Gases, prepare gases, greenhouse gases

School Science Lessons
(topic 13
2023-09-16

Gases
Contents
13.3.0 Prepare gases with gas generation apparatus
13.3.1 Gas bags
13.3.2 Hydrogen chloride
13.3.3 Hydrogen cyanide
13.3.4 Hydrogen gas
13.3.5 Hydrogen sulfide
13.3.6 Hydrochloric acid
13.3.7 Nitric acid with carbonates and bicarbonates, (Dilute nitric acid)
13.3.8 Nitric acid with copper, (Dilute nitric acid)
13.3.9 Nitric acid with copper, (Concentrated nitric acid)
13.3.10 Nitric acid with metal oxides, (Dilute nitric acid)
13.3.11 Nitric acid with metals, (Dilute nitric acid)
13.3.12 Nitric acid with sulfuric acid, (Concentrated acids)
13.3.13 Phosgene gas
13.3.27 Prepare hydrochloric acid with sodium chloride
13.3.14 Prepare hydrogen chloride / hydrochloric acid
13.3.15 Prepare hydrogen gas
13.3.16 Prepare hydrogen sulfide
13.3.17 Prepare nitric acid
13.3.18 Prepare nitric oxide
13.3.19 Prepare nitrogen dioxide
13.3.20 Prepare nitrogen gas
13.3.21 Prepare nitrous acid
13.3.22 Prepare nitrous oxide
13.3.23 Tests for dinitrogen oxide, nitrous oxide, N2O
13.2.24 Tests for hydrogen chloride
13.3.25 Tests for hydrogen gas
13.3.26 Tests for hydrogen sulfide solution

13.3.0 Prepare gases with gas generation apparatus
See diagram 1.13a: Simple fume hood
See diagram 3.32: Gas generation apparatus
See diagram 3.33: Collect gas with an angle-tube syringe
Be careful! Prepare poisonous gases only in a fume hood or fume cupboard, e.g. preparation of chlorine, hydrogen chloride, sulfur dioxide.
1. See diagram 3.32: No. 1.
Collect more dense gas by upward displacement of air (downward delivery), if molecular mass > 28.8.
The more dense gas sinks down into, and displaces, the less dense air upwards, e.g. preparation of carbon dioxide, nitrogen dioxide.
2. See diagram 3.32: No. 2.
Collect less dense gas by downward displacement of air (upward delivery), if molecular mass < 28.8.
The less dense gas rises into, and displaces, the more dense air downwards, e.g. preparation of hydrogen, ammonia.
Use a borosilicate test-tube that is not cracked.
Clamp the test-tube to a stand.
Put the solid reagent in the sidearm test-tube and the liquid reagent in the reservoir.
Add the liquid reagent very slowly drop by drop.
Keep the reservoir tap closed and the reservoir full to prevent gases blowing back.
Grease the stopper and insert it so that if an accidental sudden increase in pressure occurs, the stopper blows out of the test-tube.
Use rubber tubing to collect the sidearm to a delivery tube that leads into the receiving test-tube.
Discard the first gas coming out of the delivery tube, because it is mostly air.
Never allow a flame near the gas as it comes out of the delivery tube.
Some air probably remains in the receiving test-tube.
Use the gas bubbler to collect over water insoluble gases with similar density to air.
Some water vapour remains in the receiving test-tube.
Gases can also be collected in balloons, inflatable footballs, and plastic bags.
3. See diagram 3.32: No. 3.
Collect insoluble gas of any density over water, (by downward displacement of water, water displacement), e.g. preparation of oxygen.
Fill one third of the water trough.
Fill a test-tube with water by placing it on its side in the water trough.
Put your thumb over the end of the test-tube and invert it.
Fix the end of the gas delivery tube inside or under the test-tube.
If the gas is slightly soluble in water its solubility will be less in warmer water.
Some gases dissolve in water to produce heat and form an acid solution, e.g. HCl, SO2, NO, NO2.
Some gases dissolve in water to form a basic solution, e.g. NH3.
4. See diagram 3.32: No. 4.
Collect soluble gases in water (aqueous solution), e.g. Cl2.
5. See diagram 3.32: No. 5.
When the gas preparation equipment uses downward displacement of water (Diagram 3.) or collection in water (Diagram 4.),
water may be forced back into the equipment by atmospheric pressure i.e. "sucked back", and break hot glassware or dilute reactants.
To prevent sucking back, use an inverted glass filter funnel.
6. See diagram 3.33, No. 6.
Collect gas with an angle tube syringe that can collect and store gas of any density.

13.3.1 Gas bags
See diagram 13.01: Gas bag, cable tie
1. Party balloons can be inflated with gas only from a high pressure source, e.g. a gas cylinder.
2. Snap lock resealable polythene bags.
They can be resealed with a finger press sealing strip to give a gas tight seal.
The closure system, which reseals an opened bag includes a pressure sensitive adhesive on the front side, and a defined release surface on the back.
The top portion of the bag is folded so that the defined release surface comes into contact with the adhesive to reseal the opened bag.
3. The plastic bag used in a 2 litre wine cask can be washed out and used to store gas.
Use a cork borer to insert a glass tube through a one hole rubber stopper.
Be Careful! Leave 1 cm of glass tube to protrude from the top of the stopper.
Pull tight around the neck of the plastic bag around the rubber stopper and secure it tightly with a cable tie.
To check for leaks, close the end of the glass tube with a rubber cap, immerse the bag in water and squeeze the bag.
To fill the bag, squeeze it flat then fill it from a gas cylinder or chemical generator.
Refill the bag and squeeze out the gas more than once to ensure that any air is flushed out.
When the bag is finally filled, close the glass tube with a rubber cap.
To get a gas sample, inject the hypodermic needle of a syringe through the rubber cap and suck gas into the syringe.
A cable tie usually consists of a Nylon tape with a gear rack and a ratchet within a small open case.
When the pointed tip of the cable tie has been pulled through the case and past the ratchet, it cannot be pulled back.
But the loop formed may be pulled tighter.
Cable ties are used to bind several cables together, e.g. cables around a motor car engine.

