Preparing, Drying and Collecting Gases

There are 7 gases that you should familiarise yourself with in terms of preparation, collection and drying up:

1. hydrogen gas (or simply hydrogen), H₂
2. oxygen, O₂
3. carbon dioxide, CO₂
4. chlorine, Cl₂
5. hydrogen chloride, HCl
6. sulfur diox ide, SO₂
7. ammonia, NH₃

Preparation of Gases
Gases are prepared using chemical reactions which are represented by chemical equations. A quick recap on the structure of a chemical equation:

(left)          conditions           (right)
A + B        -------------->        C + D
reactants                                products

The substances on the left are the reactants which will be used up and be converted into the substances on the right, called the products. There could be one or more reactants and one or more products. The reaction arrow means "changing into". Conditions for the reaction such as temperature, pressure and the type of catalyst used may be written above the reaction arrow.

Experimental observations are provided so you can learn about writing them. Some items that should be described are:

• Any changes to the colour/pH of solution
• Any changes to any solid reactants
• Any effervescence being produced. If so, describe the colour and smell of the gas.
• Any precipitate being produced. If so, describe the colour and whether it dissolves in excess regents.
• If there are no changes, you will need to state that there are no visible changes.

Hydrogen, H₂

Hydrogen gas can be produced by reacting reactive metals with water or acids. The reactive metals are generally from Group I and Group II of the periodic table. The word equations and generic chemical symbol equations are provided below, along with some specific examples.

Metal + Water
metal + water -----> metal hydroxide + hydrogen
M(s) + H₂O(l) -----> M(OH)_x(aq) + H₂(g)

Example 1 - Group I Metal:
sodium + water -----> sodium hydroxide + hydrogen
2Na(s) + 2H₂O(l) -----> 2NaOH(aq) + H₂(g)

Observation: The silvery grey metal reacts vigorously, dissolving and forming a colourless alkaline solution that turns red litmus paper blue. Effervescence is observed. A colourless, odourless gas which extinguishes a lighted splint with a 'pop' sound is formed.

Notes: Sodium is a Group I metal. All Group I metals are reactive metals that can react with water to form soluble metal hydroxide. It is an alkali (soluble base). All Group I metal hydroxides are colourless solutions. Note that the reaction between Group 1 metals and water are very vigorous, sometimes even explosively. You will learn more about metals in the later topics of Periodic Table and Metals.

Example 2 - Group II Metal:
calcium + water -----> calcium hydroxide + hydrogen
Ca(s) + 2H₂O(l) -----> Ca(OH)₂(aq) + H₂(g)
                                   (limewater)

Observation: The silvery grey metal reacts readily, dissolving and forming a colourless alkaline solution that turns red litmus paper blue. Effervescence is observed. A colourless, odourless gas which extinguishes a lighted splint with a 'pop' sound is formed.

Notes: Some Group II metals do not react readily with water and one example is magnesium, Mg. However, calcium, Ca, which is below magnesium in the periodic table, is more reactive and can react with cold water. It forms a colourless alkali, calcium hydroxide. Limewater is used for the testing of carbon dioxide.

Metal + Acid
metal + acid -----> metal salt + hydrogen
M + HX -----> MX + H₂(g)

Example 1 - Group 1 Metal:
lithium + hydrochloric acid -----> lithium chloride + hydrogen
2Li(s) + 2HCl(aq) -----> 2LiCl(aq) + H₂(g)

Observation: The silvery grey metal reacts explosively, dissolving and forming a colourless solution. Effervescence is observed. A colourless, odourless gas which extinguishes a lighted splint with a 'pop' sound is formed.

Notes: Lithium is a Group I metal. All Group I metals react explosively with acids! Group I metal salts such as lithium chloride above are all colourless.

Example 2 - Group II Metal:
magnesium + hydrochloric acid -----> magnesium chloride + hydrogen
Mg(s) + 2HCl(aq) -----> MgCl₂(aq) + H₂(g)

Observation: The silvery grey metal reacts readily, dissolving and forming a colourless solution. Effervescence is observed. A colourless, odourless gas which extinguishes a lighted splint with a 'pop' sound is formed.

Notes: Magnesium is a Group II alkaline earth metal. All Group II metals react moderately fast with acids.

