School Science Lessons
2024-12-27
(Topic12F)
Sulfuric acid
Contents
12.16.0 Sulfuric acid
12.6.0.1 Acid rain, SOx, from burning sulfur or sulfur compounds
12.16.1 Concentrated sulfuric acid as oxidizing agent
12.16.2 Concentrated sulfuric acid dehydrates filter paper
12.16.3 Concentrated sulfuric acid dehydrates sucrose, (cane sugar)
12.16.4 Different forms of sulfuric acid
12.18.6.3 Reactions of sulfamic acid, NH2.SO2OH
12.16.5 Reactions of sulfuric acid
12.16.6 Sulfuric acid as a displacer
12.16.7 Sulfuric acid dehydrates copper (II) sulfate crystals
12.16.8 Sulfuric acid displaces acids from their salts
12.16.9 Sulfuric acid, ionization of hydrogen sulfate ion
12.16.10 Sulfuric acid with sodium chloride
12.18.2.0 Prepare sulfides, S2-
12.16.11 Hydrogen peroxide oxidizes sulfur dioxide to sulfuric acid
12.16.12 Prepare sulfuric acid with iron (II) sulfate
12.6.0.2 "Sulfur" in coal, iron (II) sulfide (pyrite).
12.16.13 Sulfur dioxide to sulfuric acid
12.16. Different forms of sulfuric acid
Sulfuric acid, pure, concentrated, 0.5 M, 1 M, 2 M, analytical reagent, hygroscopic, Harmful, Corrosive
Sulfuric acid, 98%
Sulfuric acid, concentrated, 18 M, 95%, 110 mL of concentrated solution for 1 litre of 2 M solution.
Sulfuric acid, slowly add acid to water with constant vigorous stirring, do NOT add water to acid
Sulfuric acid, formerly used with nitre, KNO3, for separating silver from copper
Sulfuric acid, fuming, "fuming sulfuric acid", oleum, Highly toxic by all routes, Extremely corrosive, use < 20 mL in fume cupboard
Sulfuric acid, fuming, Always very small quantities from a dropper.
Sulfuric acid, concentrated, > 10 M (90%), oil of vitriol
Sulfuric acid, Highly toxic by all routes, highly corrosive to skin, eyes, do not inhale vapour
Sulfuric acid, < 10 M (90%), > 4M (36%), Toxic if ingested, strongly acidic, corrosive
Sulfuric acid, < 4M (36%), > 0.5 M (5%), Strongly acidic, corrosive
Sulfuric acid, < 0.5 M (5%), Not hazardous
Sulfuric acid 2 M, Dilute 112 mL of 35 M in 800 mL water, then add water to 1 litre
Sulfuric acid 6 M, Dilute 168 mL of 17.8 M acid to 1 litre of water (95%sulfuric acid).
12.16.0 Sulfuric acid
Sulfuric acid, 1.7-7.8 M (15-95% wt/wt) should be handled only by an experieced teacher or a teacher with a high rate training code.
Teachers who handles it must ensure that they identify, assess and control all the risks.
Sulfuric acid, H2SO4, oleum, fuming sulfuric acid, vitriol or oil of vitriol is concentrated H2SO4, clear, colourless, heavy oily liquid, very hygroscopic, conc. 18 M, r.d. about 1.84, 18 mol per mL, r.d. fuming H2SO4: 1.92 gm cm-3 (for accumulators), battery acid, r.d. 1.25 gm cm-3, E513, dilute H2SO4, 3 M, store in cool place.
Pure sulfuric acid is called concentrated sulfuric acid.
Sulfuric acid reacts violently with water to form hydrated hydrogen ions and hydrated sulfate ions.
So the concept of dilution with water is inappropriate for the concentrated acid, because a vigorous chemical reaction occurs that releases much heat.
Do not add water to concentrated sulfuric acid.
Dilute sulfuric acid by adding the acid slowly to water with continuous stirring.
Wear eye protection when handling concentrated sulfuric acid.
Water is less dense than sulfuric acid and will layer itself on top of the acid.
At the acid / water interface, heat is generated by reaction of acid with water.
