Class 10 Science Chapter 3 Notes

Metals and Non-metals Class 10 Science Chapter 3 Notes are available here. These notes are prepared by the subject experts of our team.

Metals and Non-metals Class 10 Science Chapter 3 Notes

Metals

1. Grouping Substances by Physical Properties:

• Metals can be grouped based on physical properties such as appearance, hardness, ductility, conductivity, and sonority.

2. Metallic Lustre:

• Pure metals have a shining surface known as metallic lustre.

3. Hardness of Metals:

• Metals are generally hard, but hardness varies between metals.
• Sodium is softer and can be cut with a knife.

4. Ductility:

• Metals can be drawn into thin wires (ductility).
• Gold is the most ductile metal, capable of forming very long wires from a small amount.

5. Uses in Cooking Vessels:

• Metals like aluminium and copper are used for cooking vessels due to their properties.

6. Conductivity and Melting Points:

• Metals are good conductors of heat and have high melting points.
• Good conductors: Silver and copper.
• Poor conductors: Lead and mercury.

7. Electrical Conductivity

• Metals conduct electricity, evident when the bulb in an electric circuit glows.
• Electric wires are coated with insulating materials like PVC or rubber.

8. Sonority:

• Metals produce sound when struck on a hard surface (sonorous).
• School bells are made of metals due to this property.

Non Metals

1. Fewer Non-Metals than Metals:

• Examples: Carbon, sulphur, iodine, oxygen, hydrogen, bromine (liquid).

2. Physical Properties of Non-Metals:

• Non-metals can be solids or gases (except liquid bromine).
• Non-metals generally do not share the same physical properties as metals.

3. Exceptions in Physical Properties:

• Most metals are solid at room temperature; mercury is liquid.
• Metals generally have high melting points, but gallium and caesium have very low melting points.
• Iodine, a non-metal, is lustrous.
• Carbon, a non-metal, has allotropes (e.g., diamond is the hardest natural substance, graphite conducts electricity).
• Alkali metals (lithium, sodium, potassium) are soft and have low densities and melting points.

4. Chemical Properties for Classification:

• Elements are more accurately classified as metals or non-metals based on chemical properties.

5. Chemical Properties:

• Magnesium ribbon burns to form basic oxides.
• Sulphur powder burns to form acidic oxides.
• Non-metals typically produce acidic oxides in water.
• Metals typically produce basic oxides in water.

Chemical Properties of Metals

1. Reaction of Metals with Oxygen:

• Metals react with oxygen to form metal oxides.
• Example reactions:
  • Copper + Oxygen → Copper(II) oxide (black oxide)
  • Aluminium + Oxygen → Aluminium oxide

2. Reactivity Order:

• Metals can be arranged in decreasing order of reactivity based on their reactions with oxygen.

3. Amphoteric Oxides:

• Some metal oxides (e.g., aluminium oxide, zinc oxide) show both acidic and basic behavior.
• Reactions with acids and bases:
  • Aluminium oxide + Hydrochloric acid → Aluminium chloride + Water
  • Aluminium oxide + Sodium hydroxide → Sodium aluminate + Water

4. Solubility of Metal Oxides:

• Most metal oxides are insoluble in water.
• Some (e.g., sodium oxide, potassium oxide) dissolve in water to form alkalis:
  • Sodium oxide + Water → Sodium hydroxide
  • Potassium oxide + Water → Potassium hydroxide

5. Different Reactivities:

• Metals like potassium and sodium react very vigorously with oxygen and are stored in kerosene oil to prevent fires.
• Metals like magnesium, aluminium, zinc, and lead form protective oxide layers.
• Iron filings burn vigorously, while solid iron and copper do not burn easily.
• Silver and gold do not react with oxygen even at high temperatures.

6. Conclusion on Reactivity:

• Sodium is highly reactive.
• Magnesium is less reactive than sodium.
• Further tests needed to determine the reactivity of zinc, iron, copper, and lead.