13.3.2 Hydrogen chloride
Hydrogen chloride, HCl, Toxic, corrosive, highly irritating gas
Hydrogen chloride gas is corrosive.
Do not prepare hydrogen chloride in an open room.
Hydrogen chloride gas has a choking odour, because it combines with the water vapour in the air to form hydrochloric acid.
Concentrated sulfuric acid reacts with metal chlorides to form hydrogen chloride that dissolves in water to form hydrochloric acid.
Be careful! Prepare hydrogen chloride gas, HCl, in a fume hood!
13.3.14 Prepare hydrogen chloride / hydrochloric acid
13.3.27 Prepare hydrochloric acid with sodium chloride
13.2.24 Tests for hydrogen chloride

13.3.3 Hydrogen cyanide
See diagram 16.2.4.2.2: Hydrogen cyanide
Hydrogen cyanide, HCN, H-C≡N (note triple bond), (formerly prussic acid), hydrocyanic acid (especially as a vapour), toxic, Not permitted in schools
It is an extremely poisonous chemical, and is lethal in very small doses of a few hundred parts per million (ppm) in air.
It can kill a human in less than one hour, and if ingested can cause almost immediate death.
High concentrations of about 5.6% of HCN gas exposed to air is explosive.
Hydrocyanic acid, HCN
Hydrocyanic acid is an aqueous solution of hydrogen cyanide HCN, poisonous weak acid, used for industrial fumigating
Hydrocyanic acid, hydrochloride, CH2ClN
Salts of cyanide anion are called cyanides.

13.3.4 Hydrogen gas
Hydrogen, H, (Greek hydro water), is a colourless odourless gas at room temperature and pressure H2, lightest element.
It burns to form water, most common element in space, natural isotopes hydrogen and deuterium and manufactured isotope radioactive tritium.
It is the product of electrolysis of water.
It is used to fix nitrogen and make ammonia in the Haber process, reduction of ore oxides, manufacture of HCl, hydrogenation of oils.
Hydrogen is the elemental gas in balloons, and has possible use as hydrogen gas fuel in motor vehicles.
Atomic number: 1, Relative atomic mass: 1.0079, RD 0.070 (20 K), MP = -259oC, BP = -252oC.
Specific heat capacity: 1.43 × 104 J kg-1 K-1.
Hydrogen gas, H2, was first discovered by Theophrastus Paracelsus, (1493-1541, Swiss- German), and
Robert Boyle (1627-1691, England), who demonstrated the existence of "inflammable air"
Both chemists used sulfuric acid with iron to make hydrogen gas.
In 1766, Henry Cavendish (1731-1810, England), using iron with hydrochloric acid, showed that hydrogen is an element and a constituent of water.
Hydrogen gas, both as a compressed gas and a gas generated by experiment is highly flammable and can act as a non-toxic asphyxiant.
"Green" hydrogen gas genration
If the electricity is generated from renewable sources such as solar or wind, production of hydrogen in this way emits no greenhouse gasses.
The different "shades" of hydrogen:
"Brown hydrogen" is produced using coal where the emissions are released to the air.
"Grey hydrogen" is produced from natural gas where the associated emissions are released to the air.
"Blue hydrogen" is produced from natural gas, where the emissions are captured using carbon capture and storage.
"Green hydrogen" is produced from electrolysis powered by renewable electricity.

Experiments
13.3.4.1 Burn hydrogen gas
See diagram 13.1.4: Reduction tube
Safety
For these experiments, wear safety goggles, ear protection and have a fire extinguisher close by.
Ensure that all ignition of hydrogen occurs at least 2 metres away from any other source of hydrogen, and away from glass light fittings or glass-fronted cabinets.
Prepare hydrogen gas by reaction of an acid on an active metal, e.g. zinc with hydrochloric acid.
Hydrogen forms spontaneously combustible / explosive mixture in air at low temperature.
Hydrogen gas forms explosive mixtures with air when mixed in almost any proportions.
Pure hydrogen burns in air with a hot colourless flame without exploding.
However, mixtures of hydrogen and oxygen or air burn with explosions.
In experiments to demonstrate ignition of hydrogen, air / hydrogen mixtures are often ignited instead.
For example, when igniting a jet of hydrogen gas from a flask containing metal and acid, sufficient air may be left in the flask to form an explosive mixture so the whole flask explodes, throwing shards of glass in all directions.
Exploding soap bubbles containing hydrogen / oxygen mixture on the surface of a beaker of water is safe, if the delivery tube is at least 10 cm below the surface of the water.
13.3.4.2 Prepare hydrogen gas
Prepare hydrogen in test-tubes
Generate small quantities of gas in a test-tube only.
There may be some risk associated with popping air / hydrogen mixtures in test-tubes, because the test-tube may shatter.
Check that the test-tubes are not flawed.
Students may be injured by "popping" (burning) hydrogen in other glass containers.
Exploding about 1 mL of hydrogen / air or hydrogen / oxygen mixture is safe, if there is no possibility of broken materials being propelled outwards by the explosion.
13.3.4.3 Explode hydrogen gas
Explode a hydrogen gas balloon
Never explode hydrogen in a glass or solid container!
Exploding a hydrogen balloon is safe, but loud, because there are no hard pieces of shrapnel from the explosion.
Attach a string to the balloon clip and tether it to a rail away from any combustible material.
All other observers should be at least 5 metres away from the explosion.
Attach a small birthday candle to a long stick or thin dowel to be a safe distance from the ignition.
Extend your arms and hold the lighted candle at the end of the dowel below the balloon.
Tie the string of the 30 cm hydrogen balloon to a weight on the desk, well away from any chemicals or combustible substances.
Keep a soda acid fire extinguisher nearby.
With your safety goggles and ear protection on, light the candle and hold it under the floating balloon at arm's length.
2H2 + O2--> 2H2O + energy
13.3.4.4 Tests for hydrogen gas
Na (s) + 2(C2H5OH) (l) --> 2(C2H5ONa) (s) + H2 (g)
sodium + ethanol --> sodium ethoxide + hydrogen
Evaporate the sodium ethoxide solution to form white crystals.
Add drops of water and test for litmus that turns blue.