This causes the water just above the interface to become very hot.
When a suitable nucleating particle or surface is found, a large steam bubble forms, the middle section of the container boils suddenly and hot acid / water mixture is ejected from the container, often travelling several metres and causing serious chemical burns.
If, the acid is slowly poured into water with stirring, the same amount of heat is liberated, but the boiling point of the solution is much higher than 100oC, because of the presence of the sulfuric acid.
The temperature of the solution may rise to 150oC, but boiling still does not occur.
Sulfuric acid is highly corrosive to the skin and eyes.
Wash sulfuric acid off the skin immediately with a large volume of water.
A small amount of water simply brings about the acid water reaction and liberation of a great amount of heat, causing the remaining sulfuric acid to be even more corrosive.
Hot sulfuric acid is a strong dehydrating agent and will cause organic materials, e.g. skin, to turn black, because of the removal of water, leaving a residue of carbon.
Sulfuric acid is not volatile, so neither the pure acid nor its solutions release vapour.
However, strong heating of the acid may cause decomposition and boiling to form an aerosol of sulfuric acid droplets and sulfur dioxide gas that must not be inhaled.
Sulfuric acid forms a spontaneously explosive mixture with potassium permanganate to form liquid manganese heptoxide.
Similarly, sulfuric acid forms a spontaneously explosive mixture with chlorate salts where gaseous chlorine dioxide forms.
Dilute sulfuric acid solutions are far less dangerous than the concentrated acid, because they do not dehydrate or reducing materials and do not form much heat when further diluted.
Use sulfuric acid from motor car batteries if it is first diluted with water to four times its volume (Remember: acid to water!), but sulfuric acid collected from lead-acid batteries will be contaminated with lead.
Poisonous and corrosive if left on clothes.
In many experiments, a solution of sodium hydrogen sulfate (sodium bisulfate), can be used instead of sulfuric acid.
Sulfuric acid dissolves most metals, but not copper.
Sulfuric acid dissolves metal oxides, including copper oxide, is neutralized by alkalis, e.g. ammonia solution and forms carbon dioxide gas from metal carbonates and bicarbonates.
Sulfuric acid was formerly manufactured by heating iron (II) sulfate crystals, so it was called "oil of vitriol", a name that is still used for it today.
12.16.9 Sulfuric acid, ionization of hydrogen sulfate ion
See diagram: 12.6.0: Burn sulfur
Sulfuric acid ionizes to form hydronium ions, H3O+, and hydrogen sulfate (bisulfate) ions, HSO4.
Sulfuric acid is a colourless oily liquid available as:
A. 2.0 M (4.0 N) 1.0 M (2.0 N) and 0.5 M (1.0 N) volumetric solutions
B. Minimum assay 97% solution density 1.83 g cm-3 3. 98% "ANALAR"solution
C. "Battery acid" solution for lead cell accumulators minimum assay 30%density 1.25 g cm-3 at 20oC (battery acid).
Sulfuric acid is a strong dibasic acid that forms sulfates and hydrogen sulfates a strong oxidizing agent that dissolves copper and a strong dehydrates agent that can remove water from organic compounds.
Sulfuric acid is made by the contact process.
Sulfur is burned or the ores zinc sulfide or iron sulfide (pyrites), are heated to form sulfur dioxide.
The gases pass over vanadium (V) oxide or platinum catalyst at 450oC to form sulfur trioxide that combines with water to form sulfuric acid.
Sulfur trioxide is produced by the action of oxygen on sulfur dioxide in the presence of a catalyst, e.g. iron oxide.
12.6.0.1 Acid rain, SOx, from burning sulfur or sulfur compounds
When coal is burnt, the compounds that contain sulfur can form sulfuric acid, as in the equations below, to become components of acid rain (rainwater pH = 5.6, acid rain pH < 5).
There may be more than one pathway for the formation of sulfuric acid from sulfur dioxide.
12.6.0.2 "Sulfur" in coal, iron (II) sulfide (pyrite)
. The 1 - 3% "sulfur" in coal is usually iron (II) sulfide (pyrite).
