Study of the First Element – Hydrogen

Hydrogen, the very first element in the periodic table, earns its position due to its singular proton and electron, making it the simplest and lightest atom. Often called the most abundant element in the universe, it is a key building block of stars like our sun. However, its presence on Earth is almost entirely in combined forms, most famously as water (H₂O), which covers a majority of our planet’s surface. This chapter establishes that while hydrogen gas itself is rare in our atmosphere, its compounds are fundamental to the environment and life itself. The element is primarily prepared in laboratories through reactions such as acids with active metals like zinc or the electrolysis of water, processes that demonstrate its reactive nature.

The chemical behavior of hydrogen is dual-natured, earning it a unique spot in the periodic table. It can both donate and accept an electron, allowing it to form compounds with a vast range of elements. As a strong reducing agent, it readily gives up its electron to non-metals like oxygen, with its combustion reaction being highly exothermic. This reaction with oxygen to form water is a key area of study, highlighting its energy potential. Furthermore, its ability to act somewhat like a metal allows it to form ionic hydrides with highly reactive s-block elements, while its covalent bonding with non-metals like carbon and nitrogen is the foundation of all organic chemistry.

Beyond its chemical behavior, the biological and industrial significance of hydrogen is immense. Its role in water makes it a non-negotiable component for all known life forms, acting as the universal solvent for biological processes. In industry, hydrogen is crucial for the Haber process in manufacturing ammonia-based fertilizers, which are vital for global agriculture. It is also used in the hydrogenation of oils to produce fats like vanaspati ghee. Looking forward, hydrogen is championed as a clean fuel of the future, as its combustion produces only water vapor, offering a potential solution to fossil fuel pollution and positioning it as a cornerstone of sustainable energy research.

Exercise 6 (A)

Question 1. 

Justify the position of hydrogen in the periodic table.

Ans:

Hydrogen’s position is unique and debated because it exhibits properties of both Group 1 (Alkali Metals) and Group 17 (Halogens).

Justification for Group 1:

• Like alkali metals, hydrogen has one electron in its valence shell.
• It can lose one electron to form a positive ion (H⁺), which is why it is commonly placed above Lithium.

Justification for Group 17:

• Like halogens, hydrogen is diatomic (H₂), similar to F₂ or Cl₂.
• It needs one more electron to complete its valence shell and can form a negative ion (H⁻), called a hydride, similar to halides (F⁻, Cl⁻).
• It forms covalent compounds with non-metals (e.g., H₂O, CH₄), just as halogens do.

Conclusion:
Due to its dual behavior, hydrogen is best placed separately at the top of the periodic table, acting as a bridge between the electropositive alkali metals and the electronegative halogens. Its common placement in Group 1 is primarily based on its electron configuration, not its chemical resemblance.

Question 2. 

Why does hydrogen show dual nature?

Ans:

Hydrogen exhibits a dual nature because it is a quantum-scale object. This behavior is not unique to hydrogen but applies to all matter. However, it becomes significant and observable for tiny particles like the electron in a hydrogen atom.

At this small scale, the electron does not behave like a miniature planet orbiting a sun. Instead, its position and momentum are described by a wavefunction—a mathematical expression representing a probability wave. This wave-like nature explains the fixed energy levels and the distinct orbital shapes within the atom.

In experiments, we sometimes detect the electron as a single, localized particle (particle nature), but the probability of where we will find it is governed by its wave-like character. Therefore, hydrogen’s electron, and by extension hydrogen itself, demonstrates both wave-like and particle-like properties, a fundamental principle of quantum mechanics.

Question 3. 

1. Compare hydrogen with alkali metals on the basis of: Ion formation 

2. Compare hydrogen with alkali metals on the basis of: Reducing power 

3. Compare hydrogen with alkali metals on the basis of: Reaction with oxygen 

4. Compare hydrogen with alkali metals on the basis of: Oxide formation 

Ans:

1. Ion Formation

• Hydrogen: It can form a positive ion (H⁺) by losing its single electron, similar to alkali metals. However, the H⁺ ion is just a bare proton, making it unstable and non-existent in a free state. It immediately associates with other molecules.
• Alkali Metals: They readily and stably form positive ions (M⁺) by losing their single valence electron. These ions are stable in solutions and crystalline structures.

2. Reducing Power

• Hydrogen: It acts as a mild reducing agent, especially when heated. Its ability to donate an electron is relatively weak.
• Alkali Metals: These are among the strongest known reducing agents. They violently donate their valence electron, reacting vigorously with many substances.

3. Reaction with Oxygen

• Hydrogen: It reacts with oxygen to form a neutral oxide—water (H₂O). The reaction is highly exothermic but typically requires initiation.
• Alkali Metals: They react with oxygen to form various ionic oxides (like M₂O, M₂O₂, MO₂). The reaction is spontaneous and often vigorous or explosive, especially for heavier members like Potassium and Rubidium.

4. Oxide Formation

• Hydrogen: Its oxide is water (H₂O), a neutral, covalent compound. It is a liquid at room temperature.
• Alkali Metals: Their oxides (e.g., Li₂O, Na₂O₂) are solid, ionic compounds. When dissolved in water, they form strong alkaline (basic) solutions.

Question 4. 

1. In what respect does hydrogen differ from: alkali metals 

2. In what respect does hydrogen differ from: halogens?  

Ans:

1. In what respect does hydrogen differ from Alkali Metals?

While alkali metals have one valence electron like hydrogen, they differ significantly. Alkali metals are solid, readily lose their single electron to form positive ions (cations), and are highly reactive metals. In contrast, hydrogen is a diatomic gas, it can both lose and gain an electron, and it does not possess metallic characteristics like luster or malleability under standard conditions.

2. In what respect does hydrogen differ from Halogens?

Hydrogen differs from halogens in its electron behavior and ionic form. Halogens need only one electron to complete their octet, which they gain easily to form stable negative ions (anions). Hydrogen, while it can gain an electron to form a hydride ion (H⁻), does so very rarely and is a weak tendency. Furthermore, hydrogen exists as a diatomic gas but lacks the characteristic color, high reactivity with metals, and disinfectant properties typical of halogens like chlorine.

Question 5. 

1. Give the general group study of hydrogen with reference to valence electrons 

2. Give the general group study of hydrogen with reference to burning 

3. Give the general group study of hydrogen with reference to reducing power

Ans:

1. General Group Study of Hydrogen with Reference to Valence Electrons

Hydrogen’s single valence electron creates a unique dual identity. Like alkali metals (Group 1), it can lose this electron to form a H⁺ cation (a proton). However, it also resembles halogens (Group 17), as it can gain one electron to achieve a stable duplet, forming a H⁻ hydride ion. This single valence electron prevents a perfect fit with any one group, making hydrogen a true chemical anomaly.

2. General Group Study of Hydrogen with Reference to Burning

When burned in air, hydrogen reacts with oxygen to form water (H₂O), releasing significant energy. This combustion is clean, producing no carbon byproducts. In this behavior, it is fundamentally different from Groups 1 and 17. Alkali metals burn in air to form solid oxides, peroxides, or superoxides, while halogens are non-flammable. Hydrogen’s burning nature is unique and defines its role as a clean fuel.

3. General Group Study of Hydrogen with Reference to Reducing Power

Hydrogen is a potent reducing agent, especially when heated. It can reduce the heated oxides of many less-reactive metals (e.g., copper oxide, iron oxide) to the pure metal. In this respect, it acts similarly to Group 1 metals, which are strong reductants. However, unlike alkali metals that reduce water violently at room temperature, hydrogen’s reducing action typically requires thermal activation, placing its reducing power in a distinct, intermediate category.

Question 6. 

Why was hydrogen called ‘inflammable air’?

Ans:

The story of hydrogen’s discovery centers on the meticulous work of Henry Cavendish in the 1760s. While conducting experiments that involved reacting certain metals with acids, he isolated a new, mysterious gas. This substance didn’t fit the profile of any air known at the time, such as common air or fixed air (carbon dioxide), prompting careful investigation.

Its most defining characteristic was its incredible combustibility. Cavendish observed that this new gas would readily ignite, often producing a very faint, pale blue flame that could be difficult to see in well-lit conditions. Based directly on this dramatic and hazardous property, he logically named it ‘inflammable air,’ a perfectly descriptive term for an era before modern atomic theory.

This label, ‘inflammable air,’ stuck for some time, perfectly capturing the gas’s behavior without yet revealing its fundamental identity. It wasn’t until later, when its role in forming water was proven and its status as a primary chemical element was established, that the name ‘hydrogen,’ meaning ‘water-former,’ was adopted, leaving Cavendish’s initial name as a historical footnote.

**********Question 7. 

State some sources of hydrogen.

Ans:

Natural gas reforming (steam methane reforming) is the most common industrial source. Electrolysis of water splits it into hydrogen and oxygen using electricity. Biomass gasification converts plant or waste material into hydrogen-rich gas. Some industrial processes, like chlorine production, yield hydrogen as a by-product. Photoelectrochemical and biological methods, using sunlight or microbes, are also emerging sources.

Question 8. 

1. Compare hydrogen and halogens on the basis of: physical state 

2. Compare hydrogen and halogens on the basis of: ion formation 

3. Compare hydrogen and halogens on the basis of: valency 

4. Compare hydrogen and halogens on the basis of: reaction with oxygen

Ans:

1. Physical State: At room temperature, hydrogen exists as a colourless, odourless diatomic gas (H₂). In contrast, the halogens display a progression of physical states: fluorine (F₂) and chlorine (Cl₂) are gases, bromine (Br₂) is a volatile liquid, and iodine (I₂) is a solid. This trend highlights increasing molecular size and intermolecular forces down the halogen group, unlike hydrogen’s consistent gaseous nature.

