Hydrogen Class 9 ICSE — Chemistry Chapter 6
The hydrogen class 9 ICSE chapter, Study of the First Element Hydrogen, is Chapter 6 of your Chemistry book. It explains hydrogen’s position in the periodic table and its resemblance to both alkali metals and halogens. The chapter then moves through preparation from water and acids, the activity series, the Bosch process, properties, and uses. It closes with oxidation and reduction, where hydrogen acts as the reducing agent.
Download Study of the First Element Hydrogen Class 9 ICSE Chemistry Chapter 6 PDF
Download the Study of the First Element Hydrogen Class 9 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 18 pages, 16 sections.
What is inside Chapter 6
18 pages · 16 sections · 1 worked examples
| Section | Title |
|---|---|
| 6.1 | POSITION OF HYDROGEN IN PERIODIC TABLE |
| 6.2 | SIMILARITIES BETWEEN HYDROGEN AND ALKALI METALS |
| 6.3 | SIMILARITIES BETWEEN HYDROGEN AND HALOGENS |
| 6.4 | DISCOVERY |
| 6.5 | OCCURRENCE |
| 6.6 | PREPARATION OF HYDROGEN |
| 6.6.1 | Hydrogen from water |
| 6.7.1 | Salient features of the Activity Series |
| 6.7.2 | Displacement of hydrogen from dilute acids |
| 6.9 | MANUFACTURE OF HYDROGEN |
| 6.10.1 | Physical properties |
| 6.10.2 | Chemical properties |
| 6.11 | USES OF HYDROGEN |
| 6.12 | OXIDATION AND REDUCTION |
| 6.12.1 | Oxidation |
| 6.12.2 | Reduction |
6.1 POSITION OF HYDROGEN IN PERIODIC TABLE
Hydrogen is the first element in the periodic table. Its atomic number is 1, and it carries a single electron in its valence shell. So it belongs to the first group and the first period. You first met this arrangement in the periodic table chapter of Class 9 Chemistry.
Ever since Mendeleev’s table appeared, hydrogen’s position has been debated. Some of its properties match the Group IA alkali metals, while others match the Group VIIA halogens. Thompson suggested a separate place for hydrogen at the top of the table, which keeps the symmetry intact.
The key point is that hydrogen shows a dual nature because it resembles both families.
6.2 SIMILARITIES BETWEEN HYDROGEN AND ALKALI METALS
Hydrogen shares seven features with the alkali metals. Both have one electron in the outermost shell, one valence electron, and valency 1. Both form a positive ion by losing that electron. The pattern for hydrogen is \mathrm{H} \rightarrow \mathrm{H}^{+} + \mathrm{e}^{-}, exactly as lithium gives \mathrm{Li} \rightarrow \mathrm{Li}^{+} + \mathrm{e}^{-}.
Both act as reducing agents. Hydrogen pulls oxygen off copper(II) oxide: \mathrm{CuO} + \mathrm{H}_{2} \rightarrow \mathrm{Cu} + \mathrm{H}_{2}\mathrm{O}. Sodium does the same: \mathrm{CuO} + 2\mathrm{Na} \rightarrow \mathrm{Cu} + \mathrm{Na}_{2}\mathrm{O}.
Both burn in oxygen to form oxides. Hydrogen burns with a pop sound: 2\mathrm{H}_{2} + \mathrm{O}_{2} \rightarrow 2\mathrm{H}_{2}\mathrm{O}. The alkali metals differ among themselves. Lithium gives a monoxide, sodium a peroxide, and potassium, rubidium and caesium give superoxides of the general form \mathrm{MO}_{2}.
Both combine with non-metals like oxygen, sulphur and chlorine. Hydrogen forms \mathrm{H}_{2}\mathrm{O}, \mathrm{H}_{2}\mathrm{S} and \mathrm{HCl}, while sodium forms \mathrm{Na}_{2}\mathrm{O}, \mathrm{Na}_{2}\mathrm{S} and \mathrm{NaCl}. Alkali metals tarnish in air through oxide, hydroxide and carbonate layers. That is why they are stored under kerosene oil.