Example 3 - Transition Metal:
iron + hydrochloric acid -----> iron(II) chloride + hydrogen
Fe(s) + 2HCl(aq) -----> FeCl₂(aq) + H₂(g)

Observation: The silvery grey solid reacts slowly, dissolving and forming a pale green solution. Effervescence is observed. A colourless, odourless gas which extinguishes a lighted splint with a 'pop' sound is formed.

Notes: The pale green solution, iron(II) chloride. Iron is a metal from the transition metals block. Metals from this block are much less reactive than those from Group I and II and they usually form coloured compounds, such as the green iron(II) chloride.

Laboratory Setup for Metal + Water/Acid Reaction

Procedure:

1. Add 20cm³ (or any suitable volume) of water or acid solution into a conical flask.
2. Using a spatula, add a small amount of metal powder/metal foil/metal filings to the water or acid solution.
3. (Optional) Cover the conical flask with a rubber stopper and delivery tube to collect the gas produced.
4. Observe the effervescence produced.

Observation: Effervescence is observed. Bubbles of colourless and odourless gas are produced which extinguishes a lighted splint with a 'pop' sound.

You will learn more about testing for hydrogen gas will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the preparation, collection and drying of gas.

Notes:

• Some descriptions of metal: zinc granules, magnesium ribbon, iron filings, iron wool, copper coil
• Alternative, you can use a test tube instead of a conical flask. However, you will have to rest the test tube in a test tube rack or on a retort stand with a test tube clamp if you want to avoid holding it with a test tube holder. 
• You can also use a thistle funnel to transfer water or acid solution to the conical flask.

Industrial Production of Hydrogen
In real life, most of the hydrogen is produced by the steam reforming process of fossil fuels, which makes it not very environmentally friendly. Another method is the electrolysis of water but the use of electricity makes it an expensive process.

Oxygen, O₂

Oxygen can be produced by decomposing hydrogen peroxide, H₂O₂. At room temperature, the rate of decomposition is rather slow. So a little manganese(IV) oxide (or manganese dioxide), MnO₂, is added to speed up the reaction. Manganese(IV) oxide is a black solid and acts as a catalyst.

                               manganese(IV) oxide
hydrogen peroxide -------------------------> oxygen + water
                     MnO2
2H₂O₂(aq) ----------> O₂(g) + 2H₂O(l)

Observation: Effervescence is observed. A colourless, odourless gas which rekindles a glowing splint is formed. The black solid remains unchanged.

Notes: Manganese(IV) oxide is a catalyst, meaning that it does not get used up in the chemical reaction. It might be physically changed (broken up into smaller pieces) but does not react to form new substances.

Laboratory Setup for the Decomposition of Hydrogen Peroxide

The same setup used for Metal + Water/Acid reaction is used.

Procedure:

1. Add 20cm³ (or any suitable volume) of hydrogen peroxide solution into a conical flask.
2. Using a spatula, add a small amount of manganese(IV) oxide powder to the solution.
3. (Optional) Cover the conical flask with a rubber stopper and delivery tube to collect the gas produced.
4. Observe the effervescence produced.

Observation: Effervescence is observed. Bubbles of colourless and odourless gas are produced which rekindle a glowing splint.

You will learn more about testing for oxygen gas will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the preparation, collection and drying of gas.

Notes:

• In humans, the liver produces an enzyme called catalase to carry out the same reaction. A pork liver works in a similar way. You may want to experiment with a piece of liver.
• Alternative, you can use a test tube instead of a conical flask. However, you will have to rest the test tube in a test tube rack or on a retort stand with a test tube clamp if you want to avoid holding it with a test tube holder. 
• You can also use a thistle funnel to transfer hydroxide peroxide solution to the conical flask.

Industrial Production of Oxygen
In real life, oxygen is usually obtained from the fractional distillation of liquid air in which oxygen occupies about 21% of it.

Carbon dioxide, CO₂
Carbon dioxide can be produced by decomposing certain metal carbonates using heat or by reacting carbonates or hydrogencarbonates with acids. Most carbonates/hydrogencarbontes that you will use will be metal carbonates/hydrogencarbonates.

Thermal Decomposition of Metal Carbonates

Example 1a - Group II Metals:
calcium carbonate -----> calcium oxide + carbon dioxide
CaCO₃(s) -----> CaO(s) + CO₂(g)

Observation: A colourless, odourless gas are produced which form a white precipitate in limewater. The remaining solid remains white.