2. Ion Formation: Hydrogen exhibits unique behaviour, as it can both lose and gain a single electron. It loses an electron to form a positive ion (H⁺, a proton) or gain one to form a negative hydride ion (H⁻). Halogens, being highly electronegative, only gain one electron to form stable negative ions called halide ions (X⁻, e.g., F⁻, Cl⁻). Thus, hydrogen shares properties with both alkali metals (losing e⁻) and halogens (gaining e⁻).

3. Valency: Both hydrogen and all halogens have a valency of 1. However, the basis differs. Hydrogen achieves a stable duplet (like helium) by either sharing its single electron, donating it to form H⁺, or accepting one to form H⁻. Halogens achieve a stable octet (like noble gases) specifically by gaining or sharing one electron, making their valency consistently 1.

4. Reaction with Oxygen: Hydrogen burns in oxygen with a pale blue flame to form a stable, neutral oxide—water (H₂O). The halogens react with oxygen to form multiple, often unstable, acidic oxides. For example, chlorine forms oxides like Cl₂O and ClO₂, which are powerful oxidizing agents and decompose readily, unlike the very stable water molecule.

Question 9. 

1. Which metal is preferred for the preparation of hydrogen? from water? 

2. Which metal is preferred for the preparation of hydrogen? from acid? 

Ans:

1. Calcium is typically chosen for hydrogen production from water because it reacts moderately, ensuring safer handling compared to more reactive metals.
2. Zinc is often selected for hydrogen preparation from acid due to its consistent and manageable reaction rate, which avoids excessive speed or slowdown.

Question 10. 

1. Write the reaction of steam with a red hot iron. 

2.Why is this reaction considered a reversible reaction?

Ans:

1. Steam passed over red-hot iron produces magnetite and hydrogen gas. The reaction is:
   3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
2. This reaction is reversible because the products can revert to the original substances. When hydrogen gas is heated with iron oxide, it forms iron and steam, demonstrating the reverse process. Thus, the reaction can proceed in both directions under different conditions.

Question 11. 

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

Ans:

Zinc and aluminium are considered to have a unique nature primarily because they are amphoteric. This means they, and their oxides, can react with both acids and strong bases to form salts and water, a property not common to most metals like iron or copper.

This amphoteric behaviour is supported by their reactions. With acids, they behave like typical metals, displacing hydrogen. For example:
With Acid:
Zn + 2HCl → ZnCl₂ + H₂↑
2Al + 6HCl → 2AlCl₃ + 3H₂↑

With Base:
They react with strong alkalis like sodium hydroxide to form complex salts (zincate or aluminate) and hydrogen gas.
Zn + 2NaOH → Na₂ZnO₂ + H₂↑ (Sodium zincate)
2Al + 2NaOH + 2H₂O → 2NaAlO₂ + 3H₂↑ (Sodium aluminate)

Furthermore, their oxides also show this dual character:
Zinc Oxide:
ZnO + 2HCl → ZnCl₂ + H₂O (acidic behaviour)
ZnO + 2NaOH → Na₂ZnO₂ + H₂O (basic behaviour)

Aluminium Oxide:
Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

Question 12. 

1. Write a balanced equation for the following: Iron reacts with dil. HCl 

2. Write a balanced equation for the following: Zinc reacts with caustic soda solution 

3. Write a balanced equation for the following: Lead reacts with potassium hydroxide 4. Write a balanced equation for the following: Aluminium reacts with fused sodium hydroxide.

Ans:

1. Iron with dilute hydrochloric acid:
   Fe + 2HCl → FeCl₂ + H₂
2. Zinc with caustic soda solution (aqueous sodium hydroxide):
   Zn + 2NaOH + 2H₂O → Na₂[Zn(OH)₄] + H₂
3. Lead with potassium hydroxide:
   Pb + 2KOH → K₂PbO₂ + H₂
4. Aluminium with fused sodium hydroxide (molten, anhydrous):
   2Al + 6NaOH → 2Na₃AlO₃ + 3H₂

Question 13. 

1. Write the balanced equation and give your observation when the following metal reacts: Sodium with cold water 

2. Write the balanced equation and give your observation when the following metal reacts: Calcium with cold water 

3. Write the balanced equation and give your observation when the following metal reacts: Magnesium with boiling water 

4. Write the balanced equation and give your observation when the following metal reacts: Magnesium with steam

Ans:

1. Sodium with Cold Water
Balanced Equation: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
Observation: The sodium metal darts on the water surface, melts into a shiny ball due to the heat released, and produces a hissing sound. The reaction is vigorous, and the hydrogen gas evolved may burn with a yellow flame. The resulting solution is alkaline (sodium hydroxide).

2. Calcium with Cold Water
Balanced Equation: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)
Observation: Calcium reacts less violently than sodium. It sinks and bubbles of hydrogen gas are evolved steadily. The water becomes milky due to the formation of sparingly soluble calcium hydroxide.

3. Magnesium with Boiling Water
Balanced Equation: Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g)
Observation: In boiling water, magnesium reacts very slowly. A faint fizzing can be observed over time as hydrogen gas is produced, but no strong ignition occurs. A thin layer of magnesium hydroxide forms on the metal.

4. Magnesium with Steam
Balanced Equation: Mg(s) + H₂O(g) → MgO(s) + H₂(g)
Observation: When steam is passed over heated magnesium, it reacts vigorously. The magnesium burns with a bright white flame, producing white magnesium oxide powder and releasing hydrogen gas which may ignite with a pop.

Question 14. 

1. Under what conditions iron reacts with water. 

2. Give the balanced equation of the reaction. 

3. What is noticed if the products are not allowed to escape?

Ans:

1. Conditions for Iron Reacting with Water

Iron does not readily react with water under standard conditions, such as at room temperature or with liquid water. For a direct chemical reaction to occur, iron must be exposed to water vapor (steam) at elevated temperatures. Typically, this requires heating iron to a high temperature, often described as red-hot or above 570°C, while steam is passed over it. This process facilitates the formation of iron oxide and hydrogen gas, unlike the slow corrosion (rusting) that happens in the presence of liquid water and oxygen.

2. Balanced Equation of the Reaction

The balanced chemical equation for the reaction between iron and steam is:
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
In this equation, solid iron (Fe) reacts with gaseous water (steam, H₂O) to produce solid iron(II,III) oxide (Fe₃O₄, also known as magnetite) and hydrogen gas (H₂).

3. Observation if Products Are Not Allowed to Escape

If the hydrogen gas produced in the reaction is not permitted to escape—for instance, in a closed system—it accumulates and increases pressure. This buildup can inhibit the forward reaction because hydrogen gas acts as a reducing agent and may react with the iron oxide formed. Specifically, the reverse reaction Fe₃O₄(s) + 4H₂(g) → 3Fe(s) + 4H₂O(g) can occur, leading to an equilibrium state. As a result, the net reaction slows down or may not go to completion, and one might observe that the iron is not fully converted to iron oxide. Additionally, if the system is sealed, pressure buildup could pose safety risks, such as potential rupture.

Question 15. 

1. From the knowledge of activity series, name a metal which shows the following properties It reacts readily with cold water 

2. From the knowledge of activity series, name a metal which shows the following properties: It displaces hydrogen from hot water. 

3. From the knowledge of activity series, name a metal which shows the following properties: It displaces hydrogen from dilute HCl. 

4. From the knowledge of activity series, name a metal which shows the following properties: It forms a base which is insoluble in water.

Ans:

1. Sodium (Na) is a metal from the activity series that reacts readily with cold water. It reacts vigorously, displacing hydrogen gas and forming sodium hydroxide.
2. Magnesium (Mg) is a metal that displaces hydrogen from hot water. It reacts with steam or hot water to produce magnesium oxide or hydroxide and hydrogen gas.
3. Zinc (Zn) is a metal that displaces hydrogen from dilute hydrochloric acid. Being placed above hydrogen in the activity series, it reacts to form zinc chloride and hydrogen gas.
4. Copper (Cu) is a metal that forms a base which is insoluble in water. When its ions react with a base like sodium hydroxide, it forms copper(II) hydroxide, a characteristic blue precipitate that does not dissolve.

Question 16. 

1. Complete the following word equation: Sodium hydroxide + zinc → hydrogen + _________ 

2. Complete the following word equation:Calcium + water → calcium hydroxide + _________

Ans:

1. Sodium hydroxide + zinc → hydrogen + Sodium zincate

2. Calcium + water → calcium hydroxide + Hydrogen

Exercise 6 (B)

Question 1. 

1. Hydrogen can be prepared with the metal zinc by using: acid Give an equation. 

2. Hydrogen can be prepared with the metal zinc by using: alkali Give an equation. 

3. Hydrogen can be prepared with the metal zinc by using: water Give an equation.

Ans:

1. Zinc with Acid
   Zinc reacts with dilute sulfuric acid to yield zinc sulfate and hydrogen gas:
   Zn + H₂SO₄ → ZnSO₄ + H₂↑
2. Zinc with Alkali
   When zinc is heated with concentrated sodium hydroxide solution, sodium zincate and hydrogen are formed:
   Zn + 2NaOH + 2H₂O → Na₂[Zn(OH)₄] + H₂↑
3. Zinc with Water
   Zinc reacts with steam at high temperature to produce zinc oxide and hydrogen gas:
   Zn + H₂O (steam) → ZnO + H₂↑

Question 2. 

