The Reactivity Series of Metals Towards Oxygen and Its Application Archives

Rusting as a Redox Reaction

December 1, 2020December 1, 2020 by Veerendra

Rusting (Corrosion) as a Redox Reaction

What is corrosion of a metal?

Corrosion of metal:

When metals are exposed to their environment, they undergo corrosion. For example, after some time, a shiny aluminium pot will lose its shine, silverware will tarnish and an iron structure will rust.
Corrosion of metal is a redox reaction in which a metal is oxidised naturally to its ions, resulting in partial or complete destruction of the metal.
During corrosion, the metal atoms lose electrons to form positive ions.
M → Mⁿ⁺ + ne^–
Some metals corrode more easily than others. How easily a metal corrodes depends on two factors:
(a) Electropositivity of metals
(b) Nature of the product of corrosion
Electropositivity of metals
Nature of the product of corrosion
When a metal corrodes, it usually forms an oxide coating.
(a) The oxide coating of aluminium, for example, is tightly packed, non-porous and is firmly attached to the metal. It does not let water and air penetrate through it, protecting the aluminium underneath from further corrosion. This explains why aluminium is quite resistant to corrosion even though it is very electropositive. Other metals with similar protective oxide coating include nickel, chromium, tin, lead and zinc.
(b) The oxide coating of iron on the other hand, is not tightly packed, porous, weak and easily peels off. Thus, water and air can penetrate through the coating to further corrode the iron metal underneath it.
Table compares the resistance to corrosion of some common metals.

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Rusting as a redox reaction

Rusting is the corrosion of iron. It is the most common corrosion of metal around.
For iron to rust, oxygen and water must be present.
Rusting is a redox reaction whereby oxygen acts as the oxidising agent and iron acts as the reducing agent.
Figure shows the half-reactions of rusting.
The surface of iron at the middle of the water droplet serves as the anode, the electrode at which oxidation occurs. The iron atoms here lose electrons to form iron(II) ions.
The electrons flow to the edge of the water droplet, where there is plenty of dissolved oxygen. The iron surface there serves as the cathode, the electrode at which reduction occurs. Oxygen gains the electrons and is reduced to hydroxide ions.
The iron(II) ions produced combine with the hydroxide ions to form iron(II) hydroxide.
Thus, the overall redox reaction is as follows:
The iron(II) hydroxide is then further oxidised by oxygen to form hydrated iron(III) oxide, Fe₂O₃.xH₂O) whereby x varies. The hydrates come in various shades of brown and orange and together make up what is commonly known as rust.
In the presence of acids and salts, rusting occurs faster. These substances increase the electrical conductivity of water, making water a better electrolyte.
For example:
(a) Iron structures such as bridges, fences and cars at coastal areas rust faster due to the presence of salts in the coastal breeze.
(b) Iron structures in industrial areas rust quickly as these areas have air polluted with acidic gases such as sulphur dioxide and nitrogen oxides.

Other metals and rusting of iron

When two metals are in contact with each other, the more electropositive metal will corrode first. This is simply because the more electropositive metal can lose its electrons more readily than the less electropositive metal.
So, when iron is in contact with a more electropositive metal, rusting of iron is prevented or inhibited.
For example, magnesium is more electropositive than iron. So, when iron is in contact with magnesium, magnesium corrodes or is oxidised instead of iron.
On the other hand, when iron is in contact with a less electropositive metal, rusting of iron is speeded up.
For example, copper is less electropositive than iron. Therefore, when iron is in contact with copper, iron rusts faster.
The further apart the metals are in the electrochemical series, the faster the more electropositive metal corrodes. For example, iron rusts faster when in contact with copper than when it is in contact with tin.

What are some ways of preventing corrosion of metals?

Controlling rusting:
Generally, there are three main ways to control rusting.
Rusting as a Redox Reaction 11

1. Using protective coating
The protective coating prevents water and air from reaching the surface of iron. Various materials can be used as the protective coating, depending on the costing and usage of the iron items.

