Gonit Vigyan

 

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Physical Properties of Metals

Metals exhibit distinct physical properties that make them useful for various applications. These properties help differentiate metals from non-metals. Let's explore them in detail along with activities to observe these properties.

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1. Metallic Lustre

Activity 3.1: Observing Metallic Lustre

Materials Required:

• Samples of iron, copper, aluminium, and magnesium
• Sandpaper

Procedure:

1. Observe the appearance of each metal sample.
2. Rub the surface of each sample with sandpaper.
3. Observe the appearance again after cleaning.

Observation:

• Metals have a shiny appearance, which is called metallic lustre.
• The surface becomes shinier after removing the outer oxide layer.

Conclusion:

• Metals exhibit metallic lustre in their pure state.
• This is why metals like gold and silver are used in jewelry.

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2. Hardness

Activity 3.2: Testing Hardness of Metals

Materials Required:

• Small pieces of iron, copper, aluminium, and magnesium
• A sharp knife
• Sodium metal (Handle with care using tongs and filter paper)

Procedure:

1. Try cutting small pieces of iron, copper, aluminium, and magnesium with a knife.
2. Take sodium metal with tongs, place it on a watch glass, and try cutting it.

Observation:

• Iron, copper, aluminium, and magnesium are hard and difficult to cut.
• Sodium is soft and can be easily cut with a knife.

Conclusion:

• Most metals are hard, but some (like sodium and potassium) are soft.

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3. Malleability (Ability to be Hammered into Sheets)

Activity 3.3: Observing Malleability

Materials Required:

• Small pieces of iron, zinc, lead, and copper
• A hammer
• A block of iron

Procedure:

1. Place a piece of iron on the iron block.
2. Strike it 4-5 times with the hammer.
3. Repeat the process with zinc, lead, and copper.

Observation:

• The shape of the metals changes but they do not break.
• Some metals flatten more than others.

Conclusion:

• Metals can be hammered into thin sheets (malleability).
• Gold and silver are the most malleable metals.
• Aluminium and copper are used for making foils and utensils.

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4. Ductility (Ability to be Drawn into Wires)

Activity 3.4: Observing Ductility

Think and Answer:

• List some metals whose wires you see in daily life.
• Why are these metals used for making electrical wires?

Observation:

• Metals like copper and aluminium are used for wires.
• Gold is the most ductile metal—a 2 km wire can be drawn from 1 gram of gold.

Conclusion:

• Metals can be drawn into thin wires (ductility).
• Copper and aluminium are used in electrical wiring due to their high ductility.

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5. Thermal Conductivity

Activity 3.5: Metals as Conductors of Heat

Materials Required:

• Aluminium or copper wire
• A stand, pin, wax, candle or spirit lamp

Procedure:

1. Clamp the metal wire on a stand.
2. Attach a pin to the free end of the wire using wax.
3. Heat the wire near the clamped end using a candle or burner.

Observation:

• After some time, the pin falls off as the wax melts.

Conclusion:

• Metals are good conductors of heat.
• Silver and copper are the best conductors, while lead and mercury are poor conductors.
• This is why cooking vessels are made of metals.

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6. Electrical Conductivity

Activity 3.6: Metals as Conductors of Electricity

Materials Required:

• Battery, bulb, wires, and various metals

Procedure:

1. Set up an electric circuit with a battery, bulb, and wires.
2. Insert a metal piece between terminals A and B in the circuit.
3. Observe whether the bulb glows.

Observation:

• The bulb glows, proving that metals conduct electricity.

Conclusion:

• Metals allow electricity to pass through them due to free electrons.
• Copper and aluminium are used in electrical wiring.
• Electrical wires are coated with PVC or rubber to prevent shocks.

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7. Sonority (Production of Sound when Struck)

Think and Answer:

• What happens when you strike a metal object?
• Why are school bells made of metals?

Observation:

• Metals produce a ringing sound when struck against a hard surface.

Conclusion:

• Metals are sonorous (they produce sound).
• This is why bells, gongs, and musical instruments are made of metals.

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8. High Melting and Boiling Points

Think and Answer:

• Why do metals not melt easily?
• Which metal remains liquid at room temperature?

Observation:

• Metals generally have high melting and boiling points.
• Mercury is an exception—it is liquid at room temperature.

Conclusion:

• Most metals have high melting points, making them useful in industries.

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9. Density and Strength

• Most metals are dense due to tightly packed atoms.
• Exceptions:
  • Lithium, sodium, and potassium have low density and float on water.
• Metals are strong and can bear heavy loads.
• Iron and steel are used in construction due to their strength.

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10. State at Room Temperature

• Most metals are solids at room temperature.
• Exception: Mercury is a liquid metal.

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Summary Table of Physical Properties of Metals

Property	Description	Example
Metallic Lustre	Shiny appearance	Gold, silver, copper
Hardness	Most metals are hard	Iron, copper
Malleability	Can be hammered into sheets	Aluminium foil
Ductility	Can be drawn into wires	Copper, gold
Thermal Conductivity	Good conductors of heat	Silver, copper
Electrical Conductivity	Conducts electricity	Copper, aluminium
Sonority	Produces sound when struck	School bells
Melting & Boiling Points	Generally high	Iron, tungsten
Density	Generally high	Lead, iron
State at Room Temperature	Mostly solids	Mercury is liquid

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Conclusion

Metals possess lustre, hardness, malleability, ductility, conductivity, sonority, and high melting points. These properties make metals essential in electrical wiring, construction, cookware, machinery, and jewelry.

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3.1.2 Physical Properties of Non-Metals

Non-metals are fewer in number compared to metals and have different physical properties. They are found in all three states—solids, liquids, and gases. Some common non-metals include carbon, sulphur, iodine, oxygen, and hydrogen.

Let us explore the physical properties of non-metals and compare them with metals through activities.

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1. Physical State

• Non-metals exist in all three states at room temperature:
  • Solids: Carbon, sulphur, phosphorus, iodine
  • Liquids: Bromine (the only liquid non-metal)
  • Gases: Oxygen, hydrogen, nitrogen, chlorine

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2. Lack of Lustre (Except Iodine)

Activity 3.7: Observing Lustre in Non-Metals

Materials Required:

• Carbon (coal or graphite), sulphur, iodine
• Sandpaper

Procedure:

1. Observe the appearance of carbon, sulphur, and iodine.
2. Rub each sample with sandpaper and note the change.

Observation:

• Sulphur and carbon (coal) do not have a shiny surface.
• Iodine is an exception—it is a non-metal but has a lustrous (shiny) surface.

Conclusion:

• Most non-metals do not have a metallic lustre.
• Iodine is an exception—it is shiny like a metal.

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3. Hardness (Non-Metals are Brittle)

Activity 3.7 (continued): Testing Hardness of Non-Metals

Materials Required:

• Carbon (coal), sulphur, and iodine
• Hammer

Procedure:

1. Take small pieces of carbon, sulphur, and iodine.
2. Strike them with a hammer.

Observation:

• Non-metals break into pieces when hammered—they are brittle.
• Diamond (a form of carbon) is an exception—it is the hardest natural substance.

Conclusion:

• Non-metals are brittle and cannot be hammered into sheets.
• Diamond is an exception—it is extremely hard.

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4. Non-Malleability and Non-Ductility

• Non-metals cannot be beaten into sheets (not malleable).
• Non-metals cannot be drawn into wires (not ductile).

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5. Poor Conductors of Heat and Electricity

Activity 3.6: Testing Electrical Conductivity

Materials Required:

• Battery, bulb, wires, and non-metal samples (coal, graphite, sulphur)

Procedure:

1. Set up an electric circuit like in Activity 3.6.
2. Place different non-metals (carbon, sulphur, iodine) between terminals A and B.
3. Observe if the bulb glows.

Observation:

• Sulphur and coal do not conduct electricity.
• Graphite (a form of carbon) is an exception—it is a good conductor.

Conclusion:

• Most non-metals are poor conductors of electricity.
• Graphite (a form of carbon) is an exception—it conducts electricity.

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6. Non-Sonorous

• Non-metals do not produce a ringing sound when struck.
• This is why bells are made of metals, not non-metals.

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7. Low Melting and Boiling Points

Exceptions:

• Diamond (a form of carbon) has a very high melting point.
• Iodine sublimes (directly changes from solid to gas).

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8. Low Density

• Non-metals generally have low density.
• Examples: Hydrogen and oxygen are very light gases.