4 FeS2 (s) + 11O2 (g) --> 2Fe2O3 (s) + 8SO2 (g)
S (s) + O2 (g) --> SO2 (g) sulfur dioxide
Also, other sulfide ores may produce sulfur dioxide in the atmosphere during smelting to obtain the pure metal.
35.20.15 Galena
3PbS + 3O2 (g) --> 3Pb (s) + 3SO2 (g)
Copper (II) sulfide, chalcocite mineral
Cu2S (s) + O2 (g) --> 2Cu (s) + SO2 (g).
12.16.13 Sulfur dioxide to sulfuric acid
1. Sulfur dioxide to sulfuric acid, 1..
Sulfur dioxide is oxidized to sulfur trioxide by oxygen gas.
2SO2 (g) + O2 (g) <--> 2SO3 (g)
Sulfur dioxide is oxidized to sulfur trioxide by nitrogen dioxide.
SO2 (g) + NO2 (g) --> SO3 (g) + NO (g)
Sulfur trioxide dissolves in water to form sulfuric acid.
SO3 (g) + H2O (l) --> H2SO4 (aq) sulfuric acid.
2. Sulfur dioxide to sulfuric acid, 2.
Sulfur dioxide dissolves in water to form sulfurous acid
SO3 (g) + H2O (l) --> H2SO3 (aq)
Sulfurous acid is oxidized to sulfuric acid by ozone
H2SO3 (aq) + O3 (g) --> H2SO4 (aq) + O2 (g).
12.16.11 Hydrogen peroxide oxidizes sulfur dioxide to sulfuric acid
SO2 (g) + H2O2 (l) --> H2SO4 (aq).
12.16.1 Concentrared sulfuric acid as oxidizing agent
BE CAREFUL! YOU ARE USING HOT CONCENTRATED SULFURIC ACID!
Add hot concentrated sulfuric acid to carbon.
The reaction forms carbon dioxide and sulfur dioxide.
C (s) + 2H2SO4 (l) --> CO2 (g) + 2SO2 (g) + 2H2O (l)
Add hot concentrated sulfuric acid to sulfur.
The reaction forms sulfur dioxide and water.
S (s) + 2H2SO4 (l) --> 3SO2 (g) + 2H2O (l)
Add hot concentrated sulfuric acid to carbohydrates.
The reaction forms carbon dioxide or carbon and water.
12.16.3 Concentrated sulfuric acid dehydrates sucrose
This favourite experiment forms a large quantity of toxic gases, e.g. carbon monoxide and sulfur dioxide, and the voluminous char, (remaining carbon and other substances), should be washed thoroughly with water to remove any remaining acid before handling.
Experiments
1. Heat a mixture of 0.5 cm of sucrose and 1.0 cm of concentrated sulfuric acid gently for 2 seconds and then leave to stand.
Note the vigorous reaction and the colour change from white sugar to black carbon.
C12H22O11 (s) (H2SO4 catalyst) --> 12C (s) + 11H2O (l)
Students are very interested in this experiment if the sugar is in the shape of a volcano in a deep beaker and the sulfuric acid is poured into the "crater " of the "volcano".
This experiment may be sold as a toy product.
2. Add 35 mL of sulfuric acid to 50 g of sugar in a 100 mL beaker placed on a heat resistant mat.
3. Add 10 mL of sulfuric acid to a large test-tube almost filled with sugar.
The sugar is dehydrated to form carbon and steam, which causes the material to expand, forming a black porous column that rises out of the beaker or test-tube.
Do the experiment in a fume cupboard, because foul smelling vapours are also released.
To remove remaining acid, cool and wash the product thoroughly with water before touching it.
4. Put some sucrose (cane sugar) in a tall beaker.
Add drops of concentrated acid to the sugar.
BE CAREFUL!
The sugar turns yellow then brown then black and rises in the beaker.
It reacts with carbohydrates like sugar and cellulose charring them by removing the elements of water from them and leaving a mass of black carbon behind.
5. Roll paper into a tube and hold it in the middle of a soft plastic container, e.g. ice cream tub.
Do not use a glass jar.