Notes: All Group II metal carbonates and oxides are white. Marble chips refer to calcium carbonate.

Example 1b - Group II Metals:
magnesium carbonate -----> magnesium oxide + carbon dioxide
MgCO₃(s) -----> MgO(s) + CO₂(g)

Observation: A colourless, odourless gas are produced which form a white precipitate in limewater. The solid changes from white to black.

Notes: All Group II metal carbonates and oxides are white.

Example 2a - Transition Metals:
iron(II) carbonate -----> iron(II) oxide + carbon dioxide
FeCO₃(s) + HCl(aq) -----> FeCl₂(aq) + CO₂(g) + H₂O(l)

Observation: A colourless, odourless gas are produced which form a white precipitate in limewater. The solid changes from green to black.

Notes: Iron(II) carbonate is green while iron(II) oxide is black. Iron is a transition metal and one of the characteristics of transition metals is that they form coloured compounds.

Example 2b - Transition Metals:
zinc carbonate -----> zinc oxide + carbon dioxide
ZnCO₃(s) -----> ZnO(aq) + CO₂(g)

Observation: A colourless, odourless gas are produced which form a white precipitate in limewater. The solid changes from white to yellow. Upon cooling, the yellow solid changes to white.

Notes: Zinc oxide is yellow when hot and white when cold.

Example 2c - Transition Metals:
copper(II) carbonate -----> copper(II) oxide + carbon dioxide
CuCO₃(s) -----> CuO(aq) + CO₂(g)

Observation: A colourless, odourless gas are produced which form a white precipitate in limewater. The solid changes from green to black.

Notes: Copper(II) carbonate is green while copper(II) oxide is black. Copper is a transition metal and one of the characteristics of transition metals is that they form coloured compounds.

Laboratory Setup for the Thermal Decomposition of Metal Carbonate

Procedure:
1. Add some metal carbonate powder to a test tube.
2. Cover the test tube with a rubber stopper and a delivery tube.
3. Heat the test tube with Bunsen burner using non-luminous flame.
4. Observe the gas produced and any changes to the colour of the solid that remains in the test tube.

Observation: A colourless, odourless gas which forms a white precipitate in limewater. The colour of the solid remains unchanged/changes from <some colour> to <another colour>.

Notes:

• There are no examples of Group I metal carbonates because they do not decompose upon heating.
• Learning about the colour of substances is very useful to help you in describe changes during a reaction. You will learn more about it in the next post on the Purification of Substances.

Carbonate/Hydrogencarbonate + Acid
Both carbonate and hydrogencarbonate compounds form the same products.

carbonate/hydrogencarbonate + acid -----> salt + carbon dioxide + water

Example 1 - Group I Metal Carbonate:
potassium carbonate + hydrochloric acid -----> potassium chloride  + carbon dioxide + water
K₂CO₃(aq) + 2HCl(aq) -----> 2KCl(aq) + CO₂(g) + H₂O(l)

Observation: Effervescence is observed. Bubbles of colourless, odourless gas are produced which form a white precipitate in limewater. Solution remains colourless.

Notes: All metal carbonates exist as solid due to their high melting points. However, Group I metal carbonates are soluble in water. In the reaction above, potassium carbonate has been dissolved in water to form a solution and that is why (aq) is used to indicate its aqueous state. All Group I metal salts are soluble and form colourless solution.

Example 2 - Group II Metal Carbonate:
calcium carbonate + hydrochloric acid -----> calcium chloride + carbon dioxide + water
CaCO₃(s) + 2HCl ------> CaCl₂(aq) + CO₂(g) + H₂O(l)

Observation: White solid dissolves to form a colourless solution. Effervescence is observed. Bubbles of colourless, odourless gas are produced which form a white precipitate in limewater.

Notes: Group II metal carbonates are insoluble, unlike Group I metal carbonates, hence the (s) solid state symbol. Most Group II metal salts are soluble and form colourless solution.