Anodising

1. Anodising Definition:

• Process of forming a thick oxide layer on aluminium.

2. Natural Oxide Layer:

• Aluminium naturally develops a thin oxide layer when exposed to air, providing some corrosion resistance.

3. Enhanced Corrosion Resistance:

• The resistance to corrosion can be improved by making the oxide layer thicker.

4. Anodising Process:

• Clean aluminium article is made the anode in an electrolytic cell.
• Electrolysis is performed using dilute sulphuric acid.
• Oxygen gas evolved at the anode reacts with aluminium to form a thicker oxide layer.

5. Advantages:

• Thicker oxide layer provides better protection.
• Oxide layer can be dyed to give aluminium articles an attractive finish.

What happens when Metals react with Water?

1. Chemical Reactions:

• Metals react with water to produce metal oxides and hydrogen gas.
• Soluble metal oxides form metal hydroxides in water.
• General reactions:
  • Metal + Water → Metal oxide + Hydrogen
  • Metal oxide + Water → Metal hydroxide

2. Reactivity of Specific Metals:

• Potassium and Sodium: React violently with cold water; the hydrogen gas ignites.
• Calcium: Reacts less violently; hydrogen gas bubbles make it float.
• Magnesium: Does not react with cold water; reacts with hot water and floats.
• Aluminium, Iron, Zinc: Do not react with cold or hot water; react with steam.
• Lead, Copper, Silver, Gold: Do not react with water at all.

What happens when Metals react with Acids?

1. Reactions with Dilute Hydrochloric Acid:

• General reaction: Metal + Dilute acid → Salt + Hydrogen
• Record which metals react vigorously and produce the highest temperature.
• Write equations for reactions of magnesium, aluminium, zinc, and iron with dilute hydrochloric acid.

2. Reactions with Nitric Acid:

• Hydrogen gas is not evolved with nitric acid due to its strong oxidizing properties.
• Nitric acid oxidizes hydrogen gas to water and itself gets reduced to nitrogen oxides.
• Magnesium and manganese can react with very dilute nitric acid to produce hydrogen gas.

3. Observations and Reactivity:

• Magnesium reacts the fastest and most exothermically.
• Reactivity order with dilute hydrochloric acid: Mg > Al > Zn > Fe.
• Copper does not react with dilute hydrochloric acid, showing no bubbles or temperature change.

Aqua Regia

• Aqua regia is a mixture of concentrated hydrochloric acid and concentrated nitric acid.
• The ratio of hydrochloric acid to nitric acid is 3:1.
• It is a freshly prepared, highly corrosive, fuming liquid.
• Aqua regia can dissolve gold and platinum, which neither of the acids can do alone.

How do Metals react with Solutions of other Metal Salts?

• Reactive metals can displace less reactive metals from their compounds in solution or molten form.
• The reactivity of metals can be compared by observing displacement reactions.
• General reaction format: Metal A + Salt solution of Metal B → Salt solution of Metal A + Metal B.

Reactivity Series

1. Reactivity Series Definition: The reactivity series is a list of metals arranged in order of decreasing reactivity.

• Most Reactive:
  • K (Potassium)
  • Na (Sodium)
  • Ca (Calcium)
  • Mg (Magnesium)
  • Al (Aluminium)
  • Zn (Zinc)
• Intermediate Reactivity:
  • Fe (Iron)
  • Sn (Tin)
  • Pb (Lead)
  • [H] (Hydrogen)
• Least Reactive:
  • Cu (Copper)
  • Hg (Mercury)
  • Ag (Silver)
  • Au (Gold)

How Do Metals and Non Metals React

1. Reactivity and Electronic Configuration:

• Metals and non-metals react to attain a completely filled valence shell, similar to noble gases which are chemically inactive due to their stable electronic configuration.

