Set up of apparatus.
The apparatus is set up as shown in the diagram. On heat concentrated sulphuric acid reacts with potassium nitrate to give nitric acid vapour and potassium hydrogen sulphate
Nitric acid collects in the collecting apparatus cooled tap water. The acid collected is yellow. It is converted to colourless nitric acid by bubbling air/ oxygen through it.
Industrial preparation of nitric acid (Ostwald process)
Nitric acid is manufactured by catalytic oxidization of ammonia (oxidization of ammonia in presence of a catalyst).
Ammonia and excess air are passed over the hot platinum catalyst at about 8000C forming colourless nitrogen monoxides gas. The reaction is exothermic.
The mixture of nitrogen monoxide and excess air is cooled and nitrogen monoxide is oxidized to nitrogen dioxide.
Nitrogen dioxide together with excess air is dissolved in hot water to form nitric acid.
A VIDEO ABOUT LABORATORY PREPARATION OF NITRIC ACID
Uses of nitric acid
- It is used in the manufacture of brass.
- It is used in the manufacture of dyes and explosives.
- Used in the manufacture of fertilizers like ammonium nitrate, NH4, NO3, potassium nitrate, KNO3 and sodium nitrate, NaNO3.
Ammonium nitrate is prepared by neutralizing nitric acid with ammonia solution.
Properties of nitric acid
It behaves chemically in 2 ways
- As a strong acid
- As a powerful oxidizing agent
Nitric acid as a strong acid
Conditions under which nitric acid behaves as a strong acid
- The acid should be dilute
- Room temperature (no heating)
Reactions of nitric acid as a strong acid
- It reacts with carbonates and hydrogen carbonates to give carbon dioxide gas
It reacts with bases and alkalis forming salts and water only
Very dilute nitric acid reacts magnesium metal evolving hydrogen.
Magnesium is the only metal that liberates hydrogen with nitric acid. Other metals are oxidized by the acid to the corresponding metal nitrate.
Nitric acid as an oxidizing agent
- Must be concentrated
- Sometimes it needs to be hot.
- Concentrated nitric acid reacts with a green solution of iron (III) sulphate. The most important change is the oxidizing of iron (II) IONS.
- Concentrated nitric acid reacts with the copper metal to give reddish brown nitrogen dioxide gas. A reaction occurs at room temperature.
If the acid is diluted with its own volume of water, colourless nitrogen monoxide is formed.
The lead metal reacts with nitric acid in a similar way.
Aluminium and iron are made passive because of the formation of the oxide layer which forms a protective coating over the metal and stops further chemical reaction.
- Reaction with non-metals
- Concentrated nitric acid reacts with sulphur powder to give reddish brown fumes of nitrogen dioxide.
- When a piece of red hot charcoal is put into concentrated nitric acid. It continues to burn and reddish-brown gas of nitrogen dioxide are formed.
- When red phosphorus is heated gently with moderately dilute nitric acid, reddish brown fumes of nitrogen dioxide are formed.
Other oxidation reactions
When hydrogen sulphide gas is passed through moderate dilute nitric acid, a pale yellow precipitate of sulphur is formed.
Hydrogen sulphide is oxidized to sulphur and nitric acid is reduced to colourless nitrogen monoxide.
With concentrated nitric acid, nitrogen dioxide is formed.
All nitrates are soluble in water
The reaction of heat on nitrates
On very strong heating sodium and potassium nitrates melt to a colourless liquid which the decomposed to give a pale yellow nitrite and oxygen gas.
Lead (II) nitrate;
On heating, it makes a cracking sound and gives off white fumes
Reddish – brown fumes of NO2 and colourless gas which supports burning is given off. (oxygen).
The residue is reddish brown when hot and yellow on cooling (lead (II) nitrate)
It is white crystalline solid. On heating, t decomposes to give reddish brown nitrogen dioxide and oxygen gas and a white residue.
Magnesium nitrate; white crystalline solid. Decomposes on heating to give NO2,O2 gas and white residue of magnesium oxide.
It is a white crystalline solid and is deliquescent on heating it decomposes to give reddish – brown NO2, oxygen gas and a residue which is yellow when hot and white on cooling. (zinc (II) oxide).
Copper (II) nitrate;
It’s a green crystalline solid an is deliquescent. On heating, it decomposed to give nitrogen dioxide, oxygen gases and black residue of copper (II) oxide.
Mercury (II) nitrate and silver nitrate; they decompose on heating to give the metal-nitrogen dioxide and oxygen.
Test for nitrates
The brown ring test
To a solution of a nitrate, add an equal volume of freshly prepared iron (II) sulphate in a boiling tube- hold the tube in slanting position and very carefully poor concentrated sulphuric acid down the side of the tube.
Concentrated sulphuric acid is denser than the solution and therefore sinks to the bottom of the tube and a brown ring from where the two layers meet.
Using concentrated sulphuric acid
Add concentrated sulphuric acid to a solid nitrate in a test and heat the tube to form a nitric acid vapour. Reddish – brown fumes of nitrogen dioxide are formed.