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Comparison of Physical Properties of Metals and Non-Metals

Property	Metals	Non-Metals
Physical State	Mostly solids (except mercury)	Solids, liquids (bromine), gases
Lustre	Shiny (metallic lustre)	Dull (except iodine)
Hardness	Hard (except sodium, potassium)	Brittle (except diamond)
Malleability	Can be beaten into sheets	Cannot be beaten, breaks easily
Ductility	Can be drawn into wires	Cannot be drawn into wires
Conductivity	Good conductors of heat & electricity	Poor conductors (except graphite)
Sonority	Produces a ringing sound	Does not produce sound
Melting Points	High (except gallium, caesium)	Low (except diamond)

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9. Chemical Properties of Non-Metals

Non-metals are better classified based on their chemical properties rather than physical properties.

Activity 3.8: Reaction of Non-Metals with Oxygen

Materials Required:

• Magnesium ribbon, sulphur powder, test tubes, litmus paper, water, burner

Procedure:

1. Burn a magnesium ribbon in air.
2. Collect the ash formed and dissolve it in water.
3. Test the solution with red and blue litmus paper.
4. Burn sulphur powder in air.
5. Collect the fumes in a test tube, add water, and shake.
6. Test this solution with litmus paper.

Observations:

• Magnesium forms magnesium oxide (MgO), which is basic.
• Sulphur forms sulphur dioxide (SO₂), which is acidic.

Conclusion:

• Metal oxides are basic.
• Non-metal oxides are acidic.

Chemical Reactions:

1. Magnesium burns in oxygen:
   $2Mg + O₂ → 2MgO$
   $MgO + H₂O → Mg(OH)₂$ (Basic solution)

2. Sulphur burns in oxygen:
   $S + O₂ → SO₂$
   $SO₂ + H₂O → H₂SO₃$ (Acidic solution)

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10. Exceptions in Physical Properties of Non-Metals

• Iodine is a non-metal but has lustre.
• Diamond (carbon) is the hardest known natural substance.
• Graphite (carbon) is a good conductor of electricity.
• Bromine is a liquid non-metal.

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Conclusion

• Non-metals generally have low melting points, are brittle, poor conductors, and do not produce sound.
• There are exceptions like iodine (lustrous), graphite (conductive), and diamond (hardest substance).
• The classification of metals and non-metals based on physical properties is not perfect because of exceptions.
• Chemical properties are a better way to classify elements.

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QUESTIONS

1. Give an example of a metal which:

(i) is a liquid at room temperature.

Answer: Mercury (Hg) is the only metal that exists as a liquid at room temperature.

(ii) can be easily cut with a knife.

Answer: Sodium (Na) and Potassium (K) are soft metals that can be easily cut with a knife.

(iii) is the best conductor of heat.

Answer: Silver (Ag) is the best conductor of heat, followed by copper (Cu).

(iv) is a poor conductor of heat.

Answer: Lead (Pb) and Mercury (Hg) are poor conductors of heat compared to other metals.

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2. Explain the meanings of malleability and ductility.

(i) Malleability:

Answer: Malleability is the ability of a metal to be hammered or rolled into thin sheets without breaking.

• Example: Gold (Au) and Aluminium (Al) are highly malleable metals.
• Application: Aluminium foils used for food wrapping and gold used in jewelry.

(ii) Ductility:

Answer: Ductility is the ability of a metal to be drawn into thin wires.

• Example: Copper (Cu) and Silver (Ag) are highly ductile.
• Application: Copper is used in electrical wiring due to its ductility and good conductivity.

3.2 CHEMICAL PROPERTIES OF METALS

Metals react with various substances such as oxygen, water, acids, and bases, leading to different chemical changes. These reactions help us understand their reactivity and their properties.

We will explore these reactions in Sections 3.2.1 to 3.2.4.

3.2.1 What Happens When Metals Are Burnt in Air?

Most metals combine with oxygen when burnt in air to form metal oxides.

General Reaction:

$$\text{Metal} + \text{Oxygen} → \text{Metal Oxide}$$

Activity 3.9: Burning Metals in Air

Materials Required:

• Metal samples (Aluminium, Copper, Iron, Lead, Magnesium, Zinc, Sodium)
• Pair of tongs
• Burner

Procedure:

1. Hold a piece of metal using tongs and heat it in a flame.
2. Observe:
   • Which metals burn easily?
   • What flame colour do you see?
   • How does the metal surface appear after burning?
3. Collect the product formed and test its solubility in water.

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Observations and Reactions of Metals with Oxygen

Metal	Observation When Heated	Reaction with Oxygen	Nature of Oxide
Sodium (Na)	Burns vigorously with a bright yellow flame	$4Na + O_2 → 2Na_2O$	Basic (forms NaOH in water)
Potassium (K)	Burns with a lilac flame	$4K + O_2 → 2K_2O$	Basic (forms KOH in water)
Magnesium (Mg)	Burns with a bright white flame	$2Mg + O_2 → 2MgO$	Basic
Aluminium (Al)	Forms a thin protective oxide layer	$4Al + 3O_2 → 2Al_2O_3$	Amphoteric
Zinc (Zn)	Forms a yellowish-white oxide layer	$2Zn + O_2 → 2ZnO$	Amphoteric
Iron (Fe)	Iron filings burn in flame, solid iron does not	$3Fe + 2O_2 → Fe_3O_4$	Basic
Copper (Cu)	Does not burn, forms a black coating	$2Cu + O_2 → 2CuO$	Basic
Silver (Ag), Gold (Au)	No reaction	Does not react with oxygen	N/A

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Types of Metal Oxides

1. Basic Oxides: Dissolve in water to form alkalis.
   • Example: $$Na_2O + H_2O → 2NaOH$$ $$K_2O + H_2O → 2KOH$$
2. Amphoteric Oxides: Show both acidic and basic properties.
   • Example: Aluminium oxide (Al₂O₃), Zinc oxide (ZnO)
   • Reacts with acids and bases to form salts and water: $$Al_2O_3 + 6HCl → 2AlCl_3 + 3H_2O$$ $$Al_2O_3 + 2NaOH → 2NaAlO_2 + H_2O$$

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3.2.2 Reactivity of Metals Towards Oxygen

• Highly Reactive Metals (Sodium, Potassium): React violently with oxygen and are stored in kerosene to prevent accidents.
• Moderately Reactive Metals (Magnesium, Aluminium, Zinc, Iron): Form oxide layers that protect them from further oxidation.
• Least Reactive Metals (Copper, Silver, Gold): Do not burn in air or react with oxygen even at high temperatures.

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3.2.3 Protective Oxide Layers

• Some metals like Aluminium, Zinc, and Lead form a protective oxide layer that prevents further oxidation.
• Iron filings burn in oxygen, but iron blocks do not, due to the formation of a slow rusting layer.
• Copper forms a black coating of copper(II) oxide.
• Silver and Gold do not react with oxygen at all, making them ideal for making jewelry.

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3.2.4 Anodising – Protection of Aluminium

• Anodising is a process used to increase the thickness of the protective aluminium oxide layer.
• Process:
  1. A clean aluminium article is made the anode.
  2. It is electrolysed using dilute sulphuric acid.
  3. Oxygen gas is released at the anode and forms a thicker oxide layer.
  4. This improves resistance to corrosion and can be dyed for decorative purposes.

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Reactivity Series of Metals Based on Reaction with Oxygen

Metal	Reactivity with Oxygen
Potassium (K)	Very highly reactive, catches fire
Sodium (Na)	Highly reactive, stored in kerosene
Calcium (Ca)	Reacts quickly, forms oxide layer
Magnesium (Mg)	Burns with bright white flame
Aluminium (Al)	Forms protective oxide layer
Zinc (Zn)	Forms oxide coating
Iron (Fe)	Burns as filings, not as block
Lead (Pb)	Forms a protective oxide layer
Copper (Cu)	Forms a black coating (CuO)
Silver (Ag), Gold (Au)	Do not react

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Conclusion

1. Metals react with oxygen to form metal oxides.
2. Metal oxides can be basic or amphoteric.
3. Highly reactive metals like sodium and potassium react violently and must be stored in kerosene.
4. Less reactive metals like copper form oxide coatings but do not burn.
5. Unreactive metals like silver and gold do not react even at high temperatures.
6. Anodising protects aluminium from corrosion by thickening the oxide layer.
7. **Metals show different reactivities towards oxygen, forming the basis of the reactivity series.

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3.2 CHEMICAL PROPERTIES OF METALS

Metals show different chemical properties when they react with water, acids, and oxygen. These reactions help us understand the reactivity series of metals. In this section, we will explore how metals react with water and acids, along with key observations and reactions.

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3.2.2 What Happens When Metals React with Water?