Example 3 - Transition Metal Carbonate:
copper(II) carbonate + hydrochloric acid -----> copper(II) chloride + carbon dioxide + water
CuCO₃(s) + 2HCl(aq) -----> CuCl₂(aq) + CO₂(g) + H₂O(l)

Observation: Green solid dissolves to form a blue-green solution. Effervescence is observed. Bubbles of colourless, odourless gas are produced which form a white precipitate in limewater.

Notes: All transition metal carbonates are insoluble, unlike Group I metal carbonates. Copper is a transition metal and one of the characteristics of transition metals is that they form coloured compounds, hence the green copper(II) carbonate and blue-green copper(II) chloride solution.

Example 4 - Group I Metal Hydrogencarbonate:
Sodium hydrogencarbonate + hydrochloric acid -----> sodium chloride + carbon dioxide + water
NaHCO₃(aq) + HCl(aq) -----> NaCl(aq) + CO₂(g) + H₂O(l)

Observation: Effervescence is observed. Bubbles of colourless, odourless gas are produced which form a white precipitate in limewater. Solution remains colourless.

Laboratory Setup for Carbonate/Hydrogencarbonte + Acid Reaction

For solid/insoluble carbonate/hydrogencarbonate:

For aqueous/soluble carbonate/hydrogencarbonate:

Procedure:
1. Add a small amount of solid carbonate/hydrogencarbonate or a small volume of aqueous carbonate/hydrogencarbonate into a conical flask.
2. Add an excess of acid solution to the carbonate or hydrogencarbonate.
3. Cover the conical flask with a rubber stopper and a delivery tube.
4. Observe the effervescence produced and also any changes to the colour of the solution.

Observations: If a solid carbonate/hydrogencarbonate is used, you can mention that it dissolves to form a <colour> solution. Effervescence is observed. Bubbles of colourless, odourless gas are produced which form a white precipitate in limewater. Solution remains colourless/changes from <colour> to <colour>.

You will learn more about testing for carbon dioxide will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the collection and drying of gas.

Notes:

• Other acids other than dilute hydrochloric acid can be used. However, avoid using acids which will form another solid product. For example, avoid using dilute sulfuric acid with calcium carbonate, the product calcium sulfate forms a protective coating around calcium carbonate which prevents further reaction. You will learn more about insoluble salts under the topic of Salts.
• There are non-metal carbonates as well. For example, ammonium carbonate, (NH4)2CO3.
• You can also use a thistle funnel to transfer water or acid solution to the conical flask.

For solid/insoluble carbonate/hydrogencarbonate:

For aqueous/soluble carbonate/hydrogencarbonate:

Industrial Production of Carbon Dioxide
In real life, carbon dioxide is mainly obtained from the steam reforming process of fossil fuels (same as hydrogen). A small amount is also obtained from the fractional distillation process of liquid air.

Chlorine, Cl₂
Chlorine can be produced by heating concentrated hydrochloric acid with manganese(IV) oxide. 

conc. hydrochloric acid + manganese(IV) oxide -----> manganese(II) chloride + chlorine + water
4HCl(l) + MnO₂(s) -----> MnCl₂(aq) + Cl₂(g) + 2H₂O(l)

Laboratory Setup for Producing Chlorine Gas

Procedure:
1. Set up the apparatus as shown above.
2. Add manganese(IV) oxide to the the round-bottomed flask.
3. Using a thistle funnel secured using a rubber stopper, transfer concentrated hydrochloric acid into the round-bottomed flask.
4. Cover the other opening with a rubber stopper and a delivery tube.
5. Heat the mixture using non-luminous flame.
6. Observe the effervescence produced and any changes to the colour of the solid or liquid.

Observation: Effervescence is observed. A yellowish-green irritating pungent gas, which turns moist blue litmus paper red and bleaches it, is produced. The black solid dissolves, forming a very pale pink solution.

You will learn more about testing for chlorine gas will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the collection and drying of gas.

Notes: 

• Since this involves a concentrated acid, it should be done in a fume chamber and by a teacher.
• Ensure that the thistle funnel is immersed in the liquid acid.
• This is a redox reaction which you will be learning more about in a topic called Oxidation and Reduction. 
• The manganese(IV) oxide acts as an oxidising agent, removing hydrogen from the concentrated hydrochloric acid to produce chlorine gas. In other words, the hydrochloric acid is oxidised and managanse(IV) oxide is reduced. 
• Other oxidising agents such as potassium manganate(VII) may also be used. 
• Manganese is a transition metal and one of the characteristics of transition metals is that they form coloured compounds, hence the black manganese(IV) oxide and the pale pink manganese(II) chloride. Since water is produced, manganese(II) chloride is able to dissolve in it.