Oil and grease are used for moving parts of engine.
Paints are used for items that are not easily scratched such as cars, ships, bridges, railings and gates. For example, most modern cars have a few layers of anti-rust coating and paints on them. Some pots and plates have enamel-paint coated on them.
Plastics are used for light items such as clothes hanger and wire fences.
Galvanising (zinc plating) involves coating an iron or steel sheet with a thin layer of zinc. This is done by dipping the iron into molten zinc.
- Galvanising is used on objects that are exposed to the atmosphere, such as iron roofing, water tanks and iron rubbish bins.
- Galvanised iron is prevented from rusting in two ways. Firstly, the zinc layer provides a protective oxide coating.
- Secondly, when the galvanised iron is scratched, zinc corrodes first instead of iron because zinc is more electropositive than iron.
In tin plating, an iron or steel sheet is coated with a very thin coating of tin. This is done by dipping the iron into molten tin or by electroplating an iron sheet using tin(IV) chloride as the electrolyte.
- Tin plating is usually used for making tin cans as tin is very expensive. After tin plating, the inside of the can is coated with a thin layer of plastic.
- The tin provides a protective oxide coating to the cans. The cans do not rust as long as the tin coating remains unbroken.
- However, as soon as the can is scratched, rusting will occur quickly. This is because iron is more electropositive than tin. Thus, food in dented or scratched tin cans should not be consumed.

2. Alloying
Stainless steel is a corrosion resisting alloy of iron.

It contains carbon and a varying amount of chromium and nickel. The typical stainless steel contains about 18% chromium and 8% nickel.
Thus, stainless steel is expensive and is mainly used for small objects such as cutlery and decorative items.
The chromium and nickel provide a protective oxide coating which is firmly bonded to the iron and is not easily removed. Furthermore, the oxide coating is shiny, hence giving stainless steel an attractive, mirror-like finish.

3. Sacrificial protection

In this method, iron is attached to a more electropositive metal which acts as the sacrificial metal.
It is used for objects that are exposed to conditions that speed up rusting such as water and sea water. For example, bridge pillars and hulls of ship are usually attached to zinc blocks while underground pipelines are tied to bags of magnesium.
Being more electropositive, the sacrificial metal would act as the anode whereby it is oxidised, protecting iron from rusting. Thus, the sacrificial metal is also known as the sacrificial anode.
The sacrificial metal has to-be renewed from time to time.

Effect of other metals on rusting experiment

Aim: To investigate the effect of other metals on rusting.
Problem statment: How do different types of metals in contact with iron affect rusting?
Hypothesis: When a more electropositive metal is in contact with iron, the metal inhibits rusting. When a less electropositive metal is in contact with iron, the metal speeds up rusting.
Variables:
(a) Manipulated variable : Different metals in contact with iron
(b) Responding variable : Presence of blue colouration
(c) Controlled variables : Clean iron nails, medium in which iron nails are kept, temperature
Operational definition:
Blue colouration indicates rusting of iron.
Materials: Iron nails, magnesium ribbon, copper strip, zinc strip, tin strip, hot jelly solution containing a little
potassium hexacyanoferrate(III) solution and phenolphthalein indicator, sandpaper.
Apparatus: Test tubes, test tube rack.
Safety measure:
Potassium hexacyanoferrate(III) solution is poisonous. Thus, the hot jelly solution should be handled with care.
Procedure:

All five iron nails, magnesium ribbon, strips of copper, zinc and tin are cleaned with sandpaper.
Four iron nails are coiled tightly with magnesium ribbon, strips of copper, zinc and tin each.
All five iron nails are placed in five separate test tubes as shown in Figure.
The same amount of hot jelly solution containing potassium hexacyanoferrate(III) solution and phenolphthalein indicator is poured into the test tubes to completely cover ail the nails.
The test tubes are kept in a test tube rack and left aside for a day. Any changes are observed.

