ICSE Class 9 Chemistry Chapter 06 Study of the first Element Hydrogen

Study of the First Element - Hydrogen

Syllabus

Position of the non-metal (Hydrogen) in the periodic table and general group characteristics with reference to valency electrons, burning, ion formation applied to the above mentioned element.

(i) Hydrogen from: water, dilute acids and alkalis.

(a) Hydrogen from water.

The action of cold water on sodium, potassium and calcium.

The action of hot water on magnesium.

The action of steam on aluminium, zinc, and iron; (reversibility of reaction between iron and steam)

The action of steam on non-metal (carbon).

Application of activity series for the above mentioned reactions.

(b) Displacement of hydrogen from dilute acids:

The action of dilute sulphuric acid or hydrochloric acid on metals: Mg, Al, Zn and Fe (To understand reasons for not using other metals and dilute nitric acid)

(c) Displacement of hydrogen from alkalis:

The action of Alkalis ((NaOH, KOH) on Al, Zn and Pb - unique nature of these elements.

(ii) The preparation and collection of hydrogen by a standard laboratory method other than electrolysis. In the laboratory preparation, the reason for using zinc, the impurities in the gas, their removal and the precautions in the collection of the gas must be mentioned.

(iii) Industrial manufacture of hydrogen by Bosch process:

Main reactions and conditions. Separation of CO2 and CO from hydrogen.

(iv) Oxidation and reduction reactions

Differences in terms of addition and removal of oxygen/hydrogen.

6.1 Position of Hydrogen in Periodic Table

Hydrogen is the first element in the periodic table. Its atomic number is 1, and it has only one electron in its valence shell. Therefore, it belongs to the first group and the first period of the periodic table.

It is expected that the properties of hydrogen should be similar to those of the other members of the 1st group, but this is not the case. Ever since Mendeleev presented his periodic table, the position of hydrogen in the table has been a matter of controversy and debate. This is mainly because some of the properties of hydrogen resemble the properties of the Group I A elements (alkali metals) while others are similar to those of the Group VII A (halogens) elements. Thomson suggested a separate position for hydrogen. He puts it at the top of the periodic table; that does not disturb the symmetry of the modern periodic table.

Note: Hydrogen shows a dual nature since it resembles the alkali metals of Group IA and the halogens of Group VII A (17).

Teacher's Note

Hydrogen's unique position in the periodic table reflects its role as the most abundant element in the universe and the simplest atom, making it fundamental to understanding chemistry and stellar processes.

6.2 Similarities Between Hydrogen and Alkali Metals

1. Electronic configuration: They have only one electron in their outermost orbits.

K L M N O P (Shells)

H (1): 1

Li (3): 2, 1

Na (11): 2, 8, 1

K (19): 2, 8, 8, 1

Rb (37): 2, 8, 18, 8, 1

Cs (55): 2, 8, 18, 18, 8, 1

Fr (87): 2, 8, 18, 32, 18, 8, 1

2. Valence electrons: All elements have one electron in their outermost orbit, i.e. their valency shell and so all of them have one valence electron.

3. Valency: All alkali metals, including hydrogen, have valency 1.

4. Ion formation: Each of them can form cation, i.e. positive ion, by loss of an electron.

\[H \rightarrow H^+ + e^-\]

\[Li \rightarrow Li^+ + e^-\]

\[Na \rightarrow Na^+ + e^-\]

As such all these elements have electropositive character.

5. Reducing power: Both the alkali metals and hydrogen act as reducing agents.

\[CuO + H_2 \rightarrow Cu + H_2O\]

\[CuO + 2Na \rightarrow Cu + Na_2O\]

6. Burning: Hydrogen burns in oxygen to form its oxide (water).

\[2H_2 + O_2 \rightarrow 2H_2O\]

Hydrogen burns with a pop sound. Alkali metals also burn vigorously when heated in oxygen to form their respective oxides.

Lithium forms monoxide

\[4Li + O_2 \rightarrow 2Li_2O\]

Sodium forms peroxide

\[2Na + O_2 \rightarrow Na_2O_2\]

while potassium, rubidium and caesium form superoxides having the general formula \(MO_2\) (where M stands for metal).

\[K + O_2 \rightarrow KO_2\]

Action of alkali metals on air: Alkali metals are so reactive that they get rapidly tarnished when exposed to air. This is due to the formation of oxides and hydroxides, and finally, carbonates, on their surface.

