Frank Brothers Solutions for ICSE Class 10 Chemistry Chapter 7b Metallurgy

Chapter 7. Metallurgy

Solution 9:
(i) Sodium is more reactive than aluminium because sodium can lose electron easily to form sodium ions which can react with other substances readily while aluminium cannot lose electrons so easily as compared to sodium hence it is less reactive.
(ii) Gallium and caesium metals melt below 30°C so if room temperature is around 30°C they may also be in liquid state along with mercury.
(iii) Metals always combine with electro valency because they always have a tendency to form positive ions by losing electrons.
(iv) Metal ions are always positively charged because these are formed by loss of electrons from the metal atoms as:
M → M⁺ + e⁻
Hence the resulting metal ions have less number of electrons and more number of protons so overall these are positively charged.
(v) Since metal ions are always positively charged so they are attracted towards negatively charged cathode (negatively charged) in the electrolytic cell and are discharged or get reduced there as:
At cathode: M⁺ + e⁻ → M
(vi) Metals generally do not form hydrides because hydrogen itself behaves like metals and has the tendency to lose its electron.
(vii) Metal atoms have largest atomic sizes among all the elements. Hence due to large size they cannot combine with other similar atoms to form diatomic or triatomic molecules so metals are monoatomic.
(viii) Sodium and potassium cannot be kept exposed to air get oxidized to form their respective oxides as:
4K + O₂ → 2K₂O
4Na + O₂ → 2Na₂O
(ix) Sodium and potassium should not be treated with acids because they react with acids with explosive violence and liberates hydrogen gas:
2K + 2H₂SO₄ → 2K₂SO₄ + H₂
2Na + 2HCl → 2NaCl + H₂
(x) Hydrogen is not a metal but it has been assigned a place in the reactivity series of metals because like metals, hydrogen also loses electrons and form positive ions, H⁺.

📝 Teacher's Note: Use simple experiments like dropping sodium in water to show students why reactive metals need special storage. This makes the abstract concept of reactivity very concrete and memorable.

🎯 Exam Tip: For electron loss equations, always balance the charges on both sides - this is where most students lose marks in metallurgy questions.

Solution 10:
(i) 2Na + 2H₂O → 2NaOH + H₂
(ii) Mg + H₂O → MgO + H₂
(iii) Zn + H₂SO₄ → ZnSO₄ + H₂
(iv) 2HgO → 2Hg + O₂
(v) 2Ag₂O → 4Ag + O₂
(vi) Cu(OH)₂ → CuO + H₂O
(vii) Ca(OH)₂ → CaO + H₂O
(viii) PbCO₃ → PbO + CO₂
(ix) 2KNO₃ (s) → 2KNO₂ (s) + O₂ (g)
(x) 2Cu(NO₃)₂ (s) → 2CuO (s) + 4NO₂ (g) + O₂ (g)
(xi) Hg(NO₃)₂ (s) → Hg (l) + 2NO₂ (g) + O₂ (g)
(xii) 2Ag(NO₃)₂ (s) → 2Ag (s) + 2NO₂ (g) + O₂ (g)

📝 Teacher's Note: Practice balancing these equations step by step with students. Start with atoms, then balance charges, and finally check that mass is conserved on both sides.

🎯 Exam Tip: Always write the physical states (s), (l), (g) in decomposition reactions - examiners specifically look for this detail and award extra marks.

Solution 11:
(i) Reduction with carbon.
The oxide of metal is lead oxide (PbO).
(ii) Electrolytic reduction
The oxide of metal is disodium oxide Na₂O.
2NaCl → 2Na⁺ + Cl⁻
2Na⁺ + 2e⁻ → 2Na (at cathode)
Cl⁻ - 2e⁻ → Cl₂(anode)
(iii) Reduction with heat alone
The oxide of metal is mercuric oxide HgO
2HgO → 2Hg + O₂

📝 Teacher's Note: Explain that different metals need different extraction methods based on their position in the reactivity series - this helps students understand why we can't use one method for all metals.

🎯 Exam Tip: Mention the specific oxide compound for each method - don't just write "metal oxide" as this is too vague and loses marks.

Solution 12:
1. Aqueous solution of sodium chloride is not used for electrolytic reduction of sodium metal because sodium metal formed at cathode after discharge of sodium ions (at cathode) will react with water to form alkali NaOH.
2. For the reduction of a metal oxide a reducing agent other than carbon is carbon monoxide (CO).

📝 Teacher's Note: Demonstrate why molten salt is used instead of aqueous solution by showing students what happens when sodium reacts with water - the vigorous reaction explains everything!

🎯 Exam Tip: Always mention that the sodium would react with WATER specifically, not just "react with the solution" - this precision in language is what examiners look for.

Solution 13:
Corrosion is a process involving the conversion of a metal into an undesirable compound (usually oxide) on exposure to atmospheric conditions i.e. moisture and oxygen. In case of iron, corrosion is known as rusting.
Chemically rust is hydrated ferric oxide Fe₂O₃.xH₂O.
Reaction of corrosion: 4Fe + 3O₂ + xH₂O → 2Fe₂O₃.xH₂O

📝 Teacher's Note: Show students real examples of rusted iron objects and fresh iron to make the concept visual. Explain how water and oxygen both are needed for rusting to occur.

🎯 Exam Tip: Always include the chemical formula for rust (Fe₂O₃.xH₂O) and mention that it's specifically "hydrated" ferric oxide - this shows deeper understanding.