Results:

Test tube	Pair of metals	Intensity of dark blue colouration	Pink colouration	Inference regarding rusting
I	Fe only	Low	Present	The iron nail rusts a little.
II	Fe + Mg	None	Present	The iron nail does not rust.
III	Fe + Cu	Very high	Present	The iron nail rusts very quickly.
IV	Fe + Zn	None	Present	The iron nail does not rust.
V	Fe + Sn	High	Present	The iron nail rusts quickly.

Discussion:

During rusting, iron(II) ions are produced. These ions form dark blue colouration with potassium hexacyanoferrate(III). The more iron(II) ions are produced, the higher the intensity of the dark blue colouration.
During the corrosion of a metal, the reduction of oxygen forms hydroxide ions, thus giving rise to basic condition. The hydroxide ions give pink colouration with phenolphthalein.
The jelly is used to enable us see the blue and pink colourations clearly as diffusion occurs slowly in a solid state. Otherwise, the blue and pink colourations are mixed up and difficult to distinguish.
Since pink colouration is found in all test tubes, it is inferred that corrosion of metal has taken place in each test tube. The metal that corrodes must either be the iron nail or the metal it is in contact with.
Sometimes, the pink colouration is not clear as the hydroxide ions formed, immediately combine with the metal ions. Thus, not many free hydroxide ions are present in the jelly.
Test tube I acts as a control. The iron nail rusts a little.
Test tube II
(a) Magnesium is more electropositive than iron. This means that magnesium can lose its electrons more readily than iron. Therefore, magnesium is oxidised. Magnesium acts as the anode.

(b) The electrons flow to iron which acts as the cathode. At the cathode, the electrons are gained by
oxygen. Thus, oxygen undergoes reduction, producing hydroxide ions which give pink colouration with phenolphathalein.

(c) The iron nail does not corrode. This explains the absence of blue colouration in this test tube.
Test tube III
(a) Iron is more electropositive than copper. This means that iron can lose its electrons more readily than copper. Therefore, iron rusts or is oxidised. Iron acts as the anode.

(b) Since iron and copper have a large difference in electropositivity, the rusting of iron occurs very quickly, producing a large amount of iron(II) ions. This explains the high intensity of blue colouration in this test tube.
(c) The electrons flow to copper which acts as the cathode. At the cathode, the electrons are gained by oxygen, thus reducing oxygen to hydroxide ions. The presence of hydroxide ions is indicated by the pink colouration.
Test tube IV
Similar to test tube II, the iron nail in this test tube does not corrode, thus no dark blue colouration is found.
Zinc acts as the anode and is oxidised as zinc is more electropositive than iron.
Test tube V
Similar to test tube III, the iron nail acts as the anode and rusts. This is because iron is more electropositive than tin.

However, the rate of rusting in test tube V is lower than that in test tube III as the difference in electropositivity
between iron and tin is smaller than the difference in electropositivity between iron and copper.

Conclusions:

Rusting is inhibited when iron is in contact with a more electropositive metal.
Rusting is speeded up when iron is in contact with a less electropositive metal.

Application of the reactivity series of metals in the extraction of metals

December 1, 2020December 1, 2020 by Veerendra

Application of the reactivity series of metals in the extraction of metals

Application of the reactivity series of metals towards oxygen in the extraction of metals:

Only a few metals such as platinum, gold and silver are found free in nature. Other metals have to be extracted from their ores.
Ores are naturally occurring rocks that contain high concentration of one or more mineral of metals. The common minerals found in ores are oxides, sulphides and carbonates of metals.

The following table shows the main contents of some common ores and the metals extracted from them.

Ore	Main mineral in ore	Metal extracted from ore
Carnallite	Hydrated potassium magnesium chloride, KCl.MgCl₂.6H₂O	Potassium or magnesium
Halite (rock salt)	Sodium chloride, NaCl	Sodium
Bauxite	Aluminium oxide, Al₂O₃	Aluminium
Zinc blende (sphalerite)	Zinc sulphide, ZnS	Zinc
Smithsonite	Zinc carbonate, ZnCO₃	Zinc
Hematite	Iron(III) oxide, Fe₂O₃	Iron
Magnetite	Triiron tetroxide. Fe₃O₄	Iron
Cassiterite	Tin(IV) oxide. SnO₂	Tin
Galena	Lead(II) sulphide, PbS	Lead
Malachite	Copper(II) carbonate hydroxide, CuCO₃.Cu(OH)₂	Copper
Chalcopyrite	Copper iron sulphide, CuFeS₂	Copper
Cinnabar	Mercury(II) sulphide, HgS	Mercury

