Metals

Metals

Imagine holding a shiny coin, a copper wire, or an aluminium foil. What makes these materials so different from a lump of coal or a piece of sulphur? The answer lies in understanding metals and their unique physical properties. In this chapter, we will explore what sets metals apart from non-metals through hands-on observation and inquiry-based learning.

Introduction to Metals

Metals form the largest group of elements in the periodic table. They include familiar substances like iron, copper, aluminium, magnesium, sodium, lead, and zinc. While we encounter metals daily—in utensils, wires, jewellery, and construction materials—have you ever wondered why these elements are chosen for these specific uses?

The secret lies in their physical properties: characteristics that can be observed or measured without changing the substance's chemical identity. Let us embark on a journey of discovery through simple, investigative activities that reveal the fascinating world of metals.

Physical Properties of Metals

1. Metallic Lustre

Lustre is the property that makes metals shine. When you take a fresh piece of iron, copper, or aluminium and clean its surface with sandpaper, you'll notice a brilliant, reflective shine. This is called metallic lustre.

Why do metals shine? In their pure state, metals have tightly packed atoms that reflect light uniformly. However, when exposed to air and moisture, many metals form a dull layer of oxide on their surface. Rubbing with sandpaper removes this layer, revealing the shiny surface beneath.

{{VISUAL: photo: comparison of a dull copper coin and a freshly cleaned shiny copper surface side by side}}

{{KEY: type=definition | title=Metallic Lustre | text=The shiny appearance of metals in their pure state due to their ability to reflect light uniformly from their smooth surface.}}

Metals, in their pure state, have a shining surface — this property distinguishes them at first glance.

2. Hardness

Not all metals feel the same when you try to cut them. Take a sharp knife and attempt to cut through iron, copper, aluminium, or magnesium—you'll find them quite resistant. These metals are hard.

But here's a surprising fact: sodium metal is so soft that it can be cut with a knife like butter! This shows that hardness varies from metal to metal. While most metals are hard and strong, alkali metals (lithium, sodium, potassium) are exceptions—they are soft, have low densities, and low melting points.

{{KEY: type=points | title=Hardness in Metals | text=- Most metals are hard and cannot be easily cut with a knife.

• Hardness varies: iron is harder than aluminium.
• Alkali metals (sodium, potassium) are soft exceptions.
• Sodium metal must be handled with care and stored properly.}}

3. Malleability

Have you ever wondered how thin aluminium foils or gold ornaments are made? The answer lies in a property called malleability—the ability of metals to be beaten into thin sheets without breaking.

When you place a piece of iron, zinc, lead, or copper on a hard surface (like an iron block) and strike it repeatedly with a hammer, the metal flattens and spreads into a thin sheet instead of shattering. This is malleability in action.

Gold and silver are the most malleable metals known. Jewellers can hammer gold into sheets so thin that they are almost transparent!

{{VISUAL: diagram: illustration showing a metal piece being hammered on an anvil, demonstrating malleability with thin metal sheets below}}

{{KEY: type=definition | title=Malleability | text=The property of metals that allows them to be beaten or rolled into thin sheets without breaking. Gold and silver are the most malleable metals.}}

4. Ductility

Think about the electric wires running through your home, or the cables supporting a suspension bridge. They are made from metals drawn into long, thin wires—a property called ductility.

Ductility is the ability of metals to be stretched into thin wires. Not all materials can do this; try pulling a piece of chalk or charcoal—it will simply break. But metals like copper, aluminium, and gold can be drawn into wires of remarkable length.

Gold holds the record as the most ductile metal. Would you believe that a single gram of gold can be drawn into a wire nearly 2 kilometres long? This extraordinary property makes metals invaluable for electrical and mechanical applications.

{{KEY: type=concept | title=Ductility | text=Ductility is the ability of metals to be drawn into thin wires without breaking. Copper is commonly used for electrical wiring due to its excellent ductility and conductivity. Gold is the most ductile metal, allowing a single gram to be stretched into a 2 km wire.}}

Why are malleability and ductility important? These properties allow metals to be shaped according to our needs—into sheets, wires, utensils, tools, and intricate designs—making them incredibly versatile materials.

Conductivity of Metals

Heat Conductivity

Why are cooking vessels made from aluminium, copper, or stainless steel—and never from wood or plastic? The answer is thermal conductivity: metals are excellent conductors of heat.

