Metals and Non-metals consistently contributes 10–12 marks in CBSE Boards. The reactivity series, ionic bond formation, and extraction of metals (metallurgy) are the most-tested topics. Properties of metals vs non-metals is a guaranteed 1-mark MCQ. NTSE includes activity series reasoning and corrosion prevention questions.
Key Concept Highlights
Physical Properties of Metals and Non-metals
Chemical Properties of Metals
Reactions with Oxygen, Water, Acids, and Other Metal Salts
Reactivity Series of Metals
Ionic Bond Formation (Electron Transfer)
Properties of Ionic Compounds
Occurrence of Metals (Minerals and Ores)
Extraction of Metals (Low, Medium, High Reactivity)
Memorise the reactivity series K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au — it underpins 40% of questions. Ionic compound properties follow a standard list (hard, high melting point, conduct in solution). Extraction method depends on position in reactivity series: low = reduction with C; medium = electrolysis; high = thermite. Time investment: 4 days.
MetalsNon-metalsReactivity SeriesIonic CompoundsChemical Properties of MetalsCorrosionExtraction of MetalsElectrolysisAlloysCBSE Class 10NCERT Science
🗒️ Defiition
Metals are elements that readily lose electrons to form positive ions (cations) and exhibit properties
such as high electrical and thermal conductivity, malleability, ductility, and metallic lustre.
From a chemical perspective, metals are electropositive elements that tend to form
ionic compounds by donating electrons.
💡 Concept
Metals have 1–3 electrons in their valence shell. Due to low ionisation energy, they
lose electrons easily to form cations.
🎨 SVG Diagram
🏷️ Properties
Physical Properties of Metals
Physical Properties of Metals
Conductivity
Free electrons allow metals to conduct electricity and heat efficiently.
Malleability
Layers of atoms slide without breaking due to metallic bonding.
Ductility
Metals can be drawn into wires (e.g., copper wires).
Lustre
Metals reflect light due to free electrons.
High Density & Strength
Atoms are closely packed in metallic lattice.
High Melting & Boiling Points
Strong metallic bonds require high energy to break
Sonorous
Produce sound when struck (used in bells)
Exceptions (Very Important for Boards)
Mercury is liquid at room temperature.
Sodium and Potassium are soft and can be cut with a knife.
Gallium melts in hand (low melting point).
🏷️ Properties
Chemical Properties of Metals
Chemical Properties of Metals
Reaction with Oxygen:
Metal + Oxygen → Metal Oxide
\[ \mathrm{4Al + 3O_2 → 2Al_2O_3} \]
Reaction with Water
Metal + Water → Metal Hydroxide + Hydrogen
\[ \small\mathrm{2Na + 2H_2O → 2NaOH + H_2} \]
Reaction with Acids
Metal + Acid → Salt + Hydrogen
\[ \mathrm{Zn + 2HCl → ZnCl_2 + H_2} \]
📌 Note
Reactivity of Metals
Reactivity depends on the tendency to lose electrons. Metals higher in the reactivity series displace
those below them.
📌
Note
This concept is critical for understanding displacement reactions and extraction of metals.
Non-metals are elements that tend to gain electrons to form negative ions (anions) and generally
exhibit poor conductivity, low density, and lack of metallic lustre.
Chemically, non-metals are electronegative elements that prefer sharing or gaining
electrons to achieve stable electronic configuration.
💡 Concept
Electronic Basis of Non-Metallic Nature
Non-metals usually have 4–8 valence electrons. Due to high electronegativity, they
attract electrons rather than losing them.
This tendency explains why non-metals form anions or share electrons in covalent bonds.
🏷️ Properties
Physical Properties of Non-Metals
Physical Properties of Non-Metals
Low Density
Generally lighter due to loosely packed atoms.
Non-Lustrous
Lack shiny appearance (except iodine).
Brittle Nature
Break easily when hammered (not malleable).
Poor Conductivity
No free electrons for conduction.
Low Melting & Boiling Points
Weak intermolecular forces.
Non-Sonorous
Do not produce sound when struck.
Exceptions (Very Important for Exams)
Graphite conducts electricity.
Iodine is lustrous.
Diamond is extremely hard (hardest natural substance).
🏷️ Properties
Chemical Properties of Non-Metals
Chemical Properties of Non-Metals
Electron Gain (Reduction)
—
\[ \mathrm{X + e^- → X^- } \]
Reaction with Oxygen
Non-metal + Oxygen → Acidic Oxide
\[ \mathrm{C + O_2 → CO_2} \]
Reaction with Hydrogen
orms covalent hydrides
\[ \mathrm{H_2 + Cl_2 → 2HCl} \]
Reaction with Metals
Forms ionic compounds
\[ \mathrm{Na + Cl → NaCl} \]
📌 Note
Covalent Bonding in Non-Metals
Non-metals achieve stability by sharing electrons, forming covalent bonds.
1
Example
Example: Hydrogen molecule \( H_2 \) formed by sharing one electron each.
Non-metals are insulators because their electrons are tightly held in covalent bonds. Unlike metals, they lack a mobile "sea" of delocalised electrons required to carry an electric current.
Why are non-metals brittle?
Non-metals lack the non-directional metallic bonding that allows layers of atoms to slide over each other. Under stress, their rigid covalent structures simply snap or shatter rather than deforming.
⚡ Exam Tip
Use keywords: electronegative, anion, covalent bond.
Always mention electron gain in definitions.
Quote exceptions (graphite, iodine, diamond).
Write balanced chemical equations properly.
⚠️ Warning
Common Mistakes
Confusing ionic and covalent bonding.
Ignoring exceptions like graphite.
Writing incorrect formulas for oxides.