1. For laboratory preparation of hydrogen, give the following: materials used 

2. For laboratory preparation of hydrogen, give the following: method of collection 

3. For laboratory preparation of hydrogen, give the following: chemical equation 

4. For laboratory preparation of hydrogen, give the following: fully-labeled diagram 

Ans:

1. Materials Used
The laboratory preparation of hydrogen gas typically employs zinc granules and dilute sulfuric acid as the primary reactants. Zinc is preferred due to its moderate reactivity, which ensures a steady and controllable evolution of gas. The apparatus consists of a Woulfe’s bottle or a flat-bottomed flask fitted with a thistle funnel and a delivery tube. The acid used is usually dilute sulfuric acid (approximately 1:4 ratio with water). Very dilute or very concentrated acid is avoided, as the former yields gas too slowly and the latter can produce irritating sulfur dioxide. Small amounts of impurities in the zinc help initiate the reaction more effectively.

2. Method of Collection
Hydrogen gas is collected by the method of downward displacement of water. This technique is suitable because hydrogen is only very slightly soluble in water. The delivery tube from the reaction flask is connected and led into an inverted gas jar or burette, which is initially full of water and placed in a water trough. As the hydrogen gas bubbles into the jar, it displaces the water downwards and gets collected at the top. It can also be collected by upward displacement of air, given it is lighter than air, but the water displacement method is preferred as it yields purer gas, free from moisture and air.

3. Chemical Equation
The reaction between zinc and dilute sulfuric acid is a single displacement reaction, where zinc replaces hydrogen from the acid. The balanced chemical equation is:
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
The ionic equation highlights the redox nature of the reaction:
Zn(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂(g)
In this process, zinc sulfate is formed in solution, and hydrogen gas is liberated as colourless bubbles.

4. Fully-Labeled Diagram
(A written description of the diagram for clarity)
The setup shows a Woulfe’s bottle or flask placed on a stand. The thistle funnel is inserted through a cork, with its lower end dipped below the acid level in the bottle to prevent the escape of gas. A delivery tube is fitted to the side arm of the bottle. This tube leads into an inverted gas jar (or a graduated burette) placed in a water trough. The gas jar is initially full of water and held in place by a stand and clamp. The entire assembly for collecting gas by downward displacement of water is clearly shown. Labels indicate: Zinc granules, Dilute Sulfuric Acid, Hydrogen Gas collecting at the top of the jar, and Water Level falling as gas collects. The diagram should clearly show that the end of the delivery tube is beneath the mouth of the gas jar, allowing bubbles of gas to rise and accumulate.

Question 3. 

1. Name the impurities present in hydrogen prepared in the laboratory. 

2. How can these impurities be removed? 

Ans:

1. Impurities Present in Laboratory-Prepared Hydrogen

Hydrogen gas prepared in the laboratory, typically via the reaction of a dilute acid (like hydrochloric or sulfuric acid) with an active metal (like zinc or iron), contains the following main impurities:

• Water Vapour (H₂O): The reaction occurs in an aqueous solution, so the evolved gas is saturated with moisture.
• Acidic Vapours: Traces of the reacting acid (e.g., HCl mist or H₂SO₄ spray) can be carried over with the gas.
• Hydrogen Sulfide (H₂S): A common impurity when commercial grades of zinc or iron, which contain sulfides, are used.
• Arsine (AsH₃) and Phosphine (PH₃): These toxic, foul-smelling gases can form in trace amounts if the metal used contains arsenic or phosphorus impurities.
• Carbon Dioxide (CO₂) and Oxides of Nitrogen: These may form as minor impurities depending on the concentration and type of acid used.
• Unreacted Air: Air present in the apparatus at the start of the reaction may mix with the initial hydrogen collected.

2. Removal of These Impurities

The impurities are removed by passing the generated hydrogen gas through a series of washing or drying bottles containing specific reagents, in a logical order:

1. Acidic Vapours: The gas is first bubbled through a water scrubber or a wash bottle containing distilled water. This removes most of the entrained acid spray.
2. Hydrogen Sulfide, Arsine, and Phosphine: The gas is then passed through a wash bottle containing a solution of lead(II) acetate or silver nitrate. Lead acetate removes H₂S by forming a black precipitate of lead sulfide. Silver nitrate solution removes all three (H₂S, AsH₃, PH₃) by forming precipitates.
3. Oxidic Impurities (like Oxides of Nitrogen): The gas may be passed through a wash bottle containing an alkaline solution (e.g., potassium hydroxide, KOH) to absorb any acidic oxides.
4. Water Vapour (Drying): Finally, the gas is dried by passing it through a drying tower or U-tube containing a suitable desiccant. Common drying agents used are:
   • Concentrated sulfuric acid (in a wash bottle).
   • Anhydrous calcium chloride (granular).
   • Phosphorus pentoxide (P₂O₅) for very thorough drying.

Important Note on Order: The gas is always dried (step 4) as the final step to prevent moisture from interfering with the previous chemical scrubbing steps. The typical sequence is: Acid Removal → Chemical Scrubbers (Pb(CH₃COO)₂, KOH) → Final Drying.

Question 4. 

Which test should be made before collecting hydrogen in a gas jar?

Ans:

Here’s how to do it:

1. Allow the gas to flow initially: When you begin producing hydrogen (e.g., from a reaction between zinc and dilute acid), let the gas escape for a few seconds. This flushes out any air trapped in the apparatus.
2. Collect a small sample: Use a test tube or a similar small container to capture a bit of the gas coming from the source. Avoid using the gas jar itself at this stage.
3. Test with a flame: Hold a burning splint or a lit match near the mouth of the test tube containing the gas sample. Do not insert the splint directly into the tube; just bring it close to the opening.
4. Observe the reaction:
   • If the gas is pure hydrogen, you’ll hear a quiet “pop” sound as it ignites harmlessly.
   • If the gas is mixed with air, you’ll hear a loud bang or sharp explosion, indicating that it’s unsafe to collect.
5. Proceed only if pure: Once you confirm a quiet pop, you can safely collect hydrogen in the gas jar. If there’s a loud bang, continue flushing the apparatus and repeat the test until pure hydrogen is confirmed.

Question 5. 

Why is nitric acid not used in the preparation of hydrogen?

Ans:

Nitric acid is unsuitable for the preparation of hydrogen gas due to its inherently strong oxidizing nature, which fundamentally alters the typical acid-metal reaction. While acids like dilute hydrochloric or sulfuric acid react with reactive metals (e.g., zinc, magnesium) to produce salt and hydrogen gas, nitric acid behaves differently. It is a potent oxidizing agent, meaning it readily accepts electrons. When a metal reacts with nitric acid, the acid itself gets reduced, and the primary products are not hydrogen but various nitrogen oxides (like NO₂ or NO), water, and the corresponding metal salt.

Even with dilute nitric acid and highly reactive metals, any trace of hydrogen initially formed is immediately oxidized by the acid into water. For instance, with magnesium, very dilute nitric acid can yield some hydrogen, but the reaction is neither efficient nor pure, as it is mixed with other reduction products. With most common laboratory metals like zinc or iron, the reaction yields no hydrogen at all; instead, products like nitrous oxide, ammonia, or nitrogen monoxide are formed.

Therefore, for the specific and reliable laboratory preparation of hydrogen gas, non-oxidizing acids such as dilute hydrochloric acid or dilute sulfuric acid are employed. These acids provide the H⁺ ions needed for reduction to H₂ gas without competing oxidative side reactions, ensuring a pure and controllable yield of hydrogen.

Question 6.

Why is hot concentrated sulphuric acid not used in the preparation of hydrogen?

Ans:

Hot concentrated sulphuric acid is avoided in the preparation of hydrogen gas due to its strong oxidizing properties. Unlike dilute acids, which readily release H⁺ ions that react with metals to produce hydrogen, concentrated sulphuric acid behaves differently. When heated, it becomes an even more potent oxidizing agent. Instead of facilitating the displacement of hydrogen, it tends to undergo reduction itself, often leading to the formation of sulfur dioxide or other sulfur compounds rather than hydrogen gas.

For example, with common metals like zinc or iron, hot concentrated sulphuric acid typically produces sulfur dioxide, water, and metal sulfates. This reaction does not yield hydrogen, making it unsuitable for hydrogen generation. Additionally, the process can be hazardous due to the release of toxic gases and the risk of violent reactions. Therefore, for safe and effective hydrogen preparation, dilute acids like hydrochloric or sulphuric acid are preferred, as they provide the necessary H⁺ ions without unwanted oxidation.

Question 7. 

1. Hydrogen is manufactured by the ‘Bosch Process’. Give the equation with conditions. 

2. Hydrogen is manufactured by the ‘Bosch Process’. How can you obtain hydrogen from a mixture of hydrogen and carbon monoxide? 

Ans:

1. The Bosch Process for manufacturing hydrogen involves reacting steam with red-hot coke. The chemical equation and conditions are as follows:

Steam is passed over coke (carbon) heated to a temperature of approximately 1100-1200°C. This primary reaction produces a mixture of carbon monoxide and hydrogen, known as water gas.
Equation: C + H₂O → CO + H₂

This water gas mixture is then mixed with more steam and passed over a catalyst of ferric oxide (Fe₂O₃) at a lower temperature of around 450°C. This secondary step, known as the water-gas shift reaction, converts the carbon monoxide into carbon dioxide, releasing more hydrogen.
Equation: CO + H₂O → CO₂ + H₂

Overall, the process converts coke and steam into hydrogen and carbon dioxide under the stated conditions of high heat and the presence of a catalyst.