Example:

\[4Na + O_2 \rightarrow 2Na_2O\]

\[Na_2O + H_2O \rightarrow 2NaOH\]

\[2NaOH + CO_2 \rightarrow Na_2CO_3 + H_2O\]

Note: Due to their reactivity, alkali metals are always stored in an inert organic solvent, such as kerosene oil. This inert solvent prevents them from coming in contact with air and moisture.

7. Combination with non-metals: Hydrogen as well as alkali metals react with non-metals like oxygen, sulphur and chlorine to form respective compounds.

Hydrogen - forms H2O; H2S; HCl

Sodium - forms Na2O; Na2S; NaCl and so on.

6.3 Similarities Between Hydrogen and Halogens

1. Electronic configuration: Hydrogen and halogens have one electron less than the nearest inert gas.

H = 1 [He = 2],

F = 2, 7 [Ne = 2, 8),

Cl = 2, 8, 7 (Ar 2, 8, 8)

2. Valency: Both have valency 1. Thus, they accept one electron to attain the electronic configuration of the nearest inert gas.

3. Formation of ions: Both show a tendency to form anions since they are one electron short of the nearest inert gas configuration.

\[H + e^- \rightarrow H^-\]

\[Cl + e^- \rightarrow Cl^-\]

4. Electronegative character: Both halogens and hydrogen are non-metals. They show electronegative character.

\[H + e^- \rightarrow H^-\]

\[F + e^- \rightarrow F^-\]

5. Physical state: Like halogens (fluorine and chlorine), hydrogen too is a gas.

6. Atomicity: Hydrogen as well as halogens exist in the form of diatomic molecules (H2, F2, Cl2, Br2, I2).

Properties of Hydrogen different from those of Alkali Metals and Halogens.

1. Hydrogen atom has only one shell but alkali metals and halogens have two or more shells.

2. Oxide of hydrogen, H2O, is a neutral oxide. Oxides of halogens like Cl2O, Cl2O7, etc., are acidic in nature, while oxides of alkali metals like Na2O, K2O, etc. are basic in nature.

Teacher's Note

Understanding hydrogen's dual nature helps us appreciate why it can form both covalent bonds (like halogens) and ionic compounds (like alkali metals), making it crucial in both organic and inorganic chemistry we encounter daily.

6.4 Discovery

The credit for the discovery of hydrogen goes to Henry Cavendish (1766), although Robert Boyle had prepared the gas in 1672. Not only did Cavendish prepare the gas from iron and dilute acids, he also established its elementary nature and showed that water is formed when this gas burns in air. He called it inflammable air because of its combustible nature. It was on account of its ability to form water that Lavoisier, in 1783, named it hydrogen (Greek word meaning water former).

6.5 Occurrence

Free state

In free state, hydrogen is found in traces in the earth's crust and atmosphere. Volcanic gases contain 0.025% of it, the earth's crust 0.98%, the earth's atmosphere 0.01% and the atmospheres of the Sun and the stars contain 1.1% hydrogen.

Combined state

(i) Plant and animal tissues are made up of compounds of hydrogen with carbon, oxygen and nitrogen.

(ii) Hydrogen is the characteristic constituent of acids, alkalis, hydrocarbons and proteins. In addition to these, sugar, starch, petroleum products, proteins, carbohydrates and also fats contain hydrogen.

(iii) In water, it is 11-1% by weight.

6.6 Preparation of Hydrogen

Since water, acids and alkalis contain hydrogen, they form the cheapest raw materials for preparation and manufacture of hydrogen.

6.6.1 Hydrogen from water

Hydrogen from cold water and metals

Reactive metals like potassium, sodium and calcium react with cold water forming their corresponding hydroxides and evolving hydrogen. Their reactions are exothermic.

