Comprehensive Study Guide: Hydrogen - Properties, Reactions, and Activity Series

Position of Hydrogen in Periodic Table

• Hydrogen is the first element in the periodic table, having an atomic number of $1$ and possessing a single electron in its valence shell ($K\text{ shell}$).
• Basic Properties and Parameters of Hydrogen:
  • Symbol: $H$
  • Atomic number: $1$
  • Relative atomic mass: $1$
  • Molecular formula: $H_2$
  • Relative molecular mass: $2$
  • Valency: $1$
  • Electronic configuration: $1$ ($K\text{ shell}$)
  • Position in Periodic Table: $1\text{st Period}$, $1\text{st Group (Group IA)}$
• Position Controversy and Dual Nature:
  • Ever since Mendeleev presented his periodic table, the exact position of hydrogen has been a subject of controversy and debate.
  • Hydrogen exhibits a dual nature because its chemical and physical properties resemble both Group $1$ (Group IA, alkali metals) and Group $17$ (Group VIIA, halogens).
  • Thompson proposed a separate position for hydrogen at the top of the periodic table, which preserves the overall symmetry of the modern periodic table.

Similarities and Differences: Hydrogen vs. Alkali Metals vs. Halogens

Similarities Between Hydrogen and Alkali Metals (Group IA / Group 1)

• Electronic Configuration:
  • Both hydrogen and all alkali metals possess exactly one electron in their outermost orbit:
  • Hydrogen ($H$, atomic number $1$): $1$
  • Lithium ($Li$, atomic number $3$): $2, 1$
  • Sodium ($Na$, atomic number $11$): $2, 8, 1$
  • Potassium ($K$, atomic number $19$): $2, 8, 8, 1$
  • Rubidium ($Rb$, atomic number $37$): $2, 8, 18, 8, 1$
  • Caesium ($Cs$, atomic number $55$): $2, 8, 18, 18, 8, 1$
  • Francium ($Fr$, atomic number $87$): $2, 8, 18, 32, 18, 8, 1$
• Valence Electrons:
  • All elements in this group have $1$ electron in their valence shell.
• Valency:
  • All alkali metals, along with hydrogen, exhibit a valency of $1$
• Ion Formation:
  • Each forms a monovalent positive cation by losing its single outer electron, demonstrating electropositive character:
  • $H \rightarrow H^+ + e^-$
  • $Li \rightarrow Li^+ + e^-$
  • $Na \rightarrow Na^+ + e^-$
• Reducing Power:
  • Both hydrogen and alkali metals act as effective reducing agents by removing oxygen from metallic oxides:
  • $CuO + H_2 \rightarrow Cu + H_2O$
  • $CuO + 2Na \rightarrow Cu + Na_2O$
• Reaction with Oxygen and Burning:
  • Hydrogen burns in oxygen with a characteristic pop sound to form its neutral oxide (water):
  • $2H_2 + O_2 \rightarrow 2H_2O$
  • Alkali metals burn vigorously when heated in oxygen to yield their respective oxides:
  • Lithium forms a monoxide: $4Li + O_2 \rightarrow 2Li_2O$
  • Sodium and hydrogen form peroxides: $2Na + O_2 \rightarrow Na_2O_2$ and $H_2 + O_2 \rightarrow H_2O_2$
  • Potassium, rubidium, and caesium form superoxides with the general formula $MO_2$ (where $M$ represents the metal): $K + O_2 \rightarrow KO_2$
• Air Reactivity and Storage:
  • Alkali metals are highly reactive and tarnish rapidly when exposed to air due to the formation of surface layers of oxides, hydroxides, and carbonates:
  • $4Na + O_2 \rightarrow 2Na_2O$
  • $Na_2O + H_2O \rightarrow 2NaOH$
  • $2NaOH + CO_2 \rightarrow Na_2CO_3 + H_2O$
  • Due to this extreme reactivity, alkali metals are always stored immersed in an inert organic solvent such as kerosene oil to isolate them from atmospheric oxygen and moisture.
• Combination with Non-Metals:
  • Hydrogen and alkali metals combine with non-metals such as oxygen, sulphur, and chlorine to form analogous compounds:
  • Hydrogen compounds: $H_2O$, $H_2S$, $HCl$
  • Sodium compounds: $Na_2O$, $Na_2S$, $NaCl$

Similarities Between Hydrogen and Halogens (Group VIIA / Group 17)