Metals react with water to form metal oxides and hydrogen gas.
Some metal oxides dissolve in water to form metal hydroxides.

General Reactions:

1. Metal + Water → Metal Oxide + Hydrogen
2. Metal Oxide + Water → Metal Hydroxide (If soluble)

Activity 3.10: Reaction of Metals with Water

Materials Required:

• Metal samples: Aluminium, Copper, Iron, Lead, Magnesium, Zinc, Sodium
• Beakers with cold water, hot water, and steam setup
• Tongs

Procedure:

1. Put small pieces of metals into cold water. Observe which metals react.
2. For metals that do not react with cold water, place them in hot water.
3. If no reaction occurs, expose them to steam and observe.
4. Arrange the metals in decreasing order of reactivity with water.

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Observations: How Different Metals React with Water

Metal	Reaction with Cold Water	Reaction with Hot Water	Reaction with Steam
Potassium (K)	Reacts violently, catches fire	-	-
Sodium (Na)	Reacts violently, catches fire	-	-
Calcium (Ca)	Reacts moderately, forms bubbles	-	-
Magnesium (Mg)	No reaction	Reacts slowly	Reacts with steam
Aluminium (Al)	No reaction	No reaction	Reacts with steam
Zinc (Zn)	No reaction	No reaction	Reacts with steam
Iron (Fe)	No reaction	No reaction	Reacts slowly with steam
Lead (Pb), Copper (Cu), Silver (Ag), Gold (Au)	No reaction	No reaction	No reaction

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Reactions of Metals with Water

Highly Reactive Metals (React with Cold Water)

1. Sodium reacts violently with water and catches fire: $$2Na + 2H_2O → 2NaOH + H_2 + \text{heat}$$
2. Potassium reacts even more violently and bursts into flames: $$2K + 2H_2O → 2KOH + H_2 + \text{heat}$$
3. Calcium reacts moderately with water, forming hydrogen gas: $$Ca + 2H_2O → Ca(OH)_2 + H_2$$
   • Calcium starts floating due to bubbles of hydrogen gas sticking to its surface.

Moderately Reactive Metals (React with Hot Water/Steam)

4. Magnesium does not react with cold water but reacts with hot water: $$Mg + 2H_2O \text{(hot)} → Mg(OH)_2 + H_2$$
5. Aluminium, Zinc, and Iron do not react with cold or hot water but react with steam to form metal oxides and hydrogen gas: $$2Al + 3H_2O \text{(steam)} → Al_2O_3 + 3H_2$$ $$3Fe + 4H_2O \text{(steam)} → Fe_3O_4 + 4H_2$$

Unreactive Metals (No Reaction with Water or Steam)

6. Metals like Lead, Copper, Silver, and Gold do not react with water at all.

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Reactivity Series Based on Reaction with Water

Most Reactive → Least Reactive
K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag > Au

Note: Sodium and potassium are stored in kerosene to prevent accidental fires.

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3.2.3 What Happens When Metals React with Acids?

Metals react with dilute acids to form salt and hydrogen gas.

General Reaction:

$$\text{Metal} + \text{Dilute Acid} → \text{Salt} + H_2$$

Activity 3.11: Reaction of Metals with Dilute Acids

Materials Required:

• Metal samples: Magnesium, Aluminium, Zinc, Iron, Copper
• Test tubes with dilute hydrochloric acid (HCl)
• Thermometers

Procedure:

1. Put small pieces of metal into test tubes containing dilute HCl.
2. Observe:
   • Which metals react vigorously?
   • Which metal produces the most bubbles?
   • Record the temperature change.
3. Arrange metals in decreasing order of reactivity.

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Observations: How Different Metals React with Acids

Metal	Reaction with Dilute Acid	Rate of Bubble Formation	Heat Released?
Magnesium (Mg)	Vigorous reaction, dissolves quickly	Fastest	Yes, high
Aluminium (Al)	Reacts rapidly, forms salt and H₂	Fast	Yes
Zinc (Zn)	Moderate reaction, releases H₂	Medium	Moderate
Iron (Fe)	Slow reaction, few bubbles	Slow	Low
Copper (Cu), Silver (Ag), Gold (Au)	No reaction	No bubbles	No heat change

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Reactions of Metals with Dilute Hydrochloric Acid (HCl)

1. Magnesium reacts violently with HCl: $$Mg + 2HCl → MgCl_2 + H_2$$
2. Aluminium reacts but less vigorously: $$2Al + 6HCl → 2AlCl_3 + 3H_2$$
3. Zinc reacts moderately with HCl: $$Zn + 2HCl → ZnCl_2 + H_2$$
4. Iron reacts slowly with HCl: $$Fe + 2HCl → FeCl_2 + H_2$$
5. Copper, Silver, and Gold do not react with acids.

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Why Doesn't Copper React with Acids?

• Copper is less reactive than hydrogen, so it cannot displace hydrogen from acids.

Why Doesn't Hydrogen Gas Evolve in HNO₃?

• Nitric acid (HNO₃) is a strong oxidizing agent.
• It oxidizes hydrogen gas into water and gets reduced to nitrogen oxides (NO, NO₂, N₂O).
• However, Magnesium (Mg) and Manganese (Mn) react with very dilute HNO₃ to release H₂ gas.

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Aqua Regia – The Strongest Acid

• Aqua regia (Latin: "Royal Water") is a 3:1 mixture of concentrated HCl and HNO₃.
• It is one of the few substances that can dissolve gold and platinum.
• Neither HCl nor HNO₃ alone can dissolve gold, but together they can.

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Reactivity Series Based on Reaction with Acids

Most Reactive → Least Reactive
Mg > Al > Zn > Fe > Pb > Cu > Ag > Au

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Conclusion

1. Metals react with water to form metal oxides/hydroxides and hydrogen gas.
2. Highly reactive metals (K, Na, Ca) react violently with cold water.
3. Moderately reactive metals (Mg, Zn, Fe) react with hot water or steam.
4. Metals react with acids to produce salts and hydrogen gas.
5. Copper, Silver, and Gold do not react with acids.
6. Aqua regia is a strong acid that dissolves gold and platinum.

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3.2.4 How Do Metals React with Solutions of Other Metal Salts?

Metals vary in their reactivity. A more reactive metal can displace a less reactive metal from its salt solution. This principle is used to determine the reactivity of metals through displacement reactions.

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Displacement Reaction of Metals

A more reactive metal can replace a less reactive metal from its salt solution.

General Reaction:

$$\text{Metal A} + \text{Salt Solution of B} → \text{Salt Solution of A} + \text{Metal B}$$

For example, if iron (Fe) is more reactive than copper (Cu), then iron will displace copper from copper sulphate solution:

$$Fe + CuSO_4 → FeSO_4 + Cu$$

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Activity 3.12: Reaction of Metals with Other Metal Salts

Materials Required:

• Copper wire
• Iron nail
• Solutions of Iron Sulphate (FeSO₄) and Copper Sulphate (CuSO₄)
• Test tubes

Procedure:

1. Place a copper wire in iron sulphate solution (FeSO₄).
2. Place an iron nail in copper sulphate solution (CuSO₄).
3. Leave both setups undisturbed for 20 minutes.
4. Observe any colour change or deposition of a new metal.

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Observations

Test Tube Setup	Observation	Reaction Occurred?
Copper wire in FeSO₄ solution	No visible change	No reaction
Iron nail in CuSO₄ solution	Iron nail gets coated with reddish-brown copper	Reaction occurred

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Explanation

• Iron is more reactive than copper, so it displaces copper from CuSO₄ solution, forming iron sulphate (FeSO₄).
• The reddish-brown deposit on the iron nail is copper metal.
• Since copper is less reactive than iron, it cannot displace iron from FeSO₄ solution.

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Balanced Chemical Equation

$$Fe + CuSO_4 → FeSO_4 + Cu$$

• Iron (Fe) displaces copper (Cu) from copper sulphate solution, forming iron sulphate (FeSO₄).
• Type of Reaction: Displacement Reaction

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Conclusion from Activities 3.9, 3.10, and 3.11

• In Activity 3.9, metals reacted differently with oxygen (burning in air).
• In Activity 3.10, metals reacted differently with water.
• In Activity 3.11, metals reacted at different rates with acids.
• In Activity 3.12, iron displaced copper, proving that iron is more reactive than copper.
• Displacement reactions are the best way to compare metal reactivity.

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3.2.5 The Reactivity Series of Metals

The reactivity series is a list of metals arranged in decreasing order of reactivity. It helps predict whether a metal can displace another from its salt solution.