📋 Case Study
Carbon exists as both diamond and graphite with very different properties.
Question: Explain the reason.
Solution:
Due to different atomic arrangements (allotropes), graphite has free electrons (conductive) while
diamond has a rigid 3D structure (non-conductive and very hard).
🌟 Importance
Frequently asked in comparison-based questions.
Key concept for bonding and compounds.
Important in case-study and assertion-reason questions.
✏️ Example
Examples of non-metals
Oxygen, Nitrogen, Carbon, Sulphur, Chlorine.
🧫
Difference between Metals and Non-Metals
📌 Note
Understanding the differences between metals and non-metals is essential for mastering chemical bonding,
reactivity, and material properties. This comparison is frequently tested in CBSE board exams in
descriptive, assertion-reason, and case-study formats.
📊 Comparison Table
Difference between Metals and Non-Metals
Difference between Metals and Non-Metals
metals
Non-Metals
Mostly solids at room temperature (exception: mercury)
Exist in all three states (solid, liquid, gas)
Generally hard (exceptions: sodium, potassium are soft)
Generally soft (exception: diamond is extremely hard)
Malleable and ductile
Brittle and non-ductile
Lustrous (shiny surface)
Non-lustrous (exception: iodine)
Good conductors of heat and electricity
Poor conductors (exception: graphite)
Electropositive (lose electrons easily)
Electronegative (gain electrons)
Form basic or amphoteric oxides
Form acidic or neutral oxides
High melting and boiling points
Low melting and boiling points
High density
Low density
Sonorous (produce sound when struck)
Non-sonorous
Form ionic compounds
Form covalent compounds
🎨 SVG Diagram
Concept Map for Quick Revision
🔢 Formula
Key Electron Transfer Concept
Metals: \[ M → M^{n+} + ne^- \]
Non-metals: \[ X + e^- → X^- \]
✏️ Example
Why do metals form cations while non-metals form anions?
Metals have 1 to 3 electrons in their valence shells and low ionisation energy, making it easier to lose electrons and achieve a stable octet, forming positive cations. Non-metals have 4 to 7 electrons and high electronegativity, leading them to gain electrons and form negative anions.
Why are metals good conductors but non-metals are not?
Metals possess a "sea of delocalised electrons" that are free to move through the crystal lattice and carry electric charge. In contrast, non-metals hold their electrons tightly in covalent bonds, leaving no mobile charge carriers available for conduction.
Use keywords: electropositive, electronegative, ionic, covalent.
Draw comparison in tabular form for 3–5 mark questions.
Link properties with electron behavior for higher marks.
⚠️ Warning
Common Mistakes
Writing “all metals are hard” (incorrect).
Ignoring exceptions like graphite conductivity.
Confusing ionic vs covalent bonding.
📋 Case Study
A material is shiny, conducts electricity, and can be beaten into thin sheets.
Question: Classify it and justify.
Solution:
It is a metal because it shows lustre, conductivity, and malleability.
🌟 Importance
Most frequently asked 3–5 mark question.
Foundation for chemical bonding and reactions.
Helps in solving case-study and assertion-reason questions.
🧫
Exceptions in Properties of Metals and Non-Metals
🧠 Remember
Important Exceptions with Scientific Reason
Mercury (Hg)
The only metal that exists as a liquid at room temperature. Reason: Weak metallic bonding due to larger atomic size
reduces intermolecular attraction.
Gallium (Ga) and Caesium (Cs)
Gallium (Ga) and Caesium (Cs): Metals with very low melting points; can melt on
palm. Reason: Weak metallic bonds due to loosely held valence
electrons.
Iodine (I)
Iodine (I): A non-metal that shows lustre. Reason: Its crystal structure reflects light like metals.: A non-metal that shows lustre. Reason: Its crystal structure reflects light like metals.
Graphite (C): A non-metal that conducts electricity. Reason: Presence of free electrons between carbon layers.
Alkali Metals (Li, Na, K)
Alkali Metals (Li, Na, K): Very soft, low density, low melting points. Reason: Only one valence electron → weak metallic bonding.
🎨 SVG Diagram
✏️ Example
Why does graphite conduct electricity?
In graphite, each carbon atom is bonded to only three others, leaving one free valence electron per atom. These electrons become delocalised between the hexagonal layers, allowing them to move and conduct electricity.
Why is diamond extremely hard?
Diamond has a rigid 3D tetrahedral structure where each carbon atom is strongly bonded to four others by covalent bonds. This lack of free space or sliding layers makes it the hardest known natural substance.
Why are alkali metals (Li, Na, K) soft?
Alkali metals have large atomic sizes and only one valence electron, resulting in weak metallic bonding. This weak attraction between the kernels and the electron sea allows them to be easily cut with a knife.
⚡ Exam Tip
Always write reason with exception for full marks.
Memorize at least 5–6 key exceptions.
Link exception with bonding concept (metallic or covalent).
Used frequently in assertion-reason questions.
⚠️ Warning
Common Mistakes
Writing only exception without explanation.
Confusing graphite with diamond properties.
Ignoring alkali metals' softness.
📋 Case Study
A student observes that carbon conducts electricity in one form but not in another.
Question: Explain this behavior.
Solution:
Graphite conducts electricity due to free electrons, whereas diamond does not because all electrons
are involved in strong covalent bonds.
🌟 Importance
Frequently asked in 1–3 mark conceptual questions.
Strengthens understanding beyond rote learning.
Helps in tricky MCQs and case-based questions.
🧫
Chemical Properties of Metals
📌 Note
Reaction with Oxygen
Formation of metal oxide
Metals react with oxygen to form metal oxides, which are generally basic in nature.