Fill the container with sugar.
Pour just enough water to dampen the sugar down the tube to reach the bottom.
Leave to stand for five minutes to allow the water to spread throughout the sugar.
Remove the paper tube to leave a hole in the damp sugar.
BE CAREFUL!
Pour 30 mL of concentrated (98%) sulfuric acid down the hole and onto the top of the sugar.
The sugar starts to turn brown, and black in patches.
After some minutes bubbles of steam form.
The reaction became more vigorous as the material in the container expands.
A black cylinder rises out of the jar.
Jets of steam spurt out.
Heat is given out as the cylinder keeps rising.
The black steaming cylinder is spongy carbon.
Tap with a spatula to show it is hard, like expanded polystyrene packaging.
If the carbon solidifies to make a seal over the top of the jar and the reaction continues deeper, below the seal, pressure may build up to cause an explosion and a shower of black crumbling carbon.
12.16.2 Concentrated sulfuric acid dehydrates filter paper
Place a piece of dry filter paper top of a beaker.
Put 3 drops of concentrated sulfuric acid on the centre of the filter paper.
Note how the centre of the filter paper chars, becomes black, as the concentrated sulfuric acid withdraws water from the filter paper fibres leaving carbon.
12.16.12 Prepare sulfuric acid with iron (II) sulfate
An early method of preparing vitriol (sulfuric acid), used for bleaching, was by dry distillation of green vitriol (ferrous sulfate, iron (II) sulfate).
Put 2 cm of powdered iron (II) sulfate crystals into a hard glass test-tube.
Fit it with a stopper through which passes a right angle piece of glass tubing.
The other end of the tubing dips into a test-tube.
Clamp the hard glass test-tube loosely in a stand or hold it in a paper holder.
Keep the test-tube sloping downward or moisture condenses in the cooler part of the test-tube, runs back on to the hot glass, and cracks the test-tube.
Heat gently at first, moving the flame, then more strongly.
Water of crystallization is produced at first and the substance changes to white anhydrous iron (II) sulfate, that decomposes with stronger heat.
A thick white vapour appears and condenses in the test-tube as a colourless liquid.
When no more vapour is produced, let the apparatus cool.
The liquid collected in the test-tube is a weak solution of sulfuric acid.
Test it with blue litmus paper and a crystal of sodium carbonate (washing soda).
Carbon dioxide gas forms.
The red substance left in the hard glass test-tube is red iron oxide, called jewellers' rouge, because jewellers use it for polishing gold and silver.
12.11.1 Reactions of nitrites
Use a fume cupboard for these experiments.
The oxide N2O3 is unstable at temperatures above -21oC and decomposes into nitric oxide and nitrogen dioxide.
N2O3 --> NO + NO2.
1. Add dilute hydrochloric acid to sodium nitrite solution.
Note the solution turns pale blue with effervescence and brown fumes are given off.
Acids react with nitrites to produce unstable nitrous acid that decomposes into nitric oxide and oxygen.
The oxygen oxidizes nitrous acid to nitric acid.
3NO3- + 2H + --> NO3- + H2O + 2NO
Finally, the nitric oxide finally reacts with the oxygen of the air to form nitrogen dioxide
2NO + O2 --> 2NO2.
2. Mix iron (II) sulfate solution and sodium nitrite solution.
Add drops of dilute sulfuric acid.
The solution turns brown, because of the loose compound that iron (II) sulfate makes with nitric oxide.
This test distinguishes between a nitrite and a nitrate.
3. Add drops of concentrated hydrochloric acid to two mL of potassium iodide solution.
Pour the mixture into sodium nitrite solution.
Iodine forms as a brown colour or a black precipitate, because of the oxidation of potassium iodide to iodine.
Iodide ion is oxidized by electron loss and nitrous acid is reduced by electron gain.
2I- - 2e- --> I2
2HNO3 + 2H + + 2e- --> 2H2O + 2NO
2HNO3 + 2H + + 2I- --> 2H2O + I2 + 2NO
Finally, the nitric oxide finally reacts with the oxygen of the air to form nitrogen dioxide
2NO + O2 --> 2NO2
The nitrous acid acts as an oxidizing agent.