Solution 1:
1. Gold and Platinum.
2. Charge.
3. Gangue.
4. Flux.
5. Calcination.
6. Roasting.
7. Iron pyrites.
8. Bauxite.
9. Cryolite, aluminium fluoride, Calcium fluoride.
10. Cathode: inner lining of gas-carbon of the electrolytic cell.
Anode: Thick carbon rods dipping into the fused electrolytes.
11. Thermite welding.
12. Copper and silver.
13. Aluminium, Iron.
14. platinum and gold
15. sodium and potassium

📝 Teacher's Note: These are key metallurgy terms that students must memorize. Create flashcards or word association games to help students remember these definitions and examples.

🎯 Exam Tip: Learn the exact spelling of technical terms like "cryolite" and "bauxite" - spelling mistakes in technical terms can cost marks even if the concept is understood.

Solution 2:
1. Zinc is used in galvanization and dry cells because zinc coating protects the iron from corrosion as it is more electropositive than iron hence it would be attacked first.
2. Nitric acid can be stored in aluminium containers because it do not attack aluminium. It renders aluminium passive due to the formation of an oxide film on surface of aluminum.
3. Aluminium oxide cannot be reduced by carbon because it is comparatively high in electrochemical series hence more reactive than carbon.
4. A neutral gas other than oxygen is formed at the anode during electrolysis of fused alumina because the oxygen gas formed at the anode oxidizes the carbon of the anode to carbon dioxide.
5. Extraction of aluminium was very difficult in the beginning because it was very expensive.
6. Carbon anodes are used in the electrolytic extraction of aluminium because carbon in the form of graphite is a good conductor of electricity.
7. Galvanized metal ions should not be used for storing food as food acids may react with the zinc coating and cause food poisoning.

📝 Teacher's Note: Connect these concepts to real life - show galvanized iron sheets, aluminium foil, and discuss why certain metals are chosen for specific uses. This makes chemistry practical and relevant.

🎯 Exam Tip: For galvanization questions, always mention that zinc is "more electropositive" than iron - this is the key scientific explanation that examiners want to see.

Solution 3:
1. Mineral: The naturally occurring compounds of metals which are generally mixed with earthy such as soil, sand, limestone and rocks are known as minerals.
2. Ore: Those minerals from which a metal can be extracted profitably are called ores.
3. Gangue: The rocky impurities like (SiO₂) present in an ore are called gangue.
4. Charge: The mixture of materials fed into a furnace to extract a metal is called charge.
5. Flux: The substance added to get rid of gangue in the extraction of metal is called flux.
6. Slag: The product obtained by the combination of gangue with flux is called slag.

📝 Teacher's Note: Emphasize the relationship between these terms - mineral → ore → extraction process involving flux, charge, and slag formation. Draw a flowchart to show these connections.

🎯 Exam Tip: Remember that ores are always profitable to extract from - this economic aspect is what distinguishes ores from ordinary minerals and examiners often test this distinction.

Solution 4:
(i) The set of processes used to remove as much of gangue as possible, is known as concentration of ores. It is also known as 'ore dressing.'
(ii) The process of concentration
a. Based on densities is Gravity separation.
b. Based on magnetic nature is Electromagnetic separation.
(iii) Concentrated of an ore by froth floatation process: This process depends on preferential wettability of the ore and the gangue particles. Crushed ore is taken in a large tank containing water and certain oils. The ore particles get wetted by the oil and the gangue particles get wetted by water. The mixture is then agitated with the help of compressed air. The ore particles that get wetted with the oil form a froth on the top, and can be scooped out. This method is used for the concentration of sulphide ores.

📝 Teacher's Note: Use a simple demonstration with oil and water to show how different materials have different wetting properties. This makes the froth floatation concept much clearer to students.

🎯 Exam Tip: Always mention that froth floatation is specifically used for "sulphide ores" - this specificity shows you understand which method applies to which type of ore.

Solution 5:
1. Mercury.
2. Silver.
3. Zinc.
4. Aluminium.

📝 Teacher's Note: These appear to be answers to specific questions about metals. Help students understand the reactivity series and which metals can be found free in nature versus those that need extraction.

🎯 Exam Tip: When listing metals, think about their position in the reactivity series - this helps you remember which extraction methods are used for different metals.

Solution 6:
As we know that minerals are the naturally occurring compounds of metals which are generally mixed with earthy such as soil, sand, limestone and rocks while ores are those minerals from which a metal can be extracted profitably.
Hence "All ores are minerals but all minerals are not ores".

📝 Teacher's Note: Use Venn diagrams to show this relationship visually. Draw a large circle for minerals with a smaller circle inside for ores - this makes the concept crystal clear.

🎯 Exam Tip: This is a very common exam statement to prove or explain. Remember the key word "profitably" - this is what makes an ore different from just any mineral.

Solution 7:
1. Iron: Haematite(Fe₂O₃) and Magnetite (Fe₃O₄).
2. Zinc: Zinc blende (ZnS) and Calamine (ZnCO₃).
3. Aluminium:Bauxite(Al₂O₃) and Cryolite (AlF₃.3NaF).

📝 Teacher's Note: Create a chart with metal names, ore names, and chemical formulas. Students should memorize the most common ores for iron, zinc, and aluminium as these appear frequently in exams.

🎯 Exam Tip: Always write both the common name AND the chemical formula for ores - for example, "Haematite (Fe₂O₃)" shows complete knowledge and gets full marks.