Questions:

• Which hazard symbol(s) are suitable for this setup?
• The diagram does not show how the round-bottomed flask is being stabilised. How do you think the round-bottomed flask is being supported in the apparatus setup above?

[Corrosive due to concentrated acid, Irritating due to chlorine gas; retort stand with a suitable clamp]

Industrial Production of Chlorine
In real life, chlorine is produced from the electrolysis of aqueous sodium chloride or molten sodium chloride. You will learn more about this from the topic of electrochemistry.

Electrolysis of aqueous sodium chloride:
sodium chloride + water -----> sodium hydroxide + hydrogen + chloride
2NaCl(aq) + 2H₂O(l) -----> 2NaOH(aq) + H₂(g) + Cl₂(g)

Electrolysis of molten sodium chloride:
sodium chloride -----> sodium + chlorine
2NaCl(l) -----> 2Na(l) + Cl₂(g)

Hydrogen chloride, HCl
Hydrogen chloride gas can be produced by reacting concentrated sulfuric acid with sodium chloride salt. Depending on the temperature, the products will differ.

Room temperature (25°C or 298K):
conc. sulfuric acid + sodium chloride -----> sodium hydrogensulfate + hydrogen chloride
H₂SO4(l) + NaCl(s) -----> NaHSO₄(s) + HCl(g)

Around 200°C:
sodium hydrogensulfate + sodium chloride -----> sodium sulfate + hydrogen chloride
NaHSO₄(s) + NaCl(s) -----> Na₂SO₄(s) + HCl(g)

Laboratory Setup for Producing Hydrogen Chloride Gas

Procedure:
1. Set up the apparatus as shown above.
2. Add sodium chloride to the the round-bottomed flask.
3. Using a thistle funnel secured using a rubber stopper, transfer concentrated sulfuric acid into the round-bottomed flask.
4. Cover the other opening with a rubber stopper and a delivery tube.
5. Heat the mixture using non-luminous flame.
6. Observe the effervescence produced and any changes to the colour of the solid or liquid.

Observation: Effervescence is observed. A colourless, choking and pungent gas which turns most blue litmus paper red is produced. No visible change to the colour of the white solid.

You will learn more about testing for hydrogen chloride will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the collection and drying of gas.

Notes: 

• Since this involves a concentrated acid, it should be done in a fume chamber and by a teacher.
• Ensure that the thistle funnel is immersed in the liquid acid.
• Because concentrated sulfuric acid has very low water content, Group I metal salts such as the sodium chloride, sodium hydrogensulfate and sodium sulfate do not dissolve and exist as solid state during the reaction.

Industrial Production of Hydrogen Chloride
In real life, a lot of  hydrogen chloride gas is produced as a by-product of reactions involving chlorinated hydrocarbons. A small amount is also produced by direct reaction between hydrogen gas and chlorine gas.

hydrogen + chlorine -----> hydrogen chloride
H₂(g) + Cl₂(g) -----> 2HCl(g)

Sulfur dioxide, SO₂
Sulfur dioxide can be prepared by reacting metal sulfites with dilute sulfuric acid or by reacting copper with concentrated sulfuric acid. This is very similar to reacting carbonates with acids.

Metal Sulfite + Acid
metal sulfite + acid -----> salt + sulfur dioxide + water

Example 1 - Group 1 Metal Sulfite:
sodium sulfite + sulfuric acid -----> sodium sulfate + sulfur dioxide + water
Na₂SO₃(s) + H₂SO₄(aq) -----> Na₂SO₄(aq) + SO₂(g)+ H₂O(l)

Observation: Effervescence is observed. A colourless, choking gas is produced which turns acidified potassium dichromate(VI) green. The white solid dissolves to form a colourless solution.

Notes: Sulfur dioxide is very soluble in water, hence the amount of sulfur dioxide collected may be lesser. To reduce the solubility of the gas, the reaction mixture can be heated.