When you heat one end of a metal wire (say, aluminium or copper) clamped to a stand with a pin attached using wax at the free end, the heat travels quickly through the metal. Soon, the wax melts and the pin drops—even though the flame is far from the pin!

Silver and copper are the best conductors of heat, which is why copper-bottomed cookware is prized by chefs. However, lead and mercury are comparatively poor conductors.

{{VISUAL: diagram: labeled setup showing a metal wire clamped on a stand with a pin attached using wax at the free end and a flame heating the clamped end}}

{{KEY: type=points | title=Thermal Conductivity | text=- Metals are good conductors of heat.

• Silver and copper are the best heat conductors.
• Lead and mercury are relatively poor conductors.
• Metals have high melting points (except gallium and caesium).}}

Electrical Conductivity

Set up a simple electric circuit with a bulb, battery, and two terminals (A and B). Place different metals—copper, iron, aluminium—between the terminals. In every case, the bulb glows brightly. This proves that metals are excellent conductors of electricity.

This is why electrical wires are made from copper or aluminium. But have you noticed that these wires are always coated with materials like PVC (polyvinylchloride) or rubber? These coatings are insulators—they prevent electric shocks and short circuits by stopping the flow of electricity outside the wire.

{{KEY: type=exam | title=Why PVC-coated Wires? | text=A common CBSE question asks why electrical wires are coated with PVC. Answer: PVC is an insulator that prevents electric shocks and short circuits by not allowing electricity to pass through it.}}

5. Sonority

Strike a metal utensil, a steel plate, or a school bell with a hard object. What do you hear? A clear, ringing sound! This property of metals to produce sound when struck is called sonority.

Metals are sonorous—they ring when hit. This is why school bells, cymbals, and gongs are made from metals like bronze or brass. Non-metals, on the other hand, produce dull sounds or no sound at all.

{{KEY: type=definition | title=Sonority | text=The property of metals to produce a ringing sound when struck on a hard surface. This is why bells and musical instruments are made from metals.}}

Summary Table: Physical Properties of Metals

{{ZOOM: title=Exceptions in Metals | text=Not all metals fit the "typical" profile. Mercury is liquid at room temperature. Gallium and caesium melt in your palm. Alkali metals are soft and have low densities. Recognizing these exceptions helps us understand that properties exist on a spectrum.}}

Conclusion

Through these simple, inquiry-based activities, we have discovered that metals possess a unique set of physical properties—lustre, hardness, malleability, ductility, thermal and electrical conductivity, and sonority—that make them indispensable in our daily lives. From the wires that power our homes to the utensils that cook our food, metals are everywhere.

But physical properties alone don't tell the full story. In the next section, we will explore non-metals and see how they differ. Then, we'll dive deeper into the chemical properties that truly distinguish metals from non-metals—because chemistry is where the real magic happens!

Non-metals

Non-metals

In the previous chapter, we learned that elements can be broadly classified into metals and non-metals. While metals make up the majority of the periodic table, non-metals are fewer in number but equally important in nature and industry. Examples of non-metals include carbon, sulphur, iodine, oxygen, hydrogen, nitrogen, phosphorus, chlorine, and the noble gases.

An interesting physical fact: among all non-metals, bromine is the only one that exists as a liquid at room temperature. All other non-metals are either solids (like carbon, sulphur, iodine, phosphorus) or gases (like oxygen, hydrogen, nitrogen, chlorine).

Physical Properties of Non-metals: A Sharp Contrast

To understand how non-metals differ from metals, let us revisit the physical properties we studied for metals — appearance, hardness, malleability, ductility, conductivity, and sonority — and see how non-metals compare.

{{VISUAL: photo: collection of common non-metal samples including a lump of sulphur (yellow), iodine crystals (dark purple), and a piece of charcoal (black carbon)}}

Appearance and Lustre

Unlike metals, most non-metals do not have a shiny or lustrous surface. They appear dull. For instance, sulphur is a dull yellow powder, and carbon (in the form of coal or charcoal) is black and non-reflective.

