Curriculum
Science Grade 10
Grade 10 Science
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Chapter 1. Chemical Reactions
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Chapter 2: Acids, Bases and Salts
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Chapter 3: Metals and Non-metals
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Chapter 4: Carbon and its Compounds
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Chapter 5: Life Processes
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Chapter 6: Control and Coordination
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Chapter 7: How do Organisms Reproduce?
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Chapter 8: Heredity
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Chapter 9: Light Reflection and Refraction
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Chapter 10: The Human Eye and the Colourful World
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Chapter 11: Electricity
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Chapter 12: Magnetic Effects of Electric Current
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Chapter 13: Our Environment
Interactive Worksheets
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Chapter 3: Metals and Non-metals
Grade 10 Science · Chapter 3
⚗️ Metals and Non-metals
Understand the physical & chemical properties of metals and non-metals, ionic bonding, the reactivity series, metallurgy, corrosion, and alloys — explained simply for Grade 10 students.
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Covers Physical & Chemical Properties |
Key Topics Reactivity · Ionic Bonds · Metallurgy |
Includes 10+ Examples · Practice Sets A–D |
📋 Table of Contents
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▶ 1. Introduction ▶ 2. Physical Properties of Metals ▶ 3. Physical Properties of Non-metals ▶ 4. Chemical Properties of Metals ▶ 5. Reactivity Series ▶ 6. How Metals & Non-metals React (Ionic Bonding) |
▶ 7. Occurrence & Extraction of Metals ▶ 8. Corrosion & Alloys ▶ 9. Worked Examples (10+) ▶ 10. Practice Sets A–D with Answers ▶ 11. Chapter Summary ▶ 12. Exam Quick-Check (8 Points) |
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1. Introduction |
Elements around us are grouped into two large families — metals and non-metals. This grouping is based on how they look, conduct electricity and heat, and how they react chemically.
About 80% of all known elements are metals. Think of iron nails, copper wires, aluminium foil — all metals. Oxygen in air, carbon in pencils, sulphur in matchsticks — these are non-metals.
💡 Did You Know?
Mercury is the only metal that is liquid at room temperature. Bromine is the only non-metal that is liquid at room temperature. All other non-metals are either solid or gas.
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2. Physical Properties of Metals |
Key Physical Properties of Metals
📊 Diagram: Malleability vs Ductility
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MALLEABILITY
Metal beaten flat (e.g. aluminium foil) |
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DUCTILITY
Metal drawn into wire (e.g. copper wire) |
⚠️ Important Exceptions to Remember
• Mercury — the only metal that is liquid at room temperature
• Gallium & Caesium — melt at just above room temperature (they melt on your palm!)
• Sodium & Potassium — so soft they can be cut with a knife; stored in kerosene oil
• Iodine — a non-metal that IS lustrous (shiny) — exception to non-metal rule
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3. Physical Properties of Non-metals |
Non-metals are generally the opposite of metals. They are brittle, dull, and poor conductors — with a few important exceptions.
📊 Diagram: Allotropes of Carbon
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DIAMOND
3D tetrahedral C-C network |
vs |
GRAPHITE Layered flat sheets of carbon |
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4. Chemical Properties of Metals |
4.1 Metals + Oxygen → Metal Oxides (Basic in Nature)
Almost all metals react with oxygen to form metal oxides. Most metal oxides are basic in nature — they turn red litmus paper blue and react with acids.
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Copper + Oxygen 2Cu + O₂ → 2CuO Black copper(II) oxide forms |
Aluminium + Oxygen 4Al + 3O₂ → 2Al₂O₃ White aluminium oxide forms |
🌟 Amphoteric Oxides
Some metal oxides react with BOTH acids AND bases — they are called amphoteric oxides.
Example — Aluminium oxide (Al₂O₃):
Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O (acts as a base, reacts with acid)
Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (acts as an acid, reacts with base)
Zinc oxide (ZnO) is also amphoteric.
4.2 Metals + Water → Metal Hydroxide + Hydrogen Gas
4.3 Metals + Dilute Acids → Salt + Hydrogen Gas
Metals above hydrogen in the reactivity series react with dilute acids to produce a salt and hydrogen gas. Reactivity order: Mg > Al > Zn > Fe. Copper does NOT react with dilute HCl or H₂SO₄.