(a) Potassium:

1. Potassium when added to water, floats on water (density 0.86 g/cc).

2. It melts [at 62°C], forming a silver grey globule that darts about on the surface of the water.

3. Reaction: \(2K + 2H_2O \longrightarrow 2KOH + H_2\)

The reaction is highly exothermic and vigorous.

4. It catches fire and burns with a lilac-coloured flame.

5. Bubbles of hydrogen gas are seen, and the solution formed is colourless, soapy and alkaline.

Although pure hydrogen burns with a pale blue flame, the colour of the flame is lilac due to the presence of traces of potassium vapour.

Caution! Potassium readily catches fire in air because it reacts with water vapour. So it is kept in kerosene oil and handled carefully. Do not touch it with your hand, rather use a pair of tongs.

(b) Sodium:

1. It floats on water; density 0.97 g/cc; melting point 97°C.

2. It melts, forming a silvery globule, which darts about on the surface of the water.

3. Reaction: \(2Na + 2H_2O \longrightarrow 2NaOH + H_2\)

The reaction is less exothermic and less vigorous as compared to potassium.

(c) Calcium:

1. Calcium sinks in water.

2. Reaction: \(Ca + 2H_2O \longrightarrow Ca(OH)_2 + H_2\)

Reaction is less vigorous than sodium.

3. Bubbles of hydrogen are liberated, and the solution turns milky, turbid and alkaline. If red litmus is introduced in solution it turns blue.

Action of hot water and steam on metals

(i) Magnesium reacts slowly with boiling water and forms a base, magnesium hydroxide liberating hydrogen gas.

\[Mg + 2H_2O \rightarrow Mg(OH)_2 + H_2\uparrow\]

(ii) Magnesium burns in steam with an intense white light, liberating hydrogen gas and white ash, i.e., magnesium oxide.

\[Mg + H_2O \rightarrow MgO + H_2\uparrow\]

Magnesium oxide crumbles down due to heating. Further exposing magnesium to steam results in the liberation of hydrogen gas.

Action of steam on metals

(a) Aluminium:

1. It reacts with steam to liberate hydrogen.

2. It forms aluminium oxide, which makes it inactive, and liberates hydrogen. There is no further reaction due to the oxide coating.

\[2Al + 3H_2O \xrightarrow{800°C} Al_2O_3 + 3H_2\]

Steam

3. At high temperature, however, the coating breaks and aluminium reacts with steam, liberating hydrogen.

(b) Zinc:

1. It is even less reactive i.e. it reacts only when it is heated and steam is passed over it.

2. Hydrogen is liberated and zinc is converted to white zinc oxide.

\[Zn + H_2O \rightarrow ZnO + H_2\]

3. Zinc oxide is yellow when it is hot but white when it is cold.

(c) Iron:

1. Iron is less reactive than zinc, but red hot iron reacts with steam, forming triferic tetra-oxide and hydrogen gas.

2. This reaction is reversible. If the hydrogen formed is not removed, the iron oxide formed is reduced back to iron.

\[3Fe + 4H_2O \rightleftharpoons Fe_3O_4 + 4H_2\uparrow\]

Action of steam on non-metals

Carbon: When steam is passed over red hot coke, water gas (CO + H2) is formed.

\[C + H_2O \rightarrow CO + H_2\]

Water gas

6.7 Application of Activity Series in the Preparation of Hydrogen

Based on reaction with water, metals are arranged in decreasing order of reactivity.

K

Na

Ca

Mg

Al

Zn

Fe

Pb

[H]

Cu

Hg

Ag

Au

Arrangement of metals in decreasing order of reactivity in the form of a series is called the activity or reactivity series of metals.

Thus, potassium being the most reactive metal is placed at the top of the list, and the least reactive metal being gold, is placed at the bottom of the list.

Hydrogen, though it is a non-metal, is included in this series because it can form a positive ion. It would occupy a position based on formation of its positive ion.

The ability of metals to reduce water to hydrogen decreases on going down the series.

\[K > Na > Ca > Mg > Al > Zn ...............\]

Lead and the metals that are further below in the activity series: No reaction with water, even when the metal is hot and steam is used.

6.7.1 Salient features of the Activity Series

(i) Electropositive character decreases down the series.

(ii) Reducing power of metals decreases down the series. Thus, potassium is the strongest reducing agent.