13.3.5 Hydrogen sulfide
Acids react with sulfides to produce hydrogen sulfide, (rotten egg gas), and a salt.
Experiments
13.3.16 Prepare hydrogen sulfide
13.3.26 Tests for hydrogen sulfide solution

13.3.6 Hydrochloric acid, HCl
Hydrochloric acid, ACS reagent, 37% | Hydrogen chloride., ≥99.8% Hydrochloric acid is a solution of hydrogen chloride in water
Hydrogen chloride occurs as either a colourless liquid with a an irritating, pungent odour, or a colourless to slightly yellow gas which can be shipped as a liquefied compressed gas; highly soluble in water.
Hydrochloric acid, HCl (aq). spirits of salts, muriatic acid, clear, colourless, fuming liquid, conc. 12 M, RD 1.18 gm cm-3, 12 mol dm-3, E507, dilute HCl 6 M (solution of hydrogen chloride gas), distinct odour, (aqua regia: 3 vols concentrated HCl + 1 vol concentrated HNO3)
Concentrated hydrochloric acid releases hydrogen chloride gas if open to the atmosphere.
This gas is highly irritant to the lungs and causes coughing.
Use the concentrated gas in a fume cupboard to avoid inhalation.
If a container of the concentrated acid is broken outside a fume cupboard, evacuate the area until the fumes have dissipated.
Small amounts of hydrogen chloride gas are relatively harmless and a <0.1 M dilute acid has low toxicity.
The human stomach forms hydrochloric acid as part of the digestive process.
Concentrated hydrochloric acid is extremely damaging to the eyes.
Wear a face shield when doing any pouring experiments that might result in splashing of the liquid.
Avoid skin contact with hydrochloric acid.
Wash any area of the body that has been in contact with hydrochloric acid with copious amounts of water.
Dilution of hydrochloric acid with water does releases some heat, but not as much as sulfuric acid.
Always dilute by adding acid to water, not water to acid.
Do not mix hydrochloric acid with formaldehyde solution, because highly carcinogenic bis(chloromethyl) ether may form.
Do not store bottles of hydrochloric acid near formaldehyde, because the reaction might occur in the air.
Do not dispose of formaldehyde solutions and hydrochloric acid in the same sink system.
Common name: Muriatic acid (25% solution) (in concrete bleach, silverware cleaning solutions, drain cleaners

13.3.7 Nitric acid with carbonates and bicarbonates, (Dilute nitric acid)
Put 2 cm of chalk (calcium carbonate), washing soda (sodium carbonate), or baking soda (sodium hydrogen carbonate, bicarbonate) into a test-tube.
Add drops of dilute nitric acid.
Observe the vigorous effervescence, because of the formation of carbon dioxide.
Tests the gas with limewater.
CaCO3 (s) + 2HNO3 (aq) --> Ca(NO3)2 (aq) + CO2 (g) + H2O (l)

13.3.8 Nitric acid with copper, (Dilute nitric acid)
Nitric acid dissolves most metals, usually producing copious brown fumes of toxic nitrogen dioxide.
However, the composition of the gases produced depends on the concentration of the nitric acid.
The reaction of copper with dilute nitric acid produces colourless nitric oxide, NO.
The nitric oxide reacts with air at the top of a test-tube to form brown nitrogen dioxide, NO2.
With concentrated nitric acid, copper forms mostly brown NO2 gas.
Nitrogen dioxide is a powerful lung irritant.
All reactions that result in the formation of nitrogen dioxide should be done in a fume cupboard or a well-ventilated area.
Very dilute nitric acid may react with very active metals, e.g. magnesium to form hydrogen gas.
When nitric acid reacts with most metals, it oxidizes the hydrogen to water.
Experiment
Add drops of dilute nitric acid to copper.
Nitrogen monoxide forms, which immediately reacts with oxygen gas in the air to form nitrogen dioxide.
3Cu (s) + 8HNO3 (aq) --> 3Cu(NO3)2 (aq) + 4H2O (l) + 2NO (g)
2NO (g) + O2 (g) --> 2NO2 (g)

13.3.9 Nitric acid with copper, (Concentrated nitric acid)
Nitric acid reacts with metals above platinum in the reactivity series, but does not form hydrogen gas.
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD..
Pour drops of concentrated nitric acid on pieces of copper in a test-tube.
Put a stopper on the test-tube immediately, because brown nitrogen dioxide gas forms.
The nitric acid acts as an oxidizing agent and is reduced to nitrogen dioxide and water.
The reaction is exothermic.
Cu (s) + 4HNO3 (aq) --> Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g)

13.3.10 Nitric acid with metal oxides, (Dilute nitric acid)
Heat black copper oxide with 2 cm of dilute nitric acid in a test-tube.
The copper oxide disappears and a blue solution of copper nitrate forms.
CuO + 2HNO3 --> CuNO3 + H2O