💡 Aqua Regia — “Royal Water”
A fresh mixture of concentrated HCl and HNO₃ in the ratio 3:1 is called Aqua Regia. It is one of the very few substances that can dissolve gold and platinum — metals that no single acid can attack. Highly corrosive!
4.4 Metals + Salt Solutions → Displacement Reactions
A more reactive metal displaces a less reactive metal from its salt solution. This is the clearest test of reactivity.
Fe + CuSO₄ → FeSO₄ + Cu
Iron (more reactive) displaces copper from copper sulphate solution. The blue solution becomes pale green and a reddish copper deposit forms on the iron nail.
General rule: Metal A + Salt solution of B → Salt of A + Metal B (only if A is more reactive than B)
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5. The Reactivity Series |
The reactivity series lists metals from most reactive (top) to least reactive (bottom). It predicts which metals can displace others from salt solutions and which extraction method to use.
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MOST REACTIVE ↓ |
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🔑 Exam Tip — Reactivity Series Memory Trick
K Na Ca Mg Al Zn Fe Pb [H] Cu Hg Ag Au
Remember: “King Narendra Can Make A Zinc Factory, Please Help Create More Silver-Gold”
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6. How Metals and Non-metals React — Ionic Bonding |
When a metal reacts with a non-metal, the metal atom loses electrons to form a positive ion (cation), and the non-metal atom gains those electrons to form a negative ion (anion). The opposite charges attract strongly — this is an ionic bond.
Noble gases have a completely filled outer shell and are very stable. Atoms of other elements try to achieve this stable configuration by losing or gaining electrons.
Electronic Configurations of Key Elements
📊 Diagram: Formation of Sodium Chloride (NaCl)
Properties of Ionic Compounds
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7. Occurrence and Extraction of Metals |
Mineral: An element or compound found naturally in the earth’s crust.
Ore: A mineral from which a metal can be profitably extracted.
Gangue: Impurities (soil, sand etc.) mixed with the ore.
📊 Diagram: Extraction Method Based on Reactivity
🔥 The Thermit Reaction (Very Important!)
Aluminium (more reactive) displaces iron from iron(III) oxide. The reaction produces so much heat that iron comes out in molten form. This is used to weld railway tracks:
Fe₂O₃ + 2Al → 2Fe(l) + Al₂O₃ + Heat
Electrolytic Refining of Metals
Used to purify impure metals (e.g. copper). The impure metal is the anode, pure metal is the cathode, salt solution of the metal is the electrolyte. Pure metal deposits on the cathode; impurities collect at the bottom as anode mud.
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8. Corrosion and Alloys |
Types of Corrosion
📊 Diagram: Conditions Required for Iron to Rust
✅ Conclusion: Iron rusts ONLY when BOTH oxygen AND water are present together.
Prevention of Corrosion
Common Alloys and Their Uses
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9. Worked Examples |
EXAMPLE 1 Reaction of Sodium with Water
Question: Write the equation for sodium reacting with water and explain what you observe.
👁️ Show Solution
Equation: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + Heat
Observations: Sodium floats on water and moves about rapidly (reaction is exothermic). Hydrogen gas is released which may catch fire due to the heat produced. The solution formed is strongly alkaline (NaOH).
EXAMPLE 2 Identify the Amphoteric Oxide
Question: What is an amphoteric oxide? Give equations to show how aluminium oxide acts as both acid and base.
👁️ Show Solution
An amphoteric oxide reacts with both acids and bases to produce salt and water.
With acid (acts as base): Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
With base (acts as acid): Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O
EXAMPLE 3 Displacement Reaction
Question: When an iron nail is placed in copper sulphate solution, what happens? Write the equation.
👁️ Show Solution
Iron is more reactive than copper, so it displaces copper from the solution.
Equation: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
Observation: The blue colour of the copper sulphate solution fades and a reddish-brown copper deposit appears on the iron nail.
EXAMPLE 4 Formation of Ionic Compound MgCl₂
Question: Explain how magnesium chloride is formed from magnesium and chlorine.
👁️ Show Solution
Magnesium has 2 electrons in its outer shell (2,8,2) and loses both to become Mg²⁺.
Mg → Mg²⁺ + 2e⁻
Each chlorine atom (2,8,7) gains one electron to become Cl⁻. Two Cl atoms are needed.
2Cl + 2e⁻ → 2Cl⁻
Result: Mg²⁺ + 2Cl⁻ → MgCl₂ (one Mg²⁺ attracts two Cl⁻ ions)
EXAMPLE 5 Thermit Reaction
Question: Write the thermit reaction and state its industrial use.