• Electronic Configuration:
  • Hydrogen and halogens each require exactly one electron to attain the stable electronic structure of their nearest inert gas:
  • $H = 1$ (requires $1$ electron for $He = 2$)
  • $F = 2, 7$ (requires $1$ electron for $Ne = 2, 8$)
  • $Cl = 2, 8, 7$ (requires $1$ electron for $Ar = 2, 8, 8$)
• Valency:
  • Both have a valency of $1$, needing to gain or share one electron to complete their valence shell.
• Formation of Ions:
  • Both display a tendency to form monovalent anions by gaining one electron to achieve noble gas configuration:
  • $H + e^- \rightarrow H^-$
  • $Cl + e^- \rightarrow Cl^-$
• Electronegative Character:
  • Both halogens and hydrogen display non-metallic, electronegative behavior by accepting electrons:
  • $H + e^- \rightarrow H^-$
  • $F + e^- \rightarrow F^-$
• Physical State:
  • Like the lighter halogens (fluorine and chlorine), hydrogen exists as a gas at room temperature.
• Atomicity:
  • Hydrogen and halogens exist naturally as homonuclear diatomic molecules ($H_2$, $F_2$, $Cl_2$, $Br_2$, $I_2$).

Differences of Hydrogen from Alkali Metals and Halogens

• Number of Electron Shells:
  • A hydrogen atom possesses only $1$ electron shell ($K\text{ shell}$), whereas all alkali metals and halogens possess $2$ or more electron shells.
• Chemical Nature of Oxides:
  • Oxide of hydrogen ($H_2O$) is strictly neutral.
  • Oxides of halogens (such as $Cl_2O$, $Cl_2O_7$) are acidic in nature.
  • Oxides of alkali metals (such as $Na_2O$, $K_2O$) are strongly basic in nature.

Discovery and Occurrence of Hydrogen

Discovery and History

• Robert Boyle produced hydrogen gas in $1672$.
• The formal credit for discovering hydrogen is given to Henry Cavendish in $1766$.
  • Cavendish prepared the gas by reacting iron with dilute acids.
  • He established its elementary nature and proved that water is produced when the gas burns in air.
  • He initially named it "inflammable air" due to its highly combustible character.
• Antoine Lavoisier in $1783$ named the gas "hydrogen", derived from Greek words meaning "water former", owing to its ability to generate water upon combustion.

Occurrence in Free and Combined States

• Free State:
  • Hydrogen is found in trace amounts in the Earth's crust ($0.98\%$) and atmosphere ($0.01\%$).
  • Volcanic gases contain approximately $0.025\%$ hydrogen gas.
  • Massive quantities of hydrogen exist in the atmospheres of the Sun and stellar bodies.
• Combined State:
  • Plant and animal tissues are constructed from hydrogen compounds bound with carbon, oxygen, and nitrogen.
  • It is a defining constituent of acids, alkalis, hydrocarbons, and proteins.
  • Found in sugars, starches, petroleum products, carbohydrates, and fats.
  • In water, hydrogen constitutes $11.1\%$ of the total mass by weight.

Preparation of Hydrogen from Water

Reaction of Cold Water with Reactive Metals

• Extremely reactive metals (potassium, sodium, calcium) react directly with cold water, forming their soluble hydroxides and liberating hydrogen gas in exothermic reactions.

• Potassium ($K$):

  • Physical properties: Density of $0.86\,g\,cm^{-3}$ (floats on water); melting point of $62^\circ\text{C}$.
  • Reaction behavior: Melts into a silvery-grey globule that darts violently around the surface of the water.
  • Equation: $2K + 2H_2O \rightarrow 2KOH + H_2\uparrow$
  • Flame and observation: Reaction generates sufficient heat to ignite the hydrogen gas, which burns with a lilac-coloured flame due to traces of vaporized potassium. Solution formed is colourless, soapy to touch, and strongly alkaline.
  • Handling precaution: Potassium catches fire upon contact with moist air; it must be stored under kerosene oil and handled exclusively using tongs, never bare hands.

• Sodium ($Na$):

  • Physical properties: Density of $0.97\,g\,cm^{-3}$ (floats on water); melting point of $97^\circ\text{C}$.
  • Reaction behavior: Melts into a silvery sphere darting across the water surface. The reaction is less violent and less exothermic than potassium.
  • Equation: $2Na + 2H_2O \rightarrow 2NaOH + H_2\uparrow$
  • Flame and observation: Ignores and burns with a golden yellow flame. The resulting solution is colourless, soapy, slightly warm, and alkaline.
  • Amalgam alternative: Sodium amalgam (an alloy of sodium with mercury) and potassium amalgam react smoothly and slowly with water, offering a safer mechanism to produce hydrogen gas.