Some metal oxides react with both acids and bases. These are called amphoteric oxides.
Al2O3 + 6HCl → 2AlCl3 + 3H2O
Al2O3 + 2NaOH → 2NaAlO2 + H2O
Formation of Alkalis
Some metal oxides dissolve in water to form alkalis (strong bases).
Na2O + H2O → 2NaOH
K2O + H2O → 2KOH
Highly reactive metals like sodium and potassium
react vigorously with oxygen and are stored in kerosene.
📌 Note
Reaction with Dilute Acids
Metals react with dilute acids to produce salt and hydrogen gas.
\[\scriptsize\color{orange}\boxed{\mathrm{Metal + Dilute Acid \rightarrow Salt + H_2}}\]
\[
\mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2}
\]
Important Concept: Only metals above hydrogen in the reactivity series can displace hydrogen.
👁️
Observation
Metals like Cu, Ag, Au do not react with dilute acids.
Nitric acid (HNO₃) usually does not produce H₂ gas due to its oxidising nature.
📌 Note
Reaction with Water</h3>
Metals react with water to produce metal hydroxide/oxide and hydrogen gas.
\[
\color{orange}\boxed{
\begin{aligned}
\mathrm{Metal + Water} &\rightarrow \mathrm{Metal\ hydroxide + H_2}
\end{aligned}}
\]
🗂️ Types / Category
Types of Reactions
❄️
Cold Water
Na, K, Ca react vigorously; Hydrogen gas evolved may catch fire.
🔥
Hot Water
Mg does not react with cold water; it reacts with hot water to form Mg(OH)₂.
💨
Steam
Al, Zn, Fe do not react with liquid water; they react only with steam to form Metal Oxides.
Metals are electropositive; they donate electrons to oxygen to form ionic oxides. When dissolved in water, these oxides produce hydroxyl (OH⁻) ions, which gives them their basic character (e.g., $Na_2O + H_2O \rightarrow 2NaOH$).
Why is aluminium oxide ($Al_2O_3$) called an amphoteric oxide?
It is called amphoteric because it shows both acidic and basic behaviour. It reacts with acids to form salt and water (acting as a base) and also reacts with strong bases like $NaOH$ to form sodium aluminate and water (acting as an acid).
Why do metals like Copper, Gold, and Silver not react with dilute acids?
These metals are positioned below Hydrogen in the reactivity series. Since they are less reactive than hydrogen, they cannot displace hydrogen ions from dilute acids to evolve $H_2$ gas.
A metal X displaces copper from CuSO₄ solution but does not react with water.
Question: Identify its reactivity.
Solution:
Metal X is more reactive than Copper (as it successfully displaces it) but less reactive than Magnesium (as it fails to react with water). This places X in the middle of the reactivity series, likely being a metal like Iron (Fe) or Lead (Pb)..
🧫
Reactivity Series of Metals
📘 Definition
The reactivity series is an arrangement of metals in decreasing order of their tendency to lose electrons
and form positive ions. It helps predict chemical behaviour such as displacement reactions, corrosion,
and extraction methods.
💡 Concept
Core Chemical Principle
Reactivity depends on how easily a metal loses electrons:
Zinc is more reactive than copper, so it displaces copper from its solution.
📌 Note
Importance in Extraction of Metals
Highly reactive metals (K, Na, Ca): Extracted by electrolysis
Moderately reactive (Zn, Fe): Extracted by reduction with carbon
Least reactive (Au, Ag): Found in native state
🔤 Mnemonic
Please Send Cats, Monkeys And Zebras In Lead Cages Securely Guarded.
✏️ Example
Why can Zinc displace Copper from its sulphate solution, but Copper cannot displace Zinc?
Zinc is more reactive than Copper and is placed higher in the reactivity series. This allows Zinc to lose electrons more easily and displace the less reactive Copper ions from the solution. Copper, being lower in the series, cannot provide enough energy to displace Zinc.
Why are metals like Gold and Silver found in the native (free) state in nature?
Gold and Silver are least reactive (noble metals) and are placed at the very bottom of the reactivity series. They do not react with atmospheric oxygen, moisture, or dilute acids, allowing them to remain chemically stable in their metallic form over millions of years.
⚡ Exam Tip
Memorize series order (frequently asked).
Always justify using position in series.
Link with displacement and extraction.
⚠️ Warning
Common Mistakes
Reversing order of metals.
Forgetting hydrogen position (important reference).
Not explaining displacement properly.
📋 Case Study
A metal X displaces iron from FeSO₄ but does not react with sodium hydroxide.
Question: Where is X placed in the reactivity series?
Solution:
X is more reactive than iron but less reactive than sodium.
🧫
Reaction of Metals and Non-Metals
📘 Definition
Definition of Ionic Compounds
When a metal reacts with a non-metal, electrons are transferred from the metal to the non-metal,
resulting in the formation of ionic (electrovalent) compounds.
Metals form cations and non-metals form anions, which attract each other
due to electrostatic force.
Why do ionic compounds conduct electricity in molten or aqueous state?
In the solid state, ions are held in fixed positions by strong electrostatic forces. In the molten or aqueous state, this rigid lattice breaks down, allowing ions to become mobile charge carriers that can move freely toward electrodes to conduct electricity.
Why are ionic compounds brittle?
When a mechanical force is applied, the layers of the crystal lattice shift slightly. This brings ions of the same charge next to each other, leading to strong electrostatic repulsion that shatters the crystal structure.
⚡ Exam Tip
Always write electron transfer steps.
Explain properties with reason.
Use keywords: lattice, electrostatic force, ions.
⚠️ Warning
Common Mistakes
Writing molecule instead of lattice.
Not explaining conductivity condition.