👁️ Show Solution
Reaction: Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat
Use: The heat produced is so intense (above 2500°C) that iron comes out in molten form. This is used to weld cracked railway tracks on-site without any external power supply.
EXAMPLE 6 Why Sodium is Stored in Kerosene
Question: Why is sodium metal always stored under kerosene oil?
👁️ Show Solution
Sodium is extremely reactive. It reacts violently with air (oxygen) and even with moisture in the air or water. The reaction is so exothermic that it can catch fire spontaneously. Kerosene oil does not react with sodium and forms a protective layer that keeps sodium away from air and moisture.
EXAMPLE 7 Extraction of Mercury from Cinnabar
Question: How is mercury extracted from cinnabar (HgS)?
👁️ Show Solution
Step 1 — Roasting: 2HgS(s) + 3O₂(g) → 2HgO(s) + 2SO₂(g)
Step 2 — Reduction by heating: 2HgO(s) → 2Hg(l) + O₂(g)
Mercury is low in the reactivity series so its oxide breaks down by heating alone — no reducing agent needed.
EXAMPLE 8 Electrolytic Refining
Question: In electrolytic refining of copper, what is the anode, cathode, and electrolyte?
👁️ Show Solution
Anode: Impure copper (the impure metal to be refined)
Cathode: A thin strip of pure copper
Electrolyte: Acidified copper sulphate solution (CuSO₄)
Process: Impure copper dissolves from anode into solution. Pure copper deposits on cathode. Insoluble impurities collect as anode mud at bottom.
EXAMPLE 9 Galvanisation
Question: What is galvanisation and why does it protect iron even if the zinc coating is scratched?
👁️ Show Solution
Galvanisation is coating iron or steel with a thin layer of zinc. Even if the zinc coating is scratched or broken, the iron beneath is still protected. This is because zinc is higher in the reactivity series than iron — so zinc corrodes first, sacrificing itself to protect the iron underneath. This is called sacrificial protection.
EXAMPLE 10 Why Aluminium is Used for Cooking Utensils
Question: Aluminium is a highly reactive metal, yet it is used for making cooking utensils. Why?
👁️ Show Solution
Although aluminium is reactive, it quickly forms a thin, tough layer of aluminium oxide (Al₂O₃) on its surface when exposed to air. This oxide layer is very stable and acts as a protective coating that prevents further reaction with air or water. This is why aluminium doesn’t corrode like iron and is safe to use for cookware.
EXAMPLE 11 Gold Bangle Mystery (Real-world Detective Problem)
Question: A goldsmith dipped gold bangles in a solution and they became shiny, but lost weight drastically. What solution was used?
👁️ Show Solution
The solution used was Aqua Regia — a mixture of concentrated HCl and HNO₃ in the ratio 3:1. This is one of the only substances that dissolves gold. The outer layer of gold dissolved, making them temporarily shiny (fresh metal exposed), but the total weight of gold reduced significantly. The trickster fraudulently dissolved the gold!
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10. Practice Sets A–D |
📘 Practice Set A — Multiple Choice
1. Which of the following is NOT a physical property of metals?
(a) Lustrous (b) Malleable (c) Acidic oxide (d) Sonorous
2. Which pair will give a displacement reaction?
(a) CuSO₄ + Ag (b) ZnSO₄ + Fe (c) FeSO₄ + Zn (d) AgNO₃ + Cu
3. Food cans are coated with tin, not zinc, because:
(a) Tin is cheaper (b) Zinc is more reactive than tin (c) Tin is lighter (d) Zinc melts easily
4. Which metal is liquid at room temperature?
(a) Gallium (b) Mercury (c) Caesium (d) Sodium
✅ Show Answers
1. (c) Acidic oxide 2. (c) and (d) 3. (b) 4. (b) Mercury
📗 Practice Set B — Short Answer
1. Why do ionic compounds conduct electricity in solution but not in solid state?
2. What is calcination? Give one example with equation.
3. Differentiate between roasting and calcination.
4. Name the alloy used for welding electrical wires and give its composition.
✅ Show Answers
1. In solid state ions are fixed in a rigid lattice and cannot move. When dissolved in water, ions are free to move toward electrodes and carry charge.