13.3.11 Nitric acid with metals, (Dilute nitric acid)
1. Add slowly small pieces of copper, magnesium and zinc to small amounts of dilute nitric acid in separate test-tubes.
If no change is taking place, gently heat the mixture.
Repeat the procedure 1. with concentrated nitric acid, 2. with concentrated sulfuric acid, and 3. with concentrated hydrochloric acid.
The reactions of metals with nitric acid and concentrated sulfuric acid are different from reactions of metals with hydrochloric acid, dilute sulfuric acid and dilute acetic acid.
Although copper does not react with dilute acids or with concentrated hydrochloric acid, it does react with both dilute and concentrated nitric acids and with hot concentrated sulfuric acid, but does not produce hydrogen gas in reaction with them.
The residual mixtures contain solutions of salts, but writing equations for the reactions is difficult, because more than one reaction can occur simultaneously between copper or magnesium or zinc and nitric acid.
For example, when zinc reacts with nitric acid the reaction may produce nitrogen dioxide, nitric oxide, nitrous oxide, zinc nitrate and ammonium nitrate.
2. Put strands of copper wire from an old electric light flex into a test-tube with 2 cm of dilute nitric acid.
Heat the test-tube gently until effervescence starts and then stand the test-tube in the test-tube rack or a beaker.
The effervescence is caused by the formation of oxides of nitrogen, principally nitric oxide, NO, and nitrogen peroxide, NO2, the first of these gases is colourless, while the second consists of brown fumes.
In minutes, the copper has dissolved in the acid and a blue-green solution of copper nitrate is left.
3Cu (s) + 8HNO3 (aq) --> 3Cu(NO3)2 (aq) + 2NO (g) + 4H2O (l)
With concentrated nitric acid, the copper is oxidized to Cu2+ ions and the nitric acid is reduced to nitrogen dioxide.
Cu (s) + 4HNO3 (aq) --> CuNO3)2 (aq) + 2NO2 (g) + 2H2O (l)
The solution is dark brown while the nitrogen dioxide is forming.
When all the copper has reacted, with the addition of water, the solution turns pale blue.
Aluminium and zinc do not, or hardly, react with concentrated nitric acid, because an oxidation layer forms.
However, zinc reacts with dilute nitric acid.

13.3.12 Nitric acid with sulfuric acid, (Concentrated acids)
The stronger concentrated sulfuric acid donates a proton, H+, to the weaker concentrated nitric acid, which then loses a water molecule, leaving the NO2+.
HNO3 + H2SO4 --> NO2+ + H2O + 2HSO4-

13.3.13 Phosgene gas
Phosgene, carbonyl chloride, carbonyl dichloride, CoCl2
Phosgene is used for the synthesis of isocyanate-based polymers, carbonic acid esters, acid chlorides, dyestuff manufacture and insecticides.
It is a poisonous gas and was used as a chemical warfare agent in the First World War.
In low concentrations it smells like cut grass.
Colourless gas or volatile liquid, extremely toxic, heated containers may rupture, dye, pesticide, corrosive.
. It is manufactured by direct combination of carbon monoxide with chlorine, with carbon catalyst.
At high temperate the reverse reaction occurs.
CO + Cl2 --> CoCl2
Chloroform + UV --> phosgene, so chloroform is kept in dark glass bottles

13.3.14 Prepare hydrogen chloride / hydrochloric acid
| See diagram 3.42: Collecting hydrogen chloride
| See diagram 13.3.9: Prepare hydrochloric acid
| See diagram 1.13a: Simple fume hood
Be careful!
Prepare hydrogen chloride gas only in a fume hood or fume cupboard!
Hydrogen chloride gas has a choking odour, because it combines with the water vapour in the air to form hydrochloric acid.
Concentrated sulfuric acid reacts with metal chlorides to form hydrogen chloride that dissolves in water to form hydrochloric acid.
1. Put sodium chloride crystals in a 100 mL filter flask or sidearm test-tube.
Coarse rock salt causes less frothing than the fine salt.
Carefully add concentrated sulfuric acid down a funnel to just cover the sodium chloride crystals.
Heat the mixture if necessary.
Collect the hydrogen chloride gas in test-tubes by upward displacement of air then put a stopper in the receiving test-tube, and put the end of the delivery tube into water to absorb excess hydrogen chloride.
NaCl (s) + H2SO4 (aq)--> HCl (g) + NaHSO4 (aq)
Repeat the experiment with concentrated hydrochloric acid and concentrated sulfuric acid.
Be careful!
2. Prepare hydrogen chloride gas by gently warming hydrochloric acid in a water bath in a flask with a gas collection tube.
Collect the gas by displacement of air.
Hydrogen chloride can be used in place ammonia in the ammonia fountain in the ammonia fountain experiment.
3. Hydrogen chloride gas fumes in air, forming droplets of hydrochloric add, so be careful not to inhale it.
Mix well together, on a creased sheet of paper, a finger width of sodium hydrogen sulfate and the same quantity of sodium chloride and transfer the mixture to a test-tube fitted with stopper and delivery tube.
Heat the mixture by keeping the test-tube moving in the flame to prevent the glass cracking.
The misty fumes of the heavy gas pass down wards into the second test-tube.
When it is full, as shown by the fumes coming out at the top, stopper it, and collect another as a spare.
Hold a piece of blue litmus paper in the fumes.
The blue litmus turns red.
The misty fumes are minute droplets of hydrochloric acid.
formed by the reaction of the invisible hydrogen chloride with water vapour in the air.
It is this acid that turns the blue litmus red.
4. Concentrated sulfuric acid reacts with sodium chloride to form hydrogen chloride gas, and can be reduced with copper metal to form sulfur dioxide gas on gentle heating.
Do these experiments in a fume cupboard while wearing eye protection.

