NCERT Science Notes - Class 10
Chapter 3 - Metals and Non-metals

Welcome to AJs Chalo Seekhen. This webpage is dedicated to Class 10 | Science | Chapter 3 - Metals and Non-metals. The chapter delves into the properties and reactions of metals and non-metals. It explores their physical and chemical properties, how metals react with oxygen, water, dilute acids, and other metal salts, forming various compounds and exhibiting different behaviors. The chapter also discusses the concept of a reactivity series, where metals are arranged based on their reactivity levels. Additionally, it examines the nature of metal oxides and how metals displace one another in compounds, a process known as the displacement reaction. This chapter provides foundational knowledge of metals and non-metals, which is crucial for understanding various chemical processes and the principles of inorganic chemistry​.

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NCERT Science Notes - Class 10
Chapter 3 - Metals and Non-metals

3.1.1 - Metals

• Grouping Substances by Physical Properties: One way to categorize substances is by comparing their physical properties.
• Activities 3.1 to 3.6: These activities involve the collection and examination of various metals: iron, copper, aluminum, magnesium, sodium, lead, zinc, and any other readily available metal.
  
  

• Samples of Iron, Copper, Aluminum, and Magnesium: Obtain samples of these metals and observe their initial appearance.
• Cleaning Metal Surfaces: Clean the surface of each metal sample by rubbing them with sandpaper and note their appearance after cleaning.
  • Metallic Lustre: Metals in their pure state exhibit a shining surface, a property known as metallic lustre.

• Small Metal Pieces: Take small pieces of iron, copper, aluminum, and magnesium.
• Attempt to Cut Metals: Try to cut these metals with a sharp knife and record your observations.
• Handling Sodium Metal: Handle a piece of sodium metal with caution. Dry it by pressing between folds of a filter paper.
• Attempt to Cut Sodium Metal: Place the sodium metal on a watch-glass and try to cut it with a knife.
  • Metal Hardness: Metals are generally hard substances and their hardness can vary from one metal to another.

• Metals Tested: Take pieces of iron, zinc, lead, and copper.
• Impact Testing: Place each metal on a block of iron and strike it four or five times with a hammer.
• Observations: When striking the metals with a hammer, iron, and copper undergo a change in shape. They can be beaten into thin sheets without breaking. Zinc and lead do not change shape easily and do not exhibit malleability.
  • Malleability: Some metals can be beaten into thin sheets, a property known as malleability. Gold and silver are known to be the most malleable metals.

Activity 3.4

• Daily Life Metals: List the metals whose wires you have encountered in daily life.
  • Ductility: The ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal, and a significant length of wire can be drawn from a small amount of gold.

Activity 3.5

• Wire Experiment: Take an aluminum or copper wire and clamp it on a stand.
• Fixing Pin: Fix a pin to the free end of the wire using wax.
• Heating the Wire: Heat the wire near the clamped end using a spirit lamp, candle, or burner.
• Observations: When heating the aluminum or copper wire, it does not melt. Instead, it remains intact. This demonstrates that metals have high melting points and can withstand heat without melting under normal conditions.
  • Heat Conductivity: also known as thermal conductivity, refers to a material's ability to conduct heat efficiently. Materials with high heat conductivity transfer heat rapidly, whereas those with low heat conductivity are poor heat conductors. This activity demonstrates that metals are good conductors of heat and have high melting points. Silver and copper are excellent heat conductors.

• Electric Circuit Setup: Set up an electric circuit with terminals A and B.
• Metal Testing: Insert the metal to be tested between terminals A and B in the circuit.
• Observations: When the metal is placed between terminals A and B in the electric circuit, the bulb in the circuit glows. This indicates that metals are good conductors of electricity.
  • Electrical Conductivity: is the ability of a material to conduct an electric current. Materials that readily allow the flow of electricity are considered good conductors, while those that inhibit or resist electric current flow are insulators. Metals are known to conduct electricity. This activity confirms their electrical conductivity.