6.3 SIMILARITIES BETWEEN HYDROGEN AND HALOGENS
Hydrogen also resembles the halogens. Its shell arrangement \mathrm{H} = 1 is one electron short of helium (2). Fluorine (2, 7) is one short of neon, and chlorine (2, 8, 7) is one short of argon. So hydrogen and the halogens both have valency 1 and accept one electron to form anions:
\mathrm{H} + \mathrm{e}^{-} \rightarrow \mathrm{H}^{-}, \quad \mathrm{Cl} + \mathrm{e}^{-} \rightarrow \mathrm{Cl}^{-}Both are non-metals and show electronegative character. Fluorine and chlorine are gases, and so is hydrogen. All of them form diatomic molecules: \mathrm{H}_{2}, \mathrm{F}_{2}, \mathrm{Cl}_{2}, \mathrm{Br}_{2}, \mathrm{I}_{2}.
Hydrogen differs from both families in two ways. Its atom has only one shell, while alkali metals and halogens have two or more. Its oxide is neutral, while halogen oxides are acidic and alkali metal oxides are basic.
| Property | Hydrogen | Alkali metals | Halogens |
|---|---|---|---|
| Valency | 1 | 1 | 1 |
| Ion formed | \mathrm{H}^{+} or \mathrm{H}^{-} | \mathrm{M}^{+} cation | \mathrm{X}^{-} anion |
| Number of shells | one | two or more | two or more |
| Nature of oxide | neutral (\mathrm{H}_{2}\mathrm{O}) | basic (\mathrm{Na}_{2}\mathrm{O}) | acidic (\mathrm{Cl}_{2}\mathrm{O}_{7}) |
6.4 DISCOVERY
Henry Cavendish is credited with discovering hydrogen in 1766, though Robert Boyle had prepared the gas earlier, in 1672. Cavendish made it from iron and dilute acids, established that it was an element, and showed that burning it in air forms water.
Because the gas burned, he called it inflammable air. In 1783 Lavoisier renamed it hydrogen, from the Greek for water former.
6.5 OCCURRENCE
In the free state, hydrogen is found only in traces. Volcanic gases contain 0.025% of it, the earth’s crust 0.98%, the atmosphere 0.01%, and the Sun and stars 1.1%.
In the combined state, hydrogen is widespread. Plant and animal tissues contain it, and it is a characteristic constituent of acids, alkalis, hydrocarbons and proteins. Sugar, starch, petroleum products and fats all contain hydrogen. Water holds 11.1% hydrogen by weight.
6.6 PREPARATION OF HYDROGEN
Water, acids and alkalis all contain hydrogen, so they are the cheapest raw materials for its preparation and manufacture. The next sections build on this idea, first with water and metals, then with acids.
6.6.1 Hydrogen from water
Reactive metals react with cold water to give a hydroxide and hydrogen. Potassium floats, melts, darts about, and burns with a lilac flame: 2\mathrm{K} + 2\mathrm{H}_{2}\mathrm{O} \rightarrow 2\mathrm{KOH} + \mathrm{H}_{2}. Sodium behaves in the same way with a golden yellow flame: 2\mathrm{Na} + 2\mathrm{H}_{2}\mathrm{O} \rightarrow 2\mathrm{NaOH} + \mathrm{H}_{2}. Calcium sinks and gives a milky, alkaline solution: \mathrm{Ca} + 2\mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{Ca}(\mathrm{OH})_{2} + \mathrm{H}_{2}.
Hot water and steam need less reactive metals. Magnesium reacts slowly with boiling water and burns in steam with an intense white light. Aluminium needs steam near 800^\circ\mathrm{C}: 2\mathrm{Al} + 3\mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{Al}_{2}\mathrm{O}_{3} + 3\mathrm{H}_{2}. Zinc gives white zinc oxide, which is yellow while hot.
Red hot iron reacts with steam in a reversible reaction:
3\mathrm{Fe} + 4\mathrm{H}_{2}\mathrm{O} \rightleftharpoons \mathrm{Fe}_{3}\mathrm{O}_{4} + 4\mathrm{H}_{2}This reaction reaches chemical equilibrium at 700^\circ\mathrm{C}. At equilibrium, the forward and backward reactions run at the same speed, so the amounts of reactants and products stop changing. If the hydrogen is removed, the oxide is not reduced back to iron, and the forward reaction continues.
Carbon is the non-metal in this section. Steam over red hot coke forms water gas:
\mathrm{C} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{CO} + \mathrm{H}_{2}These reactions arrange metals in decreasing order of reactivity. That list is the activity series, which the chapter treats next.
6.7.1 Salient features of the Activity Series
The activity series lists metals in decreasing order of reactivity. Potassium sits at the top and gold at the bottom. Hydrogen is included even though it is a non-metal, because it can form a positive ion.