Confusing ionic with covalent bonding.
📋 Case Study
A compound does not conduct electricity in solid state but conducts in molten form.
Question: Identify the type of compound.
Solution:
It is an ionic compound because ions are fixed in solid state but free in molten state.
🌟 Importance
Why This Topic is Important
Core concept for bonding and compound formation.
Frequently asked in 3–5 mark questions.
Directly linked to properties and applications.
🧫
Occurrence and Extraction of Metals
📌 Note
Occurrence of Metals in Nature
Metals occur in the Earth's crust either in free (native) state or in
combined state as compounds called ores.
Least reactive metals (Au, Ag, Pt): Found in free state
Moderately reactive metals (Zn, Fe, Pb): Found as oxides, sulphides, carbonates
Highly reactive metals (K, Na, Ca): Always found in combined state
Oxygen is highly abundant and reactive, so most metals occur as oxides.
📌 Note
Important Terms
Ore: A mineral from which a metal can be extracted economically.
Gangue: Unwanted impurities like sand, clay present in ores.
🎨 SVG Diagram
Steps in Extraction of Metals
📌 Note
Extraction of Low Reactivity Metals
Metals like Hg and Cu are extracted by heating alone (self-reduction).
Why are highly reactive metals (like Na, Mg, Al) extracted by electrolytic reduction?
These metals have a higher affinity for oxygen than carbon does. Therefore, carbon cannot act as a reducing agent to displace oxygen from their oxides. They are instead extracted by electrolysis of their molten chlorides or oxides, where the metal is deposited at the cathode.
Why are sulphide and carbonate ores converted into metal oxides before reduction?
It is far easier to reduce a metal oxide to its metallic form using carbon (coke) than it is to reduce sulphides or carbonates directly. Processes like Roasting and Calcination are used to ensure this conversion for more efficient extraction.
⚡ Exam Tip
Always classify metals into 3 categories (low, medium, high reactivity).
Write correct equations with conditions.
Link extraction method with reactivity series.
⚠️ Warning
Common Mistakes
Confusing roasting and calcination.
Writing wrong electrode reactions.
Not linking extraction with reactivity.
🧫
Refining of Metals (Purification, Electrolytic Process & Applications)
📘 Definition
Definition of Refining
Refining is the process of purifying an impure metal obtained after extraction to achieve
high purity required for industrial and electrical applications.
🌟 Importance
Why Refining is Necessary
Impurities reduce strength and conductivity
Pure metals are required for electrical components
Enhances durability and corrosion resistance
🔄 Process
Methods of Refining
⚡
Electrolytic Refining
⚗️
Distillation
💧
Liquation
🌀
Zone Refining
📌 Note
Electrolytic Refining
In this method, impure metal is used as the anode, pure metal as the cathode, and a suitable salt
solution as the electrolyte.
Anode mud consists of insoluble impurities (like gold, silver, and platinum) that do not dissolve in the electrolyte and settle at the bottom below the anode during refining. It is industrially significant because the recovery of these precious metals often helps offset the cost of the refining process.
Why is electrolytic refining essential for Copper used in electrical wiring?
Even trace amounts of impurities significantly reduce the electrical conductivity of Copper. Electrolytic refining is used to obtain 99.9% pure metal, ensuring the high conductivity and low resistance required for efficient power transmission.
⚡ Exam Tip
Always write electrode reactions
Mention anode mud explicitly
Draw labeled diagram for 3–5 marks
⚠️ Warning
Common Mistakes
Confusing anode and cathode roles
Not writing half-reactions
Ignoring impurity behavior
🧫
Corrosion and Prevention of Corrosion
📘 Definition
What is Corrosion?
Corrosion is the slow deterioration of metals due to chemical or electrochemical reactions with
the environment, leading to formation of oxides, hydroxides, or sulphides.
Example: Rusting of iron → formation of hydrated iron(III) oxide.
Barrier Protection (Painting, Oiling, Greasing):
Prevents contact with air and moisture.
Galvanisation:
Coating iron with zinc. Zinc acts as a sacrificial metal.
Alloying:
Stainless steel (Fe + Cr + Ni) resists corrosion due to protective oxide layer.
Cathodic Protection:
More reactive metal (Zn, Mg) corrodes instead of iron.
🚨 Caution
Why Galvanisation Works
🚧Caution
Zinc is more reactive than iron, so it oxidises first:
\[
\mathrm{Zn \rightarrow Zn^{2+} + 2e^-}
\]
Thus, iron is protected even if the coating is damaged.
✏️ Example
Why does iron rust significantly faster in coastal (saline) areas?
Coastal air contains saline vapours (salts). Saltwater acts as an electrolyte, which increases the movement of ions and accelerates the electrochemical process of rusting compared to pure water.
How does galvanisation protect iron even if the zinc coating is scratched?
Galvanisation involves coating iron with Zinc. Since Zinc is more reactive than Iron, it acts as a sacrificial metal; it oxidises in preference to iron. Even if the coating is broken, the zinc continues to corrode first, providing sacrificial protection to the underlying iron.
⚡ Exam Tip
Always mention oxygen + water conditions
Write electrochemical reactions for 3–5 marks
Explain galvanisation with reason
⚠️ Warning
Common Mistakes
Writing incomplete rust equation
Ignoring electrochemical mechanism
Confusing galvanisation with painting
🌟 Importance
Why This Topic is Important
Frequently asked in case-study questions
Concept-based 3–5 mark questions
Real-life application based questions
🧫
Important Points – Metals and Non-Metals (Quick Revision + Exam Booster)
📌 Note
Important Points
Metals & Non-Metals: Complete Expert Guide
Comprehensive NCERT Class X Revision System
1. Classification of Elements
Dual Grouping: All 118 elements are broadly classified into Metals and Non-metals based on their electronic configuration.