Table 3.2: Reactivity Series of Metals

Metal	Reactivity
Potassium (K)	Most reactive
Sodium (Na)	↓
Calcium (Ca)	↓
Magnesium (Mg)	↓
Aluminium (Al)	↓
Zinc (Zn)	↓
Iron (Fe)	↓
Lead (Pb)	↓
[Hydrogen]	↓
Copper (Cu)	↓
Mercury (Hg)	↓
Silver (Ag)	↓
Gold (Au)	Least reactive

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Trends in the Reactivity Series

• Metals at the top (K, Na, Ca, Mg) are highly reactive.
  • React violently with water and acids.
  • Easily lose electrons to form positive ions (cations).
• Metals in the middle (Al, Zn, Fe, Pb) are moderately reactive.
  • React with acids and steam but not cold water.
• Metals at the bottom (Cu, Hg, Ag, Au) are the least reactive.
  • Do not react with water or acids.
  • Do not corrode easily.

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Uses of the Reactivity Series

1. Predicting Displacement Reactions

   • A more reactive metal can replace a less reactive metal from its salt solution.
   • Example: Iron can displace copper from CuSO₄ solution, but copper cannot displace iron from FeSO₄.

2. Determining Metal Extraction Methods

   • Highly reactive metals (K, Na, Ca) are extracted by electrolysis.
   • Moderate metals (Zn, Fe, Pb) are extracted by chemical reduction.
   • Least reactive metals (Cu, Ag, Au) are found in native (pure) form in nature.

3. Corrosion Resistance

   • Less reactive metals like gold and silver do not corrode, making them useful for jewelry.
   • Iron rusts easily because it is more reactive.

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Conclusion

1. More reactive metals can displace less reactive metals from their salt solutions.
2. Iron is more reactive than copper, so it can replace copper from CuSO₄.
3. The Reactivity Series ranks metals from most reactive to least reactive.
4. Highly reactive metals react with water, acids, and oxygen, while less reactive metals do not.
5. The Reactivity Series helps in understanding displacement reactions, metal extraction, and corrosion resistance.

QUESTIONS & ANSWERS

1. Why is sodium kept immersed in kerosene oil?

Answer:

• Sodium (Na) is highly reactive and reacts violently with oxygen and moisture in the air.
• If exposed to air, it can catch fire due to its extreme reactivity.
• To prevent accidental fires, sodium is stored under kerosene oil, which does not contain oxygen or water.

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2. Write equations for the reactions of:

(i) Iron with steam

Answer:

• Iron reacts with steam (H₂O in gaseous form) to produce iron(III) oxide (Fe₃O₄) and hydrogen gas (H₂).

$$3Fe + 4H_2O \text{(steam)} → Fe_3O_4 + 4H_2$$

(ii) Calcium and Potassium with water

Answer:

• Calcium reacts with cold water to form calcium hydroxide (Ca(OH)₂) and hydrogen gas (H₂).

$$Ca + 2H_2O → Ca(OH)_2 + H_2$$

• Potassium reacts violently with cold water, producing potassium hydroxide (KOH) and hydrogen gas (H₂), releasing heat energy.

$$2K + 2H_2O → 2KOH + H_2 + \text{heat}$$

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3. Analysis of Metal Reactivity from the Given Table

Metal	Iron(II) sulphate	Copper(II) sulphate	Zinc sulphate	Silver nitrate
A	No reaction	Displacement	-	-
B	Displacement	No reaction	-	-
C	No reaction	No reaction	No reaction	Displacement
D	No reaction	No reaction	No reaction	No reaction

(i) Which is the most reactive metal?

Answer:

• Metal B is the most reactive because it displaces iron from iron(II) sulphate solution.
• This means B is more reactive than iron.

(ii) What would you observe if B is added to a solution of Copper(II) sulphate?

Answer:

• Since B does not react with CuSO₄ solution, no displacement reaction occurs.
• No colour change or metal deposition would be observed.

(iii) Arrange the metals A, B, C, and D in decreasing order of reactivity.

Answer:

• B > A > C > D
• B is the most reactive (displaces iron).
• A is next (displaces copper).
• C is less reactive (only displaces silver).
• D is the least reactive (no reaction with any solution).

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4. Which gas is produced when dilute hydrochloric acid is added to a reactive metal? Write the chemical reaction when iron reacts with dilute H₂SO₄.

Answer:

• Hydrogen gas (H₂) is produced when a reactive metal reacts with dilute hydrochloric acid (HCl).

• Reaction of Iron with Sulfuric Acid:

$$Fe + H_2SO_4 → FeSO_4 + H_2$$

• Observation:
  • Bubbles of hydrogen gas will be seen.
  • The temperature of the solution increases due to an exothermic reaction.

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5. What would you observe when zinc is added to a solution of iron(II) sulphate? Write the chemical reaction that takes place.

Answer:

• Observation:

  • Zinc is more reactive than iron, so it displaces iron from iron(II) sulphate solution (FeSO₄).
  • The solution changes from green (FeSO₄) to colourless (ZnSO₄).
  • Iron metal gets deposited at the bottom of the test tube.

• Reaction:

$$Zn + FeSO_4 → ZnSO_4 + Fe$$

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3.3 HOW DO METALS AND NON-METALS REACT?

Metals and non-metals react in specific ways to form compounds. Their reactivity is based on their electronic configuration, and they tend to react in a way that helps them achieve a stable noble gas configuration.

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Why Do Metals and Non-Metals React?

• Noble gases have a completely filled outer shell and do not react easily.
• Other elements react to achieve a stable electronic configuration similar to noble gases.
• Metals tend to lose electrons to form positively charged ions (cations).
• Non-metals tend to gain electrons to form negatively charged ions (anions).
• The attraction between oppositely charged ions forms ionic (electrovalent) compounds.

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Electronic Configuration of Some Elements

Type of Element	Element	Atomic Number	Electronic Configuration (K, L, M, N shells)
Noble Gases	Helium (He)	2	2
	Neon (Ne)	10	2, 8
	Argon (Ar)	18	2, 8, 8
Metals	Sodium (Na)	11	2, 8, 1
	Magnesium (Mg)	12	2, 8, 2
	Aluminium (Al)	13	2, 8, 3
	Potassium (K)	19	2, 8, 8, 1
	Calcium (Ca)	20	2, 8, 8, 2
Non-Metals	Nitrogen (N)	7	2, 5
	Oxygen (O)	8	2, 6
	Fluorine (F)	9	2, 7
	Phosphorus (P)	15	2, 8, 5
	Sulphur (S)	16	2, 8, 6
	Chlorine (Cl)	17	2, 8, 7

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Formation of Ionic Compounds (Electrovalent Bonding)

Example 1: Formation of Sodium Chloride (NaCl)

1. Sodium (Na) has 1 electron in its outermost shell (M shell).
2. It loses 1 electron, achieving a stable 2,8 (Neon-like) configuration.
3. This forms a positively charged sodium ion (Na⁺).

$$Na → Na^+ + e^-$$

4. Chlorine (Cl) has 7 electrons in its outermost shell (M shell).
5. It gains 1 electron, achieving a stable 2,8,8 (Argon-like) configuration.
6. This forms a negatively charged chloride ion (Cl⁻).

$$Cl + e^- → Cl^-$$

7. Na⁺ and Cl⁻ are oppositely charged and attract each other due to electrostatic forces, forming sodium chloride (NaCl).
8. NaCl does not exist as single molecules but as a lattice of oppositely charged ions.

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Example 2: Formation of Magnesium Chloride (MgCl₂)

1. Magnesium (Mg) has 2 electrons in its outermost shell (M shell).
2. It loses 2 electrons to achieve a stable 2,8 (Neon-like) configuration, forming a Mg²⁺ ion.

$$Mg → Mg^{2+} + 2e^-$$

3. Each chlorine (Cl) atom needs 1 electron to complete its octet.
4. Two chlorine atoms take the two electrons lost by magnesium, forming two Cl⁻ ions.

$$Cl + e^- → Cl^-$$

5. The Mg²⁺ and two Cl⁻ ions attract each other, forming magnesium chloride (MgCl₂).

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Definition: Ionic (Electrovalent) Compounds

Compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds or electrovalent compounds.