Solution 8:

📝 Teacher's Note: Emphasize the key difference - calcination is WITHOUT air, roasting is WITH air. Use the memory trick: "Roasting needs air, just like roasting food needs air circulation in an oven."

🎯 Exam Tip: Always mention the type of ores used in each process - carbonate ores for calcination, sulphide ores for roasting. This distinction is crucial for full marks.

Solution 9:
Refining of metals: It is the further purification of metals obtained by reduction process to remove all the impurities.
Depending upon the nature of metal, nature of impurities and purpose for which metal is to be used.
The three methods used for refining are:
1. Liquation.
2. Distillation.
3. Electrolytic refining.

📝 Teacher's Note: Explain that even after extraction, metals are not 100% pure. Different refining methods are chosen based on the melting point and properties of the metal and its impurities.

🎯 Exam Tip: Always state that refining removes "all the impurities" - this shows you understand it's the final purification step after extraction.

Solution 10:
Electrolytic method: This method is based upon the phenomenon of electrolysis and is widely used to refine a number of metals such as copper, silver, gold etc. In this method, impure metal is made anode whereas a thin sheet of pure metal is used as cathode in an electrolytic tank. The electrolyte used in the tank is usually the acidified aqueous solution of a salt or complex salt of metal. On passing the electric current through electrodes the metal ions from anode go into the electrolyte solution. These cations gain electrons from the cathode and get deposited on it. The impurities either remain dissolved or get precipitated as anode mud.
At anode: M → M⁺⁺ + ne⁻
At cathode: M⁺⁺ + ne⁻ → M where M is the metal to be refined.

📝 Teacher's Note: Set up a simple electrolysis demonstration using copper sulfate solution with copper electrodes. Students can actually see the pure copper depositing at the cathode while impurities fall as sludge.

🎯 Exam Tip: Always mention that impure metal is the ANODE and pure metal is the CATHODE - getting this backwards is a common mistake that loses marks.

Solution 11:

(a) Bauxite.

(b) Purification of aluminium ore is done by the Baeyer's process.

Baeyer's process involves the following steps:

1. Bauxite is crushed to a fine powder.
2. The powdered bauxite is then treated with a strong solution of sodium hydroxide.
3. The mixture is then heated under pressure to 150°C to 250°C, for about 30 minutes. The heat and the pressure cause the alumina to dissolve in the sodium hydroxide to form sodium aluminate.

\( Al_2O_3.2H_2O + 2NaOH \rightarrow 2NaAlO_2 + 3H_2O \)

Alumina being amphoteric dissolves in sodium hydroxide.

4. The sodium aluminate solution is then filtered to separate the impurities.
5. The solution of sodium aluminate is then cooled lightly and sent into tanks called precipitators.
6. Crystals of aluminium hydroxide are then added when most of the sodium aluminate undergoes hydrolysis to precipitate insoluble aluminium hydroxide.

\( NaAlO_2 + 2H_2O \rightarrow NaOH + Al(OH)_3 \)

7. The solid aluminium hydroxide is separated by filtration.
8. The solid aluminium hydroxide is then washed and again filtered and dried, and then heated to about 1100°C to 1200°C, when aluminium hydroxide decomposes to form aluminium oxide.

\( 2Al(OH)_3 \rightarrow Al_2O_3 + 3 H_2O \)

(c) Aluminium is obtained from pure ore by Hall - Heroult's process as follows:

i. Electrolyte used:

a. Pure alumina
b. Cryolite
c. Aluminium fluoride
d. Calcium fluoride or Fluorspar

📝 Teacher's Note: Demonstrate the amphoteric nature of alumina by showing how it dissolves in both acids and bases. Use simple lab demonstrations with small amounts of Al₂O₃ to make the concept clearer for students.

🎯 Exam Tip: Always mention that Baeyer's process is specifically for purification of alumina, and include the key temperatures (150-250°C and 1100-1200°C) for full marks.

Solution 12:

(a) Fluorspar is \( CaF_2 \) and it helps in the mobility of the fused mixture.

(b) Cathode is the inner lining of gas-carbon of the electrolytic cell and anode is the thick carbon rods dipping into the fused electrolytes.

At cathode: Aluminium ions get reduced as:
\( Al^{3+} + 3e^- \rightarrow Al \)

At anode: oxygen gas is liberated as:
\( O^{2-} - 2e^- \rightarrow [O] \)
\( [O] + [O] \rightarrow O_2 \)

The oxygen formed at anode oxidizes the carbon of the anode to carbon dioxide.
\( C + O_2 \rightarrow CO_2 \)

As a result the anode gets oxidized and it has to be replaced periodically.

📝 Teacher's Note: Explain why the anode needs replacement but the cathode doesn't - this helps students understand the oxidation-reduction happening at different electrodes.

🎯 Exam Tip: Always mention both the electrode reactions and explain why anodes need periodic replacement due to carbon oxidation to CO₂.

Solution 13:

Cryolite acts as a solvent for the electrolytic mixture in the electrolytic reduction of alumina in the Hall's process.

📝 Teacher's Note: Emphasize that cryolite serves two main purposes - reducing melting point and making the mixture conductive. Students often forget the conductivity aspect.

🎯 Exam Tip: Mention both functions of cryolite: lowering melting point of alumina and making molten alumina a good conductor of electricity.