3.1.2 Non-metals

Discussion:

• Grouping Elements by Physical Properties: It is evident that elements cannot be grouped solely based on physical properties, as there are exceptions.
• States at Room Temperature: Most metals are solid at room temperature, except for mercury, which is a liquid. Gallium and Caesium, however, have very low melting points and can melt on contact with a human palm.
• Lustrous Non-Metal: Iodine, despite being a non-metal, exhibits luster.
• Allotropes of Carbon: Carbon exists in different forms called allotropes. Diamond is extremely hard with a high melting point, while graphite is a conductor of electricity.
• Alkali Metals: Elements like lithium, sodium, and potassium are extremely soft, have low densities, and low melting points.

Activity 3.8

• Magnesium Burning:
  • Burned magnesium ribbon.
  • Collected the ashes formed and dissolved them in water.
  • Tested the resultant solution with red and blue litmus paper.
• Sulphur Burning:
  • Burned sulphur powder.
  • Collected the fumes produced in a test tube and added water.
  • Tested the solution with blue and red litmus paper.

• The product formed on burning magnesium is basic, as indicated by its effect on red litmus paper.
• The product formed on burning sulfur is acidic, as indicated by its effect on blue litmus paper.

• Magnesium Burning: Mg + O₂ → MgO (Basic oxide)
• Sulphur Burning: S + O₂ → SO₂ → H₂O + SO₂ → H₂SO₃ (Sulphurous acid)

• Most non-metals produce acidic oxides when dissolved in water.
• Most metals produce basic oxides when dissolved in water.

3.2 CHEMICAL PROPERTIES OF METALS

Section Overview

• Examines the chemical properties of metals in Sections 3.2.1 to 3.2.4.
• Samples needed: Aluminium, copper, iron, lead, magnesium, zinc and sodium.

3.2.1 What happens when Metals are burnt in Air?

Do You Know?

Let's integrate the information from the "Do You Know?" section into our summary:

Anodising of Aluminium:

• Anodising is a specialized process to create a thick oxide layer on aluminium.
• Aluminium naturally forms a thin oxide layer when exposed to air, which protects it from corrosion.
• Anodising enhances this protective layer by making it thicker.
• During anodising, an aluminium article is used as the anode in an electrolytic cell with dilute sulphuric acid.
• Oxygen released at the anode reacts with the aluminium to thicken the oxide layer.
• The resulting oxide layer can be dyed to give the aluminium a decorative and durable finish.

Observations on Metal Reactivity:

• Activity 3.9 should have demonstrated that sodium is highly reactive, spontaneously reacting with oxygen.
• Magnesium's reaction, while visible, is less vigorous than sodium's, indicating lower reactivity.
• The activity does not provide a clear conclusion on the reactivity order for zinc, iron, copper, or lead.
• Further experiments are suggested to better understand the reactivity series of these metals.

In summary, anodising is a valuable industrial process for improving the durability and aesthetic appeal of aluminium products. Observations from educational experiments like Activity 3.9 are crucial in teaching students about the reactivity of metals, although not all reactivity levels can be conclusively determined from such simple experiments. Further investigation is needed to accurately place metals in a reactivity series.

• Sodium is the most reactive of the tested metals, as its reaction is described as "the most vigorous."
• Magnesium is less reactive than sodium, as its reaction is "less vigorous."
• The passage suggests that burning in oxygen alone is not enough to determine the reactivity of zinc, iron, copper, or lead. This may imply that the reactions of these metals in the experiment were not as distinct or conclusive as the reactions of sodium and magnesium.
• The passage indicates a need for additional reactions or experiments to determine the relative reactivity of zinc, iron, copper, and lead more precisely.

3.2.2 What happens when Metals react with Water?

Activity 3.10:

