steam reforming of natural gas.
The ammonia gas produced is then liquefied and stored for future use.