Electronic Nature: Metals are electropositive (ready to lose electrons to form +ve cations), while Non-metals are electronegative (ready to gain electrons to form -ve anions).
2. Physical Properties & Exceptions
Metals
Lustrous (shiny), Malleable (sheets), Ductile (wires), and excellent Heat/Electrical conductors.
Ex: Hg (Liquid)
Non-Metals
Brittle (break on impact), Non-lustrous (dull), and generally Poor conductors.
Ex: Graphite Conducts
3. Chemical Behaviour
Electron Transfer: Metals → Lose electrons → Form Cations. Non-metals → Gain electrons → Form Anions.
Oxide Nature: Metals form Basic Oxides (turn red litmus blue). Non-metals form Acidic or Neutral Oxides (like CO, N₂O).
Amphoteric Alert: Oxides like Al₂O₃ and ZnO show both acidic and basic properties.
4. Reactivity Series
This is the vertical arrangement of metals in decreasing order of their chemical reactivity.
K > Na > Ca > Mg > Al...
● Metals above Hydrogen can displace it from dilute acids.
● High reactivity = Vigorous displacement from salt solutions.
5. Occurrence & Extraction
Combined State: Highly reactive metals (K, Na) are always found as oxides, carbonates, or sulphides.
Free State: Noble metals (Au, Pt, Ag) are found in their pure metallic form.
Metallurgy: The entire scientific sequence of extraction and refining to obtain 99.9% pure metal.
6. Ionic Compounds (Electrovalent)
Formed by the complete transfer of electrons from a metal to a non-metal.
● High Melting/Boiling Points
● Strong Electrostatic Forces
● Soluble in Water
● Conducts in Molten/Aqueous
7. Alloys
A homogeneous mixture of two or more metals (or metal + non-metal) to enhance strength, hardness, and corrosion resistance (e.g., Brass, Steel).
8. Corrosion
The slow eating away of metal surfaces. Rusting of Iron is a specific type requiring both Oxygen and Moisture.
9. Chemical Behaviour of Non-Metals
● Acid Test: Non-metals do not displace hydrogen from dilute acids because they cannot provide electrons to H⁺ ions.
● Hydrides: They react with Hydrogen to form Stable Covalent Hydrides (like NH₃ or CH₄).
NCERT EXAM BOOSTER (Last-Minute Recall)
Master the 4 Exceptions: Hg, Graphite, Iodine, Diamond.
Always link metal reactions with Electropositivity.
Draw Electron-Dot Structures for NaCl and MgCl₂.
Keywords to use: Amphoteric, Calcination, Roasting, Galvanisation.
NCERT · Class X · Science · Chapter 3
Metals & Non-Metals
An interactive learning engine for deep conceptual mastery
8Concepts
40+Reactions
30+Questions
6Activities
📚Core Concepts
Eight foundational pillars of Chapter 3 — click any card to explore
Concept 01
Physical Properties of Metals
Metals are lustrous, malleable, ductile, and good conductors of heat and electricity. They are generally hard solids with high melting and boiling points.
Physical Chemistry
Exceptions to note:
Mercury is the only liquid metal at room temperature. Sodium and Potassium are so soft they can be cut with a knife. Gallium melts at body temperature (~29.8°C). Iridium and Osmium are the densest metals. Tungsten has the highest melting point (3422°C) among metals.
Key Property
Malleability → Beaten into thin sheets (Gold foil: ~0.0001 mm)
Ductility → Drawn into thin wires (1g Au → 2 km wire)
Concept 02
Physical Properties of Non-Metals
Non-metals are brittle, non-lustrous, poor conductors (except graphite), with low melting points. They may be solids, liquids, or gases.
Physical Chemistry
Carbon (in the form of graphite) is an exception — it conducts electricity. Diamond (another allotrope of carbon) is the hardest natural substance. Bromine is the only non-metal that is liquid at room temperature. Iodine has a metallic lustre but is not a metal.
Concept 03
Reaction of Metals with Oxygen
Most metals react with oxygen to form basic metallic oxides. The reactivity determines speed and conditions required for oxide formation.
Non-metallic oxides turn blue litmus red — they are acidic in nature.
🟢 Metal + Non-metal → Ionic Compound
2Na + Cl2 → 2NaCl Mg + Cl2 → MgCl2 2K + S → K2S
Metal donates electrons; non-metal accepts. Ionic bond formed.