Using copper powder and concentrated sulphuric acid.
Mix a solid nitrate with copper powder and then add concentrated sulphuric acid. Het the mixture, NO2 fumes are formed.
Chlorine and its compounds
Chlorine is a non-metallic element of group seven of the periodic table.
It has 17 electrons
Valency = 1
It forms a negative ion by accepting 1 electron.
Set up of the apparatus is as follows;
The apparatus is set up as shown in the figure above. On heating, concentrated hydrochloric acid reacts with manganese (IV) oxide to give chlorine gas
The gas is bubbled through water to remove hydrogen gas and then through concentrated sulphuric acid to dry it. It is collected by downward delivery
Set up of the apparatus:
The apparatus is set up as shown in the figure above. At room temperature potassium manganate (VII) reacts potassium manganate (VII) reacts with concentrated hydrochloric acid to give chlorine gas. The gas is collected over brine. (Liquid sodium chloride).
Potassium manganite (VII) and manganese (IV) oxide oxidize concentrated hydrochloric acid to chlorine.
Using bleaching powder. (Calcium hypochlorite)
At room temperature, bleaching powder reacts with dilute nitric or hydrochloric acid to give chlorine
Industrial preparation of chlorine
- Visit notes of electrolysis of brine using the mercury cathode cell.
Properties of chlorine
- It’s a greenish- yellow gas with a chocking irritating smell.
- Fairly soluble in cold water forming yellowish chlorine water which contains chlroci (I) acid. HOCL (aq) hydrochloric acid. HCl
- It turns blue litmus red and it bleaches it immediately.
- It doesn’t burn and it extinguishes a burning splint.
- It is denser than air.
- It is easily liquefied (gas- liquids) under pressure at ordinary temperature.
Test for chlorine –
It’s a greenish yellow gas which turns damp blue litmus red and bleaches it immediately.
Affinity for hydrogen-
It rarely combines with substances which contain hydrogen turpentine,
It is warmed, a filter paper deeper in it and then dropped in a gas jar full of chlorine, there is a red flash accompanied by a violent reaction and a black cloud of solid particles of carbon is formed.
Chlorine combines with hydrogen in turpentine leaving black carbon behind from the equation chlorine is more reactive than carbon.
Reaction with hydrogen
A mixture hydrogen and chlorine does not react in darkness. When a flame is applied to a mixture containing equal volumes of hydrogen and chlorine
Hydrogen burns with a white flame and clouds of steamy fumes of hydrogen chloride are formed.
The greenish – yellow colour of chlorine gradually disappears
Reaction with hydrogen sulphide
Chlorine reacts with hydrogen sulphide forming a yellow precipitate of sulphur and white fumes of hydrogen chloride.
Reaction with phosphorous
When a piece of yellow phosphorous is lowered in a glass jug full of chlorine, it burns spontaneously (without application of heat) giving off white fumes of phosphorous trichloride (phosphorous (III) chloride) and phosphorous pentachloride (phosphorous (V) chloride)
The action of chlorine as an iron
A stream of dry chlorine gas is passed over hot iron metals as shown in the diagram above. Hot iron burns spontaneously in dry chlorine to form iron (III) chloride which collects as sublimates of black crystals in the collecting apparatus.
Heating is stopped as soon as the reaction starts.
A tube containing a drying agent (anhydrous calcium chloride) is connected to the collecting apparatus as this would be absorbed by iron (III) chloride which is very deliquescent. Iron (III) chloride is formed instead of iron (III) chloride because chlorine is an oxidizing agent.
The reaction should be carried out in fume cupboard or excess chlorine dissolves in an alkaline. (chlorine is very toxic.)
Sodium chloride can be prepared y heating sodium metal in an atmosphere of chlorine
Magnesium chloride can also be prepared by heating magnesium metal in an atmosphere of chlorine.
Iron (II) chloride – prepared by reacting hot iron metal with hydrogen chloride to give iron (II) chloride and hydrogen gas.
Explanation; chlorine is an oxidizing agent. It oxidizes iron (II) sulphate (pale green) to iron (III) sulphate (yellow/ brown)
Reaction with alkaline
- Cold dilute alkalis; when chlorine gas is bubbled through cold aqueous alkalis, the hypochlorite and chloride of the metal are formed. The solution turns pale yellow and it smells of chlorine gas
The resulting solution has bleaching properties; bleaches clothes, removes ink stains from paper and cloth. This is because sodium hypochlorite is a bleaching agent. It can also be used as an antiseptic (kills germs in threat and mouth).
- Hot concentrated alkalines
When chlorine is passed into hot concentrated alkalines a mixture of the chlorate and chloride is formed.
When calcium is added t a jar of chlorine, the colour and smell of the gas soon disappears, showing that it has been.
Bleaching actions of chlorine
The bleaching is due to oxidization of chloric (I) acid; dry hydrochloric does not bleach.
Chlorine bleaches by oxidization
Effect of sunlight on chlorine water