2. To obtain pure hydrogen from the mixture of hydrogen and carbon monoxide (water gas) produced in the first step of the Bosch Process, the following method is used:

The gaseous mixture is combined with excess steam and directed over a catalyst, typically iron(III) oxide, at a temperature of about 450°C. In this catalytic step, the carbon monoxide reacts with the steam in a reversible reaction called the water-gas shift reaction, forming carbon dioxide and additional hydrogen. The resulting mixture now contains hydrogen, carbon dioxide, and any unreacted steam. This gas is then passed under pressure into water, where the carbon dioxide dissolves readily, leaving behind relatively pure hydrogen gas. For further purification, the hydrogen can be dried or passed through a solution of an alkali like caustic potash to remove any last traces of acidic carbon dioxide.

Question 8. 

1. Give an equation to express the reaction between: Steam and red hot iron 

2. Give an equation to express the reaction between: Calcium and water 

Ans:

1. When steam passes over red hot iron, the reaction yields iron(II,III) oxide and hydrogen gas. This can be represented by the balanced chemical equation:
   3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g).
2. Calcium metal reacts vigorously with water to form calcium hydroxide and hydrogen gas. The balanced chemical equation for this process is:
   Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g).

Question 9. 

1. A small piece of calcium metal is put into a small trough containing water. There is effervescence and white turbidity is formed. Name the gas formed in the reaction. How would you test the gas? 

2. A small piece of calcium metal is put into a small trough containing water. There is effervescence and white turbidity is formed. Write an equation for the reaction. 

3. A small piece of calcium metal is put into a small trough containing water. There is effervescence and white turbidity is formed. What do you observe when a few drops of red litmus solution are added to the turbid solution.

Ans:

1. Name the gas formed and how to test it
The gas formed is hydrogen.
To test the gas, bring a lighted splint close to the mouth of the container (or collect the gas in a test tube). Hydrogen will burn with a characteristic pop sound.

2. Write an equation for the reaction
The balanced chemical equation for the reaction is:
Ca (s) + 2H₂O (l) → Ca(OH)₂ (aq) + H₂ (g)
The white turbidity is due to the formation of sparingly soluble calcium hydroxide.

3. Observation when red litmus is added to the turbid solution
The turbid solution contains calcium hydroxide, which is alkaline.
When a few drops of red litmus solution are added, the litmus paper turns blue.

Question 10. 

1. Thin strips of magnesium, copper, and iron are taken. Write down what happens when these metals are treated as follows: Heated in presence of air Heated with dil.HCl Added to an aqueous solution of zinc sulphate. 

2. Thin strips of magnesium, copper, and iron are taken.Arrange these metals in descending order of reactivity.

Ans:

1. Observations for Magnesium, Copper, and Iron Strips

Magnesium (Mg)

• Heated in presence of air: Burns with an intense, bright white flame, producing a white, powdery ash (magnesium oxide, MgO).
• Heated with dilute HCl: Reacts vigorously with rapid effervescence of hydrogen gas. The strip dissolves completely, forming a colorless solution (magnesium chloride, MgCl₂).
• Added to aqueous zinc sulphate: No visible reaction. Magnesium is more reactive than zinc, but zinc ions in solution cannot displace magnesium. Therefore, no displacement occurs.

Iron (Fe)

• Heated in presence of air: Glows and forms a black, scale-like layer (tri-iron tetroxide, Fe₃O₄). It does not burn with a flame.
• Heated with dilute HCl: Reacts steadily with a moderate evolution of hydrogen gas bubbles. The strip dissolves to form a light green solution (ferrous chloride, FeCl₂).
• Added to aqueous zinc sulphate: No visible reaction. Iron is less reactive than zinc, so it cannot displace zinc from its salt solution.

Copper (Cu)

• Heated in presence of air: The shiny surface turns black, forming a coating of copper(II) oxide (CuO). No burning occurs.
• Added to aqueous zinc sulphate: No visible reaction. Copper is much less reactive than zinc and cannot displace it.

2. Descending Order of Reactivity
Based on the observed chemical behaviour:
Magnesium > Iron > Copper

Magnesium shows the most vigorous reactions, especially with acid and oxygen. Iron reacts but less aggressively. Copper shows the least reactivity, failing to react with acid or displace other metals.

Question 11. 

1. Choose the correct option:

Hydrogen is evolved by the action of cold dil. HNO3 on

1. Fe
2. Cu
3. Mg
4. Zn

2. Choose the correct option:

Which metal absorbs hydrogen?

1. Al
2. Fe
3. Pd
4. K

3. Choose the correct option:

The composition of the nucleus of deuterium is

1. 1 e- and 1P
2. 1 P and 1 A
3. 1 n and 1 e- 
4. 2 P and 1 e-

4. Choose the correct option:

Elements which show unique nature in the preparation of hydrogen are:

1. Na, K, Li
2. Mg, Ca, Ba
3. Al, Zn, Pb
4. Fe, Cu, Ag

12. 

1. Give a reason for the following: Zinc granules are used in the lab preparation of hydrogen. 

2. Give reason for the following: Purified and dried hydrogen is collected over mercury. 3.Give reason for the following: The end of the thistle funnel should be dipped under acid. 

4.Give reason for the following: Dilute sulphuric acid cannot be replaced by concentrated acid in the preparation of hydrogen.

Ans:

1. Zinc granules are used in the lab preparation of hydrogen because they offer a balanced surface area for reaction, ensuring a steady and manageable release of gas. Their granular form prevents the reaction from becoming too violent, which might occur with powdered zinc, and they react efficiently with dilute acids like sulfuric or hydrochloric acid to produce hydrogen without excessive frothing or splattering.
2. Purified and dried hydrogen is collected over mercury primarily because mercury is inert and does not react with or dissolve hydrogen. This method allows hydrogen to displace mercury without contamination, maintaining gas purity. Additionally, mercury’s high density prevents mixing with air, ensuring the collected hydrogen is free from moisture and other gases, though due to mercury’s toxicity, alternative methods are now often preferred.
3. The end of the thistle funnel should be dipped under the acid to create a liquid seal, which prevents hydrogen gas from escaping back through the funnel. If the funnel is not submerged, gas could leak out, reducing the yield and posing a risk of explosion if hydrogen mixes with air. Dipping it ensures that the acid flows into the reaction vessel while gas is directed through the delivery tube for safe collection.
4. Dilute sulphuric acid cannot be replaced by concentrated acid in the preparation of hydrogen because concentrated sulfuric acid acts as an oxidizing agent. When reacted with zinc, concentrated acid leads to the production of sulfur dioxide rather than hydrogen gas. In contrast, dilute sulfuric acid undergoes a displacement reaction with zinc to yield zinc sulfate and hydrogen, which is the desired outcome for hydrogen generation.

Exercise 6 (C)

Question 1. 

1. Where does hydrogen occur in a free state? 

2.How did the name ‘hydrogen’ originate? 

Ans:

1. Hydrogen in its free, uncombined state is exceptionally rare on Earth due to its highly reactive nature. It is found in minute traces within the upper layers of our atmosphere. However, the vast majority of free hydrogen exists in outer space. It is the most abundant element in the universe, forming the primary composition of stars like our Sun, where it undergoes nuclear fusion, and it is a major component of gas giant planets such as Jupiter and Saturn.

2. The name ‘hydrogen’ has its roots in the Greek language and was given by the French chemist Antoine Lavoisier in the late 18th century. It is derived from the Greek words ‘hydro’, meaning water, and ‘genes’, meaning forming or producer. Therefore, hydrogen literally translates to “water-former.” This name was chosen because when hydrogen gas burns in air, it reacts with oxygen to form water vapour, which was a key experiment in understanding the composition of water.

Question 2. 

1. Hydrogen can be prepared with the help of cold water. Give a reaction of hydrogen with: A monovalent metal 

2. Hydrogen can be prepared with the help of cold water. Give a reaction of hydrogen with: A divalent metal 

Ans:

For a Monovalent Metal:
Sodium (Na), a monovalent alkali metal, reacts vigorously with cold water. The process involves sodium displacing hydrogen from water, resulting in the formation of sodium hydroxide and hydrogen gas. The balanced chemical equation is:
2Na + 2H₂O → 2NaOH + H₂↑

For a Divalent Metal:
Calcium (Ca), a divalent alkaline earth metal, reacts with cold water at a moderate rate. In this reaction, calcium replaces hydrogen in water, producing calcium hydroxide and hydrogen gas. The balanced chemical equation is:
Ca + 2H₂O → Ca(OH)₂ + H₂↑

Question 3. 

1. Which metal is preferred for collecting hydrogen from Cold water? Write the balanced equation. 

2. Which metal is preferred for collecting hydrogen from: Write the balanced equation

Hot water 

3.Which metal is preferred for collecting hydrogen from steam? Write the balanced equation for the case.

Ans:

1. For collecting hydrogen from cold water, the preferred metal is calcium. While metals like sodium and potassium react more vigorously, their reactions are often too violent and unsafe for simple collection. Calcium reacts steadily with cold water, producing hydrogen gas at a manageable rate.
   Ca (s) + 2H₂O (l) → Ca(OH)₂ (aq) + H₂ (g)↑
2. For collecting hydrogen from hot water, the preferred metal is magnesium. Magnesium reacts only very slowly with cold water, but the reaction rate increases sufficiently with hot water to produce a steady stream of hydrogen gas. The balanced chemical equation is:
   Mg (s) + 2H₂O (l) → Mg(OH)₂ (aq) + H₂ (g)↑
   (Note: The water (l) is in the liquid state but is heated.)
3. For collecting hydrogen from steam, the preferred metal is iron, commonly in the form of iron filings or wool. Metals like magnesium and aluminum also react with steam, but the classic laboratory method uses red-hot iron. 