⚖️Comparison Tables
Side-by-side analysis for quick revision and conceptual clarity
Metals vs Non-Metals — Physical Properties
Property
Metals
Non-Metals
Physical State
Mostly solids (except Hg — liquid)
Solids, liquids (Br₂), or gases
Lustre
Shiny metallic lustre
Dull (except Iodine, Graphite)
Hardness
Generally hard (except Na, K)
Soft (except Diamond — hardest)
Malleability
Malleable (beaten into sheets)
Brittle; break into pieces
Ductility
Ductile (drawn into wires)
Not ductile
Conductivity
Good conductors of heat & electricity
Poor conductors (except Graphite)
Melting/Boiling Point
High (except Hg, Ga, Cs)
Low (except Diamond, Silicon)
Density
High (except Na, K, Li)
Low
Sonority
Sonorous (produce sound when struck)
Not sonorous
Metals vs Non-Metals — Chemical Properties
Reaction
Metals
Non-Metals
With Oxygen
Metal + O₂ → Basic oxide
Non-metal + O₂ → Acidic oxide
With Water
Reactive metals form H₂ + base
Halogens dissolve; Cl₂ + H₂O → HCl + HClO
With Acids
Form salt + H₂ ↑ (if above H)
Do not react with acids generally
With Chlorine
Form ionic chlorides (MCl)
Form covalent chlorides (e.g., PCl₃)
Electron Tendency
Lose electrons → form cations
Gain electrons → form anions
Bond Type Formed
Ionic bonds with non-metals
Covalent bonds with non-metals
Nature of Oxide
Basic (turns red litmus blue)
Acidic (turns blue litmus red)
Ionic Compounds vs Covalent Compounds
Property
Ionic Compounds
Covalent Compounds
Bond Formation
Electrostatic attraction between ions
Sharing of electrons
Melting/Boiling Point
High
Generally low
Solubility in Water
Generally soluble
Generally insoluble (except polar)
Electrical Conductivity
Conduct in molten/aqueous state
Generally non-conductors
State at Room Temp
Solid crystals
May be solid, liquid or gas
Example
NaCl, MgCl₂, CaO
HCl, CO₂, H₂O, CCl₄
Metals and Reactivity with Water — Summary
Metal
Condition
Product
Reactivity
Na, K
Cold water
MOH + H₂↑
Very high
Ca
Cold water
Ca(OH)₂ + H₂↑
High
Mg
Hot water / steam
MgO + H₂↑
Moderate
Al, Zn, Fe
Steam only
Metal oxide + H₂↑
Moderate–low
Pb
Steam (slowly)
PbO + H₂↑
Low
Cu, Ag, Au, Pt
No reaction
—
Negligible
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📝Concept-Building Question Bank
Original questions (not from textbook) — organised by concept, with full solutions
ConceptMedium
A student claims that "all metals conduct electricity, therefore copper wire is used in electrical circuits." Identify TWO distinct reasons why copper is specifically preferred over other conducting metals like aluminium or iron.
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Step-by-Step Solution
Understand the question: Why is Cu chosen over other conductors? We need properties, not just electrical conductance.
Property 1 — Electrical conductivity: Copper has the second highest electrical conductivity among metals (after silver), making current flow with minimal resistance. Silver is too expensive; hence copper is the practical choice.
Property 2 — Ductility: Copper is highly ductile — it can be drawn into very thin wires without breaking. This is essential for wire manufacture. Iron is less ductile and more brittle.
Bonus — Corrosion resistance: Copper does not rust like iron. It forms a thin protective green patina (CuCO₃·Cu(OH)₂) that prevents further corrosion, extending wire life.
Conclusion: Copper's exceptional ductility + high conductivity + corrosion resistance make it the optimal practical choice for wiring, even though silver is a better conductor.
ApplicationMedium
Zinc granules are placed in silver nitrate solution and in copper sulphate solution separately. Predict what happens in each case and write the balanced equations.
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Step-by-Step Solution
Recall reactivity series: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag
Case 1 — Zn in AgNO₃: Zn is MORE reactive than Ag. ∴ Zn displaces Ag from solution. Zn + 2AgNO₃ → Zn(NO₃)₂ + 2Ag↓ Colourless AgNO₃ → colourless Zn(NO₃)₂; grey-white silver deposits on zinc.
Case 2 — Zn in CuSO₄: Zn is MORE reactive than Cu. ∴ Zn displaces Cu. Zn + CuSO₄ → ZnSO₄ + Cu↓ Blue CuSO₄ solution becomes colourless; reddish-brown Cu deposits on zinc.
Generalisation: A more reactive metal always displaces a less reactive metal from its salt solution. This is the basis of displacement reactions.
AnalysisHard
Aluminium is more reactive than iron yet it is preferred for making aircraft bodies and cooking utensils. Explain this apparent contradiction using your chemical knowledge.
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Step-by-Step Solution
Apparent contradiction: If Al is more reactive, it should corrode faster — so why is it used where corrosion resistance is needed?
The key — Passive layer (Anodisation): When Al is exposed to air, it immediately reacts with oxygen to form a thin, impermeable layer of aluminium oxide: 4Al + 3O₂ → 2Al₂O₃
Why this protects: Al₂O₃ is hard, tightly adherent, and chemically inert. It acts as a physical barrier, preventing further oxygen and moisture from reaching the underlying metal — so corrosion stops after a thin film forms.
Contrast with iron: Iron oxide (Fe₂O₃·xH₂O — rust) is porous and flakes off, continuously exposing fresh iron. Iron corrodes layer by layer until the entire piece is destroyed.
Other advantages of Al: Low density (light-weight, essential for aircraft); high strength-to-weight ratio; good thermal conductivity (cooking utensils).
Conclusion: Al's self-passivating property makes it effectively corrosion-resistant in practical applications, despite being chemically reactive.
HOTSVery Hard
A piece of sodium is placed in oxygen, then the product is dissolved in water. The resulting solution is tested with litmus and then with phenolphthalein. Predict all observations and explain the chemistry at each step.
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Step-by-Step Solution
Step 1 — Na reacts with O₂: 4Na + O₂ → 2Na₂O Sodium burns with a bright yellow-orange flame. Product: Sodium oxide (Na₂O), a white solid. (Na can also form Na₂O₂ with excess O₂.)
Step 2 — Na₂O dissolves in water: Na₂O + H₂O → 2NaOH The white solid dissolves, forming sodium hydroxide — a strong alkali. Solution becomes warm (exothermic).
Step 3 — Litmus test: NaOH is strongly alkaline (pH ≈ 12–13). Red litmus paper turns blue. Blue litmus remains blue.
Step 4 — Phenolphthalein test: Phenolphthalein is colourless in neutral/acidic solutions and pink/magenta in alkaline solutions. Solution turns bright pink. This confirms strong alkalinity.