2NO2- + 4H + + 2e- --> 2H2O + 2NO (electron gain).
4. Pass hydrogen sulfide into sodium nitrite solution acidified with dilute hydrochloric acid.
Note the rapid reaction, the precipitate of sulfur and the brown fumes.
Sulfide ions from the partially ionized hydrogen sulfide lose electrons and are oxidized.
Nitrite ions accept these electrons and are reduced.
S2- - 2e- --> S
2NO2- + 4H + + 2e- --> 2H2O + 2NO
2NO2- + 4H + + S2- --> 2H2O + 2NO + S
The nitrous acid acts as an oxidizing agent.
2NO2- + 4H + + 2e- --> 2H2O + 2NO (electron gain).
5. Add a piece of copper to sodium nitrite solution acidified with dilute sulfuric acid.
Note the rapid reaction to attack the copper and the solution turns blue.
The nitrous acid acts as an oxidizing agent.
2NO2- + 4H + + 2e- --> 2H2O + 2NO (electron gain).
6. Add sodium nitrite solution to bromine water acidified with dilute sulfuric acid.
The bromine water loses its colour, because of reduction to hydrobromic acid.
The bromine is reduced to bromide ions by accepting electrons, the nitrite ion is oxidized and so acts as a reducing agent by supplying the electrons.
Br2 + 2e- --> 2Br-
NO2- + H2O --> NO3-+ 2H+ + 2e-
NO2- + H2O + Br2 --> NO3-+ 2H + + 2Br-
The nitrous acid acts as a reducing agent and is oxidized to nitric acid.
No gas is given off.
If the free nitrous acid is in excess, it will decompose into nitric acid and oxygen.
H2O + NO2- --> NO3- + 2H+ + 2e- (electron loss)
If the free nitrous acid is in excess, decomposition occurs.
7. Add sodium nitrite solution to potassium permanganate solution acidified with dilute sulfuric acid.
The solution loses its colour.
The permanganate ion oxidizes and the nitrite ion reduces and is itself oxidized.
MnO4- + 8H+ + 5e- --> Mn2+ + 4H2O (electron gain)
H2O + NO2- - 2e- --> NO3-+ 2H+ (electron loss)
5NO3- + 2MnO4- + 6H+ --> 5NO3- + 2Mn2+ + 3H2O
The nitrous acid acts as a reducing agent and is oxidized to nitric acid.
No gas is given off.
If the free nitrous acid is in excess, it will decompose into nitricacid and oxygen.
H2O + NO2- --> NO3- + 2H+ + 2e- (electron loss)
If the free nitrous acid is in excess, decomposition occurs.
8. Add ammonium chloride solution to sodium nitrite solution.
Heat the solution and observe the effervescence and the colourless odourless gas given off.
The gas gives negative results for the limewater test, litmus test, and lighted splint test.
The gas is nitrogen.
NH4+ + NO2- --> 2H2O + N2 (g).
9. Nitrosamines, produced by nitrous acid with secondary amines, can be formed in the gut when nitrites react with amino acids.
12.18.2.0 Prepare sulfides, S2-
1. Use an ignition tube with 3 cm of powdered sulfur and heat until melted.
Hook a strip of copper over the rim of the ignition tube so that its lower edge is just above the surface of the sulfur.
Heat to boil the sulfur and note the glow as copper sulfide forms on the copper.
2. Pass hydrogen sulfide into copper (II) sulfate solution.
Filter off the precipitated copper (II) sulfide.
Cu2+ + S2- --> CuS (s).
3. This reaction requires the destruction of an ignition tube.
To avoid this loss, some teachers do this experiment on tin lids in the open laboratory, but this practice is not recommended.
Put 2 g of the mixture into a borosilicate test-tube.
Insert a plug of mineral wool into the mouth of the test-tube.
Clamp the test-tube and heat the powder mixture.
When an orange glow is seen, stop heating the mixture, and leave to cool.
The test tube can be broken open using a pestle and mortar.