3Fe (s) + 4H₂O (g) → Fe₃O₄ (s) + 4H₂ (g)↑
Here, steam (H₂O in gaseous state) is passed over heated iron to produce iron(II,III) oxide (magnetite) and hydrogen gas.

Question 4. 

1. Hydrogen may be prepared in the laboratory by the action of a metal on an acid.

Which of the metals copper, zinc, magnesium or sodium would be the most suitable? 2. Hydrogen may be prepared in the laboratory by the action of a metal on an acid.

Which of the acids dilute sulphuric, concentrated sulphuric, dilute nitric acid and concentrated nitric acid would you choose? Explain why you would not use the acids you reject. 

3. Hydrogen may be prepared in the laboratory by the action of a metal on an acid.

How would you modify your apparatus to collect dry hydrogen? Which drying agent would you employ for this purpose?  

Ans:

1. Suitable Metal Selection
From the options given, zinc is the most suitable for laboratory preparation of hydrogen by the action of a metal on an acid.

• Copper is unsuitable as it lies below hydrogen in the reactivity series and does not react with dilute acids to release hydrogen.
• Zinc reacts steadily and controllably with dilute acids, producing a consistent hydrogen stream without being excessively violent.
• Magnesium reacts too vigorously, increasing the risk of uncontrolled splashing and rapid gas evolution, making it less manageable for standard laboratory setups.
• Sodium is dangerously reactive with acids, resulting in explosive reactions, and is unsafe for routine laboratory use.

2. Suitable Acid Choice and Reasons for Rejection
The most appropriate acid is dilute sulphuric acid.

• Dilute sulphuric acid reacts cleanly with zinc to produce hydrogen gas without significant byproducts.
• Concentrated sulphuric acid is unsuitable because it acts as an oxidizing agent, producing sulfur dioxide rather than hydrogen when reacting with metals like zinc.
• Dilute nitric acid is oxidizing and yields nitrogen oxides or ammonium ions instead of hydrogen, regardless of the metal used.
• Concentrated nitric acid is a strong oxidizing agent and causes metals to produce nitrogen dioxide, not hydrogen.

3. Modification for Dry Hydrogen Collection and Drying Agent
To collect dry hydrogen, the apparatus must include a drying tube or wash bottle placed between the gas generator and the collection setup. This vessel contains a suitable drying agent through which the gas is bubbled or passed.

• Drying agent: Concentrated sulphuric acid is commonly used, as it effectively removes moisture from hydrogen gas. Alternatively, calcium chloride or silica gel may be employed, but concentrated sulphuric acid is often preferred in school laboratory settings due to its high efficiency and availability.
• Modification: Connect the gas outlet tube from the generator to a drying vessel containing the drying agent, then lead the dried hydrogen to an upturned jar or gas syringe for collection.

Question 5. 

1. Why are the following metals not used in the lab preparation of hydrogen?

calcium 

2. Why are the following metals not used in the lab? preparation of hydrogen? iron 

3. Why are the following metals not used in the lab? preparation of hydrogen?

aluminium  

4. Why are the following metals not used in the lab? preparation of hydrogen?

sodium 

Ans:

1. Calcium
Calcium is not typically used in lab preparation of hydrogen due to its rapid and exothermic reaction with water or acids, which can be difficult to control. When calcium reacts with dilute acids, it produces hydrogen gas but also forms a sparingly soluble calcium salt (e.g., calcium sulfate with sulfuric acid) that can coat the metal surface, slowing down or interrupting the reaction. Additionally, calcium is more expensive and less readily available than commonly used metals like zinc, making it impractical for routine lab use.

2. Iron
Iron is avoided in lab preparation of hydrogen because its reaction with dilute acids (e.g., hydrochloric or sulfuric acid) is relatively slow and inconsistent. Iron often contains carbon or other impurities that can lead to the production of unwanted by-products, affecting the purity of hydrogen. The reaction may also be hindered by the formation of a passive oxide layer on the metal surface. For efficient and quick gas generation, labs prefer more reactive metals like zinc.

3. Aluminium
Aluminium is not used in lab preparation of hydrogen due to its protective oxide layer, which forms spontaneously in air. This layer is inert and prevents the metal from reacting readily with dilute acids unless it is removed or disrupted (e.g., by using concentrated acids or mercury salts). Such conditions are inconvenient or hazardous in a lab setting. Additionally, the reaction with dilute acids is often slow and unreliable, making aluminium unsuitable for routine hydrogen production.

4. Sodium
Sodium is excluded from lab preparation of hydrogen due to its extreme reactivity. It reacts violently with water or acids, producing hydrogen gas with intense heat and splashing, which poses significant safety risks such as fires or explosions. The reaction is too rapid to control easily in standard lab apparatus, and sodium must be handled under strict conditions (e.g., in inert atmospheres). For safe and manageable hydrogen generation, less reactive metals are preferred.

Question 6. 

Based on the reactions of water on metals, arrange the following metals in increasing order of reactivity: iron, sodium, magnesium, zinc, calcium

Ans:

Based on their observed reactions with water, the metals can be arranged from the least reactive to the most reactive as follows:

Iron < Zinc < Magnesium < Calcium < Sodium

This order is determined by the specific conditions required for a reaction to occur and the vigor of the reaction. Iron is the least reactive in this list, as it does not react with cold or hot water; it only reacts with steam (red-hot water vapor) to form iron oxide and hydrogen gas. Zinc is slightly more reactive than iron but still requires steam to produce a reaction, forming zinc oxide.

Magnesium shows a clearer reactivity with water. It reacts very slowly with cold water, but a visible reaction occurs with hot water, producing magnesium hydroxide and hydrogen gas. Calcium is significantly more reactive, reacting steadily with cold water to form calcium hydroxide and hydrogen gas, often with a noticeable fizz. Finally, sodium is the most violently reactive. It reacts explosively with cold water, generating sodium hydroxide and hydrogen gas with so much heat that the hydrogen can ignite. This progressive increase in vigor and decrease in the required reaction temperature establishes the reactivity sequence.

Question 7. 

Hydrogen is evolved when dilute HCl reacts with magnesium, but nothing happens in the case of mercury and silver. Explain.

Ans:

This behavior is explained by understanding where these metals fall on the reactivity series. Magnesium is a highly reactive metal, positioned well above hydrogen in this series. This means it readily loses its electrons. When it comes into contact with the hydrogen ions (H⁺) present in dilute hydrochloric acid, the magnesium atoms donate their electrons to these hydrogen ions. The ions then pair up to form hydrogen gas (H₂), which bubbles away. In essence, magnesium’s strong tendency to form positive ions pushes the reaction forward.

In contrast, both mercury and silver are found below hydrogen in the reactivity series. They are relatively unreactive and have a much weaker tendency to lose electrons compared to magnesium. The hydrogen ions in the acid are actually stronger oxidizing agents than these metals are reducing agents. As a result, mercury and silver cannot supply enough “push” to donate electrons to the hydrogen ions. Therefore, no electron transfer occurs, and hydrogen gas is not liberated. The metals simply do not react with the dilute acid.

A helpful way to picture this is to think of the reactivity series as a ladder, with hydrogen as a key rung in the middle. Any metal placed above hydrogen on this ladder can displace it from a dilute acid. Magnesium sits high on the ladder, so it can easily knock hydrogen off. Mercury and silver, however, are on rungs below hydrogen, so they lack the chemical “height” or driving force to do so, resulting in no observable reaction.

Question 8. 

Steam can react with metal and non-metal to liberate hydrogen. Give the necessary conditions and equations for the same.

Ans:

When steam reacts with metals, the necessary condition is that the metal must be placed above copper in the reactivity series. Highly reactive metals like sodium and potassium react explosively with cold water, so their reaction with steam is not commonly demonstrated. Instead, metals such as magnesium, aluminum, zinc, and iron require the steam to be passed over the heated metal. The general reaction yields the metal oxide and hydrogen gas. For instance, red-hot iron reacts with steam to form iron(II,III) oxide (magnetite) and hydrogen. Aluminum requires an exceptionally high temperature and often a cobalt catalyst to overcome its protective oxide layer, producing aluminum oxide and hydrogen.

The relevant chemical equations for metals are:

1. Magnesium: Mg(s) + H₂O(g) → MgO(s) + H₂(g)
2. Aluminum: 2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g) (with high heat and catalyst)
3. Zinc: Zn(s) + H₂O(g) → ZnO(s) + H₂(g)
4. Iron: 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

In contrast, the reaction of steam with non-metals to liberate hydrogen is far less common. The primary example is with red-hot carbon (coke or coal). When steam is passed over glowing carbon at very high temperatures (around 1000°C), a mixture of carbon monoxide and hydrogen, known as water gas, is produced. This is an industrial process, and while hydrogen is liberated, it is not in pure form but mixed with carbon monoxide. The equation for this endothermic reaction is: C(s) + H₂O(g) → CO(g) + H₂(g). This is a key method for hydrogen production, though it requires significant energy input and careful control of conditions. Most other non-metals do not yield hydrogen from steam under standard laboratory conditions.

Question 9. 

Hydrogen is obtained by displacement from: dilute sulphuric acid dilute hydrochloric acid Write equations using zinc and iron. Why does copper not show similar behavior?

Ans:

Hydrogen gas can be liberated through displacement reactions when certain metals react with dilute acids. Here, the reactions of zinc and iron with dilute sulphuric acid and dilute hydrochloric acid are considered, along with an explanation for copper’s lack of reactivity.