Solution 14:

(a) Aluminium is a more active metal than iron, but suffers less corrosion because of the formation of a thin, transparent, protective, non-porous adhering film of aluminium oxide on the surface of aluminium which makes it resistant to corrosion.

(b) Aluminium vessels should not be cleaned with powders containing alkalis because alunimun reacts with alkalis to form respective aluminate and hydrogen gas as:
\( 2Al + 2NaOH + 2H_2O \rightarrow 2NaAlO_2 + 3H_2 \)

(c) Food containing iron salts should not be cooked in aluminium utensils as there is a danger of aluminium toxicity.

📝 Teacher's Note: Show students an aluminium surface under magnification to demonstrate the protective oxide layer. This visual helps them understand why aluminium doesn't rust like iron.

🎯 Exam Tip: For corrosion resistance, always mention the "thin, transparent, protective oxide layer" - these are key descriptive words examiners look for.

Solution 15:

(i) \( Fe_2O_3 + 2Al \rightarrow 2Fe + Al_2O_3 \)
(ii) \( 2Al(OH)_3 \rightarrow Al_2O_3 + 3H_2O \)
(iii) \( 2Al + N_2 \rightarrow 2AlN \)
(iv) \( Al_2O_3 + 2H_2O + 2NaOH \rightarrow 2NaAlO_2 + 3H_2O \)

📝 Teacher's Note: Practice balancing these equations step by step. Students often struggle with the stoichiometry, especially in the thermite reaction.

🎯 Exam Tip: Always check if coefficients are in lowest terms and verify by counting atoms on both sides of each equation.

Solution 16:

An alloy is a homogeneous mixture of two or more metals fused together and then solidified. Alloys are made because they have many salient features:
1. Tensile strength.
2. Strength.
3. Electrical hardness.

📝 Teacher's Note: Show students actual samples of pure metals vs their alloys (like pure iron vs steel) to demonstrate the improved properties physically.

🎯 Exam Tip: Define alloy as "homogeneous mixture" and mention that alloys are generally stronger and more resistant to corrosion than pure metals.

Solution 17:

The properties of alloys which are different from constituent metals are:
1. Alloys are stronger and harder than the metals of which they are made.
2. Alloys are more resistant to corrosion.

📝 Teacher's Note: Explain this using the concept of crystal lattice disruption - when different sized atoms mix, they create a stronger structure that's harder to deform.

🎯 Exam Tip: Always mention both "strength/hardness" and "corrosion resistance" as the two main improved properties of alloys.

Solution 18:

Amalgam: A mixture or an alloy of mercury with a number of metals or non-metals is known as amalgams. An amalgam may be liquid such as Na/Hg or a solid like Zn/Hg.
1. Iron does not form amalgam.
2. Dental amalgam which is a mixture of mercury with a silver tin alloy is used for dental fillings.

📝 Teacher's Note: Explain why iron doesn't form amalgams while most other metals do - this helps students understand the special chemical behavior of iron.

🎯 Exam Tip: Remember that amalgams are specifically mercury alloys, and mention the dental application as a common example.

Solution 19:

📝 Teacher's Note: Connect each alloy's composition to its uses - for example, chromium in stainless steel prevents rusting, which is why it's used for cutlery.

🎯 Exam Tip: Learn the approximate percentages and main applications. Remember that stainless steel contains chromium for corrosion resistance.

Solution 20:

1. Galvanization protects iron from rusting because in galvanization coating of zinc is done over iron articles and zinc being more electropositive would be attacked preferably than iron.
2. Stainless steel is more useful than steel as it is harder, has high tensile strength, more lustre, more resistance to corrosion and many chemicals.
3. Aluminium is extensively used for making aircraft parts because of features like high tensile strength, corrosion resistance light but hard and tough.
4. Cold water has no action on aluminium while burning aluminium decomposes steam.

📝 Teacher's Note: Demonstrate galvanization with a simple experiment showing how zinc corrodes first when both zinc and iron are exposed to the same corrosive environment.

🎯 Exam Tip: For galvanization, always mention that zinc is "more electropositive" than iron, so it gets attacked first, protecting the iron underneath.

Solution 21:

(i) \( 4Al + 3O_2 \rightarrow 2Al_2O_3 \) (s)
(ii) \( 2Al + N_2 \rightarrow 2AlN \)
(iii) \( 2Al + 2KOH + 2H_2O \rightarrow 2KAlO_2 + 3H_2 \)
(iv) \( Fe_2O_3 + 2Al \rightarrow 2Fe + Al_2O_3 \)
(v) \( Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2 \)

1. Iron having a coating of zinc is called galvanized iron.
2. iron which cannot be easily acted upon by acids is called as passive iron. Galvanized iron is called passive iron since coating of zinc protects the iron from corrosion as zinc is more electropositive and so would be attacked first.

📝 Teacher's Note: Emphasize that "passive iron" doesn't mean iron is inactive, but rather that it's protected from reacting due to the zinc coating.

🎯 Exam Tip: Balance all equations carefully and remember that galvanized iron is called "passive iron" because it's protected from acid attack.

Solution 1991-2:

Zinc amalgam which is a mixture of zinc and mercury.

📝 Teacher's Note: Remind students that amalgams always involve mercury as one component, and zinc amalgam is commonly used in laboratories.

🎯 Exam Tip: Clearly state that zinc amalgam is a "mixture of zinc and mercury" for complete definition.

Solution 1992-1:

\( 2Al + 2NaOH + 2H_2O \rightarrow 2NaAlO_2 + 3H_2 \)

📝 Teacher's Note: Show students that this reaction produces hydrogen gas, which is why alkali cleaners shouldn't be used on aluminium vessels.