See diagram FeS
4. Prepare sulfides of iron, cobalt and nickel.
Prepare solutions of iron (II) cobalt and nickel salts.
Pass hydrogen sulfide into each solution Note a slight precipitate of dark iron (II) sulfide, but no precipitate with the cobalt or nickel salts.
Add ammonia solution, NH3 (aq) ("ammonium hydroxide") and pass more hydrogen sulfide through the three solutions.
All three solutions form a black precipitate of the metallic sulfide, however only iron (II) sulfide dissolves in dilute hydrochloric acid.
The sulfides of cobalt and nickel dissolve in concentrated hydrochloric acid in the presence of potassium chlorate or in aqua regia.
Transfer the sulfides to evaporating basins, add concentrated hydrochloric acid and a crystal of potassium chlorate.
Heat until the crystals dissolve.
The cobalt salt becomes is pink in solution.
The nickel salt becomes yellow-green.
12.18.3.1 Prepare sulfur monochloride, S2Cl2
This experiment may not be allowed in some school systems.
Put 10 cc of sulfur in a distilling flask.
Use a one-hole stopper fitted with a delivery tube to reach the level of the sulfur.
Connect the delivery tube to a supply of dry chlorine.
Heat the sulfur on a gauze and pass in chlorine.
Stand the flask in a water bath of cold water.
Collect the liquid product of sulfur monochloride, S2Cl2, in a dry test-tube.
Heat drops of the product in water and tests for sulfur dioxide and hydrochloric acid.
Note the deposit of sulfur.
2S2Cl2 + 2H2O --> 4HCl + 3S (s) +SO2 (g).
12.16.5 Reactions of sulfuric acid
Sulfuric acid acts as a dehydrating agent, removing water, or the elements of water, from another substance.
1. Add 2 cm of concentrated sulfuric acid to 1 cm of copper sulfate crystals.
After ten minutes, note the colour change from blue copper sulfate crystals to white anhydrous copper sulfate.
CuSO4.5H2O + (H2SO4) --> CuSO4 + (H2SO4.5H2O).
2. Heat a mixture of 0.5 cm of sucrose and 1.0 cm of concentrated sulfuric acid gently for 2 seconds and then leave to stand.
Note the vigorous reaction and the colour change from white sugar to black carbon.
C12H22O11 + (H2SO4)--> 12C + (H2SO4.11H2O)
Sulfuric acid acts as a displacer of acids from their salts, sulfuric acid being much less volatile than most other acids.
3. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium chloride.
Test the fuming gas with silver nitrate solution on a glass rod to form white silver chloride with hydrogen chloride.
The less volatile sulfuric acid displaces the hydrogen chloride.
NaCl + H2SO4 --> HCl + NaHSO4.
4. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium acetate.
Note the smell of the displaced the acetic acid.
5. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium formate.
Note the displacement of formic acid followed by dehydration.
Sulfuric acid acts as a displacer of acids from their salts, sulfuric acid being much less volatile than most other acids.
Also, sulfuric acid acts as an oxidizing agent.
6. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium bromide.
A fuming gas first forms then a brown gas.
Hydrogen bromide is displaced then partially oxidized to bromine.
Hydrogen bromide turns silver nitrate on a glass rod to pale yellow silver bromide.
Potassium permanganate solution on a glass rod decolorizes the sulfur dioxide, formed by reduction of sulfuric acid.
KBr + H2SO4 --> HBr + KHSO4
2HBr + H2SO4 --> Br2 (g) + 2H2O + SO2
4H + + 2Br- + SO42- --> Br2 (g) + 2H2O + SO2 (g).
7. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium iodide.
A fuming gas first forms then a brown gas.
Hydrogen iodide is displaced then oxidized to iodine.
Note the greater extent of oxidation compared with the previous experiment.
Much of the hydrogen iodide is oxidized to iodine.
Heat the test-tube and note the violet vapour of iodine.
4H + + 2I- + SO42- --> I2+ 2H2O + SO2.
8. Reactions of dilute sulfuric acid as an acid.
Add 2 cm of dilute sulfuric acid to 1 cc of zinc powder.