\mathrm{K} \gt \mathrm{Na} \gt \mathrm{Ca} \gt \mathrm{Mg} \gt \mathrm{Al} \gt \mathrm{Zn} \gt \mathrm{Fe} \gt \mathrm{Ni} \gt \mathrm{Sn} \gt \mathrm{Pb} \gt \mathrm{H} \gt \mathrm{Cu} \gt \mathrm{Hg} \gt \mathrm{Ag} \gt \mathrm{Au}To recall the order, use the mnemonic ‘Kings Never Come Marching After Zebras; Foxes Never Sneak Past Houses; Cats Hunt And Gorillas.’
Nine features follow from this order:
- Electropositive character decreases down the series.
- Reducing power decreases, so potassium is the strongest reducing agent.
- The tendency to lose electrons and get oxidized decreases.
- Metals above hydrogen displace it from water and dilute acids; metals below it do not.
- A higher metal displaces a lower one from its salt solution, as zinc displaces copper from \mathrm{CuSO}_{4}.
- Oxides of K, Na, Ca, Mg and Al are reduced only by electrolysis.
- Oxides below aluminium are reduced by \mathrm{H}_{2}, CO or carbon.
- Oxides and nitrates of Hg, Ag and Au decompose on strong heating.
- Metals below copper do not rust easily.
6.7.2 Displacement of hydrogen from dilute acids
Metals above hydrogen displace it from dilute hydrochloric acid and dilute sulphuric acid. How violently they react depends on their position in the activity series.
- K, Na, Ca — displace hydrogen with explosive violence.
- Mg, Al, Zn — displace it vigorously, forming their salts.
- Fe, Ni, Sn, Pb — displace it gently, forming their salts.
- Cu, Hg, Ag, Au — no reaction.
The general pattern is metal + dilute acid \rightarrow salt + hydrogen. Iron, for example, gives \mathrm{Fe} + 2\mathrm{HCl} \rightarrow \mathrm{FeCl}_{2} + \mathrm{H}_{2}.
Zinc is the preferred metal for laboratory preparation. The other candidates fail for specific reasons:
- Sodium and potassium react with explosive violence.
- Calcium and magnesium are expensive.
- Aluminium forms a protective \mathrm{Al}_{2}\mathrm{O}_{3} coating and stops reacting.
- Iron must be heated, and the gas it gives is impure.
- Lead forms an insoluble coating of lead chloride or lead sulphate.
- Copper and mercury sit below hydrogen and cannot displace it.
Zinc, lead and aluminium have a unique nature. They react with acids and also with hot concentrated alkalis, releasing hydrogen. For example, \mathrm{Zn} + 2\mathrm{NaOH} \rightarrow \mathrm{Na}_{2}\mathrm{ZnO}_{2} + \mathrm{H}_{2}. Their oxides and hydroxides are amphoteric, meaning they react with both acids and bases.
In the laboratory, granulated zinc reacts with dilute \mathrm{HCl} or \mathrm{H}_{2}\mathrm{SO}_{4} in a flat-bottom flask. Hydrogen is collected by the downward displacement of water. This works because hydrogen is only slightly soluble in water, and it cannot be collected over air since it forms an explosive mixture with it.
The gas carries impurities: \mathrm{H}_{2}\mathrm{S}, \mathrm{SO}_{2}, oxides of nitrogen, phosphine (\mathrm{PH}_{3}), arsine (\mathrm{AsH}_{3}), carbon dioxide and water vapour. Silver nitrate solution removes phosphine and arsine. Lead nitrate solution removes hydrogen sulphide. Caustic potash solution removes sulphur dioxide, carbon dioxide and oxides of nitrogen. Drying agents such as fused calcium chloride remove water vapour.
Concentrated sulphuric acid is not used for drying because it reacts with hydrogen: \mathrm{H}_{2}\mathrm{SO}_{4} + \mathrm{H}_{2} \rightarrow 2\mathrm{H}_{2}\mathrm{O} + \mathrm{SO}_{2}.
Precautions matter because hydrogen–air mixtures explode. The apparatus must be airtight, no flame may burn nearby, and gas is collected only after all the air has escaped. Nitric acid is avoided because it oxidizes the hydrogen to water: 3\mathrm{Zn} + 8\mathrm{HNO}_{3} \rightarrow 3\mathrm{Zn}(\mathrm{NO}_{3})_{2} + 4\mathrm{H}_{2}\mathrm{O} + 2\mathrm{NO}. Concentrated sulphuric acid is avoided because it gives sulphur dioxide.