Key Concept Link: Metal → Metal oxide (basic) → Metal hydroxide (alkali) in water. This confirms that metallic oxides are basic in nature.
ApplicationMedium
Explain why ionic compounds like NaCl have high melting points and can conduct electricity in molten state but not in solid state.
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Step-by-Step Solution
Structure of ionic compounds: NaCl has Na⁺ and Cl⁻ ions arranged in a regular 3D lattice. Each Na⁺ is surrounded by 6 Cl⁻ ions and vice versa (face-centred cubic structure).
High melting point: Millions of strong electrostatic attractions (Na⁺ ← → Cl⁻) hold the lattice together. Enormous energy is needed to break these bonds simultaneously → very high melting point (801°C for NaCl).
No conductivity in solid state: In solid NaCl, ions are held rigidly in fixed positions in the lattice. They cannot move. No movement of charge = no current flow.
Conductivity in molten/aqueous state: On melting (or dissolving), the lattice breaks down. Ions become free to move. When electrodes are placed, Na⁺ migrates to cathode and Cl⁻ to anode — current flows.
Conclusion: Conductivity in ionic compounds requires mobile ions. Solid state = immobile ions = insulator. Liquid/aqueous = mobile ions = conductor.
ConceptEasy
Three metals A, B and C react with dilute HCl as follows: A reacts vigorously with cold HCl; B reacts only on heating; C does not react at all. Arrange A, B and C in decreasing order of reactivity and give one example each.
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Step-by-Step Solution
Interpret observations: Ease of reaction with acids indicates position in reactivity series. More reactive metals react faster under milder conditions.
Analysis: • A reacts vigorously with cold HCl → Very reactive → Above Mg in series → e.g., Na or Mg • B reacts only on heating → Moderate reactivity → e.g., Fe or Zn • C does not react → Below H in series → e.g., Cu
Order: A > B > C in reactivity.
Representative equations: Mg + 2HCl → MgCl₂ + H₂↑ (A) Fe + 2HCl → FeCl₂ + H₂↑ (B, on heating) Cu + HCl → No reaction (C)
HOTSVery Hard
Stainless steel is rust-resistant, yet pure iron rusts rapidly. Explain both the chemistry of rusting and how alloying prevents it. Also suggest two other methods of preventing rusting in bridges.
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Step-by-Step Solution
Chemistry of rusting (electrochemical process): 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ → 2Fe₂O₃·3H₂O (rust) Rust is hydrated iron(III) oxide — orange-brown, porous, and non-adherent. It flakes off, exposing fresh iron.
Why pure iron rusts so fast: Fe readily loses electrons (Fe → Fe²⁺) in the presence of moisture. Water acts as an electrolyte, accelerating the electrochemical reaction. Rust is non-protective and keeps falling off.
How stainless steel prevents rust: Adding chromium (Cr, ~10–18%) creates an alloy where Cr reacts with O₂ to form Cr₂O₃ — a thin, invisible, adherent, non-porous passive film — preventing Fe from reacting with air/moisture. Nickel adds toughness.
Other prevention methods for bridges: ① Galvanisation: Coat iron with zinc. Zn is more reactive — it acts as a sacrificial anode; even if coating chips, Zn corrodes first, protecting Fe. ② Painting / Cathodic Protection: Apply anti-rust paints (containing zinc chromate or red lead). For bridges, connect to an external DC source making bridge the cathode (cathodic protection).
AnalysisHard
Electrolysis of molten NaCl produces sodium at the cathode and chlorine at the anode. Explain why: (a) electrolysis cannot be done in aqueous NaCl for obtaining sodium, and (b) Na is collected at the cathode, not the anode.
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Step-by-Step Solution
Part (a) — Why not aqueous NaCl? In aqueous solution, both Na⁺ and H⁺ (from water ionisation: H₂O ⇌ H⁺ + OH⁻) are present at the cathode. Since H⁺ is easier to reduce than Na⁺ (H has lower reduction potential), H₂ gas is produced at cathode, NOT sodium metal. Na is far too reactive; it would immediately react with water: 2Na + 2H₂O → 2NaOH + H₂↑
Part (b) — Na collected at cathode: Cathode is the negative electrode; it attracts positive ions (cations). Na⁺ ions migrate to cathode and gain electrons: Na⁺ + e⁻ → Na (reduction) At anode (positive), Cl⁻ ions lose electrons: 2Cl⁻ → Cl₂↑ + 2e⁻ (oxidation)
Mnemonic:OILRIG — Oxidation Is Loss, Reduction Is Gain of electrons. Anode = oxidation; Cathode = reduction. Positive ions (Na⁺) go to cathode (negative electrode).
💡Tips & Tricks for Exams
Memory aids, shortcuts, and exam strategies curated from the chapter
🔑 Memory Aids & Mnemonics
🔠
Reactivity Series Mnemonic
"King Narendra Can Make A Zebra Feel Pretty Hot, Considering His Actual Greatness Particularly" K · Na · Ca · Mg · Al · Zn · Fe · Pb · H · Cu · Hg · Ag · Au · Pt
⚡
OILRIG — Electrochemistry Core
Oxidation Is Loss, Reduction Is Gain (of electrons). Anode = Oxidation (A–O), Cathode = Reduction. Works for electrolysis and displacement reactions alike.
🏷️
Oxide Nature → Litmus Trick
Metal oxides are Basic → turns Red litmus Blue (B comes after R — basic follows the metal). Non-metal oxides are Acidic → turns Blue litmus Red.
🎯
Exceptions to Memorise (1 mark questions!)