When a metal is extracted from its ore, it involves reduction.
(a) A metal has a positive oxidation number in its ore.
(b) Hence, when it is extracted from its ore, its oxidation number decreases to zero.
The method of reduction depends on the reactivity of the metals. More often than not, the cost is also taken into consideration. For example, a metal can be extracted from its ore using more reactive metals as the reducing agents. However, this is usually not done on a large scale as it is expensive.
Figure summarises the methods of extracting metals from their ores.
In extracting metals, carbon in the form of coke is used. Carbon is widely used in the extraction of zinc, iron, tin and lead for a number of reasons:
(a) First of all, carbon is more reactive than zinc, iron, tin and lead. Therefore, carbon can easily reduce the oxide of these metals.
(b) Furthermore, carbon is cheap and easily available.
(c) The carbon dioxide gas produced is not poisonous and thus can be released directly into the air.
Two main processes in the extraction of metals:
(a) Concentrating the ores (this is done by removing the impurities in the ores)
(b) Reducing the ores into metals

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Extraction of iron

There are a number of iron ores. However, it is uneconomical to extract iron from most of the ores. Two important ores used in extracting iron are hematite, Fe₂O₃ and magnetite, Fe₃O₄.
First, these ores undergo several processes to remove the impurities.
Then, the concentrated ores are reduced by carbon in the form of coke in a very large and hot furnace called blast furnace. Its temperature can reach up to 2000°C.
A small charge is introduced from the top of the blast furnace at intervals of 10 to 15 minutes. The charge consists of concentrated iron ores, coke and limestone.
The following flow chart outlines the reduction of the iron ores in the blast furnace.
The molten iron is collected at the bottom of the furnace. It is drained off periodically into moulds and is allowed to cool. The product is called pig iron or cast iron.
At the same time, the intense heat in the blast furnace causes the limestone to decompose.
CaCO₃(s) → CaO(s) + CO₂(g)
(a) The calcium oxide then reacts with the impurities in the ores, which consist mostly of sand, SiO₂, to form calcium silicate, CaSiO₃ or slag.
CaO(s) + SiO₂(s) → CaSiO₃(l)
(b) As the molten slag is less dense than molten iron, it floats on the molten iron, protecting the molten iron from oxidation by the hot air.
(c) Like the molten iron, the slag is also drained off periodically. The slag can be used as a building material and for the manufacture of cement.

Extraction of tin

The main ore of tin is cassiterite which contains tin(IV) oxide, SnO₂. The following summarises the steps in the extraction of tin.
The ore is first crushed, grounded and washed.
Then, the ore is concentrated by mixing it with oil and water. In this flotation method, the tin minerals, which are less dense, are trapped in the floating foam. The impurities such as soil and sand, which are denser, sink to the bottom.
The concentrated ore is then roasted in the air. This converts the sulphide of tin to oxide. At the same time, impurities such as sulphur and oil are burnt off.
Similar to iron, the reduction of tin(IV) oxide takes place in the blast furnace by carbon monoxide and coke.
Calcium oxide from the limestone eliminates the remaining impurities to slag.
The molten tin is drained off into moulds to become tin blocks.

Contribution of the metal extraction industry

Malaysia has one of the largest reserves of tin ore in the world and was once the world’s largest producer of tin. Tin played a predominant role in Malaysian economy in the 19th and 20th centuries until the collapse of the tin market in the early 1980s.
Other than tin ore, significant amount of iron ore, copper ore, bauxite and gold are also mined in our country.
Besides providing jobs to our people, the local metal extraction industry has contributed greatly to our economy and has supported many local metal-related industries such as the food canning industry, alloy manufacturing and the jewellery industry.