Test for hydrogen: it burns with a pale blue flame and does not support combustion. When premixed with air or oxygen, it explodes with a pop sound.
6.9 MANUFACTURE OF HYDROGEN
Large-scale hydrogen is made by the Bosch process. In step one, steam passes over hot coke at about 1000^\circ\mathrm{C} to form water gas:
\mathrm{C} + \mathrm{H}_{2}\mathrm{O} \xrightarrow{1000^\circ\mathrm{C}} \mathrm{CO} + \mathrm{H}_{2}This step is endothermic. In step two, water gas mixes with excess steam and passes over ferric oxide as catalyst, with chromic oxide as promoter, at about 450^\circ\mathrm{C}:
\mathrm{CO} + \mathrm{H}_{2} + \mathrm{H}_{2}\mathrm{O} \xrightarrow[katex display="true"]450^\circ\mathrm{C}{\mathrm{Fe}_{2}\mathrm{O}_{3}} \mathrm{CO}_{2} + 2\mathrm{H}_{2}[/katex]
This step is exothermic. Step three removes carbon dioxide by forcing the mixture through cold water under pressure, or through caustic potash solution: 2\mathrm{KOH} + \mathrm{CO}_{2} \rightarrow \mathrm{K}_{2}\mathrm{CO}_{3} + \mathrm{H}_{2}\mathrm{O}. Step four removes leftover carbon monoxide with ammoniacal cuprous chloride solution.
Hydrogen is also obtained by electrolysing acidulated water. A little sulphuric acid makes the water conduct electricity. Hydrogen collects at the cathode and oxygen at the anode:
2\mathrm{H}_{2}\mathrm{O} \xrightarrow{\text{electric current}} 2\mathrm{H}_{2} + \mathrm{O}_{2}The main advantage is that both gases are produced pure.
6.10.1 Physical properties
Hydrogen is colourless, odourless, tasteless and non-toxic. Its density is 0.0899\ \text{g/mL}, making it the lightest known gas, 14.4 times less dense than air. It is very slightly soluble in water; one litre dissolves about 20 mL of the gas.
Hydrogen liquefies only below -240^\circ\mathrm{C} at about 20 atmospheres. Finely divided palladium takes up 900 volumes of hydrogen at room temperature. This adsorption is called occlusion.
One simple experiment proves the lightness. Hydrogen passed through soap solution forms bubbles that fly upwards and burst. Gas behaviour like this connects to the study of gas laws in this Class 9 Chemistry book.
6.10.2 Chemical properties
Hydrogen reacts with non-metals, metals, metallic oxides and organic compounds, but it does not support combustion. With oxygen it burns with a pop sound: 2\mathrm{H}_{2} + \mathrm{O}_{2} \rightarrow 2\mathrm{H}_{2}\mathrm{O}. Equal volumes of hydrogen and chlorine react slowly in diffused sunlight but explosively in direct sunlight.
With nitrogen it forms ammonia in the Haber process: \mathrm{N}_{2} + 3\mathrm{H}_{2} \rightleftharpoons 2\mathrm{NH}_{3}. The conditions are 450\text{-}500^\circ\mathrm{C}, 200–900 atm, with finely divided iron as catalyst and molybdenum as promoter. With molten sulphur, hydrogen gives hydrogen sulphide. The hydrogen–oxygen mixture is called a detonating mixture.
Dry hydrogen reacts with heated potassium, sodium and calcium to give hydrides, for example 2\mathrm{Na} + \mathrm{H}_{2} \rightarrow 2\mathrm{NaH}. These hydrides react with water to release hydrogen again.
Hydrogen is a strong reducing agent. It removes oxygen from the oxides of less active metals and is itself oxidized to water. Watch the colour changes, which the chapter lists:
- \mathrm{Fe}_{2}\mathrm{O}_{3} + 3\mathrm{H}_{2} \rightarrow 2\mathrm{Fe} + 3\mathrm{H}_{2}\mathrm{O} — brown to grey.
- \mathrm{PbO} + \mathrm{H}_{2} \rightarrow \mathrm{Pb} + \mathrm{H}_{2}\mathrm{O} — yellow to silvery grey.
- \mathrm{CuO} + \mathrm{H}_{2} \rightarrow \mathrm{Cu} + \mathrm{H}_{2}\mathrm{O} — black to reddish brown.