• Liquid metal at RT → Mercury (Hg) • Liquid non-metal at RT → Bromine (Br₂) • Non-metal conductor → Graphite (Carbon) • Metal cut by knife → Na, K, Li (very soft) • Lustrous non-metal → Iodine • Hardest natural substance → Diamond (Carbon)
🔗
Alloy Composition Quick-Recall
Brass = Cu + Zn (Z = Zinc; Brass rhymes with "class" → elegant, used in musical instruments) Bronze = Cu + Sn (Bronze age = old; Sn = ancient symbol for tin) Solder = Pb + Sn (think "Pbond + Snap together" — welding)
📐
Exam Strategy: Reaction with Water Questions
Always state: (1) the condition (cold water / hot water / steam / no reaction), (2) the products, (3) write the balanced equation. For full marks, mention the observation (e.g., vigorous effervescence, Na floats and moves).
🧪
Fe₃O₄ vs Fe₂O₃ — Don't Mix Them Up!
Fe₃O₄ (black magnetic oxide) — formed when iron reacts with steam or burns in oxygen: 3Fe + 4H₂O → Fe₃O₄ + 4H₂ Fe₂O₃ (hydrated form = rust) — formed by slow corrosion with air + moisture: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ → Fe₂O₃·3H₂O
💰
Why Gold & Platinum are "Noble" Metals
They are at the bottom of the reactivity series — lowest tendency to lose electrons. They do not react with air, water, dilute acids, or most chemicals. This is why gold jewellery doesn't tarnish and platinum is used in catalytic converters.
⚠️Common Mistakes & Misconceptions
Frequently made errors in exams — learn what's wrong and why
❌
Confusing Malleability with Ductility
Malleability = drawn into wires; Ductility = beaten into sheets
Malleability = beaten into thin sheets; Ductility = drawn into thin wires
Rust = Fe₂O₃·xH₂O (hydrated iron(III) oxide) — water is essential!
Rust requires both oxygen AND water. Neither alone causes rusting. Always write the hydrated form.
❌
Thinking Copper Reacts with Dilute HCl
Cu + 2HCl → CuCl₂ + H₂↑
Cu + HCl → No reaction (Cu is below H in reactivity series)
Cu cannot displace H from dilute HCl/H₂SO₄. It does react with concentrated HNO₃ and hot conc. H₂SO₄ — but those are not dilute acids.
❌
Saying "All metals are solid at room temperature"
Mercury is a gas at room temperature
Mercury (Hg) is a LIQUID metal at room temperature (melting point = −39°C)
Gallium (Ga) melts at ~29.8°C and becomes liquid on your palm — also an exception!
❌
Confusing Anodising with Galvanisation
Galvanisation = coating with aluminium
Galvanisation = coating with ZINC (Zn). Anodising = treating aluminium to thicken Al₂O₃ layer.
Galvanisation is named after Luigi Galvani, uses zinc because Zn is more reactive than Fe — acts as sacrificial metal.
❌
Wrong Product: Iron + Steam
3Fe + 4H₂O → 3FeO + 4H₂
3Fe + 4H₂O → Fe₃O₄ + 4H₂↑ (Iron reacts with STEAM, not cold water)
Fe₃O₄ is a mixture of FeO and Fe₂O₃ — it's the black magnetic oxide formed at high temperature. Iron does NOT react with cold or hot water.
❌
Thinking Ionic Compounds Conduct in Solid State
NaCl conducts electricity in solid state because it has ions
Solid NaCl does NOT conduct — ions are locked in lattice and cannot move. Free ions are needed for conductivity.
Conductivity requires mobile charge carriers. In solid ionic compounds, ions are stationary. Only molten or dissolved ionic compounds conduct.
❌
Mixing Brass and Bronze
Bronze = Cu + Zn; Brass = Cu + Sn
Brass = Cu + Zn; Bronze = Cu + Sn
Trick: Brass has a 'Z' (Zinc). Bronze has an 'n' → Tin (Sn comes from Latin Stannum). Bronze Age weapons were made of Cu+Sn — the stronger alloy.
🎮Interactive Learning Modules
Six activity types for active recall and concept reinforcement
🎯
MCQ Quiz
10 randomised multiple-choice questions with instant feedback and explanations
🔗
Match the Columns
Connect metals to their reactions, properties, or alloys
✏️
Fill in the Blanks
Complete the chemical equations and statements
📊
Reactivity Sorter
Drag-and-rank metals by reactivity; verify against the actual series
✅
True or False
15 quick statements — decide, then see the explanation
🃏
Flash Cards
Flip cards for key terms, reactions, and alloy compositions
Match the Columns
Click an item in Column A, then click the matching item in Column B.
Column A — Metal / Alloy
Column B — Property / Composition
Fill in the Blanks
Complete each statement or equation with the correct term.
Reactivity Sorter
Drag the metals to arrange them from MOST reactive (top) to LEAST reactive (bottom).
True or False
Flash Cards — Click to Flip
📚
ACADEMIA AETERNUMतमसो मा ज्योतिर्गमय · Est. 2025
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Complete Guide: Metals and Non-Metals Class 10 | Notes, Reactions, NCERT Q&A
Complete Guide: Metals and Non-Metals Class 10 | Notes, Reactions, NCERT Q&A — Complete Notes & Solutions · academia-aeternum.com
Chapter 3 "Metals and Non-Metals" explores the fundamental differences between these two essential groups of elements. Metals, characterized by their shiny appearance, malleability, ductility, and excellent conductivity of heat and electricity, are widely used in daily life and industries. Non-metals, on the other hand, are generally dull, brittle, poor conductors, and exist in various states—solid, liquid, or gas. This chapter delves into their physical and chemical properties, their reactions…
🎓 Class 10📐 Science📖 NCERT✅ Free Access🏆 CBSE · JEE