Examples of Ionic Compounds:

• Sodium chloride (NaCl)
• Magnesium chloride (MgCl₂)
• Calcium oxide (CaO)
• Potassium bromide (KBr)

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Cations and Anions in Ionic Compounds

Ionic Compound	Cation (Positive Ion)	Anion (Negative Ion)
NaCl (Sodium chloride)	Na⁺ (Sodium ion)	Cl⁻ (Chloride ion)
MgCl₂ (Magnesium chloride)	Mg²⁺ (Magnesium ion)	Cl⁻ (Chloride ion)
CaO (Calcium oxide)	Ca²⁺ (Calcium ion)	O²⁻ (Oxide ion)
KBr (Potassium bromide)	K⁺ (Potassium ion)	Br⁻ (Bromide ion)

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Key Features of Ionic Compounds

1. Formed by transfer of electrons

   • Metals lose electrons (form cations).
   • Non-metals gain electrons (form anions).
   • Oppositely charged ions attract, forming ionic bonds.

2. High melting and boiling points

   • Strong electrostatic forces hold ions together.
   • Requires a lot of energy to break bonds.

3. Soluble in water

   • Water molecules break electrostatic attractions, dissolving the compound.

4. Conduct electricity in molten or dissolved state

   • In solid form, ions are fixed and cannot conduct electricity.
   • In molten or dissolved state, ions move freely and conduct electricity.

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Conclusion

1. Metals lose electrons to form positive ions (cations).
2. Non-metals gain electrons to form negative ions (anions).
3. The electrostatic attraction between cations and anions forms ionic compounds.
4. Ionic compounds have high melting points, are soluble in water, and conduct electricity in molten form.

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3.3.1 Properties of Ionic Compounds

Ionic compounds have distinct physical and chemical properties due to the strong electrostatic forces between their positively and negatively charged ions. Let’s explore these properties through Activity 3.13 and other observations.

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Activity 3.13: Observing the Properties of Ionic Compounds

Materials Required:

• Samples of sodium chloride (NaCl), potassium iodide (KI), barium chloride (BaCl₂), calcium chloride (CaCl₂)
• Metal spatula, Bunsen burner, water, kerosene, petrol
• Electrical circuit with electrodes and battery

Procedure & Observations:

Step	Procedure	Observation	Inference
1. Physical State	Observe the appearance and texture of the samples	All are solids, hard but brittle	Ionic compounds are solid at room temperature
2. Heating the sample	Heat a small amount on a metal spatula	Compounds do not melt easily	Ionic compounds have high melting points
3. Solubility Test	Dissolve samples in water, petrol, and kerosene	Soluble in water, insoluble in kerosene & petrol	Ionic compounds dissolve in polar solvents (water) but not in non-polar solvents
4. Electrical Conductivity Test	Make a circuit with electrodes dipped in a salt solution	The bulb glows, indicating electricity conduction	Ionic compounds conduct electricity in solution but not in solid form

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Properties of Ionic Compounds

(i) Physical Nature

• Ionic compounds are solids at room temperature.
• They are hard due to strong electrostatic forces between oppositely charged ions.
• However, they are brittle, meaning they break into pieces when force is applied.

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(ii) High Melting and Boiling Points

• Ionic compounds have very high melting and boiling points because a large amount of energy is needed to break the strong electrostatic attraction between ions.
• The melting and boiling points of some ionic compounds are given in Table 3.4:

Ionic Compound	Melting Point (K)	Boiling Point (K)
Sodium chloride (NaCl)	1074 K	1686 K
Lithium chloride (LiCl)	887 K	1600 K
Calcium chloride (CaCl₂)	1045 K	1900 K
Calcium oxide (CaO)	2850 K	3120 K
Magnesium chloride (MgCl₂)	981 K	1685 K

• Conclusion: The higher the melting and boiling point, the stronger the electrostatic attraction in the compound.

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(iii) Solubility in Water and Other Solvents

• Ionic compounds are soluble in water but insoluble in non-polar solvents like kerosene or petrol.
• Why?
  • Water is a polar solvent and helps break the electrostatic attraction between ions, allowing them to dissolve.
  • Non-polar solvents (kerosene, petrol) cannot break ionic bonds, so they remain insoluble.

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(iv) Electrical Conductivity

State of Ionic Compound	Conducts Electricity?	Reason
Solid State	❌ No	Ions are fixed in a rigid lattice and cannot move freely.
Molten State (Liquid Form)	✅ Yes	Heat breaks electrostatic forces, allowing free movement of ions.
Aqueous Solution (Dissolved in Water)	✅ Yes	Water separates ions, allowing them to move and carry charge.

• Ionic compounds do not conduct electricity in solid form because ions are held in place and cannot move.
• However, in the molten state or aqueous solution, ions are free to move, allowing the flow of electricity.

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Summary of Properties of Ionic Compounds

Property	Observation	Reason
Physical Nature	Hard but brittle solids	Strong electrostatic forces between ions
Melting & Boiling Points	High	A large amount of energy is required to break ionic bonds
Solubility	Soluble in water, insoluble in petrol & kerosene	Water is a polar solvent, petrol and kerosene are non-polar
Electrical Conductivity	Conducts in molten state and solution, but not in solid state	In solid form, ions are fixed; in molten/solution form, ions move freely

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Conclusion

1. Ionic compounds are hard, brittle solids with high melting and boiling points.
2. They dissolve in water but not in non-polar solvents like petrol and kerosene.
3. They do not conduct electricity in solid form but conduct in molten state or aqueous solution.

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QUESTIONS & ANSWERS

1. (i) Write the electron-dot structures for sodium, oxygen, and magnesium.

Answer:

• Electron-dot structures (Lewis structures) represent valence electrons of atoms using dots.

Sodium (Na) Electron-Dot Structure:

• Atomic Number of Na = 11, Electronic Configuration: 2, 8, 1
• Valence Electrons = 1

$$\text{Na} \cdot$$

Oxygen (O) Electron-Dot Structure:

• Atomic Number of O = 8, Electronic Configuration: 2, 6
• Valence Electrons = 6

$$\cdot \cdot O \cdot \cdot$$

Magnesium (Mg) Electron-Dot Structure:

• Atomic Number of Mg = 12, Electronic Configuration: 2, 8, 2
• Valence Electrons = 2

$$\cdot \cdot Mg \cdot \cdot$$

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(ii) Show the formation of Na₂O and MgO by the transfer of electrons.

Answer:

Formation of Sodium Oxide (Na₂O)

• Sodium (Na) loses 1 electron to form Na⁺ ion.
• Oxygen (O) gains 2 electrons to form O²⁻ ion.
• Two sodium atoms transfer their electrons to one oxygen atom.

$$2Na \rightarrow 2Na^+ + 2e^-$$ $$O + 2e^- \rightarrow O^{2-}$$ $$2Na^+ + O^{2-} \rightarrow Na_2O$$

Formation of Magnesium Oxide (MgO)

• Magnesium (Mg) loses 2 electrons to form Mg²⁺ ion.
• Oxygen (O) gains these 2 electrons to form O²⁻ ion.

$$Mg \rightarrow Mg^{2+} + 2e^-$$ $$O + 2e^- \rightarrow O^{2-}$$ $$Mg^{2+} + O^{2-} \rightarrow MgO$$

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(iii) What are the ions present in these compounds?

Answer:

• In Na₂O:
  • Cation: Na⁺ (Sodium ion)
  • Anion: O²⁻ (Oxide ion)
• In MgO:
  • Cation: Mg²⁺ (Magnesium ion)
  • Anion: O²⁻ (Oxide ion)

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2. Why do ionic compounds have high melting points?

Answer:

• Ionic compounds have high melting points because of strong electrostatic forces of attraction between positively and negatively charged ions.
• A large amount of energy is required to break these strong bonds.
• Example:
  • Sodium chloride (NaCl) melts at 1074 K
  • Calcium oxide (CaO) melts at 2850 K
• Thus, ionic compounds have high melting points due to strong ionic bonds.

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3.4 OCCURRENCE OF METALS

Metals are found in nature in different forms, depending on their reactivity. Some are found in the free state, while others occur as compounds in ores. The extraction of metals from ores involves several steps, including enrichment of ores, reduction, and refining.

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3.4.1 Minerals and Ores

• Minerals: Naturally occurring elements or compounds in the Earth's crust.
• Ores: Minerals that contain a high percentage of a particular metal and can be profitably extracted.

Examples of Minerals and Ores:

Metal	Mineral	Ore
Iron (Fe)	Hematite (Fe₂O₃)	Hematite (Fe₂O₃)
Aluminium (Al)	Bauxite (Al₂O₃.2H₂O)	Bauxite (Al₂O₃.2H₂O)
Copper (Cu)	Malachite (CuCO₃.Cu(OH)₂)	Copper pyrite (CuFeS₂)
Zinc (Zn)	Zincite (ZnO)	Zinc blende (ZnS)
Silver (Ag)	Argentite (Ag₂S)	Argentite (Ag₂S)
Gold (Au)	Native gold	Native gold

Sources of Metals

1. Earth’s crust: The primary source of metals.
2. Seawater: Contains soluble metal salts like sodium chloride (NaCl) and magnesium chloride (MgCl₂).