Worked pattern: pass dry hydrogen over 4 g of black copper(II) oxide in a heated tube. The black solid slowly turns reddish brown as copper forms, and colourless water droplets collect at the cooler end. The remaining solid weighs less than 4 g because the oxygen has left with the hydrogen as water. This is the reduction half of a redox reaction.
With a nickel catalyst, hydrogen adds across double and triple bonds. This is hydrogenation: \mathrm{H}_{2}\mathrm{C}=\mathrm{CH}_{2} + \mathrm{H}_{2} \rightarrow \mathrm{C}_{2}\mathrm{H}_{6}.
6.11 USES OF HYDROGEN
Hydrogen has a high heat of combustion, so it is used as a fuel in coal gas, water gas and liquid hydrogen. The oxy-hydrogen torch burns hydrogen and oxygen to reach about 2500^\circ\mathrm{C}. It cuts and welds metals, melts platinum and quartz, and fuses alumina into synthetic rubies and sapphires. The atomic hydrogen torch reaches about 2800^\circ\mathrm{C} and welds alloys of tungsten, manganese and chromium.
Hydrogen is used to make ammonia by the Haber process, hydrogen chloride and methyl alcohol: \mathrm{CO} + 2\mathrm{H}_{2} \rightarrow \mathrm{CH}_{3}\mathrm{OH}. It hydrogenates vegetable oils into vanaspati ghee at about 200^\circ\mathrm{C}, with nickel, platinum or palladium as catalyst. It also makes artificial petrol from coal and reduces heated metallic oxides in metal extraction: \mathrm{Fe}_{2}\mathrm{O}_{3} + 3\mathrm{H}_{2} \rightarrow 2\mathrm{Fe} + 3\mathrm{H}_{2}\mathrm{O}. Oxides of very active metals like sodium and potassium are not reduced by hydrogen.
Meteorological balloons once used hydrogen. Helium is now preferred because hydrogen is inflammable.
6.12 OXIDATION AND REDUCTION
In the hydrogen class 9 ICSE chapter, oxidation and reduction are introduced as a pair. Oxidation is the loss of electrons, and reduction is the gain of electrons. The two always occur together, because the electrons lost by one substance are gained by another at the same moment. This idea links many reactions across the Class 9 Chemistry book.
6.12.1 Oxidation
In the electronic concept, oxidation is the loss of electrons. Zinc loses two: \mathrm{Zn} \rightarrow \mathrm{Zn}^{2+} + 2\mathrm{e}^{-}.
Oxidation also happens in four other ways:
- Addition of oxygen: \mathrm{C} + \mathrm{O}_{2} \rightarrow \mathrm{CO}_{2}.
- Addition of an electronegative ion: 2\mathrm{FeCl}_{2} + \mathrm{Cl}_{2} \rightarrow 2\mathrm{FeCl}_{3}.
- Removal of hydrogen: \mathrm{H}_{2}\mathrm{S} + \mathrm{Cl}_{2} \rightarrow 2\mathrm{HCl} + \mathrm{S}.
- Removal of an electropositive ion: 2\mathrm{KI} + \mathrm{H}_{2}\mathrm{O}_{2} \rightarrow 2\mathrm{KOH} + \mathrm{I}_{2}.
An oxidizing agent accepts electrons, or provides oxygen or an electronegative ion, or removes hydrogen or an electropositive ion. Examples include solids like \mathrm{KMnO}_{4}, liquids like hydrogen peroxide, and gases like chlorine.
Tests for oxidizing agents: heating liberates oxygen that rekindles a glowing splinter; hydrogen sulphide gives a yellow sulphur precipitate; warming with concentrated \mathrm{HCl} liberates chlorine; and potassium iodide paper turns brown.
Worked pattern: in \mathrm{H}_{2}\mathrm{S} + \mathrm{Cl}_{2} \rightarrow 2\mathrm{HCl} + \mathrm{S}, hydrogen sulphide is oxidized because it loses hydrogen, while chlorine gains hydrogen and is reduced.
6.12.2 Reduction
In the electronic concept, reduction is the gain of electrons. Copper(II) ions gain two: \mathrm{Cu}^{2+} + 2\mathrm{e}^{-} \rightarrow \mathrm{Cu}.
Reduction also happens in four other ways:
- Removal of oxygen: \mathrm{CuO} + \mathrm{H}_{2} \rightarrow \mathrm{Cu} + \mathrm{H}_{2}\mathrm{O}.
- Addition of an electropositive ion: 2\mathrm{HgCl}_{2} + \mathrm{SnCl}_{2} \rightarrow \mathrm{Hg}_{2}\mathrm{Cl}_{2} + \mathrm{SnCl}_{4}.