Equations for Zinc:

• With dilute sulphuric acid: Zinc displaces hydrogen to form zinc sulfate and hydrogen gas.
  Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
• With dilute hydrochloric acid: Zinc reacts to produce zinc chloride and hydrogen gas.
  Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Equations for Iron:

• With dilute sulphuric acid: Iron displaces hydrogen, yielding iron(II) sulfate and hydrogen gas.
  Fe(s) + H₂SO₄(aq) → FeSO₄(aq) + H₂(g)
• With dilute hydrochloric acid: Iron reacts to form iron(II) chloride and hydrogen gas.
  Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

Why Copper Does Not Behave Similarly:
Copper does not displace hydrogen from dilute acids like sulphuric or hydrochloric acid because it is less reactive than hydrogen. In the reactivity series, copper lies below hydrogen, meaning it cannot reduce hydrogen ions (H⁺) to hydrogen gas. Consequently, copper remains unreactive in such acids, and no displacement occurs. This inertness is due to copper’s higher reduction potential compared to hydrogen, making it unable to undergo oxidation easily in these acidic conditions.

Question 10. 

1. Account for the following fact: Though lead is above hydrogen in the activity series, it does not react with dilute hydrochloric acid or dilutes sulphuric acid. 

2. Give reason for the following: Potassium and sodium are not used for reaction with dilute hydrochloric acid or dilute sulphuric acid in the laboratory preparation of hydrogen. 

Ans:

1. Account for the following fact: Though lead is above hydrogen in the activity series, it does not react with dilute hydrochloric acid or dilute sulphuric acid.

While lead is indeed placed above hydrogen in the activity series, indicating it should theoretically displace hydrogen from dilute acids, this reaction is not observed in practice due to a phenomenon called passivation. When lead comes into contact with acids like HCl or H₂SO₄, an immediate and insoluble layer of lead chloride (PbCl₂) or lead sulphate (PbSO₄) forms on the metal’s surface. This layer is tightly adherent and non-porous, acting as a protective shield. It physically prevents the fresh acid from reaching the underlying metallic lead, thereby stopping the reaction almost as soon as it begins. Consequently, despite its favorable position in the series, the reaction becomes practically insignificant, and no appreciable amount of hydrogen gas is liberated.

2. Give reason for the following: Potassium and sodium are not used for reaction with dilute hydrochloric acid or dilute sulphuric acid in the laboratory preparation of hydrogen.

Potassium and sodium are extremely reactive metals, placed at the very top of the activity series. Their reaction with dilute acids is violently exothermic and far too dangerous for a standard laboratory setting. When these metals contact acids, they react with explosive violence, producing hydrogen gas with such intense heat that the hydrogen often ignites immediately, posing a severe fire and explosion hazard. Furthermore, the reaction is extremely difficult to control or moderate. Therefore, for a safe and manageable laboratory preparation of hydrogen, moderately reactive metals like zinc or iron are preferred, as they react steadily and controllably with dilute acids without such uncontrollable and hazardous outbursts.

Question 11. 

Name two alkalies that can displace hydrogen. Give balanced equations for the same. Why are the metals you have used considered to have a unique nature?

Ans:

Two alkalies that can displace hydrogen are sodium hydroxide (NaOH) and potassium hydroxide (KOH), which are formed when their respective metals react with water. The alkali metals themselves—sodium (Na) and potassium (K)—are the actual substances that displace hydrogen from water.

The balanced chemical equations for these reactions are:

1. 2Na (s) + 2H₂O (l) → 2NaOH (aq) + H₂ (g)
2. 2K (s) + 2H₂O (l) → 2KOH (aq) + H₂ (g)

Sodium and potassium are considered to have a unique nature primarily because they are members of the alkali metals group (Group 1) in the periodic table. Their uniqueness stems from having just a single electron in their outermost valence shell. This solitary electron is held very loosely due to the large atomic size, making it exceptionally easy to lose. This results in an extremely vigorous and exothermic reaction with cold water, unlike most other metals. Furthermore, they are so soft they can be cut with a knife, have remarkably low densities (lithium, sodium, and potassium float on water), and are never found in nature in their pure, metallic state due to their high reactivity. This combination of intense chemical reactivity and distinctive physical properties sets them apart from typical metals.

Question 12. 

1. Complete and balance the following reaction. Na + H2O →_____________ +___________ 2. Complete and balance the following reaction. Ca + H2O →_____________ +___________ 3. Complete and balance the following reaction. Mg + H2O →_____________ +___________ 4. Complete and balance the following reaction. Zn + H2O →_____________ +___________ 5. Complete and balance the following reaction. Fe + H2O →_____________ +___________ 6. Complete and balance the following reaction. Zn + HCl →_____________ +___________ 7. Complete and balance the following reaction. Al + H2SO4 →_____________ +___________ 

8. Complete and balance the following reaction. Fe + HCl →_____________ +___________ 9. Complete and balance the following reaction. Zn + NaOH →_____________ +___________ 

10. Complete and balance the following reaction. Al + KOH + H2O→_____________ +___________

Ans:

1. 2Na + 2H₂O → 2NaOH + H₂
2. Ca + 2H₂O → Ca(OH)₂ + H₂
3. Mg + H₂O → MgO + H₂ (requires steam/heat)
4. Zn + H₂O → ZnO + H₂ (requires steam/heat)
5. 3Fe + 4H₂O → Fe₃O₄ + 4H₂ (requires steam/heat)
6. Zn + 2HCl → ZnCl₂ + H₂
7. 2Al + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂
8. Fe + 2HCl → FeCl₂ + H₂
9. Zn + 2NaOH → Na₂ZnO₂ + H₂
10. 2Al + 2KOH + 6H₂O → 2K[Al(OH)₄] + 3H₂

Question 13. 

1. Metal in the powdered form reacts very slowly with boiling water, but it decomposes in steam. Name the metal. 

2. Write a balanced equation for the reaction named metal with (i) boiling water (ii) steam.

Ans:

1. The metal is magnesium.
2. Balanced equations:
   (i) Reaction with boiling water:
   Mg + 2H₂O → Mg(OH)₂ + H₂
   (ii) Reaction with steam:
   Mg + H₂O → MgO + H₂

Question 14. 

What do you observe when hydrogen gas is passed through a soap solution?

Ans:

When hydrogen gas is directed into a soap solution, you immediately notice the formation of bubbles within the liquid. These bubbles are filled with hydrogen gas and, due to the soap film stabilizing them, they detach and float into the air. Since hydrogen is much lighter than air, the bubbles rise upward remarkably fast, often appearing to zoom away if released in still conditions.

If you bring a lighted splint or a flame close to these ascending bubbles, they ignite and burn with a faint blue flame, producing a distinct popping sound. This pop occurs because the hydrogen inside mixes rapidly with oxygen in the air and combusts. The overall effect is a swift display of rising, flammable bubbles that vanish with a sharp pop when ignited, demonstrating hydrogen’s low density and explosive nature in air.

Question 15. 

1. Under what conditions can hydrogen be made to combine with nitrogen? Name the products and write the equation for the reaction. 

2. Under what conditions can hydrogen be made to combine with chlorine? Name the products and write the equation for the reaction. 

3.Under under what conditions can hydrogen be made to combine with sulphur? Name the products and write the equation for the reaction. 

4. Under what conditions can hydrogen be made to combine with oxygen? Name the products and write the equation for the reaction.

Ans:

1. Hydrogen combining with Nitrogen
Hydrogen combines with nitrogen under conditions of high pressure (around 200 atmospheres), a moderately high temperature (approximately 450°C), and in the presence of a finely divided iron catalyst, often promoted with molybdenum. This process is known as the Haber process. The product formed is ammonia gas. 

N₂ + 3H₂ ⇌ 2NH₃

2. Hydrogen combining with Chlorine
Hydrogen combines with chlorine either in the presence of direct ultraviolet light (sunlight) or when ignited. The reaction is initiated by the energy from light or heat, which breaks the chlorine molecules into reactive atoms. The product is hydrogen chloride gas. The chemical equation for this combination is:
H₂ + Cl₂ → 2HCl

3. Hydrogen combining with Sulphur
Hydrogen combines with sulphur upon direct heating. Sulphur must be heated to its vapour state to react with hydrogen gas passed through it. The product of this reaction is hydrogen sulphide gas, which has a characteristic rotten egg smell. The chemical equation is:
H₂ + S → H₂S

4. Hydrogen combining with Oxygen
Hydrogen combines with oxygen upon ignition, meaning it requires an initial spark or flame to begin the reaction. The combination is highly exothermic and explosive when the mixture is in a 2:1 volume ratio. The product formed is water, usually as steam under the reaction conditions. The chemical equation is:
2H₂ + O₂ → 2H₂O

Question 16. 

When hydrogen is passed over a black solid compound A, the products are a ‘colorless liquid’ and a ‘reddish-brown metal B’.

Substance B is divided into two parts each placed in separate test tubes.

Dilute HCl is added to one part of substance B and dilute HNO3 to the other.

1. Name the substances A and B.
2. Give two tests for the colourless liquid formed in the experiment.
3. What happens to substance A when it reacts with hydrogen? Give reasons for your answer.
4. Write an equation for the reaction between hydrogen and substance A.
5. Is there any reaction between substance B and dilute hydrochloric acid? Give reasons for your answer.

Ans:

Based on the given information, the substances can be identified through logical deduction. The black solid compound A is copper(II) oxide (CuO). When hydrogen gas is passed over hot CuO, it reduces the compound. The reddish-brown metal B formed is copper (Cu), and the colorless liquid produced is water (H₂O). This is a standard reduction reaction where hydrogen removes oxygen from the metal oxide.