Close the test-tube with the thumb until enough hydrogen forms to give a mild explosion when the mouth of the test-tube is held in a flame.
2H + + Zn (s) --> Zn2+ + H2 (g).
9. Add 2 cm of sodium carbonate to 1 cc of zinc powder.
Test for carbon dioxide by passing the gas given off to pass into limewater that turns milky due to the fine precipitate of calcium carbonate.
2H + + CO32- --> H2O +CO2 (g)
Ca(OH)2 + CO2 (g) --> CaCO3 (s) +H2O.
10. Reactions of dilute sulfuric acid as a sulfate.
Add an equal volume of barium chloride solution to 3 cm of dilute sulfuric acid.
Note the white precipitate of barium sulfate.
Allow the precipitate to settle, filter, wash and leave to dry.
SO42- + Ba2+ --> BaSO4 (s).
12.16.7 Sulfuric acid dehydrates copper (II) sulfate crystals
1. Sulfuric acid removes water, or the elements of water, from another substance
Add 2 cm of concentrated sulfuric acid to 1 cm of copper (II) sulfate crystals.
After ten minutes, note the colour change from blue copper (II) sulfatecrystals to white anhydrous copper (II) sulfate.
CuSO4.5H2O + (H2SO4) --> CuSO4 + (H2SO4.5H2O).
2. Add drops of sulfuric acid to blue copper (II) sulfate crystals.
The crystals turn white as they lose water.
Concentrated sulfuric acid combines so readily with water that it can be used as a dehydrates agent, It can remove water from hydrated copper (II) sulfate crystals and from other hydrated salts.
CuSO4.5H2O (s) <--> CuSO4 (s)+ 5H2O (l).
12.16.6 Sulfuric acid as a displacer
Sulfuric acid is much less volatile than most other acids.
Also, sulfuric acid acts as an oxidizing agent.
1. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium bromide.
A fuming gas first forms then a brown gas.
Hydrogen bromide is displaced then partially oxidized to bromine.
Hydrogen bromide turns silver nitrate on a glass rod to pale yellow silver bromide.
Potassium permanganate solution on a glass rod decolorizes the sulfur dioxide, formed by reduction of sulfuric acid.
KBr + H2SO4 --> HBr + KHSO4
2HBr + H2SO4 --> Br2 (g) + 2H2O + SO2
4H + + 2Br- + SO42- --> Br2 (g) + 2H2O + SO2 (g).
2. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium iodide.
A fuming gas first forms then a brown gas.
Hydrogen iodide is displaced then oxidized to iodine.
Note the greater extent of oxidation compared with the previous experiment.
Much of the hydrogen iodide is oxidized to iodine.
Heat the test-tube and note the violet vapour of iodine.
4H + + 2I- + SO42- --> I2 + 2H2O + SO2.
3. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium chloride.
Test the fuming gas with silver nitrate solution on a glass rod to form white silver chloride with hydrogen chloride.
The less volatile sulfuric acid displaces the hydrogen chloride.
NaCl + H2SO4 --> HCl + NaHSO4.
4. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium acetate.
Note the smell of the displaced the acetic acid.
5. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium formate.
Note the displacement of formic acid followed by dehydration.
12.16.10 Sulfuric acid with sodium chloride
In absence of water
NaCl + H2SO4 --> NaHSO4 + HCl (at room temperature)
NaCl + NaHSO4 --> HCl + Na2SO4 (if reaction proceeds above 200oC).
12.18.6.3 Reactions of sulfamic acid
Sulfamic acid, NH2.SO2OH
A tautomer is when an atom, e.g. hydrogen, moves backwards and forwards between different places on a molecule, the new and original molecules form a tautomeric pair.
(H3NSO3) (tautomer: NH2.SO2OH) (amidosulfonic acid, amidosulfuric acid, aminosulfonic acid, sulfamidic acid)
Sulfamic acid is used in preparations to clean stainless steel and copper utensils to remove hard water scale.
It is also used to make sweeteners.
1. Dissolve 1 cc of sulfamic acid in 2 cm of water.
Note the high solubility of the acid.