13.3.15 Prepare hydrogen gas
See diagram 3.41: Collecting hydrogen gas
Hydrogen, H2, is a colourless odourless diatomic gas with the lowest density of any element.
Hydrogen does not change the colour of moist litmus.
The hydrogen ion, H+, is a proton.
Do not allow direct combination of hydrogen and chlorine in bright light or ignition of the mixture by lighted taper or electric spark.
You can ignite a jet of hydrogen issuing from a delivery tube.
Hydrogen reduces metal oxides.
Hydrogen, H2, is a colourless odourless diatomic gas with the lowest density of any element.
Hydrogen does not change the colour of moist litmus.
The hydrogen ion, H+, is a proton.
Experiments
1. Prepare hydrogen gas by reaction of an acid on an active metal, e.g. zinc with hydrochloric acid.
See diagram 3.2.33: Zinc with hydrochloric acid
Do not use a container bigger than a test-tube.
Check that the test-tubes are not chipped or cracked.
Put granulated zinc in a small borosilicate test-tube and cover it with water.
Add a crystal of copper (II) sulfate to act as a catalyst.
The copper sulfate first reacts with zinc in a displacement reaction to form metallic copper, which has a catalytic action.
Slowly add dilute hydrochloric acid through a funnel, or through a syringe.
Bubbles of hydrogen appear on the surface of the zinc.
The test-tube feels hot, because the reaction is exothermic.
Use a lighted taper to "pop"' the air / hydrogen mixture.
Collect hydrogen gas by downward displacement or over water.
Let the reaction continue for some minutes to drive out all the air from the test-tube.
Discard the first two test-tubes of hydrogen, because they will contain displaced air.
Collect test-tubes of the gas and apply stoppers.
Zn (s) + 2HCl (aq)--> ZnCl2 (s) + H2 (g)
2. Prepare hydrogen with iron filings and citric acid or sulfuric acid or sodium hydrogen sulfate
Put 1 cm depth of iron filings in a test-tube.
Just cover the iron filings with a dilute acid solution.
Warm the test-tube until frothing starts.
Hydrogen is colourless and odourless, but any impurities in the iron filings give a nasty smell.
To test for hydrogen gas, remove from heating, place your thumb over the end of the test-tube, count to five, apply a lighted paper to the end of the test-tube, the hydrogen gas explodes with a loud pop sound.
Never test more than one test-tube full of hydrogen gas!
3. Prepare hydrogen with aluminium foil and sodium carbonate
Cut into small pieces aluminium foil or aluminium milk bottle tops and put them into a test-tube.
Add 5 mL of sodium carbonate solution (Na2CO3.10H2O, washing soda).
Heat until effervescence occurs.
4. Prepare hydrogen with calcium and hydrochloric acid
Use forceps to transfer about 0.1 g of calcium metal turnings to dilute hydrochloric acid in a test-tube.
Ca (s) + 2HCl (aq)--> CaCl2 (aq) + H2 (g)
5. Prepare hydrogen with iron filings and potash alum
Put 5 g of iron filings in a 1 cm depth of alum solution in a test-tube.
Heat the solution until effervescence occurs.
Potash alum, "alum" has the formula: Al2(SO4)3.K2(SO4).24H2O, also shown as KAl(SO4)2.12H2O.
6. Prepare hydrogen with iron filings and ammonium chloride
Put an equal volumes mixture of iron filings and ammonium chloride in a dry test-tube and heat.
Hydrogen gas and ammonia are given off.
7. Prepare hydrogen with aluminium kitchen foil and caustic soda drain cleaner (sodium hydroxide)
Fill a clear glass bottle one third full with tap water.
Do not use a plastic bottle.
Cut a plastic bottle, with diameter greater than the glass bottle at half its length, to make a beaker.
Half fill the plastic beaker and stand the glass bottle in it.
With safety glasses and gloves, use a funnel to add three heaped teaspoons of caustic soda to the water in the glass bottle.
Keep the bottle open and swirl the bottle gently without spilling any reagents, then return the bottle to the plastic beaker.
The dissolving is exothermic so the contents of the bottle will get hot.
Roll three 20 × 30 cm sheets of kitchen aluminium foil into cylinders and then drop them into the bottle.
Inflate a 30 cm (helium quality) balloon to stretch the rubber.
Deflate the balloon and attach it to the mouth of the bottle, with another person holding the bottle.
Swirl the bottle contents again and return the bottle to the plastic beaker.
Let the balloon swell to 30 cm diameter then remove from the bottle, with another person holding the bottle, and attach a balloon clip with attached string.
If the balloon does not rise up when you hold the attached string, either the balloon is not yet at full capacity or the reaction has occurred too quickly, and condensed water vapour inside the balloon is weighing it down.
Wash the remaining reagents down the laboratory sink.
2Al (s) + 2 NaOH (aq) + 2H2O--> 2NaAlO2 + 3H2 +energy
aluminium + sodium hydroxide + water--> sodium aluminate + hydrogen gas + energy
8. Prepare hydrogen gas bubbles
Hydrogen is much lighter than air and was formerly used in airships, dirigible balloons.
It has now been replaced by helium, because hydrogen ignites easily.
Pass hydrogen through soapy water to form soap bubbles full of hydrogen.
Shake the bubbles gently to make them float up.
The hydrogen bubbles rise into the air, showing the low relative density of hydrogen gas.
Try to ignite the bubbles with a lighted splint.
9. Prepare hydrogen gas with citric acid
Put 2 mL of citric acid crystals and an equal amount of iron filings in a test-tube.
Add drops of water then heat the test-tube.
When the effervescence becomes brisk, remove the test-tube from the flame.
To test for hydrogen, put the thumb over the end of the test-tube, count to five, then apply a glowing splint to the end of the test-tube.
Repeat the experiment with small pieces of zinc instead of iron filings.
In both experiments the formation of the hydrogen is increased by adding drops of copper sulfate solution.
10. Aluminium with acids
Dissolve aluminium in heated dilute hydrochloric acid and note that hydrogen gas forms.
2Al + 6H + --> 2A13+ + 3H2 (g).
Hot concentrated sulfuric acid will attack aluminium with the production of sulfur dioxide.
Dilute or concentrated nitric acid acts only very slowly on aluminium.
11. Hydrogen gas generator from a boiling tube
See diagram 3.2.34: Small hydrogen generator
Hydrogen generators are not allowed in many school systems.
Make holes in the bottom of a boiling tube.
Heat the bottom of a boiling tube and a glass rod to red heat in a Bunsen burner flame.
Fuse the glass rod on to the bottom of the boiling tube then pull it away to form a shred of glass pulled out from the boiling tube.
Break off the shred of glass and form smooth rounded edges to the hole in a hot flame.
Make three holes.
Put granulated zinc into the boiling tube and fix a one-hole stopper with a delivery tube, rubber tube extension and screw clip.
If the zinc is in very small pieces put glass wool in the bottom of the boiling tube before adding the zinc.
Put the apparatus in a jar containing M sulfuric acid and drops of copper sulfate solution.
Open the clip to allow acid to enter the boiling tube and react with the zinc to form hydrogen gas.
Close the clip to allow pressure inside the boiling tube to prevent acid reacting with the zinc.
Zn (s) + H2SO4 (aq) --> ZnSO4 (aq) + H2 (g)