Example 2 - Group 2 Metal Sulfite:
magnesium sulfite(s) + sulfuric acid -----> magnesium sulfate + sulfur dioxide + water
MgSO₃(s) + H₂SO₄(aq) -----> MgSO₄(aq) + SO₂(g) + H₂O(l)

Observation: Effervescence is observed. A colourless, choking gas is produced which turns acidified potassium dichromate(VI) green. The white solid dissolves to form a colourless solution.

Notes: Sulfur dioxide is very soluble in water, hence the amount of sulfur dioxide collected may be lesser. To reduce the solubility of the gas, the reaction mixture can be heated.

Laboratory Setup for Metal Sulfite + Acid Reaction

Procedure:
1. Add a small amount of metal sulfite into a conical flask.
2. Add an excess of acid solution to the sulfite.
3. Cover the conical flask with a rubber stopper and a delivery tube.
4. Observe the effervescence produced and also any changes to the colour of the solution.

Observation: Effervescence is observed. A colourless, choking gas is produced which turns acidified potassium dichromate(VI) green. The <color> solid dissolves to form a <colour> solution.

You will learn more about testing for sulfur dioxide will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the collection and drying of gas.

Notes:

• Group I metal sulfite salts are all soluble in water and may be used in aqueous form. You may use (aq) as the state symbol if the experiment is carried out using the solution form.
• Other acids such as dilute hydrochloric acid and diulte nitric acid can be used.
• Sulfur dioxide is soluble in water. Hence this setup may not yield less sulfur dioxide as some is dissolved in the reaction mixture as it is being produced. Heating the reaction mixture can help to reduce the solubility and yield more sulfur dioxide. You may add a Bunsen burner to the diagram if this is so.
• You can also use a thistle funnel to transfer water or acid solution to the conical flask.
• Sulfur dioxide is a strong reducing agent and reacts with a strong oxidising agent such as acidified potassium manganate(VII) solution or acidified potassium dichromate(VI) solution.

Copper + Concentrated Sulfuric Acid
copper + conc. sulfuric acid -----> copper(II) sulfate + sulfur dioxide + water
Cu(s) + H₂SO₄(l) -----> CuSO₄(aq) + SO₂(g) + H₂O(l)

Laboratory Setup for Copper + Concentrated Sulfuric Acid Reaction

Procedure:
1. Set up the apparatus as shown above.
2. Add copper metal to the the round-bottomed flask.
3. Using a thistle funnel secured using a rubber stopper, transfer concentrated sulfuric acid into the round-bottomed flask.
4. Cover the other opening with a rubber stopper and a delivery tube.
5. Heat the mixture using non-luminous flame.
6. Observe the effervescence produced and any changes to the colour of the solid or liquid.

Observation: Effervescence is observed. A colourless, choking gas is produced which turns acidified potassium dichromate(VI) green. The <color> solid dissolves to form a <colour> solution.

You will learn more about testing for sulfur dioxide will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the collection and drying of gas.

Notes: 

• Since this involves a concentrated acid, it should be done in a fume chamber and by a teacher.
• Ensure that the thistle funnel is immersed in the liquid acid.
• Concentrated sulfuric acid has very low water content. However, since water is produced, copper(II) sulfate may be able to dissolve to form a blue solution.
• Sulfur dioxide is a strong reducing agent and can be tested with a strong oxidising agent such as acidified potassium manganate(VII) or acidified potassium dichromate(VI).
  

Industrial Production of Sulfur Dioxide
Sulfur dioxide can be produced by heating a mineral ore, pyrite, that contains substantial amount of iron(II) sulfide with oxygen or by burning sulfur directly in oxygen.

Combusting Iron(II) Sulfide
iron(II) sulfide + oxygen -----> iron(II) oxide + sulfur dioxide
2FeS(s) + 3O₂(g) -----> 2FeO(s) + 2SO₂(g)

Combusting Sulfur
sulfur + oxygen -----> sulfur dioxide
S(s) + O₂(g) -----> SO₂(g)

Ammonia, NH₃
This is the only alkaline gas you have to know. In laboratories, ammonia is usually prepared by warming ammonium salts (NH⁺₄ salts) with strong bases such as aqueous sodium hydroxide or aqueous calcium hydroxide (limewater or slaked lime).