Conserving metals

Metals are non-renewable and the worlds reserves of metals are depleting. Thus, we need to conserve metals. This can be done through the 3Rs (reduce, reuse and recycle).
We should reduce the use of metals in whatever way we can. For example, we can use pencil cases made of cloth instead of metal.
Tin and aluminium cans can be reused as pencil holders. Broken metal furniture can be welded, painted and reused again.
Two most common recycled metals are aluminium and iron.
Recycling of metals not only conserves metals, but it also conserves energy and the environment. For example, recycling aluminium uses only 5% of the energy required to produce aluminium from its ore and reduces air pollution by 95%.

The Reactivity Series of Metals Towards Oxygen

December 1, 2020December 1, 2020 by Veerendra

The Reactivity Series of Metals Towards Oxygen

The reactivity of metals differs from one metal to another. In fact, the form in which a metal occurs in nature depends on its reactivity.
Gold has very low reactivity and therefore can be found in its metallic state in nature.
Aluminium, potassium and sodium have very high reactivity, and therefore exist as compounds in nature.
One of the common compounds formed by metals is metal oxide. The formation of metal oxide is a redox reaction.
Metal + oxygen → metal oxide
(a) Metal undergoes oxidation to form positive ions. Its oxidation number increases from zero to a positive value.
(b) Oxygen undergoes reduction to form oxide ions. Its oxidation number decreases from 0 to -2.
(c) Thus, metal acts as the reducing agent while oxygen acts as the oxidising agent in the formation of metal oxide.
The more reactive a metal is towards oxygen, the more vigorously it burns in oxygen.
Hence, by observing how vigorously the metals react with oxygen, we can arrange the metals according to their reactivity towards the oxygen.
The reactivity series of metals towards oxygen is a list of metals according to their reactivity with oxygen. This series of metals is quite similar to the electrochemical series because the reactivity of a metal with oxygen is closely linked to its ability to lose electrons.

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Reactivity Series of Metals with Oxygen Experiment

Aim: To investigate the reactivity of metals with oxygen.
Materials: Magnesium powder, zinc powder, iron filings, lead powder, copper powder, solid potassium manganate(VII), asbestos paper, glass wool.
Apparatus: Boiling tube, retort stand and clamp, Bunsen burner, spatula, forceps.
Safety measure:
Asbestos paper and glass wool are hazardous and should be handled with care.
Procedure:

One spatulaful of solid potassium manganate(VII) is put into a boiling tube.
Some glass wool is pushed into the tube. The tube is clamped horizontally as shown in Figure.
One spatulaful of magnesium powder is placed on a piece of asbestos paper and is put into the tube.
The magnesium powder is heated strongly. Then, the solid potassium manganate(VII) is heated. How vigorously magnesium reacts with oxygen and the colour of the residue when it is hot and when it is cold are observed.
Steps 1 to 4 are repeated using zinc powder, iron filings, lead powder and copper powder, one at a time, in place of magnesium powder.

Results:

Metal	Observation	Inference
Magnesium	Magnesium burns brightly with a very brilliant white flame. The residue is white when hot and cold.	Magnesium has a very high reactivity with oxygen. Magnesium oxide is formed.
Zinc	Zinc burns fairly bright. The residue is yellow when hot and white when cold.	Zinc has a high reactivity with oxygen. Zinc oxide is formed.
Iron	Iron glows very brightly. The residue is reddish- brown when hot and cold.	Iron has a medium reactivity with oxygen. Iron(III) oxide is formed.
Lead	Lead glows brightly. The residue is brown when hot and yellow when cold.	Lead has a low reactivity with oxygen. Lead(II) oxide is formed.
Copper	Copper glows faintly. The residue is black when hot and cold.	Copper has a lower reactivity with oxygen than lead. Copper(II) oxide is formed.