13.3.16 Prepare hydrogen sulfide
See diagram 1.13a: Simple fume hood
Hydrogen sulfide, Extremely Toxic, Highly flammable
Hydrogen sulfide, gas, < 1% Not hazardous
Hydrogen sulfide water, solution, Toxic if ingested
Hydrogen sulfide gas is both an irritant and an asphyxiant.
Collapse, coma and death from respiratory failure may come within a few seconds after one or two inspirations.
Do not prepare hydrogen sulfide in an open room.
Hydrogen sulfide is extremely toxic so use a fume cupboard.
You can ignite a jet of hydrogen sulfide issuing from a delivery tube.
Be careful! Hydrogen sulfide is an extremely poisonous colourless flammable gas with an unpleasant smell of rotten eggs.
At less than 1% concentration the smell disappears.
So a student may be breathing in this poisonous gas without being aware of it.
Do NOT use a Kipp's apparatus for generating hydrogen sulfide.
Hydrogen sulfide has high acute short-term toxicity to aquatic life, birds, and animals.
It is soluble in water and organic solvents and will corrode metals.
Hydrogen sulfide is used in the manufacture of pulp and paper (digesting agent), in tanneries and in sulfide ores.
It is the main offensive smell in flatus produces after a diet of certain types of beans.
Experiment
1. This experiment is regarded as the safest method to prepare hydrogen sulfide gas.
Do this experiment in a fume cupboard, fume hood.
Prepare only a very small quantity of this gas.
Have a beaker full of a weak alkali ready to stop the reaction.
Add dilute hydrochloric acid to iron (II) sulfide.
Collect the gas over warm water by downward displacement.
FeS (s) + 2HCl (aq)--> FeCl2 (aq) + H2S (g)
Ignite the gas as it leaves the delivery tube.
2H2S (g) + 3O2 (g)--> 2SO2 (g) + 2H2O (l)
2. Warning! Do this experiment in a fume cupboard, fume hood.
Put 5 sodium thiosulfate crystals in a metal screw cap.
Heat the metal screw cap gently by holding it with pincers in a Bunsen burner flame, until the crystals have melted and solidified again with steam given off.
Be careful! Do NOT inhale gas directly from the metal screw cap.
With more careful heating, note the "rotten egg" smell of hydrogen sulfide.
Allow the metal screw cap to cool.
Moisten the white residue with a weak acid, e.g. vinegar.
The smell of hydrogen sulfide gas becomes stronger.
Dip a strip of clean newspaper in the copper (II) sulfate solution and hold it over the meal screw cap.
The paper turns black.

13.3.17 Prepare nitric acid
Add concentrated sulfuric acid to sodium nitrate.
Be Careful! Heat gently, nitric acid vapour forms.
NaNO3 (s) + H2SO4 (l) --> HNO3 (l) + NaHSO4 (s)
Ionization of hydrogen sulfate ion
HSO4- + H2O <--> H3O+ + SO42-

13.3.18 Prepare nitric oxide
Nitric oxide (nitrogen monoxide, NO)
See diagram 13.3.18: Preparing nitrogen monoxide | See diagram 1.13a: Simple fume hood
A pinch tap, B one-hole stopper, Copper pieces, D rubber gasket with holes, E dilute nitric acid
Be careful!
Prepare nitric oxide only in a fume hood or fume cupboard, or use small quantities near an open window or a well-ventilated area.
Use eye and skin protection when splashes are possible.
Nitrogen monoxide, NO, is a colourless toxic gas that reacts immediately with oxygen in the air to form brown fumes of nitrogen dioxide, NO2.
Nitric oxide has many physiological functions in the body, e.g. dilation of blood vessels.
Nitrogen monoxide is a colourless gas that forms brown fumes of nitrogen dioxide, NO2 in air.
1. Do this experiment in a fume cupboard, fume hood or near an open window.
Add drops of dilute nitric acid to copper.
Nitrogen oxide forms that immediately reacts with oxygen in the air to form nitrogen dioxide.
2. Do this experiment in a fume cupboard, fume hood.
Fit each end of a glass tube, diameter 2 cm and length 15 cm, and with two one-hole stoppers fitted with short delivery tubes.
Connect one delivery tube to a short rubber tube with a pinch tap on it.
Fix a 2 cm rubber gasket with holes drilled in it in the middle part of the glass tube.
Drop pieces of copper on the rubber gasket.
Remove the stopper from the other end of the glass tube and add dilute nitric acid, i.e. 1 to 3 ratio of concentrated nitric acid to water.
Add enough diluted nitric acid so that no space is left for an air bubble after replacing the stopper.
Invert the apparatus and clamp it vertically with a container under the glass tube.
During the reaction the colourless gas produced gradually presses the solution out of the tube.
When the copper pieces are no longer in contact with the nitric acid solution the reaction stops leaving nitrogen oxide in the upper portion of the glass tube and a blue solution in the small container.
Open the pinch tap to let air enter the tube and note the oxidation to form brown fumes of nitrogen dioxide.
The level of the solution rises to show that nitrogen dioxide dissolves in water.
To prevent air contamination, absorb the nitrogen dioxide in the tube with an alkali solution.
3Cu (s) + 8HNO3 (aq) --> 3Cu (NO3)2 (aq) + 4H2O (l) + 2NO (g)
2NO (g) + O2 (g) --> 2NO2 (g), [The O2 comes from the air.]
3. Repeat the experiment with cold 7M nitric acid
4. This experiment may be done in a syringe
Add 5 mL of 1.2 M solution of FeSO4 in 1.8 M H2SO4 to 0.27 g of solid NaNO3
The reaction mixture turns black and nitric acid gas is produced
Wash the gas through deionized water
NO2-+ Fe2+ (aq) + 2H+ --> NO (g) + Fe3+ (aq) + H2O