Using sodium hydroxide
ammonium chloride + sodium hydroxide -----> sodium chloride + ammonia + water
NH₄Cl(aq) + NaOH(aq) -----> NaCl(aq) + NH₃(g) + H₂O(l)

Using calcium hydroxide (limewater or slaked lime)
ammonium chloride + calcium hydroxide -----> calcium chloride + ammonia + water
NH₄Cl(aq) + Ca(OH)₂(aq) -----> CaCl₂(aq) + NH₃(g) + H₂O(l)

Laboratory Setup for Producing Ammonia

Procedure:
1. Add a small amount of ammonium salt (eg. ammonium chloride) into a test tube.
2. Add a strong alkali (eg. aqueous sodium hydroxide) to the ammonium salt.
3. Swirl/stir to dissolve ammonium salt completely.
4. Cover the test tube with a rubber stopper and a delivery tube.
5. Warm the mixture gently.
6. Observe.

Observation: A colourless pungent gas is produced which turns moist red litmus paper blue.

You will learn more about testing for ammonia will be covered in a later topic called Qualitative Analysis. For this post, we will be focusing on the collection and drying of gas.

Notes:

• Heating with sodium hydroxide is also a test for ammonium cation, NH⁺₄. Ammonia gas is colourless and pungent and turns moist red litmus paper blue.
• Ammonia is highly soluble in water. You may not be able to observe any effervescence. You may test the presence of it by using a moist red litmus red and by detecting a pungent smell.
• Ammonium chloride can be produced by combining two gases that you have learned here: hydrogen chloride and ammonia.

Industrial Production of Ammonia
In real life, ammonia is manufactured using the Haber Process (also known as Haber-Bosch process) using a temperature of 450°C, 200 atm, and an iron catalyst. Note that this is a reversible reaction.

nitrogen + hydrogen <-----> ammonia
N₂(g) + 3H₂(g) <-----> 2NH₃(g)

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Properties of gases
You will need to know whether the solubility of the gases in water and their relative densities compared to air in order to choose a suitable way of collecting the gases produced in a reaction for further testing and analysis. You will also need to know their pH properties to select a suitable drying agent.

Below is a table that summarises the properties of the above list of gases:

Air is made up of nitrogen, N₂ mainly (78%) and oxygen, O₂ (21%), and the rest consists of water vapour, carbon dioxide and other inert gases (1%). Assuming that the 1% constitutes wholly of water,

The average relative mass of air is (using statistical knowledge):

78% x Mr(N₂) + 21% x Mr(O₂) + 1% x Mr(H₂O)
= 0.78 x 28 + 0.21 x 32 + 0.01 x 18
= 21.84 + 6.72 + 0.18
= 28.74

Since density = mass / volume, and all gases at room temperature and pressure has a molar volume of 24dm³ or 24000cm³, we can say that density is directly proportional to molar mass (which has the same value as relative molecular mass).

Note: Do not worry about the calculation above. You will learn more about it at a later topic called Chemical Calculation. Otherwise, if you have studied statistics in your mathematics class, you can review the section on how to calculate average from a distribution.

Therefore, gases which have relative molecular mass less than 28.72 is less dense and hence lighter than air. These gases rise upwards and displace air downwards. Gases whose relative molecular mass are greater than 28.72 are denser and hence heavier than air. They sink to the bottom and displace air upwards.

Gas	pH	Solubility in water	Relative molecular mass	Density
hydrogen gas, H2	neutral	insoluble	2	lighter than air
oxygen gas, O2	neutral	slightly soluble	32	slightly heavier than air, downward delivery is not effective
carbon dioxide, CO2	acidic	slightly soluble	44	heavier than air
chlorine, Cl2	acidic	very soluble	71	heavier than air
hydrogen chloride, HCl	acidic	very soluble	36.5	heavier than air
sulfur dioxide, SO2	acidic	very soluble	64	heavier than air
ammonia, NH3	alkaline	extremely soluble	15	lighter than air

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Collecting gases

There are 3 ways of collecting gases:

• Displacement of water
• Displacement of air
• Upward delivery
• Downward delivery
• Gas syringe system

The method to use depends on the solubility of the gas in water or its density (or relative molecular mass) compared to that of air.