- Addition of hydrogen: \mathrm{Cl}_{2} + \mathrm{H}_{2}\mathrm{S} \rightarrow 2\mathrm{HCl} + \mathrm{S}.
- Removal of an electronegative ion: 2\mathrm{FeCl}_{3} + \mathrm{H}_{2}\mathrm{S} \rightarrow 2\mathrm{FeCl}_{2} + 2\mathrm{HCl} + \mathrm{S}.
A reducing agent provides electrons, hydrogen or an electropositive ion, or removes oxygen or an electronegative ion. Examples include carbon, zinc, hydrogen sulphide and carbon monoxide. Common tests: brown fumes with nitric acid, reduction of heated CuO to red copper, decolourising \mathrm{KMnO}_{4}, turning acidified potassium dichromate from orange to green, and changing iron(III) salts from light yellow to light green.
When zinc meets copper sulphate, oxidation and reduction occur side by side:
\mathrm{Zn} \rightarrow \mathrm{Zn}^{2+} + 2\mathrm{e}^{-} \quad (\text{oxidation})
\mathrm{Cu}^{2+} + 2\mathrm{e}^{-} \rightarrow \mathrm{Cu} \quad (\text{reduction})
These are the two half-reactions of a redox reaction. Because electrons move from zinc to copper(II) ions in one step, the two halves cannot be separated. A reaction with simultaneous oxidation and reduction is a redox reaction.
Common mistakes to avoid:
- Addition of hydrogen is reduction of the substance that gains it, not oxidation. Removal of hydrogen is oxidation.
- In \mathrm{CuO} + \mathrm{H}_{2}, hydrogen is the reducing agent and CuO is the oxidizing agent. Do not swap them.
- Nitric acid and concentrated sulphuric acid do not liberate hydrogen from zinc. They oxidize it instead.
What to practise from Hydrogen Class 9 ICSE
Work through the exercise blocks in the order the book gives them. The question counts below come from this chapter’s own exercises.
- Exercise 6(A) — 15 questions on hydrogen’s position, dual nature, comparisons, and reactions with water and acids.
- Exercise 6(B) — 10 questions on laboratory preparation, impurities, the Bosch process, and metal–water reactions.
- Exercise 6(C) — 20 questions covering the whole chapter, including the activity series, acid reactions, and reasons.
- End-of-chapter redox questions — 13 items on oxidation, reduction, oxidizing and reducing agents, half-reactions, and multiple-choice checks.
Equations for water gas, the Bosch process, zinc with dilute acid, and the reduction of copper(II) oxide recur across all three exercises. Definitions of oxidation, reduction, oxidizing agent and reducing agent are tested in the end-of-chapter block. For the official syllabus notice, see the CISCE website.
Hydrogen Class 9 ICSE frequently asked questions
Why is hydrogen class 9 ICSE placed in both group 1 and group 17?
Hydrogen has one valence electron, like the alkali metals of group 1. It is also one electron short of the nearest inert gas, like the halogens of group 17. This dual nature has made its position a subject of debate since Mendeleev’s time.
How is hydrogen prepared in the laboratory for Class 9 ICSE and why is zinc preferred?
Granulated zinc reacts with dilute hydrochloric acid or dilute sulphuric acid in a flat-bottom flask. Zinc is preferred because sodium and potassium react violently, calcium and magnesium are costly, and iron needs heating and gives impure gas.
What is the Bosch process in this ICSE chapter?
Steam over hot coke gives water gas. The water gas, with excess steam, passes over ferric oxide with chromic oxide promoter to give carbon dioxide and more hydrogen. Carbon dioxide is removed under pressure or by caustic potash, and leftover carbon monoxide by ammoniacal cuprous chloride.
What is the test for hydrogen gas in this ICSE chapter?
Hydrogen burns with a pale blue flame and does not support combustion. When premixed with air or oxygen, it explodes with a pop sound. That pop is the standard confirmation test.
How do I tell oxidation from reduction in Class 9 ICSE Chemistry?
Oxidation is loss of electrons, addition of oxygen or an electronegative ion, or removal of hydrogen or an electropositive ion. Reduction is the reverse: gain of electrons, removal of oxygen, addition of hydrogen, or addition of an electropositive ion. The two always occur together in a redox reaction.
Reference: ICSE Class 9 Chemistry textbook material, chapter 6: Study of the First Element Hydrogen — CISCE-aligned.