13.3.19 Prepare nitrogen dioxide
Nitrogen dioxide is a brown gas with a choking smell that may irritate the lungs and lead to death
It is commonly formed by the reaction of concentrated nitric acid with copper and the decomposition of metal nitrates and nitrites
See diagram 1.13a: Simple fume hood
Be careful! Prepare nitrogen dioxide only in a fume hood or fume cupboard
Liquid dinitrogen tetroxide, N2O4, BP 21.2oC, dissociates as a gas to form nitrogen dioxide, NO2
1. Pour drops of concentrated nitric acid on pieces of copper in a test-tube
Fix a stopper in the test-tube immediately, because the brown noxious gas, nitrogen dioxide, forms with a pungent irritating odour
The nitric acid acts as an oxidizing agent and is reduced to nitrogen dioxide and water
Collect the gas by upward displacement of air in a fume hood
Be careful! The reaction is exothermic
Cu (s) + 4HNO3 (aq) --> Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g)
2. Heat lead (II) nitrate crystals
The decomposition may be noisy
Nitrogen dioxide and oxygen form, leaving yellow lead oxide
Pb(NO3)2 (s) --> 4NO2 (g) + 2PbO (s) + O2 (g)
lead (II) nitrate --> nitrogen dioxide + lead oxide + oxygen
3. Pass nitrogen dioxide through water
Nitrogen dioxide is decomposed by water to form a mixture of nitric acid HNO3 and nitrous acid HNO2
Note the colour and odour of the water
Test the solution with litmus paper
2NO2 + H2O --> HNO3 + HNO2

13.3.20 Prepare nitrogen gas
Atmospheric nitrogen cannot be used directly by the body
Liquid nitrogen is used to freeze tissues for microscopic examination
Nitrogen is colourless, odourless, tasteless, neutral and nonreactive
Nitrogen does not support combustion, but magnesium and calcium will continue to burn in nitrogen to form nitrides
Nitrogen is manufactured by fractional distillation of air
Air contains about 78% of nitrogen
Ammonium nitrite is unstable so the reaction of saturated solutions of sodium nitrite with ammonium chloride can be used to prepare nitrogen
Be careful! This reaction can explode without warning
1. Heat gently a test-tube containing 5 mL of saturated ammonium chloride solution
Add saturated sodium nitrite solution drop by drop
The reaction is exothermic, so stop heating when some gas forms
The ammonium nitrite breaks down into nitrogen and water
NH4NO2 (s) --> N2 (g) + 2H2O (l)
2. Dissolve 20 g of calcium hypochlorite (bleaching powder) in 100 mL water by shaking then filter the mixture
Add 10 mL of concentrated ammonia solution to the calcium hypochlorite solution and heat the mixture
Collect the gas by upward displacement of air
Be careful! Chloramine and explosive nitrogen trichloride may be produced!
2NH3 + 3CaOCl2 --> N2 + 3H2O + 3CaCl2
3. Invert a bell jar over a stand in water and insert lighted phosphorus into the bell jar
Oxygen and phosphorus combine to form phosphorus pentoxide, (P4 + 5O2 --> P4O10)
The phosphorus pentoxide is absorbed by the water in what may be a violent reaction to form phosphoric acid, leaving the nitrogen
P4O10 + 6H2O --> 4H3PO4

13.3.21 Prepare nitrous acid
Weak acid prepared by acids on nitrites
barium nitrate + sulfuric acid --> barium sulfate + nitrous acid
Ba(NO2)2 + H2SO4 --> BaSO4 + 2HNO2
sodium nitrite + dilute hydrochloric acid --> sodium chloride + nitrous acid
NaNO3 (aq) + HCl (aq) --> NaCl (aq) + HNO2 (aq)
Heat nitrous acid
Heated nitrous acid decomposes to form nitric acid and nitrogen monoxide (nitric oxide).
nitrous acid --> nitric acid + nitrogen monoxide.
2HNO2 --> HNO3 + NO

13.3.22 Prepare nitrous oxide
(N2O, nitrous oxide, dinitrogen oxide, nitrogen (I) oxide, dinitrogen monoxide, laughing gas)
Be careful!
Prepare nitrous oxide only in a fume hood or fume cupboard
Dinitrogen oxide is a colourless gas that is soluble in water, with a sweet smell that does not change the colour of moist litmus
It was previously used as an anaesthetic in dentistry ("laughing gas") and as a propellant in aerosol sprays
Be careful! Do NOT inhale nitrous oxide
It can cause death
Be careful! If you heat ammonium nitrate to dryness, it may decompose with an explosion
So heat the ammonium nitrate to between 170oCand 240oC, and no higher, to avoid detonation
1. Put 2 g of ammonium nitrate in a dry large test-tube
Fit a stopper with a delivery tube in the test-tube
Heat the test-tube slowly and pass the gas formed through 5% iron (II) sulfate solution to remove any nitrogen monoxide
Collect the gas in a receiving test-tube
2. Heat 10 g sodium nitrate powder with 9 g ammonium sulfate
Collect the gas by upward displacement of air
NH4NO3 --> N2O + 2H2O

13.3.23 Tests for dinitrogen oxide, nitrous oxide, N2O
Put a glowing splint down in the receiving test-tube and note the gas relights the splint