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3.4.1 Extraction of Metals

Classification Based on Reactivity

The extraction of a metal depends on its position in the reactivity series.

Category	Metals	Occurrence
Highly Reactive Metals	K, Na, Ca, Mg, Al	Always found in combined form (oxides, carbonates, sulphides). Never found in the free state.
Moderately Reactive Metals	Zn, Fe, Pb	Found as oxides, sulphides, and carbonates. Rarely found in the free state.
Low Reactive Metals	Cu, Ag, Au, Pt	Found mostly in free state (native form). Can also be found as sulphides or oxides.

Why Are Most Metal Ores Oxides?

• Oxygen is a highly reactive element and is abundant on Earth.
• Many metals react with oxygen over time to form metal oxides.
• Examples:
  • Iron exists as Fe₂O₃ (Hematite).
  • Aluminium exists as Al₂O₃ (Bauxite).

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3.4.2 Steps in Metal Extraction

Metal extraction involves the following steps (Fig. 3.10):

1. Enrichment of Ores (Ore Dressing or Beneficiation)

• Ores are contaminated with impurities (gangue) such as sand, clay, and soil.
• These impurities must be removed before extraction using physical or chemical methods.

Common Methods of Enrichment:

Method	Principle	Example
Hydraulic Washing (Gravity Separation)	Difference in density between ore and gangue.	Used for oxide ores like Haematite (Fe₂O₃).
Magnetic Separation	Difference in magnetic properties between ore and gangue.	Used for iron ores (Fe₃O₄, Fe₂O₃).
Froth Flotation	Difference in wettability of ore and gangue.	Used for sulphide ores (ZnS, PbS, CuFeS₂).
Leaching (Chemical Separation)	Ore is dissolved in a chemical solution to remove impurities.	Used for bauxite (Al₂O₃.2H₂O) extraction using NaOH.

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2. Extraction of Metals from Concentrated Ores

After enrichment, metals are extracted from their compounds using appropriate methods based on reactivity.

Metal Type	Extraction Method	Example
Highly Reactive Metals (K, Na, Ca, Mg, Al)	Electrolysis of molten ores	Na from NaCl, Al from Al₂O₃
Moderately Reactive Metals (Zn, Fe, Pb, Cu)	Reduction using Carbon (C) or Carbon Monoxide (CO)	Zn from ZnO, Fe from Fe₂O₃
Low Reactive Metals (Cu, Ag, Au, Pt)	Found in Free State or Extracted by Heating	Ag from Ag₂S, Au found as native gold

Examples of Reduction Reactions:

• Zinc Extraction (Using Carbon Reduction): $$ZnO + C \rightarrow Zn + CO$$
• Iron Extraction (Blast Furnace Reaction): $$Fe_2O_3 + 3C \rightarrow 2Fe + 3CO$$
• Aluminium Extraction (Electrolysis of Alumina): $$Al_2O_3 \rightarrow 2Al + 3O_2$$

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3. Refining of Metals

The extracted metal often contains impurities and must be purified using refining techniques.

Method	Used for	Example
Electrolytic Refining	Metals like Cu, Ag, Al, Zn	Copper refining
Distillation	Low-boiling metals like Zinc (Zn) and Mercury (Hg)	Zinc refining
Liquation	Metals with low melting points like Tin (Sn)	Tin refining
Zone Refining	High-purity metals like Silicon (Si) and Germanium (Ge)	Semiconductor purification

Example of Electrolytic Refining (Copper):

• Anode = Impure copper
• Cathode = Pure copper
• Electrolyte = Copper Sulphate (CuSO₄) solution

$$Cu^{2+} + 2e^- \rightarrow Cu \text{ (Deposited at cathode)}$$

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Summary of Metal Extraction Steps

1. Metals exist in nature as minerals or ores.
2. Highly reactive metals are extracted by electrolysis.
3. Moderately reactive metals are extracted by carbon reduction.
4. Less reactive metals occur in a free state or require mild heating.
5. Ores contain impurities (gangue) that must be removed through enrichment.
6. The extracted metal is purified using refining techniques like electrolytic refining.

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Conclusion

1. Metals are found in the Earth's crust as minerals or ores.
2. Metals are classified as highly, moderately, or low reactive based on the reactivity series.
3. Ores must be purified using enrichment techniques before metal extraction.
4. Highly reactive metals are extracted by electrolysis, moderately reactive metals by reduction, and low reactive metals by heating.
5. After extraction, metals are refined to obtain pure metal for industrial use.

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3.4.3 Extracting Metals Low in the Activity Series

• Metals low in the activity series (e.g., gold, silver, mercury, copper) are unreactive.
• They are often found in a free state or as sulphides or oxides in nature.
• Their oxides can be reduced to metals simply by heating.

Example: Extraction of Mercury from Cinnabar (HgS)

Step 1: Heating Mercury Sulphide (HgS) in Air

• HgS is first converted to mercuric oxide (HgO) by roasting in the presence of oxygen.

$$2HgS(s) + 3O_2(g) \rightarrow 2HgO(s) + 2SO_2(g)$$

Step 2: Further Heating of Mercuric Oxide (HgO)

• HgO decomposes to give liquid mercury (Hg).

$$2HgO(s) \xrightarrow{\text{Heat}} 2Hg(l) + O_2(g)$$

Example: Extraction of Copper from Copper Sulphide (Cu₂S)

Step 1: Heating Cu₂S in Air

$$2Cu_2S + 3O_2 \xrightarrow{\text{Heat}} 2Cu_2O + 2SO_2$$

Step 2: Further Heating of Cu₂O with Cu₂S

$$2Cu_2O + Cu_2S \xrightarrow{\text{Heat}} 6Cu(s) + SO_2(g)$$

Conclusion:

• Mercury and copper can be extracted from their ores by simple heating.
• This method is suitable for low-reactivity metals that do not strongly combine with oxygen.

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3.4.4 Extracting Metals in the Middle of the Activity Series

• Metals in the middle of the reactivity series (e.g., iron, zinc, lead, copper) are found as sulphides or carbonates.
• It is easier to extract metals from oxides than sulphides or carbonates.
• Thus, sulphides and carbonates are first converted to oxides before reduction.

Conversion of Ores to Oxides

Process	Definition	Example
Roasting	Strong heating of sulphide ores in excess air to form oxides	$2ZnS + 3O_2 \xrightarrow{\text{Heat}} 2ZnO + 2SO_2$
Calcination	Strong heating of carbonate ores in limited air to form oxides	$ZnCO_3 \xrightarrow{\text{Heat}} ZnO + CO_2$

Reduction of Oxides to Metals

• Once the metal oxide is formed, it is reduced to metal.
• Carbon (coke) is used as a reducing agent for moderately reactive metals.

Example: Extraction of Zinc from Zinc Oxide (ZnO)

$$ZnO + C \rightarrow Zn + CO$$

Displacement Reduction Method

• Highly reactive metals like Al, Na, Ca can also be used as reducing agents.
• These metals displace less reactive metals from their oxides.

Example: Extraction of Manganese (Mn) from Manganese Dioxide (MnO₂)

$$3MnO_2 + 4Al \rightarrow 3Mn + 2Al_2O_3 + \text{Heat}$$

• This reaction is highly exothermic, producing molten manganese.

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3.4.5 Extracting Metals Towards the Top of the Activity Series

• Highly reactive metals (K, Na, Ca, Mg, Al) have strong affinity for oxygen.
• They cannot be reduced by carbon because they form more stable oxides.
• These metals are extracted by electrolytic reduction of their molten salts.

Example: Extraction of Sodium from Sodium Chloride (NaCl)

At Cathode (Negative Electrode):

$$Na^+ + e^- \rightarrow Na$$

At Anode (Positive Electrode):

$$2Cl^- \rightarrow Cl_2 + 2e^-$$

• Sodium metal is deposited at the cathode.
• Chlorine gas is released at the anode.

Example: Extraction of Aluminium from Aluminium Oxide (Al₂O₃)

• Aluminium is extracted by the electrolysis of Al₂O₃ in molten cryolite.

At Cathode:

$$Al^{3+} + 3e^- \rightarrow Al$$

At Anode:

$$2O^{2-} \rightarrow O_2 + 4e^-$$

• Molten aluminium is deposited at the cathode.
• Oxygen gas is released at the anode.

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3.4.6 Refining of Metals

• The extracted metals are not pure and contain impurities.
• The most commonly used refining method is electrolytic refining.