Tests for sulfate ion by adding dilute hydrochloric acid and drops of barium chloride solution.
At first there is little action, but leave to stand and white suspension of barium sulfate forms.
Boil the mixture and the barium sulfate becomes more apparent as the sulfamic acid hydrolyses.
NH2.SO2.OH + H2O --> NH4HSO4.
2. Dissolve 1 cc of sulfamic acid in 2 cm of water.
Dissolve 1 cc of sodium nitrite in 2 cm of water.
Mix the solutions.
Note the vigorous effervescence as nitric oxide, nitrogen dioxide and nitrogen are given off.
Sulfamic acid is a strong fully ionized acid that reacts with the nitrites to give oxides of nitrogen and its -NH2 group.
Sulfamic acid also reacts with the nitrite to give nitrogen.
2H + + 2HNO2 + 2e- --> 2H2O + 2NO
2NO + O2 --> 2NO2
NH2.SO2.O- + H+ + NO2- --> N2 (g) + HSO4- + H2O
Tests for sulfate ion by adding dilute hydrochloric acid and drops of barium chloride solution.
A white suspension of barium sulfate forms.
3. Add sodium hydroxide solution to 1 mL of sulfamic acid to a depth of 2 cm for an excess of sodium hydroxide.
Heat the solution and test the gas formed for ammonia with damp red litmus paper.
NH2.SO2.O- + 2OH- --> NH3 (g) + SO42- + H2O.
4. Heat 1 cc of sulfamic acid in a dry test-tube.
Tests for sulfur dioxide with a spot of potassium permanganate on a filter paper.
Test for sulfur trioxide by allowing the white fumes to flow into a test-tube containing barium chloride solution acidified with hydrochloric acid.
Note the crystalline sublimate and dissolve the crystals in 2 cm of sodium hydroxide solution.
Heat the solution then tests for ammonia with damp red litmus paper.
Acidify the remaining solution with hydrochloric acid and add barium chloride solution to tests for sulfate ion.
5. Sulfamic acid solution in water is unstable and forms ammonium bisulfate.
However, the colourless crystalline solid is stable, not hygroscopic, and is very soluble in water to form the zwitterion (H3N+SO3-).
On heating, the solution produces ammonia.
6. Sulfamic acid reacts with nitrous acid to form nitrogen
HNO2 + (NH2)HSO3 --> H2SO4 + N2 + H2O.
7. Sulfamic acid reacts with nitric acid to form nitrous oxide
HNO3 + (NH2)HSO3 --> H2SO4 + N2O + H2O.
8. Sodium hydroxide solution is standardized by titration with primary standard sulfamic acid solution
NaOH + (NH2)HSO3 --> NaNH2SO3 + H2O.
Table 12.19.1.0
Element ->
 Fluorine Chlorine Bromine Iodine
Description ->
-
 yellow gas,
irritating smell
 green-yellow gas
irritating smell
 red liquid,
irritating smell
 black solid,
irritating smell
14.19.1.0
Prepare
halogens
 electrolysis
of KHF2
-
 heat mixture of
chloride, MnO2
& conc. H2SO4		 heat mixture of
bromide, MnO2
& conc. H2SO4		 heat mixture of
iodide, MnO2
& conc. H2SO4
Activity
-
-
 very reactive, combines
with most metals
& non-metals
 very reactive, combines
with most metals
& non-metals
 very reactive, combines
with most metals
& non-metals
 reactive, combines
with most metals
& few non-metals
Replacing action
-
 replaces all
halogens
 replaces Br2 & I2 from
bromides & iodides
 replaces iodine from
iodides
 -
-
Oxidizing action
-
 very powerful
oxidizing agent
 very powerful
oxidizing agent		 powerful
oxidizing agent		 weak
oxidizing agent
Action with alkalis
 forms fluoride
 forms hypochlorite
 forms hypobromite
 forms hypoiodite
Halogen aciid
-
 fuming gas,
weak acid
 fuming gas,
strong acid
 fuming gas,
strong acid
 fuming gas,
strong acid