{{KEY: type=definition | title=Lustre | text=Lustre is the property of a substance to have a shiny or reflective surface. Most non-metals are non-lustrous, meaning they appear dull and do not reflect light like metals do.}}

However, there is a notable exception: iodine is a non-metal but it is lustrous. Iodine crystals have a shiny, metallic-looking surface, demonstrating that classification based purely on physical properties can be misleading.

Hardness

Non-metals are generally not hard. Sulphur and phosphorus can be easily powdered. However, there is a striking exception: diamond, an allotrope of carbon, is the hardest natural substance known. This shows the enormous diversity even within a single element.

{{KEY: type=concept | title=Allotropes | text=Allotropes are different structural forms of the same element. Carbon exists in several allotropes — diamond (extremely hard), graphite (soft and slippery), and fullerenes. Each allotrope has vastly different physical properties despite being pure carbon.}}

Malleability and Ductility

Non-metals are neither malleable nor ductile. If you try to hammer sulphur or iodine, they will break or shatter into smaller pieces rather than flatten into sheets. Similarly, non-metals cannot be drawn into wires. This brittle nature is one of the clearest differences between metals and non-metals.

Electrical Conductivity

Most non-metals are poor conductors of electricity. Substances like sulphur, phosphorus, and iodine do not allow electric current to pass through them. They are called insulators.

However, once again there is an important exception: graphite, an allotrope of carbon, is a good conductor of electricity. Graphite is used in pencils and also as electrodes in batteries precisely because it can conduct current.

{{ZOOM: title=Why does graphite conduct electricity? | text=In graphite, carbon atoms are arranged in layers. Each carbon atom forms three bonds, leaving one free electron per atom. These free electrons can move between layers, allowing electric current to flow — unlike diamond, where all four electrons are locked in rigid bonds.}}

Thermal Conductivity

Non-metals are generally poor conductors of heat. This is why materials like wood (which contains carbon compounds) and plastic are used as handles for cooking utensils — they do not transfer heat easily and protect our hands from burns.

Sonority

Non-metals are not sonorous. If you strike a piece of coal or sulphur against a hard surface, it will not produce a ringing sound. Instead, it may crumble or break. This is in sharp contrast to metals like iron or copper, which produce a clear, ringing sound when struck.

{{VISUAL: chart: comparison table showing physical properties of metals versus non-metals across lustre, hardness, malleability, ductility, conductivity, and sonority}}

Exceptions: The Grey Area in Classification

As we have seen, classifying elements purely on physical properties is not foolproof. There are several important exceptions on both sides of the fence.

{{KEY: type=points | title=Key Exceptions in Physical Properties | text=- Iodine is a non-metal but is lustrous.

• Diamond (carbon) is a non-metal but is the hardest natural substance.
• Graphite (carbon) is a non-metal but conducts electricity.
• Mercury is a metal but is liquid at room temperature.
• Alkali metals (lithium, sodium, potassium) are metals but are soft enough to be cut with a knife.
• Gallium and caesium are metals but have very low melting points; they can melt on your palm.}}

Why Do These Exceptions Exist?

The physical properties of elements depend on their atomic structure and the type of bonding between atoms. For example:

• Diamond has a rigid three-dimensional network of carbon atoms, making it extremely hard.
• Graphite has a layered structure with weak forces between layers, making it soft and slippery, yet conductive due to free electrons.
• Alkali metals have only one electron in their outermost shell and weak metallic bonding, making them soft and reactive.

These structural differences explain why elements in the same category can behave so differently.

A Better Basis for Classification: Chemical Properties

Because physical properties have so many exceptions, elements are more reliably classified as metals or non-metals based on their chemical properties — especially how they react with oxygen, acids, and bases.

For instance, when we burn magnesium (a metal) in air, it forms magnesium oxide, which dissolves in water to give a basic solution (turns red litmus blue). On the other hand, when we burn sulphur (a non-metal) in air, it forms sulphur dioxide, which dissolves in water to give an acidic solution (turns blue litmus red).

{{KEY: type=concept | title=Oxides of Metals and Non-metals | text=Most metal oxides are basic in nature and turn red litmus blue when dissolved in water. Most non-metal oxides are acidic in nature and turn blue litmus red. This chemical behaviour is a more reliable way to distinguish metals from non-metals than physical properties alone.}}

We will explore these chemical properties in detail in the next section, where we study how metals react with oxygen, water, acids, and salt solutions.