Electrolytic Refining Process

1. Anode (Impure metal): The impure metal dissolves into the solution.
2. Cathode (Pure metal): The pure metal is deposited here.
3. Electrolyte: A solution of the metal salt (e.g., CuSO₄ for copper refining).

Example: Electrolytic Refining of Copper
At Anode:

$$Cu \rightarrow Cu^{2+} + 2e^-$$

At Cathode:

$$Cu^{2+} + 2e^- \rightarrow Cu$$

• Pure copper is deposited at the cathode.
• Impurities settle as anode mud.

Metals Purified by Electrolytic Refining

• Copper (Cu), Zinc (Zn), Tin (Sn), Nickel (Ni), Silver (Ag), Gold (Au).

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Summary of Metal Extraction Methods

Metal Reactivity	Examples	Extraction Method
Low Reactive Metals	Ag, Au, Cu, Hg	Heating in air
Medium Reactive Metals	Zn, Fe, Pb	Reduction with Carbon / Displacement Reaction
Highly Reactive Metals	Na, K, Ca, Mg, Al	Electrolysis of molten ores

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Conclusion

1. Metals low in the reactivity series can be extracted by simple heating.
2. Moderately reactive metals require roasting, calcination, and reduction.
3. Highly reactive metals must be extracted using electrolysis.
4. Refining is done through electrolytic refining to remove impurities.

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QUESTIONS & ANSWERS

1. Define the following terms:

(i) Mineral

Answer:

• A mineral is a naturally occurring substance found in the Earth's crust, containing metals or their compounds.
• Example: Bauxite (Al₂O₃·2H₂O) is a mineral of Aluminium.

(ii) Ore

Answer:

• An ore is a type of mineral that contains a high percentage of a metal and can be extracted profitably.
• Example: Hematite (Fe₂O₃) is an ore of Iron.

(iii) Gangue

Answer:

• Gangue refers to the unwanted impurities like sand, soil, and clay that are present in an ore.
• These impurities must be removed before metal extraction.

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2. Name two metals which are found in nature in the free state.

Answer:

• Gold (Au)
• Silver (Ag)
• These metals are very unreactive and do not easily combine with other elements.

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3. What chemical process is used for obtaining a metal from its oxide?

Answer:

• Metals are obtained from their oxides by reduction.
• The method used depends on the metal’s reactivity:

Metal Type	Method of Reduction	Example
Low Reactive Metals (Hg, Cu)	Simple heating	$2HgO \xrightarrow{\text{Heat}} 2Hg + O_2$
Moderately Reactive Metals (Zn, Fe, Pb)	Reduction with Carbon (C) or Carbon Monoxide (CO)	$ZnO + C \rightarrow Zn + CO$
Highly Reactive Metals (Na, Al, Mg, Ca)	Electrolysis of molten oxides	$Al_2O_3 \rightarrow 2Al + 3O_2$

Conclusion: The process of reduction depends on the metal’s reactivity in the activity series.

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3.5 Corrosion and Its Prevention

Corrosion is the gradual destruction of metals due to their reaction with air, water, and other environmental factors.

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3.5.1 What is Corrosion?

Corrosion occurs when metals react with oxygen, moisture, or other chemicals, forming undesirable compounds on their surface.

Examples of Corrosion:

1. Rusting of Iron (Fe)
   • When iron is exposed to moist air (oxygen + water), it forms brown flaky rust (iron oxide).
   • Reaction: $$4Fe + 3O_2 + 6H_2O → 4Fe(OH)_3$$ $$2Fe(OH)_3 → Fe_2O_3·xH_2O \text{ (Rust)}$$
2. Tarnishing of Silver (Ag)
   • Silver reacts with sulphur compounds in air to form black silver sulphide (Ag₂S).
   • Reaction: $$2Ag + H_2S → Ag_2S + H_2$$
3. Formation of Green Coating on Copper (Cu)
   • Copper reacts with moist carbon dioxide in air to form green basic copper carbonate.
   • Reaction: $$2Cu + H_2O + CO_2 + O_2 → Cu(OH)_2·CuCO_3$$

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Activity 3.14: Investigating Rusting of Iron

Experiment Setup:

1. Take three test tubes (A, B, C) and place clean iron nails in each.
2. Test Tube A: Contains water and air (uncovered).
3. Test Tube B: Contains boiled distilled water + oil (to prevent air from dissolving).
4. Test Tube C: Contains anhydrous calcium chloride (absorbs moisture, keeping air dry).
5. Leave for a few days and observe.

Observations & Conclusion:

Test Tube	Conditions	Observation	Inference
A	Air + Water	Iron nails rusted	Both air and water are required for rusting.
B	Water (No Air)	No rusting	Air is necessary for rusting.
C	Dry Air (No Water)	No rusting	Water is necessary for rusting.

• Rusting of iron requires both air (oxygen) and water (moisture).
• In the absence of either air or water, rusting does not occur.

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3.5.2 Prevention of Corrosion

Corrosion can be prevented using various methods, such as coating, galvanization, and alloying.

Methods to Prevent Rusting of Iron:

Method	Description
Painting	Coating with paint prevents direct contact with air and moisture.
Oiling/Greasing	A thin oil layer protects the metal surface from moisture and oxygen.
Galvanization	Coating iron/steel with zinc protects against rusting, even if the zinc layer is broken.
Chrome Plating	A thin chromium layer prevents iron from direct exposure to the environment.
Anodizing	A protective oxide layer is formed on metals like aluminium.
Alloying	Mixing iron with other metals (e.g., nickel, chromium) improves resistance to corrosion.

Why Does Galvanization Work?

• Zinc is more reactive than iron.
• When the zinc layer is damaged, zinc reacts instead of iron, protecting it from rusting (sacrificial protection).

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3.5.3 Alloying: A Method to Improve Metal Properties

• Alloying is the process of mixing two or more metals (or a metal with a non-metal) to improve properties.
• It enhances strength, hardness, corrosion resistance, and electrical properties.

Example: Iron and Its Alloys

• Pure iron is soft and weak.
• Adding carbon (0.05%) makes it stronger.
• Mixing iron with nickel and chromium forms stainless steel, which does not rust.

Examples of Alloys and Their Properties

Alloy	Composition	Properties	Uses
Stainless Steel	Iron (Fe) + Chromium (Cr) + Nickel (Ni)	Hard, does not rust	Cutlery, surgical instruments, machinery
Brass	Copper (Cu) + Zinc (Zn)	Corrosion-resistant, decorative	Electrical fittings, musical instruments
Bronze	Copper (Cu) + Tin (Sn)	Hard, corrosion-resistant	Statues, medals
Solder	Lead (Pb) + Tin (Sn)	Low melting point	Joining electrical wires
Amalgam	Mercury (Hg) + Other metals	Used in dental fillings	Dental fillings, electrical switches

Gold Alloy (Jewellery Gold)

• Pure gold (24-carat) is very soft.
• 22-carat gold is mixed with silver or copper to make it hard for jewellery.
• Composition: 22 parts gold + 2 parts silver/copper.

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3.5.4 The Iron Pillar of Delhi: A Wonder of Ancient Metallurgy

• The Iron Pillar near Qutub Minar (Delhi) was built over 1600 years ago.
• It has not rusted, despite being exposed to air and moisture.
• Ancient Indian metallurgists used a special iron-making technique that prevented rusting.
• Height: 8 meters, Weight: 6000 kg (6 tonnes).
• Scientists from around the world study its rust-resistant properties.

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Summary of Corrosion and Its Prevention

Topic	Key Points
Corrosion	Metals react with air, moisture, and other chemicals to form undesirable compounds.
Examples	Rusting of iron (Fe₂O₃·xH₂O), tarnishing of silver (Ag₂S), green coating on copper (Cu(OH)₂·CuCO₃).
Rusting Conditions	Requires both air (oxygen) and water (moisture).
Prevention Methods	Painting, greasing, galvanizing, anodizing, alloying.
Galvanization	Zinc coating protects iron from rusting (even if scratched).
Alloying	Mixing metals to improve strength, hardness, and corrosion resistance.
Examples of Alloys	Stainless steel (Fe + Cr + Ni), Brass (Cu + Zn), Bronze (Cu + Sn), Solder (Pb + Sn).
Iron Pillar of Delhi	1600-year-old structure, rust-resistant, studied worldwide.

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Conclusion

1. Corrosion is the gradual destruction of metals due to environmental reactions.
2. Rusting of iron requires both air and water.
3. Corrosion can be prevented by coatings (painting, greasing), galvanization, and alloying.
4. Alloying enhances metal properties, making them stronger and corrosion-resistant.
5. The Iron Pillar of Delhi is an example of ancient rust-resistant metallurgy.