General rules
1. Check the solubility. If the gas is insoluble or slightly soluble in water, the water displacement method can be used.
2. For water-soluble gases, consider displacement of air method. If the gas is less dense than air, it will rise up and so, upward delivery is more suitable. Otherwise, if the gas is dense, it will sink and so, downward delivery is more suitable.
3.  If you intend to measure the volume of gas, the gas syringe system is the best and can be used to collect all gases.
4. Note: Gases that are suitable for the water displacement method can also be collected using displacement of air method (except oxygen).

Here are the laboratory setup diagrams for different methods of gas collection:

Apparatus Setup	Suitable gases
	hydrogen oxygen carbon dioxide
	hydrogen ammonia
	carbon dioxide chlorine hydrogen chloride sulfur dioxide
	all gases able to measure the volume as well

Preparing and collecting hydrogen gas - an example of water displacement method

Preparing and collecting hydrogen gas - an example of upward delivery method

Preparing and collecting oxygen gas - an example of gas syringe system

Preparing and collecting carbon dioxide - an example of downward delivery method

Preparing and collecting chlorine - an example of downward delivery method

Preparing and collecting hydrogen chloride

Preparing and collecting sulfur dioxide

Preparing and collecting ammonia

Notes:

• As long as your experiment setup is reasonable, your choice of apparatus can be flexible.
• You may notice that both gas jars and test tubes have been used for collecting gases
• You may also have noticed that test tube can be supported by test tube clamp on a retort stand or resting on a test tube rack.
• It is optional to use a thistle funnel to transfer a liquid chemical into another container, such as a conical flask.
• So far, the gas collected will have moisture but most examination questions will require you to obtain a pure, dry sample of gas. In the next section, you will learn more about how to dry gases.
• If you need to collect dry gas, the water displacement method cannot be used. Otherwise there will be water vapour in the collected gas.

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Drying gases

Testing for the presence of water/water vapour

• copper(II) sulfate crystals are white when anhydrous (dry) and blue when hydrated (have water of crystallisation).
• cobalt(II) sulfate paper is blue when dry and pink when wet.

3 Drying agents	State at rtp	Remarks
concentrated sulfuric acid [conc. H2SO4(l)]	liquid	Cannot be used to dry alkaline gases, such as NH3 Sulfuric acid reacts with ammonia to form a salt, ammonium sulfate. H2SO4(l) + 2NH3(s) -----> (NH4)2SO4(aq)
calcium oxide [quicklime, CaO(s)]	solid	Cannot be used to dry acidic gases such as HCl and CO2 CaO(s) + HCl(g) ------> CaCl2(s) + H2O(l)
fused calcium chloride [CaCl2(s)]	solid	Can be used for most gases except  NH3 as they will react to form an additional complex compound CaCl2(s) + 8NH3(g) -----> CaCl2.8NH3(s)

Note: (l) means liquid state and (s) means solid state. Fused means that the calcium chloride has been heated in a unique way to remove moisture from it. r.t.p means room temperature (25 C or 298K) and 1 atmospheric pressure (1.02 x 105 Pa).

Apparatus for different drying agent and gas
(Source: minichemistry.com)

Concentrated sulfuric acid
Notice that the two delivery tubes used are positioned differently. The incoming gas has to be bubbled into the acid. The other delivery tube is not dipped into the acid. Otherwise, if the tube comes into contact with the acid, it will be wet and will not be able to deliver dry gas.

Calcium oxide (quicklime)
Below shows the use of a vertical column with delivery tubes and rubber stoppers and allowing the ammonia gas to move upwards. This is known as upward delivery, which is introduced in the previous section.

Fused calcium chloride
Below shows a U-tube used with delivery tubes and rubber stoppers.

Preparing, drying and collecting hydrogen gas

Preparing, drying and collecting oxygen gas

Preparing, drying and collecting carbon dioxide

Preparing, drying and collecting chlorine

Preparing, drying and collecting hydrogen chloride

Preparing, drying and collecting sulfur dioxide

Preparing, drying and collecting ammonia

Notes:

• As long as your experiment setup is reasonable, your choice of apparatus can be flexible.
• You may notice that both gas jars and test tubes have been used for collecting gases
• You may also have noticed that test tube can be supported by test tube clamp on a retort stand or resting on a test tube rack.
• It is optional to use a thistle funnel to transfer a liquid chemical into another container, such as a conical flask.
• Experiments involving concentrated acids should be carried out in a fume chamber.

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