1. Safety and Preparation:
   • Requires teacher assistance.
   • Use the same metal samples as in Activity 3.9.
2. Experiment Steps:
   • Place small pieces of metals in separate beakers filled halfway with cold water.
   • Observe which metals react with cold water and arrange them by reactivity.
   • Check for any metals that produce fire or start floating.
   • Place unreactive metals in hot water and observe.
   • For metals unreactive with hot water, set up an apparatus to expose them to steam and observe reactions.
3. Observations and Arrangement:
   • Identify metals that do not react even with steam.
   • Arrange all tested metals by their reactivity with water (decreasing order).
4. Chemical Reactions:
   • Metals reacting with water typically form metal oxide and hydrogen gas.
     Metal + Water → Metal oxide + Hydrogen
     Metal oxide + Water → Metal hydroxide
   • Soluble metal oxides further react with water to form metal hydroxides.
   • Not all metals react with water.
5. Examples of Reactions:
   • Potassium and sodium react violently with cold water, producing heat that ignites hydrogen:
     2K(s) + 2H₂​O(l) → 2KOH(aq) + H₂​(g) + heat energy
     2Na(s) + 2H₂​O(l) → 2NaOH(aq) + H₂​(g) + heat energy
   • Calcium's reaction with water is less violent, not igniting the hydrogen:
     Ca(s) + 2H₂​O(l) → Ca(OH)₂​(aq) + H₂​(g)
   • Magnesium reacts with hot water to form magnesium hydroxide and hydrogen:
     Mg(s) + 2H₂​O(l) → Mg(OH)₂​(aq) + H₂​(g)
   • Aluminium, iron, and zinc react with steam but not with cold or hot water:
     2Al(s) + 3H₂​O(g) → Al₂​O3​(s) + 3H₂​(g) 
     3Fe(s) + 4H₂​O(g) → Fe₃​O₄​(s) + 4H₂​(g)
   • Lead, copper, silver, and gold do not react with water at all.

3.2.3 What happens when Metals react with Acids?

Activity 3.11

1. Objective:
   • To investigate how different metals react with dilute acids.
2. Safety and Preparation:
   • Exclude sodium and potassium due to their vigorous reactions with water.
   • Clean tarnished metal samples with sandpaper.
3. Procedure:
   • Place the clean metal samples in separate test tubes containing dilute hydrochloric acid.
   • Insert thermometers into the test tubes to measure temperature changes.
   • Observe the rate of bubble formation (indicating hydrogen gas production).
4. Observations:
   • Identify which metals react vigorously with the acid.
   • Record the temperature to determine the most exothermic reaction.
   • Arrange the metals in order of decreasing reactivity with dilute hydrochloric acid.
5. Chemical Reactions with Hydrochloric Acid:
   • Write balanced chemical equations for reactions of magnesium, aluminium, zinc, and iron with dilute hydrochloric acid, reflecting the production of respective salts and hydrogen gas.
6. Exception with Nitric Acid:
   • Normally, hydrogen gas is not produced when metals react with nitric acid due to its oxidizing nature.
   • However, magnesium and manganese can produce hydrogen gas when reacting with very dilute nitric acid.
7. Observational Conclusions:
   • The fastest bubble formation and most exothermic reaction occur with magnesium.
   • The reactivity order observed is Mg > Al > Zn > Fe.
   • Copper does not react with dilute hydrochloric acid, as indicated by the absence of bubbles and no change in temperature.

Interesting Fact (Aqua Regia):

• Aqua regia is a potent mixture of hydrochloric acid and nitric acid in a 3:1 ratio.
• It can dissolve gold and platinum, despite these metals being resistant to individual acids.
• Aqua regia is highly corrosive and fuming.

3.2.4 - How do Metals react with Solutions of other Metal Salts?

Activity 3.12

1. Objective:
   • To investigate how metals react with solutions of salts of other metals.
2. Experiment Setup:
   • Use a clean copper wire and an iron nail.
   • Place the copper wire in a solution of iron sulphate and the iron nail in a solution of copper sulphate in separate test tubes.
3. Procedure:
   • Observe the reactions over 20 minutes.
4. Observations and Inferences:
   • Identify the test tube where a reaction has occurred.
   • Determine the evidence of reaction, typically by observing color changes or metal deposition.
   • Compare observations to previous activities to understand reactivity trends.
5. Chemical Equations and Reaction Types:
   • Write balanced chemical equations for the observed reactions.
   • Determine the type of reaction, likely to be a single displacement (or substitution) reaction.
6. Conceptual Understanding:
   • Displacement reactions allow us to compare the reactivity of different metals.
   • The reactivity series can be established by observing which metals can displace others from their compounds.
   • The general equation for displacement reactions is: $\text{Metal A}+\text{Salt solution of B}\to \text{Salt solution of A}+\text{Metal B}$
     
7. Conclusions from Activity 3.12:
   • The metal that displaces another from its salt solution is more reactive.
   • Based on the activity, students will conclude whether copper or iron is more reactive.