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QUESTIONS & ANSWERS

1. Metallic oxides of zinc, magnesium, and copper were heated with the following metals: Zinc, Magnesium, and Copper.

• Which cases will show displacement reactions?

Answer:

• A more reactive metal can displace a less reactive metal from its oxide.
• Magnesium (Mg) is the most reactive, followed by Zinc (Zn), and then Copper (Cu).

Reaction	Expected Outcome	Reaction Type
Zinc + Zinc Oxide (ZnO)	❌ No reaction	Zinc cannot displace itself.
Magnesium + Zinc Oxide (ZnO)	✅ Displacement	$Mg + ZnO → MgO + Zn$
Copper + Zinc Oxide (ZnO)	❌ No reaction	Copper is less reactive than zinc.
Zinc + Magnesium Oxide (MgO)	❌ No reaction	Zinc is less reactive than magnesium.
Magnesium + Magnesium Oxide (MgO)	❌ No reaction	Magnesium cannot displace itself.
Copper + Magnesium Oxide (MgO)	❌ No reaction	Copper is less reactive than magnesium.
Zinc + Copper Oxide (CuO)	✅ Displacement	$Zn + CuO → ZnO + Cu$
Magnesium + Copper Oxide (CuO)	✅ Displacement	$Mg + CuO → MgO + Cu$
Copper + Copper Oxide (CuO)	❌ No reaction	Copper cannot displace itself.

• Displacement reactions occur when magnesium is added to zinc oxide and copper oxide.
• Zinc can displace copper from copper oxide but not magnesium from magnesium oxide.

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2. Which metals do not corrode easily?

Answer:

• Metals that are unreactive or form a protective oxide layer do not corrode easily.
• Examples:
  • Gold (Au) and Silver (Ag) (Very unreactive)
  • Platinum (Pt) (Highly resistant to corrosion)
  • Aluminium (Al) (Forms a protective Al₂O₃ oxide layer)
  • Stainless Steel (Alloy of Iron, Chromium, and Nickel)

Why?

• Gold, silver, and platinum are noble metals and do not react with air or water.
• Aluminium and stainless steel form a protective oxide layer that prevents further corrosion.

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3. What are Alloys?

Answer:

• An alloy is a homogeneous mixture of two or more metals or a metal and a non-metal.
• It is made by melting the primary metal and mixing it with other elements in fixed proportions.
• Alloys improve strength, corrosion resistance, and hardness.

Examples of Alloys:

Alloy	Composition	Properties	Uses
Stainless Steel	Iron (Fe) + Chromium (Cr) + Nickel (Ni)	Rust-resistant, strong	Kitchenware, surgical tools
Brass	Copper (Cu) + Zinc (Zn)	Hard, corrosion-resistant	Electrical fittings, musical instruments
Bronze	Copper (Cu) + Tin (Sn)	Hard, corrosion-resistant	Medals, statues
Solder	Lead (Pb) + Tin (Sn)	Low melting point	Joining electrical wires
Amalgam	Mercury (Hg) + Other Metals	Liquid at room temperature	Dental fillings, electrical switches

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EXERCISES

1. Which of the following pairs will give displacement reactions?

(a) NaCl solution and copper metal ❌ (Copper is less reactive than sodium)
(b) MgCl₂ solution and aluminium metal ❌ (Aluminium is less reactive than magnesium)
(c) FeSO₄ solution and silver metal ❌ (Silver is less reactive than iron)
(d) AgNO₃ solution and copper metal ✅ (Copper is more reactive than silver)

Correct Answer: (d) AgNO₃ solution and copper metal

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2. Which of the following methods is suitable for preventing an iron frying pan from rusting?

(a) Applying grease ✅
(b) Applying paint ✅
(c) Applying a coating of zinc ✅
(d) All of the above ✅

Correct Answer: (d) All of the above

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3. An element reacts with oxygen to give a compound with a high melting point. This compound is also soluble in water. The element is likely to be

(a) Calcium ✅ (Calcium forms calcium oxide, which dissolves in water to form calcium hydroxide)
(b) Carbon ❌
(c) Silicon ❌
(d) Iron ❌

Correct Answer: (a) Calcium

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4. Food cans are coated with tin and not with zinc because

(a) Zinc is costlier than tin ❌
(b) Zinc has a higher melting point than tin ❌
(c) Zinc is more reactive than tin ✅ (Zinc would react with acids in food, making it unsafe)
(d) Zinc is less reactive than tin ❌

Correct Answer: (c) Zinc is more reactive than tin

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5. You are given a hammer, a battery, a bulb, wires, and a switch.

(a) How could you use them to distinguish between samples of metals and non-metals?

• Hammer Test: Metals are malleable (flatten without breaking), whereas non-metals are brittle.
• Electricity Test: Metals conduct electricity (bulb glows), whereas non-metals do not.

(b) Assess the usefulness of these tests in distinguishing between metals and non-metals.

• Malleability Test: Effective for most metals and non-metals.
• Electric Conductivity Test: Good, but some non-metals (like graphite) can conduct electricity.

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6. What are amphoteric oxides? Give two examples of amphoteric oxides.

Answer: Amphoteric oxides are oxides that react with both acids and bases to form salts and water.

Examples:

1. Aluminium oxide (Al₂O₃)
2. Zinc oxide (ZnO)

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7. Name two metals which will displace hydrogen from dilute acids, and two metals which will not.

✅ Metals that displace hydrogen from acids: Zinc (Zn), Magnesium (Mg)
❌ Metals that do not displace hydrogen: Copper (Cu), Silver (Ag)

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8. In the electrolytic refining of a metal M, what would you take as the anode, the cathode, and the electrolyte?

• Anode: Impure metal M
• Cathode: Pure metal M
• Electrolyte: Aqueous solution of a salt of metal M

Example: Copper Refining

• Anode: Impure copper
• Cathode: Pure copper
• Electrolyte: Copper sulphate (CuSO₄) solution

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9. Pratyush took sulphur powder on a spatula and heated it.

(a) What will be the action of gas on
(i) Dry litmus paper? → No change
(ii) Moist litmus paper? → Turns blue litmus red (Sulphur dioxide is acidic)

(b) Write a balanced chemical equation for the reaction taking place.

$$S + O_2 → SO_2$$

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10. State two ways to prevent the rusting of iron.

1. Painting, oiling, or greasing
2. Galvanization (coating with zinc)

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11. What type of oxides are formed when non-metals combine with oxygen?

• Acidic oxides (e.g., CO₂, SO₂, NO₂)
• Neutral oxides (e.g., CO, N₂O)

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12. Give reasons:

(a) Platinum, gold, and silver are used to make jewellery.

• These metals do not corrode and have lustrous appearance.

(b) Sodium, potassium, and lithium are stored under oil.

• They are highly reactive and react violently with air and water.

(c) Aluminium is a highly reactive metal, yet it is used to make utensils for cooking.

• Aluminium forms a protective oxide layer (Al₂O₃), preventing further reaction.

(d) Carbonate and sulphide ores are usually converted into oxides during the process of extraction.

• Metal oxides are easier to reduce to pure metal than carbonates or sulphides.

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13. Why are tarnished copper vessels cleaned with lemon or tamarind juice?

• Copper reacts with moist air to form a green basic copper carbonate layer.
• Lemon/tamarind juice contains acids (citric/tartaric acid), which react with the copper carbonate and remove the tarnish.

$$CuCO_3 + 2HCl → CuCl_2 + H_2O + CO_2$$

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14. Differentiate between metal and non-metal on the basis of their chemical properties.

Property	Metals	Non-Metals
Reaction with oxygen	Forms basic oxides	Forms acidic oxides
Reaction with acids	Produces hydrogen gas	No reaction
Reaction with bases	No reaction (except amphoteric metals)	Forms salt and water
Conductivity	Good conductors of heat & electricity	Poor conductors (except graphite)

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15. The case of the fake goldsmith

• The man used Aqua Regia (a mixture of HCl and HNO₃ in 3:1 ratio).
• Aqua Regia dissolves gold, reducing its weight and making it shine like new.

$$Au + HNO_3 + HCl → \text{Gold dissolves}$$

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16. Why is copper used to make hot water tanks instead of steel (iron alloy)?

1. Copper is corrosion-resistant, whereas steel (iron alloy) can rust.
2. Copper is a better conductor of heat, making water heat up faster.
3. Copper does not react with water, whereas iron may react to form rust.

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