2. Electronic Configurations:

• Noble Gases: Have completely filled valence shells (e.g., Neon: 2,8; Argon: 2,8,8).
• Metals: Tend to lose electrons to achieve a stable configuration (e.g., Sodium: 2,8,1 becomes Na⁺ with 2,8).
• Non-Metals: Tend to gain electrons to achieve a stable configuration (e.g., Chlorine: 2,8,7 becomes Cl⁻ with 2,8,8).

3. Reaction Examples:

• Sodium and Chlorine:
  • Sodium loses one electron to form Na⁺.
  • Chlorine gains one electron to form Cl⁻.
  • Sodium chloride (NaCl) forms from the electrostatic attraction between Na⁺ and Cl⁻ ions.
• Magnesium and Chlorine:
  • Magnesium loses two electrons to form Mg²⁺.
  • Chlorine gains one electron each to form two Cl⁻ ions.
  • Magnesium chloride (MgCl₂) forms from the electrostatic attraction between Mg²⁺ and Cl⁻ ions.

4. Ionic Compounds:

• Compounds formed by the transfer of electrons from metals to non-metals are called ionic or electrovalent compounds.
• Example of ionic compounds: Sodium chloride (NaCl) and Magnesium chloride (MgCl₂).

5. Ions in MgCl₂:

• Cation: Mg²⁺
• Anion: Cl⁻

Properties of Ionic Compounds

1. Physical Nature:

• Ionic compounds are solid and somewhat hard due to strong attraction between ions.
• They are generally brittle and break under pressure.
• Ionic compounds have high melting and boiling points, indicating strong inter-ionic attraction.
• Ionic compounds are generally soluble in water.
• Insoluble in solvents like kerosene and petrol.
• Ionic compounds in solid form do not conduct electricity due to the rigid structure preventing ion movement.
• They conduct electricity in molten state as ions move freely when electrostatic forces are overcome by heat.

Extraction of Metals

1. Metals are found either in a free state or as compounds.

2. Metal Categories Based on Reactivity:

• Least Reactive Metals:
  • Found in a free state in the earth’s crust.
  • Examples: Gold, silver, platinum, copper.
• Moderately Reactive Metals:
  • Found mainly as oxides, sulphides, or carbonates.
  • Examples: Zinc, iron, lead.
• Highly Reactive Metals:
  • Never found in nature as free elements due to their reactivity.
  • Examples: Potassium, sodium, calcium, magnesium, aluminum.

3. Occurrence of Metals:

• Free State: Least reactive metals like gold and silver.
• Combined State: Copper and silver as sulphides or oxides.
• Oxide Ores: Many metals form oxides due to the abundance and reactivity of oxygen.

4. Extraction Techniques:

• Low Reactivity Metals: Found in native state, extracted with minimal effort.
• Medium Reactivity Metals: Reduced using carbon.
• High Reactivity Metals: Extracted via electrolysis.

Enrichment of Ores

• Ores mined from the earth contain impurities like soil and sand, known as gangue.
• Impurities must be removed before metal extraction.
• The processes depend on the physical or chemical property differences between the gangue and the ore.
• Various techniques are used to separate the gangue from the ore based on these differences.

Extracting Metals in the Middle of the Activity Series

1. Moderately Reactive Metals:

• Examples: Iron, zinc, lead, copper.
• Usually present as sulphides or carbonates in nature.

2. Conversion to Oxides:

• Easier to extract metals from oxides than from sulphides or carbonates.
• Roasting: Converts sulphide ores to oxides by heating in excess air.
  • Example: 2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g)
• Calcination: Converts carbonate ores to oxides by heating in limited air.
  • Example: ZnCO3(s) → ZnO(s) + CO2(g)

3. Reduction of Metal Oxides:

• Metal oxides are reduced to metals using suitable reducing agents like carbon.
  • Example: ZnO(s) + C(s) → Zn(s) + CO(g)
• Reduction involves obtaining metals from their compounds.