• Calcium ($Ca$):

  • Physical properties: Sinks in water (denser than water).
  • Reaction behavior: Less vigorous than sodium.
  • Equation: $Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2\uparrow$
  • Flame and observation: Bubling of hydrogen gas occurs; the solution becomes milky, turbid, and strongly alkaline (turns red litmus paper blue).

Reaction of Hot Water and Steam with Metals

• Magnesium ($Mg$):

  • Reaction with boiling water: Reacts slowly with boiling water to form a base, magnesium hydroxide, releasing hydrogen:
  • $Mg + 2H_2O \rightarrow Mg(OH)_2 + H_2\uparrow$
  • Reaction with steam: Burning magnesium ribbon inserted into steam burns intensely with a brilliant white light, producing hydrogen gas and a white ash of magnesium oxide:
  • $Mg + H_2O \rightarrow MgO + H_2\uparrow$
  • Magnesium oxide crumbles upon heating, continually exposing fresh metal surface until the metal is consumed.

• Aluminium ($Al$):

  • Reacts with steam to produce aluminium oxide and hydrogen gas:
  • $2Al + 3H_2O \rightarrow Al_2O_3 + 3H_2\uparrow$
  • Mechanism and obstruction: The reaction initially stops because the produced aluminium oxide ($Al_2O_3$) forms a tight, protective surface layer over the metal. At higher temperatures ($800^\circ\text{C}$), this oxide layer breaks down, allowing the reaction to proceed.

• Zinc ($Zn$):

  • Less reactive than aluminium; reacts with steam only when heated continuously while steam is passed over it:
  • $Zn + H_2O \rightarrow ZnO + H_2\uparrow$
  • Physical observation: Zinc oxide ($ZnO$) displays a distinct color change—it is yellow when hot and turns white when cold.

• Iron ($Fe$):

  • Red-hot iron reacts with steam to form triferric tetra-oxide ($Fe_3O_4$) and hydrogen gas:
  • $3Fe + 4H_2O \rightleftharpoons Fe_3O_4 + 4H_2\uparrow$
  • Reversibility and Equilibrium:
  • The reaction is reversible. Initially, the forward rate is rapid while the backward rate is slow. As reactants are consumed and products accumulate, the backward reaction accelerates until both speeds match, establishing chemical equilibrium at $700^\circ\text{C}$.
  • Dynamic state: Chemical equilibrium is reached when the rates of forward and backward reactions are equal, causing the relative concentrations of reactants and products to remain constant.
  • Continuous preparation: To force the reaction forward to completion, liberated hydrogen gas must be continuously swept away from the reaction zone to prevent it from reducing $Fe_3O_4$ back to metallic iron.

• General Rate and Passivation Trends with Steam:

  • Among metals reacting with steam, the initial rate of reaction follows the order: $Mg > Al > Zn > Fe$
  • For magnesium, aluminium, and zinc, the steam reactions eventually halt prematurely because their respective metal oxides adhere firmly to the metallic surface, preventing fresh steam from contacting unreacted metal.

Reaction of Steam with Non-Metals (Carbon)

• Passivating steam over red-hot coke ($C$) yields water gas, an equimolar mixture of carbon monoxide ($CO$) and hydrogen gas ($H_2$):
  • $C + H_2O \rightarrow CO + H_2$
• Separation of Pure Hydrogen:
  • The resulting water gas mixture is bubbled through an ammoniacal cuprous chloride solution.
  • Carbon monoxide dissolves into the solution, leaving behind isolated hydrogen gas.

Activity Series of Metals and Hydrogen Preparation

Salient Features of the Activity Series

• Definition: The arrangement of metals in decreasing order of their chemical reactivity is defined as the activity series or reactivity series of metals.
• Standard Reactivity Ranking:
  • $K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au$
  • Potassium ($K$) is placed at the top as the most reactive metal; gold ($Au$) sits at the bottom as the least reactive.
• Position of Hydrogen: Although hydrogen is a non-metal, it is included in the metal activity series because it loses an electron to form a positive cation ($H^+$).
• Core Features of the Activity Series:
  • Electropositive character decreases progressively down the series.
  • Reducing power of metals decreases down the series ($K$ is the strongest reducing agent).
  • Tendency to lose valence electrons (undergo oxidation) decreases down the series ($K$ oxidizes most readily).
  • Metals positioned above hydrogen displace hydrogen from water and dilute acids; metals situated below hydrogen cannot displace it.
  • Displacement from salt solutions: A higher metal displaces any metal below it from its aqueous salt solution:
  • $CuSO_4(aq) + Zn \rightarrow ZnSO_4(aq) + Cu$
  • The greater the difference in position between two metals in the activity series, the more rapid the displacement reaction.
  • Reduction behavior of metal oxides:
  • Oxides of $K, Na, Ca, Mg, Al$ are extremely stable and cannot be reduced by common chemical reducing agents ($H_2, CO, C$); they can only be reduced by electrolytic processes.
  • Metal oxides below aluminium ($Fe_2O_3, ZnO, PbO, CuO$) can be reduced by heating with carbon, carbon monoxide, or hydrogen gas.
  • Oxides and nitrates of low-reactivity metals ($Hg, Ag, Au$) decompose directly upon strong heating to yield free metal:
    • $2HgO \rightarrow 2Hg + O_2\uparrow$
    • $2AgNO_3 \rightarrow 2Ag + 2NO_2 + O_2\uparrow$
  • Corrosion resistance: Metals located below copper ($Cu$) resist oxidation and do not rust easily.
  • Natural occurrence: Highly electropositive metals at the top occur primarily as chlorides ($Na, K$) or oxides ($Al$), whereas intermediate/lower metals ($Zn, Pb, Cu, Hg, Ag$) occur naturally as sulphides.