{{VISUAL: diagram: flowchart showing classification of elements into metals and non-metals based on physical properties (with exceptions noted) and chemical properties (more reliable)}}

Summary Table: Metals vs. Non-metals

{{KEY: type=exam | title=Common Exam Question | text=Be prepared to explain exceptions like iodine (lustrous non-metal), graphite (conducting non-metal), and alkali metals (soft metals). Questions often ask you to justify why physical properties alone are insufficient for classification.}}

The diversity in properties—both within metals and non-metals—reminds us that nature does not fit neatly into boxes. Understanding the exceptions deepens our grasp of atomic structure and bonding.

What happens when Metals are burnt in Air?

What happens when Metals are burnt in Air?

When you strike a matchstick or light a candle, you witness combustion — a reaction with oxygen. Metals, too, can react with oxygen in the air, and the nature of this reaction tells us a great deal about a metal's chemical reactivity. Some metals burn with brilliant flames, some form protective coatings, and some refuse to react even when heated intensely. Understanding how metals interact with oxygen is the first step in arranging them on a reactivity series, a tool that chemists use to predict and explain reactions.

Observing Metals Burn

In the laboratory, when we heat different metals in air, we observe strikingly different behaviours. Magnesium, for instance, burns with a dazzling white flame so bright that it can temporarily dazzle your eyes — this is why safety goggles are essential. The product left behind is a white powder, quite different from the shiny grey ribbon you started with.

Other metals show more subdued reactions. Copper, when heated, does not burst into flames but instead develops a black coating on its surface. Iron filings, when sprinkled into a flame, burn vigorously with bright sparks, yet a thick iron nail might only glow red without catching fire. Sodium and potassium, by contrast, are so reactive that they must be stored under kerosene oil to prevent them from reacting with moisture and oxygen in the air — if exposed, they can ignite spontaneously.

{{VISUAL: photo: comparison of metals burning in air showing magnesium burning with bright white flame, copper turning black, and iron filings sparking}}

{{KEY: type=concept | title=Metal Combustion in Air | text=When metals are heated in the presence of oxygen, they undergo oxidation to form metal oxides. The vigour and ease of this reaction vary widely among metals, reflecting their differing chemical reactivities. Highly reactive metals like sodium and potassium react violently, while noble metals like gold and silver do not react at all.}}

Formation of Metal Oxides

The general reaction that occurs when a metal burns in air can be written as:

Metal + Oxygen → Metal Oxide

For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black compound:

2Cu + O₂ → 2CuO

Similarly, aluminium reacts with oxygen to form aluminium oxide, a white powder:

4Al + 3O₂ → 2Al₂O₃

These metal oxides are the products of combustion. Notice that the chemical formula of the oxide depends on the valency of the metal — copper forms CuO (valency 2), while aluminium forms Al₂O₃ (valency 3).

{{KEY: type=definition | title=Metal Oxide | text=A metal oxide is a chemical compound formed when a metal reacts with oxygen. Metal oxides are generally basic in nature, meaning they react with acids to form salts and water. However, some metal oxides can exhibit both acidic and basic properties.}}

The Basic Nature of Metal Oxides

Most metal oxides are basic in nature. Recall from Chapter 2 that bases are substances that react with acids to produce salt and water. For instance, when copper(II) oxide reacts with hydrochloric acid:

CuO + 2HCl → CuCl₂ + H₂O

This confirms that CuO is a basic oxide.

Some metal oxides, such as sodium oxide (Na₂O) and potassium oxide (K₂O), are soluble in water. When they dissolve, they form alkalis (soluble bases):

Na₂O(s) + H₂O(l) → 2NaOH(aq)

K₂O(s) + H₂O(l) → 2KOH(aq)

Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are strong alkalis, commonly used in laboratories and industries.

{{VISUAL: diagram: flowchart showing metal combustion leading to metal oxide formation, then branching into soluble oxides forming alkalis and insoluble oxides remaining as basic oxides}}

Amphoteric Oxides: The Exception

Not all metal oxides behave in a straightforward basic manner. Some metal oxides, such as aluminium oxide (Al₂O₃) and zinc oxide (ZnO), can react with both acids and bases to produce salts and water. Such oxides are called amphoteric oxides.