Through Activity 3.12, students can witness firsthand the displacement reactions between metals and metal salt solutions. This activity helps in understanding the concept of reactivity series, which ranks metals based on their ability to displace others from their compounds. The reactivity series is an important concept in chemistry that explains the outcomes of various reactions, including those with oxygen, water, and acids, as explored in previous activities (3.9, 3.10, and 3.11).

3.2.5 - The Reactivity Series

The reactivity series is a list of metals arranged in the order of their decreasing activities. After performing displacement experiments (Activities 1.9 and 3.12), the following series, known as the reactivity or activity series, has been developed. 

 Table 3.2 Activity Series: Relative Reactivities of Metals

• Potassium (K) is listed as the most reactive metal.
• The reactivity decreases in the following order: Potassium (K), Sodium (Na), Calcium (Ca), Magnesium (Mg), Aluminium (Al), Zinc (Zn), Iron (Fe), Lead (Pb), [Hydrogen (H)], Copper (Cu), Mercury (Hg), Silver (Ag), and Gold (Au) as the least reactive.
• The position of hydrogen (H) in the series is significant as it allows us to predict whether a metal can displace hydrogen from an acid.
• Metals above hydrogen in the series can displace hydrogen from dilute acids, whereas those below cannot.
• The reactivity series is useful for predicting the outcomes of reactions involving metals and their compounds.

This series is fundamental in understanding the behavior of metals in various chemical reactions, including their interactions with water, acids, and other metal salts.

3.3 How Do Metals and Non-Metals React?

• The reactivity of elements is due to their tendency to achieve a completely filled valence shell.
• Noble gases have a full valence shell, which is why they exhibit minimal chemical activity.
• Metals and non-metals react to attain stable electronic configurations similar to noble gases.
  
  

Electronic Configurations:

• A sodium atom has one electron in its outermost shell (M shell). By losing this electron, its next shell (L shell) becomes the outermost, achieving a stable octet configuration.
• After losing an electron, the sodium atom becomes a cation (Na+) with a net positive charge due to having more protons (11) than electrons (10).
• Chlorine, with seven electrons in its outermost shell, requires one more electron to complete its octet.
  
  

Formation of Sodium Chloride (NaCl):

• Sodium can donate its outer electron to chlorine, which needs an extra electron.
• The chlorine atom gains an electron and becomes a chloride anion (Cl⁻) with a net negative charge.
• Sodium ions (Na⁺) and chloride ions (Cl⁻), being oppositely charged, attract each other and form sodium chloride.
• Sodium chloride exists as an aggregate of ions rather than discrete molecules.
  
  

Formation of Magnesium Chloride (MgCl₂):

• Magnesium (Mg) loses two electrons to form a magnesium cation (Mg²⁺).
• Each chlorine atom gains one electron to form chloride anions (Cl⁻).
• The compound formed is magnesium chloride (MgCl₂), an ionic compound where the magnesium ion is the cation and the chloride ions are the anions.
  
  

Ionic Compounds:

• Ionic or electrovalent compounds are formed by the transfer of electrons from metals to non-metals.
• These compounds consist of cations (positive ions) and anions (negative ions) held together by electrostatic forces.
  
  

Electronic Configurations of Some Elements (Table 3.3):

• Lists the electronic configurations of noble gases (Helium, Neon, Argon) and various metals (Sodium, Magnesium, Aluminium) and non-metals (Nitrogen, Oxygen, Fluorine, Phosphorus, Sulfur, Chlorine).
• The electronic configurations are noted in terms of the number of electrons in the K, L, M, and N shells.

3.3.1 - Properties of Ionic Compounds

You may have observed the following general properties for ionic compounds —

1. Physical nature: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. These compounds are generally brittle and break into pieces when pressure is applied.
2. Melting and Boiling points: Ionic compounds have high melting and boiling points (see Table 3.4). This is because a considerable amount of energy is required to break the strong inter-ionic attraction.
   
3. Solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.
4. Conduction of Electricity: The conduction of electricity through a solution involves the movement of charged particles. A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution. Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure. But ionic compounds conduct electricity in the molten state. This is possible in the molten state since the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

3.4.1 - Extraction of Metals

• The reactivity series of metals helps in understanding the extraction process.
• Metals found in the free state are typically at the bottom of the reactivity series (e.g., gold, silver, platinum, copper).
• Highly reactive metals (e.g., potassium, sodium, calcium, magnesium, and aluminum) are never found free in nature.
• Metals of medium reactivity (e.g., zinc, iron, lead) are commonly found as oxides, sulphides, or carbonates.
• Oxygen's abundance and reactivity mean that many metal ores are oxides.
  
  
Categories of Metals Based on Reactivity:
1. Metals of Low Reactivity: These metals like gold (Au) and silver (Ag) are often found in their native state.
2. Metals of Medium Reactivity: Metals like zinc (Zn), iron (Fe), and lead (Pb) are found as oxides, sulphides, or carbonates and are usually reduced using carbon.
3. Metals of High Reactivity: Metals such as potassium (K), sodium (Na), calcium (Ca), magnesium (Mg), and aluminum (Al) are extracted by electrolysis of their molten ores.
   
   
Steps Involved in the Extraction of Metals:
• Concentration of Ore: The first step in the extraction process is the removal of impurities to increase the metal content.
• Reduction to Metal:
  • For metals of high reactivity, electrolysis of molten ore is used to obtain the pure metal.
  • For metals of medium reactivity, reduction of metal oxide is achieved through processes such as calcination (for carbonate ores) and roasting (for sulphide ores).
  • For metals of low reactivity, reduction may not be necessary as they are often found in their native state. If found as compounds, roasting is used followed by refining.
• Purification of Metal: The final step is the purification of the metal to obtain its pure form.
In summary, the method of extraction and processing depends on the metal's reactivity and the form in which it is found in the ore. This systematic approach from concentration to purification enables the efficient production of metals for various uses.

3.4.2 - Enrichment of Ores

1. Ore and Gangue: Ores are sources of metals, contaminated with impurities known as gangue, like soil and sand.
2. Need for Removal of Gangue: These impurities must be removed before metal extraction for efficiency.
3. Separation Techniques: This is done by exploiting differences in physical or chemical properties between the ore and gangue using various techniques.

3.4.3 - Extracting Metals Low in the Activity Series

1. Low Reactivity Metals: Metals at the bottom of the activity series are less reactive and easily reduced.
2. Example - Mercury (Cinnabar - HgS): The process involves heating cinnabar to form mercuric oxide, then mercury.
   • 2HgS(s) + 3O₂(g) → 2HgO(s) + 2SO₂(g)
   • 2HgO(s) → 2Hg(l) + O₂(g)
3. Example - Copper (Cu₂S): Copper is extracted by heating its sulfide ore in air.
   • 2Cu₂S + 3O₂(g) → 2Cu₂O(s) + 2SO₂(g)
   • 2Cu₂O + Cu₂S → 6Cu(s) + SO₂(g)

3.4.4 - Extracting Metals in the Middle of the Activity Series

1. Moderately Reactive Metals: Metals like iron, zinc, lead, found as sulfides or carbonates.
2. Conversion to Oxides:
   • Roasting: Sulfide ores are converted into oxides by heating in excess air.
     • Example Reaction (Roasting of Zinc Ore): 2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)
   • Calcination: Carbonate ores are converted into oxides by heating in limited air.
     • Example Reaction (Calcination of Zinc Carbonate): ZnCO₃(s) → ZnO(s) + CO₂(g)
3. Reduction to Metals: Metal oxides are reduced to metals, like zinc oxide to zinc metal.
   • Example Reaction (Reduction of Zinc Oxide): ZnO(s) + C(s) → Zn(s) + CO(g)
4. Oxidation and Reduction: The process of obtaining metals from their compounds is a reduction process.
5. Displacement Reactions: Highly reactive metals can displace less reactive metals from their compounds.
• Example Reaction (Using Aluminium): 3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + Heat
• Example of Thermit Reaction: Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat

Notes for Understanding:

• Roasting and Calcination: These are preparatory processes that convert sulphides and carbonates into oxides, which are easier to reduce to metals.
• Reduction Process: The metal oxides obtained are then reduced to pure metals. This reduction is often done using carbon.
• Displacement Reactions: Highly reactive metals like aluminium can displace less reactive metals from their compounds. These reactions are typically very exothermic, producing metals in a molten state.
• Thermit Reaction: This is a specific type of displacement reaction used in practical applications like welding railway tracks, where iron oxide is reduced by aluminium.