Discussion:

When solid potassium manganate(VII) is heated, it decomposes to give out oxygen gas.
Other than solid potassium manganate(VII), oxygen gas can also be provided by:
(a) Heating a mixture of potassium chlorate(V) with manganese(IV) oxide as a catalyst
The glass wool separates the metal powder from the solid potassium manganate(VII). If the substances are mixed, the mixture of metal powder and solid potassium manganate(VII) will explode when heated.
Based on the vigour of the reactions, the metals can be arranged according to their reactivity with oxygen.
The following equations represent the reactions between the metals and oxygen.

Conclusion:
The descending order of reactivity of metals with oxygen is Mg, Zn, Fe, Pb, Cu.

Position of Carbon in the Series of Reactivity of Metals

The position of carbon in the series can be determined based on:
(a) The ability of carbon to remove oxygen from metal oxides
(b) The ability of metals to remove oxygen from carbon dioxide
Ability of carbon to remove oxygen from metal oxides
(a) Carbon is strongly heated with a metal oxide.
(b) If carbon is more reactive than the metal, it can remove oxygen from the metal oxide. In other words, carbon can reduce the metal oxide to metal.

(c) Conversely, if carbon is less reactive than the metal, it cannot remove oxygen from the metal oxide.
Ability of metals to remove oxygen from carbon dioxide
(a) A heated metal is placed in carbon dioxide.

(b) If the metal is more reactive than carbon, the metal can remove oxygen from carbon dioxide. In other words, the metal can reduce carbon dioxide to carbon.

(c) On the other hand, if the metal is less reactive than carbon, it is unable to remove oxygen from carbon dioxide.
(d) (i) For example, when a piece of magnesium ribbon is heated and placed in a gas jar filled with carbon dioxide, the magnesium ribbon burns brightly, producing a white residue of magnesium oxide. A lot of black powder of carbon is also formed on the wall of the gas jar.
(ii) The redox reaction can be represented as follows.

As magnesium is more reactive than carbon, it is able to remove oxygen from carbon dioxide. Magnesium reduces carbon dioxide to carbon and magnesium itself is oxidised to magnesium oxide.

Position of Carbon in the Series of Reactivity of Metals Experiment

Aim: To determine the position of carbon in the reactivity series of metals towards oxygen.
Materials: Carbon powder, solid copper(II) oxide, solid magnesium oxide, solid aluminium oxide, solid zinc oxide.
Apparatus: Crucible, spatula, Bunsen burner, pipe-clay triangle, tripod stand.
Procedure:

A spatulaful of carbon powder and a spatulaful of solid copper(II) oxide are mixed thoroughly in a crucible.
The apparatus is set up as shown in Figure.
The mixture is heated strongly. Any changes that occur are observed.
Steps 1 to 3 are repeated using solid zinc oxide, solid aluminium oxide and solid magnesium oxide, one at a time, in place of solid copper(II) oxide.

Results:

Mixture	Observation	Inference
Carbon + copper(II) oxide	A flame spreads to the whole mixture. A brown residue is formed.	Metallic copper is formed. Carbon has reduced copper(II) oxide to copper. Thus, carbon is more reactive than copper.
Carbon + zinc oxide	A glow spreads to the whole mixture. A grey residue is formed.	Metallic zinc is formed. Carbon has reduced zinc oxide to zinc. Thus, carbon is more reactive than zinc.
Carbon + aluminium oxide	No change	Carbon is unable to reduce aluminium oxide. Thus, carbon is less reactive than aluminium.
Carbon + magnesium oxide	No change	Carbon is unable to reduce magnesium oxide. Thus, carbon is less reactive than magnesium.

Discussion:

Carbon is more reactive than copper and zinc. Therefore, carbon can reduce copper(II) oxide and zinc oxide to their respective metals.
Carbon is less reactive than aluminium and magnesium. Thus, carbon is unable to reduce aluminium oxide and magnesium oxide.

Conclusion:
Carbon is positioned between aluminium and zinc in the reactivity series of metals towards oxygen.