For example, aluminium oxide reacts with hydrochloric acid (an acid):

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

It also reacts with sodium hydroxide (a base):

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

In the second reaction, sodium aluminate (NaAlO₂) is formed. This dual behaviour makes amphoteric oxides unique and important in industrial processes such as aluminium extraction.

{{KEY: type=definition | title=Amphoteric Oxide | text=An amphoteric oxide is a metal oxide that can react with both acids and bases to produce salts and water. Examples include aluminium oxide and zinc oxide. This property distinguishes them from typical basic metal oxides.}}

{{KEY: type=exam | title=Commonly Tested | text=CBSE exams frequently ask students to identify amphoteric oxides and write balanced equations showing their reactions with both acids and bases. Be able to name Al₂O₃ and ZnO as examples and write their reactions with HCl and NaOH.}}

Reactivity and Protective Oxide Layers

Why do some metals burn vigorously while others barely react? The answer lies in their reactivity — a measure of how readily a metal gives up electrons to form positive ions.

• Highly reactive metals (e.g., sodium, potassium) react instantly and violently with oxygen, even at room temperature. They must be stored under oil to prevent contact with air.
• Moderately reactive metals (e.g., magnesium, aluminium, zinc) react readily when heated, but at ordinary temperatures they form a thin protective oxide layer on their surface. This layer prevents further oxidation, acting like a shield. Aluminium, for instance, is widely used in construction because its oxide layer makes it corrosion-resistant.
• Low reactivity metals (e.g., copper, lead) react slowly, forming oxide coatings only when heated.
• Noble metals (e.g., silver, gold) do not react with oxygen even at high temperatures, which is why they remain shiny and are prized for jewellery.

The table below summarises the reactivity of common metals with oxygen:

{{VISUAL: chart: reactivity series bar chart showing decreasing reactivity from sodium to gold based on reaction with oxygen}}

{{ZOOM: title=Anodising Aluminium | text=Anodising is an industrial process that artificially thickens the oxide layer on aluminium. The aluminium article is made the anode in an electrolytic cell with dilute sulphuric acid. Oxygen evolved at the anode reacts with aluminium, forming a thicker, more durable Al₂O₃ layer. This layer can be dyed to give aluminium articles vibrant, corrosion-resistant finishes used in cookware, window frames, and mobile phone bodies.}}

Arranging Metals by Reactivity

Based on how readily metals burn in air, we can start arranging them in order of decreasing reactivity:

Sodium > Magnesium > Aluminium > Zinc > Iron > Copper > Silver > Gold

This sequence is the foundation of the reactivity series of metals, a powerful concept that helps predict how metals will behave in various chemical reactions — not just with oxygen, but also with water, acids, and salt solutions.

"The order in which metals react with oxygen is the first clue to their position in the reactivity series — a ladder that ranks metals from the most eager to lose electrons to the most reluctant."

{{KEY: type=points | title=Key Observations from Metal Combustion | text=- Metals react with oxygen to form metal oxides, which are generally basic in nature.

• Some metal oxides like Al₂O₃ and ZnO are amphoteric, reacting with both acids and bases.
• Highly reactive metals burn vigorously, while noble metals do not react at all.
• A protective oxide layer prevents further oxidation in metals like aluminium and zinc.
• The ease of combustion helps arrange metals in a reactivity series.}}

In the next section, we will explore what happens when these same metals are brought into contact with water — another window into their reactivity and a way to refine our understanding of the reactivity series.

What happens when Metals react with Water?

What happens when Metals react with Water?

We have seen that metals react with oxygen to form oxides. But reactivity with oxygen alone does not paint the full picture of how different metals behave. To truly understand the reactivity series, we need to observe how metals interact with water — one of the most common substances on Earth.

Water can exist in three states for our experiments: cold water, hot water, and steam (water vapour). Different metals respond differently to each of these states, and this variation helps us build a clear hierarchy of metal reactivity.

The Experiment: Testing Metals with Water

When conducting Activity 3.10 from your NCERT textbook, you would collect samples of metals like sodium (Na), potassium (K), calcium (Ca), magnesium (Mg), aluminium (Al), zinc (Zn), iron (Fe), lead (Pb), copper (Cu), and silver (Ag).