Two reliable tests for the colorless liquid, water, are as follows. First, it turns white anhydrous copper sulphate powder into a bright blue crystalline hydrate, indicating the presence of water. Second, it changes blue cobalt chloride paper to a distinct pink colour, which is a characteristic and sensitive test for moisture.

When substance A (CuO) reacts with hydrogen, it undergoes a chemical reduction. In this reaction, hydrogen acts as a reducing agent by removing the oxygen from copper(II) oxide, converting it into metallic copper and water. The reason for this reaction is that hydrogen has a higher affinity for oxygen under the applied heat, forming a more stable product (water) and leaving the pure metal behind.
H₂ + CuO → Cu + H₂O.
Regarding the reaction of substance B (copper) with dilute hydrochloric acid, there is no reaction. The reason is that copper is placed below hydrogen in the metal reactivity series. Metals below hydrogen cannot displace hydrogen from dilute acids. Since copper is less reactive than hydrogen, it does not react with dilute HCl, and no gas is evolved. In contrast, dilute nitric acid, being a strong oxidizing agent, does react with copper, but that is a separate observation not asked for in this specific part of the question.

Exercise 6 (D)

Question 1. 

Describe briefly the ionic concept of oxidation and reduction. Give an equation to illustrate.

Ans:

The ionic concept of oxidation and reduction, often termed redox reactions, moves beyond the addition or loss of oxygen. Instead, it focuses entirely on the transfer of electrons between reacting species. From this perspective, oxidation is defined as the loss of electrons from an atom or ion. 

A classic equation that perfectly illustrates this concept is the reaction between zinc metal and copper(II) sulfate solution:
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Written as an ionic equation to highlight the electron transfer, it becomes:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

In this reaction, the zinc atom (Zn) loses two electrons to form a zinc ion (Zn²⁺). Therefore, zinc is oxidized. The copper ion (Cu²⁺) gains those two electrons to become a neutral copper atom (Cu). Thus, the copper ion is reduced. The zinc acts as a reducing agent by causing reduction, and the copper ion acts as an oxidizing agent by causing oxidation. This direct exchange of electrons is the cornerstone of the ionic redox theory.

Question 2. 

Is it essential that oxidation and reduction must occur side by side in a chemical reaction? Explain

Ans:

In any chemical reaction involving electron transfer, oxidation and reduction are inseparable processes. These processes must occur simultaneously because electrons cannot exist in isolation; they must be transferred from one species to another. If a substance loses electrons (oxidation), another substance must accept those electrons (reduction) to maintain charge balance and adhere to the principle of conservation of charge.

For instance, when zinc metal reacts with copper sulfate, zinc atoms lose electrons to form zinc ions (oxidation), and copper ions gain those electrons to form copper metal (reduction). Even in electrochemical cells, where half-reactions are physically separated, oxidation at the anode and reduction at the cathode are connected through a circuit, ensuring simultaneous occurrence.

Thus, in redox reactions, oxidation and reduction are inherently coupled. While non-redox reactions (like acid-base or precipitation reactions) do not involve electron transfer, for any reaction classified as redox, it is essential that oxidation and reduction proceed side by side.

Question 3. 

1. State, giving reason, whether the substances printed in bold letters have been oxidized or reduced. PbO + CO → Pb + CO2 

2. State, giving reason, whether the substances printed in bold letters have been oxidized or reduced. Mg + 2HCl → MgCl2 + H2 

3. State, giving reason, whether the substances printed in bold letters have been oxidized or reduced. H2S + Cl2→ 2HCl + S 

4. State, giving reason, whether the substances printed in bold letters have been oxidized or reduced. Cl2 + H2S → 2HCl + S 

Ans:

1. In the reaction PbO + CO → Pb + CO₂, the substance lead (Pb) has been reduced. This is because its oxidation state decreases from +2 in lead(II) oxide (PbO) to 0 in the elemental form (Pb). Reduction involves a gain of electrons or a decrease in oxidation state, which is clearly observed here as the lead ion accepts electrons.
2. In the reaction Mg + 2HCl → MgCl₂ + H₂, the substance magnesium (Mg)—which forms MgCl₂—has been oxidized. The oxidation state of magnesium increases from 0 in the elemental metal to +2 in magnesium chloride. Oxidation is defined by a loss of electrons or an increase in oxidation state, which occurs as magnesium loses two electrons to form Mg²⁺ ions.
3. For the reaction H₂S + Cl₂ → 2HCl + S, the substance sulfur (S) in H₂S has been oxidized. The oxidation state of sulfur rises from -2 in hydrogen sulfide to 0 in the elemental sulfur that is produced. This increase in oxidation number signifies the loss of electrons, which is the defining characteristic of an oxidation process.
4. In the reaction Cl₂ + H₂S → 2HCl + S, the substance chlorine (Cl) in Cl₂ has been reduced. The oxidation state of chlorine decreases from 0 in the diatomic molecule to -1 in hydrogen chloride (HCl). This decrease indicates a gain of electrons, fulfilling the criteria for a reduction reaction.

Question 4. 

1. State whether the following conversion is oxidation or reduction: PbO2 + SO2→ PbSO4  

2. State whether the following conversion is oxidation or reduction: Cu2+ + 2 e-→ Cu 3. State whether the following conversion is oxidation or reduction: K → K+ + e-  

4. State whether the following conversion is oxidation or reduction: 2Cl- – e-→ Cl2 

Ans:

1. In the conversion PbO₂ + SO₂ → PbSO₄, the oxidation state of lead decreases from +4 in PbO₂ to +2 in PbSO₄, indicating a gain of electrons or reduction. Therefore, this conversion is reduction.
2. In the conversion Cu²⁺ + 2 e⁻ → Cu, copper ions gain electrons to form copper metal, which is a decrease in oxidation state. Thus, this conversion is reduction.
3. In the conversion K → K⁺ + e⁻, potassium loses an electron to form a potassium ion, which is an increase in oxidation state. Hence, this conversion is oxidation.
4. In the conversion 2Cl⁻ – e⁻ → Cl₂, chloride ions lose electrons to form chlorine gas, as represented by the implied loss of electrons (typically written as 2Cl⁻ → Cl₂ + 2e⁻). This increase in oxidation state classifies the conversion as oxidation.

Question 5. 

In the following reaction: A+ + B → A + B+. Write half-reactions for this reaction and name:

1. oxidizing agent
2. substance oxidized
3. reducing agent 

Ans:

Based on the reaction A⁺ + B → A + B⁺, the analysis of electron transfer is as follows.

The reaction can be split into two half-reactions. The first shows B losing an electron: B → B⁺ + e⁻. This is an oxidation half-reaction because electron loss occurs. The second shows A⁺ gaining that electron: A⁺ + e⁻ → A. This is a reduction half-reaction because electron gain occurs.

Identifying the key components, the oxidizing agent is A⁺. It accepts electrons from another substance, thereby causing oxidation and becoming reduced itself to A. The substance oxidized is B, as it directly loses electrons. Consequently, because B provides the electrons that reduce A⁺, it also acts as the reducing agent.

Question 6. 

1. Divide the following reactions into oxidation and reduction half-reaction: Zn + Pb2+→ Pb + Zn 2+ 

2. Divide the following reactions into oxidation and reduction half-reaction: Zn + Cu2+ → Cu + Zn 2+ 

3. Divide the following reactions into oxidation and reduction half-reaction: Cl2 + 2Br- → Br2 + 2Cl-

Ans:

1. For the reaction: Zn + Pb²⁺ → Pb + Zn²⁺

In this displacement reaction, zinc atoms lose electrons while lead ions gain them. The oxidation half-reaction shows zinc undergoing oxidation: Zn → Zn²⁺ + 2e⁻. Here, the neutral zinc atom loses two electrons to form a zinc ion. Concurrently, the reduction half-reaction is Pb²⁺ + 2e⁻ → Pb. In this step, the lead(II) ion gains the two electrons lost by zinc, forming a neutral lead atom. The transfer of two electrons from zinc to lead ions is what drives the reaction forward.

2. For the reaction: Zn + Cu²⁺ → Cu + Zn²⁺

This is another classic metal displacement reaction. The oxidation half-reaction involves zinc losing electrons: Zn → Zn²⁺ + 2e⁻. The zinc metal is oxidized as its oxidation state increases from 0 to +2. The complementary reduction half-reaction is Cu²⁺ + 2e⁻ → Cu. The copper(II) ions in solution accept the two electrons provided by zinc, reducing their oxidation state from +2 to 0 and depositing as solid copper metal. The flow of electrons from Zn to Cu²⁺ is the key process.

3. For the reaction: Cl₂ + 2Br⁻ → Br₂ + 2Cl⁻

This reaction involves the displacement of one halogen by another. The reduction half-reaction is Cl₂ + 2e⁻ → 2Cl⁻. Each chlorine atom in the Cl₂ molecule gains one electron, reducing its oxidation state from 0 to -1. The oxidation half-reaction is 2Br⁻ → Br₂ + 2e⁻. Here, two bromide ions each lose one electron, resulting in the formation of a bromine molecule and increasing the oxidation state of bromine from -1 to 0. The electrons released by the bromide ions are consumed by the chlorine molecules.

Question 7. 