Position of Hydrogen in the Reactivity Series of Metals

We can determine the position of hydrogen in the series based on the ability of hydrogen to remove oxygen from metal oxides.
To do this, a metal oxide is heated in the presence of hydrogen.
(a) If hydrogen is more reactive than the metal, hydrogen is able to remove the oxygen from the metal oxide. In other words, hydrogen is able to reduce the metal oxide to its metal while hydrogen itself is oxidised to water.
Hydrogen + metal oxide → metal + water
(b) On the other hand, if hydrogen is less reactive than the metal, hydrogen is unable to remove the oxygen from the metal oxide. Thus, no reaction will take place.
Using the same way, we can predict the position of other metals in the reactivity series.
The reactivity series of metals towards oxygen can assist us in predicting reactions involving metals.
(a) Reaction of metal oxides with carbon or hydrogen
Carbon + metal oxide → metal + carbon dioxide
Hydrogen + metal oxide → metal + water
A reaction will take place if carbon or hydrogen is more reactive than the metal. The carbon or hydrogen will remove the oxygen from the metal oxide.
(b) Reaction of metal oxides with other metals
Metal X + oxide of metal Y → oxide of metal X + Metal Y
This reaction occurs if metal X is more reactive than metal Y.
(c) Reaction of metals with water or steam.
Metal + water/steam → metal oxide + hydrogen
This reaction occurs if the metal is more reactive than hydrogen.
(d) Reaction of metals with carbon dioxide
Metal + carbon dioxide → metal oxide + carbon
This reaction occurs if the metal is more reactive than carbon.

Position of Hydrogen in the Reactivity Series of Metals Experiment

Aim: To determine the position of hydrogen in the reactivity series of metals towards oxygen.
Materials: 2 mol dm^-3 sulphuric acid, 1 mol m^-3 copper(II) sulphate solution, zinc granules, solid copper(II) oxide, solid zinc oxide, solid lead(II) oxide, solid iron(III) oxide, anhydrous calcium chloride.
Apparatus: Combustion tube, porcelain dish, flat-bottomed flask, U-tube, thistle funnel, delivery tubes, Bunsen burner, retort stand and clamps, stoppers with one hole, stopper with two holes.
Safety measure:
A mixture of hydrogen and air will explode when lighted.
Ensure that the flow of hydrogen is continuous throughout the activity.
Procedure:

A spatulaful of solid copper(II) oxide is placed in a porcelain dish.
The porcelain dish is placed in a combustion tube and the tube is clamped horizontally.
The apparatus is set up as shown in Figure.
Dry hydrogen gas is passed through the combustion tube for 5 to 10 minutes to remove all the air in the tube.
A sample of gas is collected from the small hole at the end of the combustion tube.
The gas is tested with a lighted wooden splint.
If the gas burns quietly without a squeaky ‘pop’, all the air in the tube has been removed. Otherwise, steps 5 and 6 are repeated until all the air in the tube has been removed.
The excess hydrogen gas that comes out of the end of the combustion tube is lighted.
Solid copper(II) oxide is strongly heated. Any change is observed. The flow of hydrogen gas should be continuous throughout this activity.
Steps 1 to 9 are repeated using solid zinc oxide, solid lead(II) oxide and solid iron(III) oxide, one at a time, in place of solid copper(II) oxide.

Results:
Discussion:

Hydrogen gas is produced when the zinc granules react with sulphuric acid with the presence of copper(II) sulphate solution as a catalyst.

The hydrogen gas produced is dried by passing it through anhydrous calcium chloride. Another drying agent that can be used is concentrated sulphuric acid.
The following precautions must be taken to prevent any explosion from happening.
(a) All connections to delivery tubes and stoppers should be tight.
(b) All the air in the combustion tube must be removed before lighting up the hydrogen gas that comes out of the end of the combustion tube. Otherwise, a mixture of hydrogen and air will explode when lighted.
(c) The flow of hydrogen gas should be continuous throughout the activity.
Hydrogen is more reactive than copper, lead and iron. Therefore, hydrogen can reduce copper(II) oxide, lead(II) oxide and iron(III) oxide to their respective metals.
Hydrogen is less reactive than zinc. Therefore, hydrogen is unable to reduce zinc oxide.

Conclusion:
Hydrogen is positioned between zinc and iron in the reactivity series of metals towards oxygen.