The method is systematic:

1. First, test each metal with cold water in separate beakers
2. If a metal does not react with cold water, try hot water
3. If there is still no reaction, expose the metal to steam using the apparatus shown below

{{VISUAL: diagram: laboratory setup showing a boiling flask with water, delivery tube passing steam over a heated metal sample in a combustion tube, with a collection tube for hydrogen gas}}

{{KEY: type=concept | title=General Reaction Pattern | text=When metals react with water, they produce a metal oxide and hydrogen gas. If the metal oxide is soluble in water, it further reacts to form a metal hydroxide. The general equations are: Metal + Water → Metal oxide + Hydrogen, and Metal oxide + Water → Metal hydroxide.}}

Highly Reactive Metals: Violent Reactions with Cold Water

Potassium and Sodium

Potassium (K) and sodium (Na) are alkali metals — the most reactive metals in the reactivity series. When even a small piece of sodium is dropped into cold water, the reaction is instantaneous and violent.

The reactions are:

2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heat energy

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat energy

What you would observe:

• The metal darts rapidly across the water surface
• It melts into a silvery ball due to the intense heat released (exothermic reaction)
• The evolved hydrogen gas catches fire, producing a characteristic yellow-orange flame (for sodium) or lilac flame (for potassium)
• The solution becomes strongly alkaline (metal hydroxide forms)

The reaction is so violent that potassium and sodium must be stored under kerosene to prevent contact with moisture in the air.

{{KEY: type=exam | title=Common Exam Question | text=You may be asked to explain why hydrogen catches fire during sodium's reaction with water. Answer: The reaction is highly exothermic, releasing enough heat to ignite the hydrogen gas produced. This is a 2-3 mark favourite question.}}

Calcium

Calcium (Ca) also reacts with cold water, but the reaction is less violent than sodium or potassium:

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

Key observations:

• Calcium reacts steadily, producing bubbles of hydrogen gas
• The metal starts floating because hydrogen bubbles stick to its surface, making it buoyant
• The heat evolved is not sufficient for the hydrogen to catch fire
• A white precipitate of calcium hydroxide (lime water) is formed, making the solution milky and alkaline

{{VISUAL: photo: calcium metal floating on water surface with hydrogen bubbles forming around it in a beaker}}

Moderately Reactive Metals: Hot Water and Steam Needed

Magnesium

Magnesium (Mg) does not react with cold water, but reacts with hot water (around 50-60°C):

Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g)

Like calcium, magnesium also floats due to hydrogen bubbles clinging to its surface. The reaction is slow and steady, producing a slightly alkaline solution of magnesium hydroxide.

However, magnesium reacts vigorously with steam, producing a bright white powder of magnesium oxide:

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

The reaction with steam is so energetic that magnesium burns with a brilliant white light — the same principle used in old photographic flash bulbs and fireworks.

{{KEY: type=points | title=Magnesium's Dual Behaviour | text=- Reacts slowly with hot water, producing Mg(OH)₂ and H₂.

• Reacts vigorously with steam, producing MgO and H₂.
• Does NOT react with cold water at all.
• This illustrates that temperature and the state of water matter in reactivity.}}

Aluminium, Zinc, and Iron: Steam Only

Metals like aluminium (Al), zinc (Zn), and iron (Fe) do not react with cold or hot water. They react only with steam when heated strongly.

The reactions are:

2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)

Zn(s) + H₂O(g) → ZnO(s) + H₂(g)

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Important observations:

• These reactions require continuous heating to keep producing steam
• A metal oxide (not hydroxide) is formed as a residue
• Hydrogen gas can be collected and tested with a burning splint (it produces a characteristic pop sound)
• Iron forms Fe₃O₄ (magnetic iron oxide), not FeO or Fe₂O₃

{{ZOOM: title=Why does iron form Fe₃O₄ and not Fe₂O₃? | text=Fe₃O₄ (magnetite) is a mixed oxide containing both Fe²⁺ and Fe³⁺ ions. At high temperatures (steam reaction), iron preferentially forms this stable mixed oxide rather than pure ferric oxide (Fe₂O₃). This is why the black magnetic oxide forms instead of the red rust we see in everyday corrosion.}}

Unreactive Metals: No Reaction with Water

Metals such as lead (Pb), copper (Cu), silver (Ag), and gold (Au) do not react with water in any form — cold, hot, or steam.