Displacement of Hydrogen from Dilute Acids

• Metals preceding hydrogen in the activity series displace hydrogen from dilute hydrochloric acid ($HCl$) and dilute sulphuric acid ($H_2SO_4$).
• Reactivity Categories with Dilute Acids:
  • Potassium ($K$), Sodium ($Na$), Calcium ($Ca$): React violently and explosively with dilute acids; unsafe for hydrogen collection.
  • Magnesium ($Mg$), Aluminium ($Al$), Zinc ($Zn$), Iron ($Fe$): Displace hydrogen vigorously to gently, providing controlled sources of hydrogen gas:
  • $Mg + 2HCl \rightarrow MgCl_2 + H_2\uparrow$
  • $Mg + H_2SO_4 \rightarrow MgSO_4 + H_2\uparrow$
  • $2Al + 6HCl \rightarrow 2AlCl_3 + 3H_2\uparrow$
  • $2Al + 3H_2SO_4 \rightarrow Al_2(SO_4)_3 + 3H_2\uparrow$
  • $Zn + 2HCl \rightarrow ZnCl_2 + H_2\uparrow$
  • $Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2\uparrow$
  • $Fe + 2HCl \rightarrow FeCl_2 + H_2\uparrow$
  • $Fe + H_2SO_4 \rightarrow FeSO_4 + H_2\uparrow$
  • Lead ($Pb$), Copper ($Cu$), Mercury ($Hg$), Silver ($Ag$), Gold ($Au$): Do not displace hydrogen from dilute acids.
  • Example: $Cu + HCl\text{ (dil.)} \rightarrow \text{No reaction}$

Reasons for Selecting Zinc in Laboratory Preparation

• Zinc ($Zn$) is the preferred metal for the standard laboratory preparation of hydrogen. Alternative metals are excluded due to specific chemical and practical limitations:
  1. Sodium ($Na$) and Potassium ($K$): React explosively with acids, making the process dangerous.
  2. Calcium ($Ca$) and Magnesium ($Mg$): Economically prohibitive and expensive for routine routine laboratory use.
  3. Aluminium ($Al$): Develops a dense, protective surface coating of aluminium oxide ($Al_2O_3$) due to high oxygen affinity, preventing further acid attack.
  4. Iron ($Fe$): Requires external heating to react efficiently, and the liberated hydrogen is contaminated with foul-smelling impurities such as hydrogen sulphide ($H_2S$) and sulphur dioxide ($SO_2$).
  5. Lead ($Pb$): Reacts initially but quickly forms an insoluble crust of lead sulphate ($PbSO_4$) or lead chloride ($PbCl_2$) over the metal, terminating further chemical action.
  6. Metals Below Hydrogen ($Cu, Hg, Ag, Au$): Possess lower electropositivity than hydrogen and cannot displace it from acid solutions.

Displacement of Hydrogen from Alkalis (Amphoteric Nature)

• Unique Nature of $Zn, Pb, Al$:

  • Metallic zinc, lead, and aluminium possess amphoteric properties. They react with acids as well as with hot, concentrated alkalis (such as sodium hydroxide $NaOH$ and potassium hydroxide $KOH$) to generate soluble salts and hydrogen gas.