🎯 Exam Tip: Remember this equation shows why aluminium reacts with alkalis - it's important for explaining why alkali cleaners damage aluminium utensils.

Solution 1992-3:

1. Nitrogen.
2. Iodine
3. Bromine
4. Carbon in the form of graphite

📝 Teacher's Note: Explain that these elements are either unreactive (like nitrogen and carbon as graphite) or only moderately reactive under normal conditions.

🎯 Exam Tip: Remember that carbon must be specified as "graphite" form when discussing its use in electrodes.

Solution 1992:

1. Cryolite is \( Na_3AlF_6 \) and its chemical name is Sodium aluminium fluoride.
2. Cryolite is used in the electrolysis of alumina. The function of cryolite is to
• Reduce melting point of alumina
• Make molten alumina a good conductor of electricity

📝 Teacher's Note: Emphasize that without cryolite, the process would require much higher temperatures and wouldn't conduct electricity well enough for electrolysis.

🎯 Exam Tip: Always mention both functions of cryolite - lowering melting point AND improving conductivity - for complete marks.

Solution 1993-1:

Gold.

📝 Teacher's Note: Explain that gold is so unreactive that it's found in pure form in nature, unlike most other metals which exist as compounds.

🎯 Exam Tip: Gold is the classic example of a metal found free in nature due to its extreme chemical inertness.

Solution 1993-2:

An ore of zinc is Zinc blende (ZnS).

(a) \( 2ZnS + 3O_2 \rightarrow 2ZnO + 3SO_2 \)
(b) \( ZnO + C \rightarrow Zn + CO \)
(c) In addition to zinc oxide carbon is put in the furnace to reduce it to zinc metal, large scale use of zinc is that it is used for electroplating

📝 Teacher's Note: Connect the roasting and reduction steps - first sulfide is converted to oxide, then oxide is reduced to metal using carbon.

🎯 Exam Tip: Always mention that zinc blende is roasted first to convert ZnS to ZnO, then reduced with carbon to get pure zinc.

Solution 1994-1:

Reactivity of metals with water: Sodium, calcium, magnesium, iron

📝 Teacher's Note: This appears to be an incomplete answer. The full reactivity series with water would show sodium and calcium reacting vigorously, magnesium slowly, and iron only with steam.

🎯 Exam Tip: When listing metal reactivity, arrange them from most reactive to least reactive and mention the conditions needed for reaction.

Solution 1994-2:

i. Reaction of sodium with water:
\( 2Na + H_2O \rightarrow 2NaOH + H_2 \)

ii. Reaction of calcium with water:
\( Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2 \)

iii. Reaction of magnesium with water:
\( Mg + H_2O \rightarrow MgO + H_2 \)

iv. Reaction of iron with water:
\( 3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2 \)

📝 Teacher's Note: Demonstrate these reactions safely in class, starting with the least reactive (iron with steam) and working up to the most reactive (sodium with water).

🎯 Exam Tip: Remember that sodium and calcium react with cold water, magnesium with hot water, and iron only with steam at high temperatures.

Solution 1995-1:

(a) Ore of iron is haematite and ore of aluminium is bauxite.

(b) Reduction of the oxide is the most important chemical process in the extraction of any metal.

(i) In case of iron:
\( Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2 \)

(ii) In case of aluminium:
\( Al_2O_3 \) cannot be easily reduced; hence it is subjected to electrolysis. Aluminium is collected at the cathode.

(c)

(d) Carbon.

📝 Teacher's Note: Highlight that iron extraction uses chemical reduction with carbon monoxide, while aluminium requires electrolysis because it's too reactive for chemical reduction.

🎯 Exam Tip: Remember that aluminum requires electrolysis because Al₂O₃ cannot be easily reduced chemically, unlike Fe₂O₃ which reduces easily with CO.

Solution 1995-2:

Zinc amalgam.

📝 Teacher's Note: Zinc amalgam is commonly used in laboratory preparations and as a reducing agent in organic chemistry.

🎯 Exam Tip: Remember that zinc amalgam is specifically a mixture of zinc and mercury, used as a reducing agent.

Solution 1996-1:
Answer: (a) Bauxite: It is the principle ore from which aluminium is extracted.
(b) Sodium hydroxide: It reacts with alumina to form sodium aluminate which is then further filtered to separate impurities.
\( Al_2O_3.2H_2O + 2NaOH \rightarrow 2NaAlO_2 + 3H_2O \)
Alumina → Sodium aluminate
(c) Cryolite is used in the electrolysis of alumina. The function of cryolite is to
(1) Reduce melting point of alumina
(2) Make molten alumina a good conductor of electricity
(d) Graphite: It acts as electrodes in the electrolytic extraction of aluminium.
In simple words: Bauxite is the main source of aluminium, sodium hydroxide helps purify it, cryolite makes the process easier by lowering temperature, and graphite acts as electrical conductors.

📝 Teacher's Note: Use a simple analogy - cryolite works like salt on ice, making alumina melt at lower temperatures. Show students the chemical equation step by step to build understanding.

🎯 Exam Tip: Always mention both functions of cryolite - reducing melting point AND increasing conductivity - for full marks.