This is why:

• Copper pipes are used for water supply systems
• Silver and gold vessels can hold water indefinitely without reacting
• Ancient bronze (copper-tin alloy) and brass (copper-zinc alloy) utensils remain intact even after centuries

{{VISUAL: chart: table comparing reactivity of metals with water, showing metal names in decreasing order of reactivity, type of water they react with (cold/hot/steam/no reaction), and products formed}}

{{KEY: type=definition | title=Reactivity with Water | text=The tendency of a metal to react with water (cold, hot, or steam) to produce metal oxide or hydroxide and hydrogen gas. Highly reactive metals like Na and K react violently with cold water, while less reactive metals like Cu and Ag show no reaction at all.}}

Building the Reactivity Series

Based on the observations from Activity 3.10, we can now arrange metals in decreasing order of reactivity with water:

K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag > Au

This order is a crucial part of the reactivity series — a ranking that helps predict how metals will behave in displacement reactions, extraction processes, and corrosion. The next section will explore how metals react with acids, adding another layer to this reactivity map.

{{KEY: type=exam | title=Diagram-Based Questions | text=CBSE often asks you to draw and label the steam-on-metal apparatus (Fig. 3.3). Practice labeling: boiling flask, delivery tube, combustion tube, metal sample, heat source, and gas collection tube. This is a common 2-mark diagram question.}}

What happens when Metals react with Acids?

What happens when Metals react with Acids?

You have already learnt that metals react with acids to give a salt and hydrogen gas. This is one of the most important and frequently observed chemical reactions in everyday life — from the fizzing of antacid tablets in water to the corrosion of metal structures. But do all metals react in the same manner? Let us explore the reactivity of different metals with dilute acids through careful observation and experimentation.

The general equation for the reaction between a metal and a dilute acid is:

{{FORMULA: expr=Metal + Dilute acid → Salt + Hydrogen | symbols=Metal:reactive metal element, Dilute acid:HCl or H₂SO₄, Salt:metal chloride or sulphate, Hydrogen:H₂ gas}}

{{KEY: type=concept | title=Reaction of Metals with Dilute Acids | text=When a reactive metal is added to a dilute acid (such as hydrochloric acid or sulphuric acid), it displaces hydrogen from the acid to form a salt and hydrogen gas. The salt formed depends on the metal used and the acid: hydrochloric acid produces chlorides, while sulphuric acid produces sulphates.}}

Activity 3.11: Investigating Metal Reactivity with Dilute Hydrochloric Acid

To understand the varying reactivity of metals with acids, let us perform a systematic experiment. This activity will help us arrange metals in order of their reactivity.

Procedure

1. Collect metal samples — magnesium, aluminium, zinc, iron, and copper. Clean tarnished surfaces with sandpaper to expose fresh metal.

2. Safety first — Do NOT use sodium or potassium, as they react vigorously and dangerously even with cold water, let alone acids.

3. Set up test tubes — Place each metal sample in a separate test tube containing dilute hydrochloric acid (HCl).

4. Measure temperature change — Suspend thermometers in each test tube so that the bulbs are dipped in the acid. Record the initial temperature.

5. Observe bubble formation — Watch carefully and note the rate at which bubbles appear in each test tube.

6. Record observations — Note which metals reacted vigorously, which reacted slowly, and which showed no reaction at all.

{{VISUAL: photo: five test tubes containing different metals (Mg, Al, Zn, Fe, Cu) in dilute HCl, showing varying rates of bubble formation and thermometers suspended in each}}

{{KEY: type=points | title=Observations from Activity 3.11 | text=- Magnesium reacted most vigorously, producing bubbles rapidly and showing the highest temperature rise (most exothermic).

• Aluminium, zinc, and iron also reacted, but at progressively slower rates.
• Copper showed NO reaction — no bubbles formed and temperature remained unchanged.
• The reactivity order observed: Mg > Al > Zn > Fe > Cu.}}

Writing the Chemical Equations

Let us write balanced chemical equations for the reactions that occurred:

Magnesium with dilute hydrochloric acid:

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

Aluminium with dilute hydrochloric acid:

2Al(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂(g)

Zinc with dilute hydrochloric acid:

Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Iron with dilute hydrochloric acid:

Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

Copper with dilute hydrochloric acid:

No reaction occurs — copper is less reactive than hydrogen and cannot displace it from acids.