2. Calcination = heating carbonate ores in limited air. E.g.: ZnCO₃ → ZnO + CO₂
3. Roasting = heating sulphide ore in excess air. Calcination = heating carbonate ore in limited air. Both convert ore to metal oxide.
4. Solder — Lead (Pb) + Tin (Sn). Has a low melting point, ideal for welding without damaging electronic components.
📙 Practice Set C — Equations and Reactions
1. Write the balanced equation for the reaction of iron with steam.
2. Write equations showing how Na₂O and K₂O dissolve in water.
3. Write the equation for the roasting of zinc sulphide.
4. What happens when zinc is added to iron sulphate solution? Write the equation.
✅ Show Answers
1. 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
2. Na₂O + H₂O → 2NaOH K₂O + H₂O → 2KOH
3. 2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)
4. Zinc is more reactive than iron. Zn + FeSO₄ → ZnSO₄ + Fe. The pale green colour of the solution fades as ZnSO₄ forms. Iron deposits.
📕 Practice Set D — Reasoning and Application
1. Why is platinum used in jewellery even though it is expensive?
2. Why are copper and not steel tanks used to store hot water?
3. Copper vessels turn green over time. Explain this with a chemical reaction.
4. Why is carbon unable to reduce the oxides of sodium, calcium and aluminium?
✅ Show Answers
1. Platinum, gold, and silver are used for jewellery because they are very unreactive (at the bottom of the reactivity series). They do not corrode, tarnish, or react with air, moisture or acids — making them long-lasting and visually appealing.
2. Copper does not react with hot water. Steel (iron alloy) would rust due to water + oxygen. Copper is also a better conductor of heat, making it ideal for hot water systems.
3. Copper reacts with CO₂ and moisture in air: 2Cu + CO₂ + H₂O → Cu(OH)₂·CuCO₃ (basic copper carbonate — green patina)
4. Na, Ca, Mg, Al are very high in the reactivity series. They have a stronger affinity for oxygen than carbon does. Carbon cannot “pull” oxygen away from these metal oxides. Electrolysis is the only way to extract them.
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11. Chapter Summary |
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Physical Properties Metals: lustrous, malleable, ductile, sonorous, good conductors Chemical Properties Metals form basic oxides. Non-metals form acidic oxides. Amphoteric oxides: Al₂O₃, ZnO react with both. Ionic Bonding Metal loses e⁻ → cation. Non-metal gains e⁻ → anion. Opposite ions attract = ionic compound. High MP, brittle, conducts in solution. |
Reactivity Series K>Na>Ca>Mg>Al>Zn>Fe>Pb>[H]>Cu>Hg>Ag>Au. More reactive displaces less reactive from salt solution. Metallurgy High reactivity → electrolysis. Medium → carbon reduction. Low → heating alone. Refining → electrolytic refining. Corrosion & Alloys Rusting needs both O₂ and water. Prevented by galvanising, painting, alloying. Alloys: Brass, Bronze, Stainless Steel, Solder. |
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12. Exam Quick-Check — 8 Must-Know Points |
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✅ Point 1 Iron rusts only when BOTH O₂ and H₂O are present. Remove either and rusting stops. |
✅ Point 2 Graphite is the only non-metal that conducts electricity. Iodine is the only lustrous non-metal. |
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✅ Point 3 Amphoteric oxides (Al₂O₃, ZnO) react with both acids AND bases — neither fully acidic nor basic. |
✅ Point 4 Ionic compounds conduct electricity in molten state or solution — NOT in solid state (ions are fixed). |
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✅ Point 5 Thermit reaction — Al reduces Fe₂O₃ to molten iron. Used to weld railway tracks on-site. |
✅ Point 6 Aqua Regia (3:1 HCl:HNO₃) dissolves gold and platinum — even though neither acid alone can do so. |
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✅ Point 7 Reactivity order: K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au |
✅ Point 8 Alloys have LOWER electrical conductivity and melting point than their parent pure metals. |
This Grade 10 Science Chapter 3 guide on Metals and Non-metals covers all key topics for board exam preparation including physical properties, chemical reactions, the reactivity series, ionic bonding, metallurgy, corrosion, and alloys. Original educational content for Grade 10 students aligned with standard science curriculum. Topics include: lustre, malleability, ductility, conductivity, amphoteric oxides, displacement reactions, ionic compounds, electrolytic refining, galvanisation, thermit reaction, and prevention of corrosion. Practice questions with answers and worked examples included for complete exam readiness.