(iii) Tendency of metals to lose valence electrons, i.e. tendency of metals to get oxidized, decreases down the series. Thus, potassium is the most readily oxidized metal.

(iv) Metals above hydrogen displace hydrogen from water and dilute acid, but metals below hydrogen do not.

\[2Na + 2H_2O \rightarrow 2NaOH + H_2\]

\[Cu + H_2O \rightarrow\] No reaction

\[Zn + 2HCl \text{(dil.)} \rightarrow ZnCl_2 + H_2\]

\[Cu + HCl \text{(dil.)} \rightarrow\] No reaction

(v) A metal higher up in the series displaces a metal below it from the salt solutions of the latter.

\[CuSO_4(aq) + Zn \rightarrow ZnSO_4(aq) + Cu\]

Larger the difference in the positions of the metals in the activity series, more rapidly the displacement occurs.

(vi) Oxides of metals K, Na, Ca, Mg and Al cannot be reduced by common reducing agents like H2, CO or C. They can be reduced only by electrolysis. Metal oxides below aluminium can be reduced by heating them in the presence of reducing agents.

\[ZnO + C \rightarrow Zn + CO\]

\[PbO + C \rightarrow Pb + CO\]

\[PbO + CO \rightarrow Pb + CO_2\]

\[PbO + H_2 \rightarrow Pb + H_2O\]

(vii) Oxides and nitrates of less reactive metals like Hg, Ag and Au decompose on strong heating to give metals.

\[2HgO \rightarrow 2Hg + O_2\]

\[2AgNO_3 \rightarrow 2Ag + 2NO_2 + O_2\]

(viii) Metals below copper do not rust easily.

(ix) Since the reactivity series lists metals in decreasing order of electropositivity, it is expected that those at the top will combine with more electronegative elements in their minerals. For example, sodium and potassium occur mainly as their chlorides, aluminium as its oxide, and zinc, lead, copper, mercury and silver as their sulphides.

6.7.2 Displacement of hydrogen from dilute acids

Hydrogen is displaced from dilute hydrochloric acid and dilute sulphuric acid when the latter react with some metals which are more reactive than hydrogen.

The extent to which these reactions occur for a given metal is also based on the activity series of metals.

K | These displace hydrogen from dilute acids (HCl or H2SO4) with explosive violence

Na |

Ca |

Mg | These displace hydrogen from dilute acids (HCl or H2SO4) vigorously, but not violently forming their respective salts

Al |

Zn |

F | These displace hydrogen from dilute acids (HCl or H2SO4) gently forming their respective salts

Ni |

Sn |

Pb |

H |

Cu | These do not displace hydrogen from dilute acids at all

Hg |

Ag |

Au |

When dilute sulphuric acid or dilute hydrochloric acid reacts with metals above hydrogen in the activity series, hydrogen is produced. But in the case of metals below hydrogen in the activity series, like copper, silver, etc., this does not happen.

Magnesium, aluminium, zinc and iron react with dilute HCl or dilute H2SO4, liberating hydrogen and forming their respective salts.

Zinc is the most preferred metal in laboratory preparation of hydrogen because of the following reasons:

(i) Sodium and potassium react violently with acid.

(ii) Calcium and magnesium are expensive.

(iii) Aluminium forms a protective coating of Al2O3 due to its great affinity for oxygen. So, it does not give hydrogen with acid after the reaction has occurred for some time.

(iv) Iron has to be heated, but then the hydrogen thus produced contains impurities like hydrogen sulphide and sulphur dioxide.

(v) Lead reacts with dilute sulphuric or hydrochloric acid and forms an insoluble coating of lead sulphate or lead chloride. Therefore, further reaction is prevented.

(vi) Hydrogen cannot be prepared from metals that are below it in the activity series of metals, such as copper and mercury, since only metals that are more reactive than hydrogen can displace it from acids.

e.g., \(Cu + HCl \rightarrow\) No reaction (dil.)

Teacher's Note

Understanding why zinc is preferred in laboratory hydrogen production connects to real-world chemistry labs where safety and efficiency matter - it reacts steadily without extreme violence or requiring expensive resources.

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