{{KEY: type=exam | title=Common Exam Question | text=Students are frequently asked to write balanced equations for metal-acid reactions and arrange metals in order of reactivity based on experimental observations. Remember to balance the equations carefully and state the physical states: (s) for solid, (aq) for aqueous, (g) for gas.}}

Why is Hydrogen Gas NOT Evolved with Nitric Acid?

An important exception to the general behaviour of metals with acids involves nitric acid (HNO₃). When most metals react with dilute or concentrated nitric acid, hydrogen gas is NOT produced. Why?

Nitric acid is a strong oxidising agent. It oxidises the hydrogen gas (H₂) that would normally be produced to water (H₂O), and in the process, nitric acid itself gets reduced to various nitrogen oxides such as:

• Nitrous oxide (N₂O)
• Nitric oxide (NO)
• Nitrogen dioxide (NO₂)

The reactions are more complex and depend on the concentration of the acid and the nature of the metal.

{{ZOOM: title=Exception — Very Dilute HNO₃ | text=Magnesium (Mg) and manganese (Mn) are exceptions. When treated with very dilute nitric acid, they DO evolve hydrogen gas because the oxidising power of very dilute HNO₃ is significantly reduced.}}

{{VISUAL: diagram: flowchart showing the reaction pathway of metals with HNO₃, illustrating how H₂ is oxidised to H₂O and HNO₃ is reduced to NO, NO₂, or N₂O depending on concentration}}

{{KEY: type=definition | title=Aqua Regia | text=Aqua regia (Latin for 'royal water') is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1. It is a highly corrosive, fuming liquid capable of dissolving noble metals like gold and platinum, which neither acid can dissolve alone.}}

Understanding Reactivity Trends

The experiment clearly demonstrates that metals have different reactivities when treated with dilute acids. This difference arises from the tendency of metal atoms to lose electrons and form positive ions (cations).

From the observations, we can deduce the following reactivity series for the metals tested:

Mg > Al > Zn > Fe > Cu

This order will be crucial when we study the complete Activity Series in the next section, which ranks all common metals by their reactivity with various reagents.

{{KEY: type=concept | title=Exothermic Nature of Metal-Acid Reactions | text=The reaction of reactive metals with dilute acids is exothermic, meaning it releases heat energy. The more reactive the metal, the greater the heat released. Magnesium showed the highest temperature rise because it is the most reactive metal in the series tested, releasing the most energy during the reaction.}}

Test of Hydrogen Gas

When metals react with dilute acids, the gas evolved is hydrogen (H₂). How can we confirm this?

Test: Bring a burning matchstick or splint near the mouth of the test tube containing the gas.

Observation: Hydrogen burns with a pop sound — this is the characteristic test for hydrogen gas.

{{VISUAL: photo: a test tube with bubbles of hydrogen gas rising, and a lit splint near the mouth producing a small flame with a popping sound}}

Chemical equation for hydrogen combustion:

2H₂(g) + O₂(g) → 2H₂O(l)

This simple test confirms that the gas produced during metal-acid reactions is indeed hydrogen.

Practical Applications and Real-World Connections

Understanding how metals react with acids has several real-world applications:

• Antacid tablets contain metal carbonates or bicarbonates that react with stomach acid (HCl) to neutralise excess acidity and produce carbon dioxide gas (fizzing).

• Metal etching and cleaning — dilute acids are used to remove oxide layers and clean metal surfaces in industry.

• Hydrogen production — reactive metals like zinc reacting with dilute acids are sometimes used in laboratory-scale hydrogen generation.

• Corrosion prevention — understanding reactivity helps engineers choose appropriate metals for construction in acidic environments (e.g., coastal areas, chemical plants).

{{KEY: type=exam | title=Exam Strategy | text=In CBSE exams, you may be asked to compare the reactivity of metals based on experimental observations, write balanced equations, or explain why certain metals do not react with acids. Always mention the test for hydrogen gas (pop sound) when describing metal-acid reactions.}}

By systematically investigating how different metals react with dilute acids, we have built a foundation for understanding the Activity Series — a powerful tool that ranks metals by their chemical reactivity and predicts which reactions will occur. This knowledge will help us explain displacement reactions, extraction of metals from ores, and everyday phenomena like rusting and corrosion in the sections ahead.