Solution 1996-2:
Answer: (a) Reaction at cathode:
\( Al^{3+} + 3e^- \rightarrow Al \)
(b) Oxygen gas is liberated at anode as:
\( O^{2-} - 2e^- \rightarrow [O] \)
\[ [O] + [O] \rightarrow O_2 \]
The oxygen formed at anode oxidizes the carbon of the anode to carbon dioxide.
\( C + O_2 \rightarrow CO_2 \)
As a result the anode gets oxidized and it has to be replaced periodically.
In simple words: At the cathode, aluminium metal forms, while at the anode oxygen gas is produced which burns away the carbon electrodes, so they need regular replacement.

📝 Teacher's Note: Emphasize that anode replacement is costly and a major challenge in aluminium production. Use the analogy of a candle burning away.

🎯 Exam Tip: Write the cathode and anode reactions separately with proper electron transfer equations for full marks.

Solution 1996-3:
Answer: An alloy is a homogeneous mixture of two or more metals fused together and then solidified.
1. The special property of duralumin is:
• Light but hard
• Resistant to corrosion
• Ductile
2. Type metal = Hard
In simple words: Alloys are like metal cocktails - mixing different metals to get better properties than any single metal alone.

📝 Teacher's Note: Give examples of common alloys students see daily - brass in door handles, steel in utensils. This makes the concept relatable.

🎯 Exam Tip: Always mention "homogeneous mixture" and "fused together" when defining alloys - these are key marking points.

Solution 1997-1:
Answer: Chromium and nickel is added to steel to make it stainless steel.
In simple words: Adding chromium and nickel to steel makes it rust-proof, just like adding ingredients to food changes its taste and properties.

📝 Teacher's Note: Show students actual stainless steel items and regular iron items to demonstrate the difference in corrosion resistance.

🎯 Exam Tip: Remember both elements - chromium AND nickel - are needed for stainless steel, not just one.

Solution 1997-2:
Answer: Ore: Those minerals from which a metal can be extracted profitably are called ores. For example bauxite ore is used to extract aluminium metal, hematite ore is used to extract iron metal.
In simple words: Ores are like treasure rocks - they contain metals that can be extracted in a way that makes business sense.

📝 Teacher's Note: Emphasize the word "profitably" - many minerals contain metals but only ores make extraction economically viable.

🎯 Exam Tip: Always include the word "profitably" in your definition of ore - it's the key distinguishing factor from mere minerals.

Solution 1998-1:
Answer: 1. good, poor.
2. non-malleable.
3. form negative ions.
4. basic oxides.
In simple words: These are basic properties that distinguish metals from non-metals in terms of conductivity, shape-changing ability, ion formation, and oxide nature.

📝 Teacher's Note: Use simple demonstrations - bend a copper wire vs. break a pencil lead to show malleability differences.

🎯 Exam Tip: For fill-in-the-blanks about metal properties, remember the opposite characteristics of non-metals.

Solution 1998-2:
Answer: 1. Mercury.
2. Graphite.
In simple words: Mercury is the only liquid metal at room temperature, and graphite is the only non-metal that conducts electricity well.

📝 Teacher's Note: Show students a thermometer (mercury) and pencil lead (graphite) as real-world examples of these exceptional cases.

🎯 Exam Tip: These are classic exceptions to general metal/non-metal properties - memorize them as they appear frequently in exams.

Solution 1998-3:
Answer: Metals have 1, 2, 3 valence electrons while non-metals have 4, 5, 6 or 7 valence electrons.
In simple words: Metals have few outer electrons (1-3) that they can easily lose, while non-metals have many outer electrons (4-7) that they prefer to gain or share.

📝 Teacher's Note: Draw electronic configurations on the board to show this pattern clearly. Use sodium (1) and chlorine (7) as classic examples.

🎯 Exam Tip: Remember the exact numbers - metals: 1,2,3 and non-metals: 4,5,6,7 valence electrons.

Solution 1999-1:
Answer: Magnesium oxide, iron (II) oxide, lead (II) oxide and then copper (II) oxide.
In simple words: This is the order from easiest to hardest to reduce - magnesium oxide needs the most energy while copper oxide needs the least.

📝 Teacher's Note: Relate this to the reactivity series - more reactive metals form oxides that are harder to reduce back to metal.

🎯 Exam Tip: The order follows reactivity series - more reactive metal = harder to reduce its oxide.

Solution 1999-2:
Answer: (a) Reduction of copper oxide by hydrogen:
\( CuO + H_2 + \text{heat} \rightarrow Cu + H_2O \)
(b) Reduction of iron (III) oxide by carbon monoxide:
\( Fe_2O_3 + 3CO + \text{heat} \rightarrow 2Fe + 3CO_2 \)
(c) Reduction of lead (II) oxide by carbon:
\( PbO + C + \text{heat} \rightarrow Pb + CO \)
In simple words: These reactions remove oxygen from metal oxides using different reducing agents - hydrogen, carbon monoxide, and carbon respectively.

📝 Teacher's Note: Emphasize that reduction means removal of oxygen. Show students how different reducing agents work for different metals.

🎯 Exam Tip: Always balance the chemical equations and mention heat as a condition for these reactions.

Solution 1999-3:
Answer: 1. Blue
2. Red
3. Hydrogen
4. acidic, acidic
5. graphite.
In simple words: These are standard answers for common chemistry fill-in-the-blanks about color changes, gas evolution, and oxide properties.

📝 Teacher's Note: Connect these individual facts to broader concepts - like why certain reactions produce specific colors or gases.

🎯 Exam Tip: These appear to be standard textbook answers - ensure you understand the context of each blank.