1. Write the equation in the ionic form CuSO4(aq)  + Fe(s)→ FeSO4(aq) + Cu(s) 

2. Divide the above equation into oxidation and reduction half-reactions. 

Ans:

1. The given equation is CuSO₄(aq) + Fe(s) → FeSO₄(aq) + Cu(s). In ionic form, the soluble ionic compounds dissociate into their ions. Copper(II) sulfate dissociates into Cu²⁺ and SO₄²⁻ ions, and iron(II) sulfate dissociates into Fe²⁺ and SO₄²⁻ ions. Thus, the full ionic equation is:
   Cu²⁺(aq) + SO₄²⁻(aq) + Fe(s) → Fe²⁺(aq) + SO₄²⁻(aq) + Cu(s)
   The sulfate ion (SO₄²⁻) is a spectator ion present on both sides. Canceling it gives the net ionic equation:
   Cu²⁺(aq) + Fe(s) → Fe²⁺(aq) + Cu(s)
2. To divide the net ionic equation into half-reactions, identify the oxidation and reduction processes. Iron metal (Fe) loses electrons and is oxidized to Fe²⁺, while copper ions (Cu²⁺) gain electrons and are reduced to copper metal (Cu).
   • Oxidation half-reaction: Fe(s) → Fe²⁺(aq) + 2e⁻
   • Reduction half-reaction: Cu²⁺(aq) + 2e⁻ → Cu(s)
3. These half-reactions are balanced for mass and charge, showing the transfer of two electrons.

Question 8. 

1. Give a reason for the following. Hydrogen is collected by the downward displacement of water and not of air, even though it is lighter than air. 

2. Give reason: A candle brought near the mouth of a jar containing hydrogen gas starts burning but is extinguished when pushed inside the jar. 

3. Give reason: Apparatus for laboratory preparation of hydrogen should be airtight and away from a naked flame.

Ans:

1. Hydrogen is collected via downward displacement of water primarily for safety, not simply because of its density. Although hydrogen is lighter than air and could be collected over air, this method would allow it to mix with atmospheric oxygen inside the gas jar. Since a mixture of hydrogen and air is highly explosive, even a small spark could trigger a dangerous reaction. Collecting it over water isolates the hydrogen from oxygen, providing a much safer method of preparation and collection, even if it is slightly less convenient.

2. When a burning candle is brought near the mouth of the hydrogen-filled jar, the hydrogen at the mouth mixes with the oxygen in the surrounding air. This mixture supports combustion, so the hydrogen ignites and burns. However, when the candle is pushed inside the jar, it enters an atmosphere of almost pure hydrogen. Combustion cannot occur in the absence of oxygen, so the flame is immediately extinguished. This demonstrates that hydrogen itself is not a supporter of combustion; it only burns in the presence of an oxidizer like oxygen.

3. The apparatus for preparing hydrogen must be airtight to prevent any leakage of the gas into the laboratory atmosphere. Since hydrogen is highly flammable and forms explosive mixtures with air, even a minor leak could create a dangerous buildup of gas. Keeping the setup away from any naked flame is a direct safety precaution for the same reason. A single spark could ignite any escaped hydrogen, causing a fire or explosion. These strict measures are essential to prevent accidents during the experiment.

Question 9. 

1. Select the odd one out and justify your answer. Zn, Fe, Mg, and Na 

2. Select the odd one out and justify your answer. SO2, H2S, NH3, and CO3 

3. Select the odd one out and justify your answer. Fe, Zn, Cu and Mg 

4. Select the odd one out and justify your answer. Fe, Pb, Al and Zn

Ans:

1. Zn, Fe, Mg, Na
Odd one out: Na
Justification: Sodium (Na) is the only metal in the list that reacts violently with cold water, producing hydrogen gas and a strong alkali. Zinc (Zn), iron (Fe), and magnesium (Mg) do not react with cold water; they react only with steam or hot water.

2. SO₂, H₂S, NH₃, CO₃
Odd one out: CO₃
Justification: CO₃ (carbonate ion) is a polyatomic ion, while SO₂ (sulfur dioxide), H₂S (hydrogen sulfide), and NH₃ (ammonia) are all discrete molecular compounds under standard conditions. Additionally, CO₃ carries a -2 charge, unlike the neutral molecules.

3. Fe, Zn, Cu, Mg
Odd one out: Cu
Justification: Copper (Cu) is the only metal listed that does not react with dilute hydrochloric acid to produce hydrogen gas. Iron (Fe), zinc (Zn), and magnesium (Mg) readily displace hydrogen from dilute acids.

4. Fe, Pb, Al, Zn
Odd one out: Al
Justification: Aluminum (Al) is amphoteric, meaning it reacts with both acids and strong bases to produce hydrogen gas. Iron (Fe), lead (Pb), and zinc (Zn) are not amphoteric in the same way—they react with acids but not with alkalis under normal conditions.

Question 10. 

1. Helium is preferred to hydrogen for filling balloons because it is:

1. lighter than air
2. almost as light as hydrogen
3. non-combustible
4. Inflammable

2. Reacting with water, an active metal produces

1. oxygen
2. nitric acid
3. a base
4. none of these

3. A metal oxide that is reduced by hydrogen is

1. Al2O3
2. CuO
3. CaO
4. Na2O

4. Which of the following statements about hydrogen is incorrect?

1. It is an inflammable gas
2. It is the lightest gas.
3. It is not easily liquefied
4. It is a strong oxidizing agent.

5. For the reaction PbO + H2→ Pb + H2O, which of the following statements is wrong?

1. H2 is the reducing agent.
2. PbO is the oxidizing agent.
3. PbO is oxidized to Pb.
4. H2 is oxidized to H2O.

6. Which metal gives hydrogen with all of the following: water, acids, alkalis?

1. Fe
2. Zn
3. Mg
4. Pb

7. Which of the following metals does not give hydrogen with acids?

• Iron
• Copper
• Lead
• Zinc

Question 11. 

1. Choose terms from the options given in brackets to complete this sentence.

When CuO reacts with hydrogen,………………… is reduced and ……………….is oxidized to …………………. 

(CuO, H2, Cu, H2O) 

2. Choose terms from the options given in brackets to complete this sentence.

Hydrogen is ………………… soluble in water.

(sparingly, highly, moderately) 

3. Choose terms from the options given in brackets to complete this sentence.

Metals like …………….. , ……………… and ……………… give H2 with steam.

(iron, magnesium, aluminium, sodium, calcium) 

4. Choose terms from the options given in brackets to complete this sentence.

Sodium ………………. reacts smoothly with cold water.

(metal, amalgam, in the molten state) 

5. Choose terms from the options given in brackets to complete this sentence.

A metal …………….. hydrogen in the activity series gives hydrogen with …………… acid or …………… acid.

(above, below, dilute hydrochloric, concentrated hydrochloric, dilute sulphuric).

Ans:

1. Choose terms from the options given in brackets to complete this sentence.

When CuO reacts with hydrogen , CuO is reduced and H2 is oxidized to H2O.

2. Choose terms from the options given in brackets to complete this sentence.

Hydrogen is sparingly soluble in water.

3. Choose terms from the options given in brackets to complete this sentence.

Metals like iron , magnesium and aluminium give H2 with steam.

4. Choose terms from the options given in brackets to complete this sentence.

Sodium amalgam reacts smoothly with cold water.

5. Choose terms from the options given in brackets to complete this sentence.

A metal above hydrogen in the activity series gives hydrogen with dilute hydrochloric acid or dilute sulphuric acid.

Question 12. 

1. Correct the following statement: Hydrogen is separated from CO by passing the mixture through caustic potash solution.

2. Correct the following statement: All metals react with acids to give hydrogen. 

3. Correct the following statement: Hydrogen is dried by passing it through conc. H2SO4. 

4. Correct the following statement: Very dilute nitric acid reacts with iron to produce hydrogen. 

5. Correct the following statement: Conc. H2SO4 reacts with zinc to liberate hydrogen. 

Ans:

1. The given statement is incorrect. Carbon monoxide (CO) is not absorbed by caustic potash (potassium hydroxide). The correct process is that hydrogen is separated from carbon dioxide (CO₂) by passing the mixture through caustic potash solution. Caustic potash readily absorbs acidic CO₂, leaving hydrogen gas unaffected.
2. This statement is inaccurate. It is not true that all metals react with acids to give hydrogen. The correct generalization is that only metals above hydrogen in the reactivity series react with dilute acids to liberate hydrogen gas. Metals like copper, silver, and gold, which are below hydrogen, do not displace it from acids.
3. The statement “Hydrogen is dried by passing it through conc. H₂SO₄” is actually correct. Concentrated sulphuric acid is an excellent drying agent for hydrogen because it is highly hygroscopic and does not react chemically with hydrogen gas under normal conditions.
4. The statement is false. Nitric acid, even when dilute, is a strong oxidizing agent. Very dilute nitric acid reacts with iron, but it does not produce hydrogen. The oxidizing nature of nitric acid leads to the formation of nitrogen oxides (like nitric oxide) instead of hydrogen gas.
5. This statement is incorrect. Concentrated sulphuric acid reacts with zinc, but it does not liberate hydrogen gas. The correct observation is that dilute sulphuric acid reacts with zinc to liberate hydrogen. Concentrated sulphuric acid, being an oxidizing agent, reacts with zinc to produce sulphur dioxide gas.

Question 13. 

1. Name: an oxidizing agent that does not contain oxygen. 

2. Name: a substance that oxidizes concentrated HCl to chlorine. 

3. Name: a substance that will reduce aqueous Iron(III)ions to iron(II)ions. 

4. Name: a liquid that is an oxidizing agent as well as a reducing agent. 

5. Name: a gas that is oxidizing as well as a reducing agent. 

6. Name: a solid that is an oxidizing agent.

Ans:

1. Chlorine
2. Manganese dioxide
3. Sulfur dioxide
4. Hydrogen peroxide
5. Nitric oxide
6. Potassium permanganate