13.2.24 Tests for hydrogen chloride
1. Ammonia test for hydrogen chloride
Hold a piece of cotton wool soaked in ammonia solution, NH3 (aq) at the mouth of a receiving test-tube.
Note the white cloud of ammonium chloride above the hydrochloric acid.
2. Fountain test for hydrogen chloride 1.
This test is similar to the ammonia fountain test.
Heat the end of a delivery tube and draw it out to form a fine jet.
Fill a flask with hydrogen chloride and close the flask with a one-hole stopper with a delivery tube.
Add litmus to alkaline water in a beaker.
Warm the flask gently to expand the gas and then hold the flask upside down with the lower end of the delivery tube in alkaline water.
Water soon sprays into the flask through the fine jet as the hydrogen chloride gas dissolves in the water, and the pressure of hydrogen chloride in the flask decreases.
The litmus in the water changes from blue to red.
2. Fountain test for hydrogen chloride 2.
Fill a beaker with litmus solution.
Fit a glass jet tube into the stopper of a flask.
Remove the stopper and jet, and start filling the dry flask with hydrogen chloride.
When the flask is full of gas, replace the stopper and jet, and quickly invert the flask with the other end of the jet tube in the litmus solution.
With the spirit burner at a safe distance, pour a finger width of methylated spirit on the flask and blow on it.
This causes the spirit to evaporate and thereby cool the flask and the gas inside it.
The contraction of the gas reduces its pressure, and atmospheric pressure forces litmus solution up the glass tube and out of the jet.
The fountain from the jet suddenly increases and the litmus changes colour.
The fountain from the jet suddenly increases for the reason given above.
The litmus changes from blue to red, because the water in the litmus solutionreacts with the hydrogen chloride to form hydrochloric acid.
3. Lighted splint test for hydrogen chloride
Hydrogen chloride extinguishes a lighted splint.
Hydrogen chloride neither burns nor supports combustion.
4. Litmus paper test for hydrogen chloride
* Test the solution in the receiving test-tube with moist litmus paper.
Red litmus paper turns blue.
* Hold a piece of blue litmus paper in the fumes.
The blue litmus turns red.
The misty fumes are minute droplets of hydrochloric acid, formed by the reaction of the invisible hydrogen chloride with water vapour in the air.
It is this acid that turns the blue litmus red.
5. Magnesium ribbon test for hydrogen chloride
Shake a receiving test-tube with water to form a solution of hydrogen chloride, hydrochloric acid.
Put a piece of magnesium ribbon in the solution.
Collect any gas formed and test for hydrogen with the glowing splint test.
6. Solubility test for hydrogen chloride
Remove the stopper from the receiving test-tube under water.
Note the solubility of hydrogen chloride.
Invert a receiving test-tube over water.
The gas dissolves immediately to form hydrochloric acid.
The water rises almost to the top, because collection by upward displacement of air results in some residual air remaining in the test-tube.
7. To show the extreme solubility of hydrogen chloride, remove the stopper from the test-tube, and quickly put your thumb or finger over the mouth of the test-tube.
Invert the test-tube of gas in a dish of water, removing your thumb only when the mouth of the test-tube is under the water.
Water rushes up into the test-tube.
Hydrogen chloride is so soluble that it dissolves almost at once in the water at the mouth of the test-tube.
Atmospheric pressure forces the water into the empty test-tube.

13.3.25 Tests for hydrogen gas
1. Lighted splint test for hydrogen gas
Be careful! A dangerous explosion may occur if you use anything bigger than a small test-tube when igniting the gas, particularly if the gas is mixed with air.
Never test more than a test-tube full of hydrogen gas.
Never dry hydrogen gas with concentrated sulfuric acid.
Hold a lighted splint or burning taper to the mouth of a test-tube.
The gas explodes with a squeaky pop sound.
The splint is extinguished.
The squeaky pop shows rapid combustion of hydrogen to form water vapour.
Look for vapour on the sides of the test-tube.
However, as 2 litres of gas forms only about 1 mL of liquid, the liquid on the sides of the test-tube may just show that test-tube was already wet before the experiment.
2H2 (g) + O2 (g)--> 2H2O (l)
hydrogen gas + oxygen gas--> water
2. Na (s) + 2(C2H5OH) (l) --> 2(C2H5ONa) (s) + H2 (g)
sodium + ethanol --> sodium ethoxide + hydrogen
Evaporate the sodium ethoxide solution to form white crystals.
Add drops of water and test for litmus that turns blue.
3. Pouring test for hydrogen gas
Test whether hydrogen is lighter than air by "pouring" the gas into a test-tube, held either above the first test-tube or below it.
Use a lighted taper to investigate where the hydrogen has gone.
4. Litmus test for hydrogen gas
Hydrogen does not change the colour of moist litmus.

13.3.26 Tests for hydrogen sulfide solution
See diagram 1.13a: Simple fume hood
Be careful!
The gas is soluble in water, so use a solution of hydrogen sulfide in water instead of the gas.
1. Odour test
Hydrogen sulfide has the odour of rotten eggs.
Be careful! Do NOT inhale gases directly from the test-tube.
Fan the gas towards the nose with the hand and sniff cautiously.
If you detect no odour, move closer and try again.
2. Lead (II) nitrate test
Hydrogen sulfide solution turns lead (II) nitrate solution test paper black.
3. Litmus test
Hydrogen sulfide solution turns blue litmus slightly pink-red.
4. Copper (II) sulfate test
Hydrogen sulfide solution turns copper (II) sulfate solution black.
Ionization of hydrogen sulfide
H2S + H2O--> H3O+ + HS-
HS- + H2O--> H3O+ + S2-5.
Hydrogen sulfide with lead ethanoate paper forms a black precipitate of lead sulfide.
H2S (g) + Pb (C2H3O2)2 (aq)--> PbS (s) + 2HC2H3O2 (aq)
lead (II) acetate + hydrogen sulfide--> lead (II) sulfide + acetic acid

13.3.27 Prepare hydrochloric acid with sodium chloride
1. Mix together 2 g of sodium chloride and 2 g of powdered alum (hydrated potassium aluminium sulfate), or sodium hydrogen sulfate (sodium bisulfate).
Put the mixture into a dry test-tube and have ready a damp blue litmus paper and a bottle of strong ammonia.
Heat the mixture over a medium Bunsen burner flame, holding the test-tube in a paper holder and moving the test-tube in the flame.
Hydrochloric acid gas, or hydrogen chloride forms as steam-like fumes.
Sniff the gas cautiously and put the blue litmus paper into the fumes.
Tests the gas, also, by removing the stopper from the bottle of strong ammonia and blowing the steamy fumes across the top of the bottle.
A dense white smoke forms.
The white smoke consists of ammonium chloride.
2. Repeat the above experiment.
Lead the delivery tube into a 500 mL bottle, half full of water.
Keep the end of the tube clear of the water to prevent sucking back.
As you heat the test-tube, hold the bottle in your other hand, and keep the water swirling to dissolve the hydrogen chloride gas.
Continue heating until no more gas forms.
Recharge the test-tube and repeat the procedure many times.
Label the bottle of dilute hydrochloric acid.