Solution 2000-1:
Answer: 1. Blue
2. Red
3. Hydrogen
4. acidic, acidic
5. graphite.
In simple words: This appears to be the same set of standard fill-in-the-blank answers as the previous solution.

📝 Teacher's Note: Use these consistent patterns to help students remember common chemistry facts and color changes.

🎯 Exam Tip: Notice the repetition - these are likely important standard facts that appear regularly in exams.

Solution 2001-1:
Answer: 1. Copper
2. Iron
3. Zinc
4. Magnesium
In simple words: This is a list of common metals, likely arranged in some order of reactivity or properties.

📝 Teacher's Note: Help students understand what property or characteristic determines this particular order of metals.

🎯 Exam Tip: Without the original question, memorize this sequence as it likely represents an important metal ordering.

Solution 2001-2:
Answer: Sodium > magnesium > Zinc > Iron > Copper
In simple words: This shows the reactivity series order - sodium is most reactive and copper is least reactive among these metals.

📝 Teacher's Note: Use the memory aid "Some More Zinc In Copper" to help students remember this reactivity order.

🎯 Exam Tip: This is the classic reactivity series order - more reactive metals displace less reactive ones from their compounds.

Solution 2002-1:
Answer:

In simple words: This is a matching exercise showing which metal (numbered 1-5) corresponds to which use (lettered A-E).

📝 Teacher's Note: Explain to students how different properties of metals make them suitable for specific applications.

🎯 Exam Tip: In matching questions, cross out each option as you use it to avoid repetition errors.

Solution 2002-2:
Answer: (a) (i) Aluminium oxide being an amphoteric oxide reacts with sodium hydroxide solution to form soluble sodium meta aluminate.
\( Al_2O_3.2H_2O + 2NaOH + 2H_2O \rightarrow 2NaAlO_2 + 3H_2O \)
Alumina → Sodium aluminate
(ii) Iron(III) oxide remains undissolved in the sodium hydroxide solution and settles down.
(b) (i) Baeyer's process or Hall's process is used for the purification of bauxite.
(ii) \( 2Al(OH)_3 + \text{Heat} \rightarrow Al_2O_3 + 3H_2O \)
(c) (i) \( Na_3AlF_6 \). It is also known as Sodium aluminium fluoride.
(ii) Conducting solution is produced.
(iii) Because it reacts with \( O_2 \) produced at anode and gets consumed to form \( CO_2 \).
\( O^{2-} - 2e^- \rightarrow [O] \)
\( C + 2[O] \rightarrow CO_2 \)
So carbon rods are replaced from time to time.
(iv) \( 2Al^{3+} + 6e^- \rightarrow 2Al \)
(d) Duralumin is light, strong and more resistant to corrosion than aluminium.
In simple words: This covers the complete process of aluminium extraction - from purifying bauxite to electrolysis, including why electrodes need replacement.

📝 Teacher's Note: Break this complex process into steps. Show students how amphoteric means reacting with both acids and bases.

🎯 Exam Tip: Remember that aluminium oxide is amphoteric - this property is key to the purification process.

Solution 2003-1:
Answer:

In simple words: This table shows the main differences between metals and non-metals in terms of electron arrangement, oxide behavior, chemical reactions, and physical properties.

📝 Teacher's Note: Use examples for each property - sodium for metals and chlorine for non-metals to make comparisons clearer.

🎯 Exam Tip: Learn these four key differences - they form the basis of most metal vs non-metal comparison questions.

Solution 2004-1:
Answer: Iodine is a non-metal that has a metallic luster and sublimes on heating.
In simple words: Iodine is special among non-metals because it shines like a metal and directly changes from solid to gas when heated.

📝 Teacher's Note: Show students iodine crystals if available - they can see the metallic shine and sublimation process.

🎯 Exam Tip: Iodine is the classic exception - a non-metal with metallic luster. Remember both properties for full marks.

Solution 2004-2:
Answer: \( 2Al + 2NaOH + 2H_2O \rightarrow 2NaAlO_2 + 3H_2 \)
In simple words: This reaction shows aluminium metal reacting with sodium hydroxide and water to produce sodium aluminate and hydrogen gas.

📝 Teacher's Note: Point out that this reaction produces hydrogen gas, which can be tested with a burning splint.

🎯 Exam Tip: Balance this equation carefully - notice the 3H₂ on the product side, not 2H₂.

Solution 2004-3:
Answer: Zinc blende (ZnS)
In simple words: Zinc blende is the chemical name for zinc sulfide, which is an important ore of zinc.

📝 Teacher's Note: Explain that blende is an old mining term, and students should know both the common name and chemical formula.

🎯 Exam Tip: Remember the chemical formula ZnS along with the name zinc blende for complete answers.

Solution 2004-4:
Answer: (i) \( Zn + 2NaOH \rightarrow Na_2ZnO_2 + H_2 \)
(ii) \( Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2 \)
(iii) \( Zn + CuSO_4 \rightarrow ZnSO_4 + Cu \)
In simple words: These reactions show zinc reacting with a base, an acid, and a salt respectively, producing different products in each case.

📝 Teacher's Note: Use these reactions to show the versatile nature of zinc and different types of chemical reactions.

🎯 Exam Tip: Notice the pattern - zinc displaces copper from copper sulfate because zinc is more reactive than copper.

Solution 2004-5:
Answer: Galvanization.
In simple words: Galvanization is the process of coating iron with zinc to prevent rusting.