• Chemical Equations with Hot Concentrated Alkalis:

  • Reaction with Zinc:
  • $Zn + 2NaOH \rightarrow Na_2ZnO_2 + H_2\uparrow$ (Sodium zincate)
  • $Zn + 2KOH \rightarrow K_2ZnO_2 + H_2\uparrow$ (Potassium zincate)
  • Reaction with Lead:
  • $Pb + 2NaOH \rightarrow Na_2PbO_2 + H_2\uparrow$ (Sodium plumbite)
  • $Pb + 2KOH \rightarrow K_2PbO_2 + H_2\uparrow$ (Potassium plumbite)
  • Reaction with Aluminium:
  • $2Al + 2NaOH + 2H_2O \rightarrow 2NaAlO_2 + 3H_2\uparrow$ (Sodium aluminate)
  • $2Al + 2KOH + 2H_2O \rightarrow 2KAlO_2 + 3H_2\uparrow$ (Potassium aluminate)

• Amphoteric Nature of Oxides and Hydroxides:

  • The oxides and hydroxides of zinc, lead, and aluminium are amphoteric—they react with both acids and strong bases to yield salt and water:
  • Reaction with acid: $ZnO + 2HCl \rightarrow ZnCl_2 + H_2O$
  • Reaction with alkali: $ZnO + 2NaOH \rightarrow Na_2ZnO_2 + H_2O$

Questions & Discussion

• Justify the position of Hydrogen in the periodic table.

  • Response Summary: Hydrogen is placed in Period $1$, Group IA ($1$) due to its atomic number $1$ and single valence electron ($1s^1$). However, because it shows properties of both alkali metals (losing $1$ electron, acting as a reducing agent) and halogens (gaining $1$ electron, diatomic molecule $H_2$), its dual nature led Thompson to suggest placing it distinctly at the top of the table to preserve periodic symmetry.

• Why does hydrogen show dual nature?

  • Response Summary: Hydrogen shows a dual nature because its single valence shell ($K\text{ shell}$) needs one electron to complete its duplet. It can lose one electron like alkali metals to form $H^+$, or gain one electron like halogens to form $H^-$.

• Compare hydrogen with alkali metals on the basis of:

  • (i) Ion formation: Both lose one electron ($H \rightarrow H^+ + e^-$, $Na \rightarrow Na^+ + e^-$).
  • (ii) Reducing power: Both reduce metal oxides ($CuO + H_2 \rightarrow Cu + H_2O$, $CuO + 2Na \rightarrow Cu + Na_2O$).
  • (iii) Reaction with oxygen: Both burn in oxygen to form oxides ($2H_2 + O_2 \rightarrow 2H_2O$, $4Li + O_2 \rightarrow 2Li_2O$).
  • (iv) Oxide formation: Hydrogen forms neutral oxide ($H_2O$); alkali metals form strongly basic monoxides, peroxides, or superoxides ($Na_2O_2$, $KO_2$).

• In what respect does hydrogen differ from:

  • (i) Alkali metals: Hydrogen has only $1$ electron shell; alkali metals have $2$ or more. Hydrogen forms a neutral oxide, while alkali metals form basic oxides.
  • (ii) Halogens: Hydrogen has only $1$ shell and forms neutral oxide; halogens have $2$ or more shells and form acidic oxides ($Cl_2O_7$).

• Why was hydrogen called "inflammable air"?

  • Response Summary: Henry Cavendish named it "inflammable air" because of its highly combustible nature and its flame production when ignited in air.

• State some sources of hydrogen.

  • Response Summary: Water ($11.1\%$ by weight), organic tissues, volcanic gases ($0.025\%$), Earth's crust ($0.98\%$), atmosphere ($0.01\%$), hydrocarbons, acids, and alkalis.

• Compare hydrogen and halogens on the basis of:

  • (i) Physical state: Both are gases ($H_2, F_2, Cl_2$).
  • (ii) Ion formation: Both gain one electron to form anions ($H^-$, $Cl^-$).
  • (iii) Valency: Both have a valency of $1$
  • (iv) Reaction with oxygen: Hydrogen forms neutral oxide ($H_2O$); halogens form acidic oxides.

• Which metal is preferred for preparation of hydrogen:

  • (i) From water? Sodium/Potassium amalgams for controlled reactions with cold water; iron or zinc with steam.
  • (ii) From acid? Granulated zinc ($Zn$).

• Why are zinc and aluminium considered to have a unique nature? Give balanced equations to support your answer.

  • Response Summary: They are amphoteric—reacting with both acids and alkalis to form hydrogen. Equations: $Zn + 2HCl \rightarrow ZnCl_2 + H_2$ and $Zn + 2NaOH \rightarrow Na_2ZnO_2 + H_2$

• Complete the following word equations:

  • (a) Sodium hydroxide + zinc $\rightarrow$ hydrogen + sodium zincate ($Na_2ZnO_2$)
  • (b) Calcium + water $\rightarrow$ calcium hydroxide + hydrogen ($Ca(OH)_2 + H_2$)