3.4.5 Extracting Metals towards the Top of the Activity Series

Note: This section discusses the extraction of highly reactive metals, which cannot be reduced using carbon due to their high affinity for oxygen.

Electrolytic Reduction

Metals like sodium, magnesium, and calcium are obtained through the electrolysis of their molten chlorides. During this process:

• The metals are deposited at the cathode (negatively charged electrode).
• Chlorine is liberated at the anode (positively charged electrode).

Example Reactions:

At cathode: Na⁺ + e^– → Na

At anode: 2Cl^– → Cl₂ + 2e^–

Aluminium Extraction

Aluminium is obtained by the electrolytic reduction of aluminium oxide.

3.4.6 Refining of Metals

Note: Metals obtained through various reduction processes contain impurities and require refining to achieve purity.

Electrolytic Refining

This is a widely used method for refining metals like copper, zinc, tin, nickel, silver, and gold. The process involves:

• The impure metal as the anode and a thin strip of pure metal as the cathode.
• Using a solution of the metal salt as an electrolyte.

During electrolysis:

• Pure metal from the anode dissolves into the electrolyte.
• An equivalent amount of pure metal is deposited on the cathode.
• Soluble impurities dissolve in the solution, while insoluble impurities settle as anode mud.

Notes for Understanding:

• Highly Reactive Metals: Metals at the top of the activity series are too reactive to be reduced by carbon. They require a different method, typically electrolytic reduction.
• Electrolytic Reduction: This involves using an electric current to decompose the compound of the metal. For metals like sodium, magnesium, and calcium, this is done by electrolyzing their molten chlorides.
• Extraction of Aluminium: Aluminium is extracted by the electrolysis of aluminium oxide, a process different from other metals due to its unique properties.
• Refining Process: The metals obtained from these processes often contain impurities. Electrolytic refining is a common method to purify these metals. This involves using the impure metal as the anode and a pure metal strip as the cathode, with a suitable electrolyte.

3.5 CORROSION

Overview from Chapter 1: Corrosion is a natural process where metals deteriorate due to reactions with substances in the environment.

• Silver Corrosion: Silver articles become black due to the formation of silver sulphide from the reaction with sulphur in the air.
• Copper Corrosion: Copper loses its shiny brown surface and forms a green coat of basic copper carbonate when reacting with moist carbon dioxide.
• Iron Rusting: Iron develops a coating of rust, a brown flaky substance, when exposed to moist air over time.
  
  

Activity 3.14 : Understanding Iron Rusting

Experiment to determine the conditions for iron rusting:

1. Setup: Place clean iron nails in three separate test tubes labeled A, B, and C.
2. Test Tube A: Add water and cork it. This exposes the nails to both air and water.
3. Test Tube B: Add boiled distilled water, 1 mL of oil, and cork it. The oil layer prevents air from dissolving in the water, exposing the nails only to water.
4. Test Tube C: Add anhydrous calcium chloride and cork it. The calcium chloride absorbs moisture, exposing the nails to dry air.

After a few days, observe the changes:

• Observation: Rust forms in Test Tube A, but not in B and C.
• Conclusion: This indicates that both air and water are necessary for iron to rust.

Notes for Understanding:

• Corrosion of Different Metals: Each metal corrodes differently due to its unique chemical properties. Silver tarnishes, copper forms a green patina, and iron rusts.
• Activity 3.14 - Experiment on Iron Rusting:
  • Purpose: To find out the conditions under which iron rusts.
  • Method: Iron nails are placed in three different conditions involving air, water, and oil or desiccants.
  • Results: Rusting occurs only in the presence of both air and water, as seen in Test Tube A. Test Tubes B and C, which lack either air or water, do not show rusting.
  • Conclusion: The experiment demonstrates that both oxygen (from air) and water are necessary for the rusting of iron.

3.5.1 - Prevention of Corrosion

CBSE Class 10 Science Chapter 3 - Metals and Non-metals Notes