📝 Teacher's Note: Show students galvanized iron sheets or buckets to demonstrate this practical application.

🎯 Exam Tip: Galvanization uses zinc coating - remember this is different from other coating methods like tinning.

Solution 2004-6:
Answer: (a) \( Na_3AlF_6 \rightleftarrows 3Na^+ + Al^{3+} + 6F^- \)
\( CaF_2 \rightleftarrows Ca^{2+} + 2F^- \)
\( Al_2O_3 \rightleftarrows 2Al^{3+} + 3O^{2-} \)
(b) Fluorspar and cryolite act as solvent. The percentage by weight composition is as follows:
Alumina - 20% by mass
Cryolite - 60% by mass
Fluorspar - 20% by mass
(c) \( O^{2-} - 2e^- \rightarrow [O] \)
\[ [O] + [O] \rightarrow O_2 \]
(d) As a reducing agent: Aluminium has a high affinity for oxygen. It readily removes oxygen from oxides of less reactive metals.
In simple words: This explains the electrolysis setup for aluminium extraction, including the role of different compounds and why aluminium can act as a reducing agent.

📝 Teacher's Note: Emphasize the percentage composition - students often forget these specific numbers in exams.

🎯 Exam Tip: Remember the 20:60:20 ratio for alumina:cryolite:fluorspar - this is a common exam question.

Solution 2005-1:
Answer: 1. (i) B, D F
(ii) A, C E
2. (i) Sodium hydroxide solution
(ii) Cryolite
3. \( Na_3AlF_6 \)
In simple words: This appears to be matching exercises and short answers related to aluminium extraction and related processes.

📝 Teacher's Note: Use these matching exercises to reinforce the connections between chemicals and their roles in metal extraction.

🎯 Exam Tip: In matching questions, ensure each option is used only once unless specified otherwise.

Solution 2005-2:
Answer: 1. For stainless steel: iron, chromium
2. For brass: Copper and zinc.
In simple words: These are the main components of two important alloys - stainless steel contains iron and chromium, while brass contains copper and zinc.

📝 Teacher's Note: Bring examples of these alloys to class so students can see and feel the difference in properties.

🎯 Exam Tip: Don't forget to mention both components for each alloy - partial answers get partial marks.

Solution 2006-1:
Answer: 1. Mercury.
2. Cryolite.
3. Roasting.
4. Calcium silicate.
5. Zone of heat absorption.
In simple words: These are standard one-word or short answers covering various metallurgy and chemistry concepts.

📝 Teacher's Note: Connect each of these terms to broader concepts in metallurgy and explain their significance.

🎯 Exam Tip: These appear to be standard textbook answers - ensure you understand the context of each term.

Solution 2007-1:
Answer: (a) Sodium hydroxide.
(b) \( 2Al(OH)_3 + \text{Heat} \rightarrow Al_2O_3 + 3H_2O \)
(c) Carbon
(d) \( [Al^{3+} + 3e^- \rightarrow Al] \times 2 \)
(e) \( [O^{2-} - 2e^- \rightarrow O_2] \)
In simple words: These equations and answers cover the complete aluminium extraction process from bauxite purification to electrolysis.

📝 Teacher's Note: Show how each step connects to the next in the overall process of aluminium production.

🎯 Exam Tip: Remember to balance equations properly and show electron transfer clearly in electrolysis reactions.

Solution 2008-1:
Answer: (b)
In simple words: This is a multiple choice answer, likely for a question about metallurgy or chemical processes.

📝 Teacher's Note: Without the original question, ensure students understand why option (b) is correct in context.

🎯 Exam Tip: Always read MCQ options carefully and eliminate obviously wrong answers first.

Solution 2009-1:
Answer: 1. Carbon as it forms very large number of compounds while the rest do not.
2. Mercury as it is a liquid metal while the rest are solid.
In simple words: These identify the odd one out - carbon forms many compounds unlike other elements listed, and mercury is liquid unlike other solid metals.

📝 Teacher's Note: Use these as examples of how to identify patterns and exceptions in chemistry.

🎯 Exam Tip: In "odd one out" questions, clearly state the property that makes the chosen element different.

Solution 2009-2:
Answer: 1. Copper reacts with concentrated nitric acid to produce nitrogen dioxide.
2. Bauxite is the chief ore of aluminium.
In simple words: These are key facts about copper's reaction with nitric acid and the main source of aluminium metal.

📝 Teacher's Note: Demonstrate the copper-nitric acid reaction if possible - the brown gas formation is quite dramatic.

🎯 Exam Tip: Remember that copper reacts with concentrated nitric acid, not dilute - concentration matters in these reactions.

Solution 2009-3:
Answer: 1. A is cathode and B is anode.
2. Molten fluorides of Al, Na and Ba.
3. Graphite rods.
In simple words: This describes an electrolysis setup with electrode identification, electrolyte composition, and electrode material.

📝 Teacher's Note: Draw a clear diagram of the electrolysis cell showing cathode, anode, and the direction of current flow.

🎯 Exam Tip: Remember that in electrolysis, reduction occurs at cathode and oxidation at anode.

Solution 2009-4:
Answer:

In simple words: This table connects the specific uses of metals to the properties that make them suitable for those applications.

📝 Teacher's Note: Emphasize how material properties determine their practical applications in industry.

🎯 Exam Tip: Always link the property to the application - explain WHY that property makes the metal suitable for that use.