4. Alternative Reduction Methods:

• Displacement Reactions: Highly reactive metals (e.g., sodium, calcium, aluminum) used as reducing agents to displace metals of lower reactivity.
  • Example: 3MnO2(s) + 4Al(s) → 3Mn(l) + 2Al2O3(s) + Heat
  • Identifying oxidation and reduction: Aluminum is oxidized, manganese dioxide is reduced.

5. Thermit Reaction:

• Highly exothermic displacement reaction producing molten metal.
• Used in applications like joining railway tracks or repairing machine parts.
  • Example: Fe2O3(s) + 2Al(s) → 2Fe(l) + Al2O3(s) + Heat

Extracting Metals Towards the Top of the Activity Series:

1. High Reactivity Metals:

• Examples: Sodium, magnesium, calcium, aluminum.
• These metals are very reactive and cannot be reduced using carbon.

2. Limitation of Carbon Reduction:

• Carbon cannot reduce the oxides of these metals due to their higher affinity for oxygen compared to carbon.

3. Electrolytic Reduction:

• Metals are extracted by electrolytic reduction of their compounds.
• Example: Sodium, magnesium, and calcium are obtained by electrolysis of their molten chlorides.

4. Electrolysis Process:

• At the Cathode (Negatively Charged Electrode):
  • Metal ions gain electrons and are deposited as metals.
  • Example reaction: (Na+) + (e-) → Na
• At the Anode (Positively Charged Electrode):
  • Non-metal ions lose electrons and are liberated.
  • Example reaction: 2(Cl-) → Cl2 + 2(e^-)

5. Example: Aluminium Extraction:

• Aluminium is extracted through the electrolytic reduction of aluminum oxide (Al₂O₃).

Refining of Metals

1. Impurities in Metals:

• Metals obtained from reduction processes contain impurities that need to be removed for purity.

2. Electrolytic Refining:

• The most common method for refining impure metals.
• Used for metals like copper, zinc, tin, nickel, silver, and gold.

3. Process of Electrolytic Refining:

• Anode: Impure metal is used as the anode.
• Cathode: A thin strip of pure metal is used as the cathode.
• Electrolyte: A solution of the metal salt is used as the electrolyte.
• When electric current passes through the electrolyte:
  • Pure metal dissolves from the anode into the electrolyte.
  • An equivalent amount of pure metal is deposited on the cathode.
  • Soluble impurities remain in the solution, while insoluble impurities settle as anode mud.

4. Example: Electrolytic Refining of Copper:

• Electrolyte: Acidified copper sulfate solution.
• Anode: Impure copper.
• Cathode: Pure copper strip.
• Pure copper is deposited on the cathode as the current is passed.

Corrosion

• Silver: Becomes black over time when exposed to air due to the formation of silver sulfide from sulfur in the air.
• Copper: Reacts with moist carbon dioxide in the air, losing its shiny brown surface and forming a green coat of copper carbonate.
• Iron: Develops a brown flaky substance called rust when exposed to moist air for a prolonged period.

Experiment to Determine Rusting Conditions:

1. Setup:

• Test Tube A: Contains iron nails with water and air.
• Test Tube B: Contains iron nails in boiled distilled water, covered with oil to prevent air dissolution.
• Test Tube C: Contains iron nails with anhydrous calcium chloride to absorb moisture, creating dry air.

2. Observations:

• Test Tube A: Nails rust, exposed to both air and water.
• Test Tube B: Nails do not rust, exposed only to water without air.
• Test Tube C: Nails do not rust, exposed to dry air without moisture.

3. Conclusion:

• Iron rusts in the presence of both air (oxygen) and water (moisture).
• Rusting does not occur if either air or moisture is absent.

Prevention of Corrosion

1. Methods to Prevent Rusting:

• Painting, oiling, greasing, galvanizing, chrome plating, anodizing, and alloying.

2. Galvanization:

• Protects steel and iron by coating with a thin layer of zinc.
• Even if the zinc coating is broken, the galvanized article remains protected from rusting.

Alloying

1. Benefits of Alloying:

• Improves the properties of metals.
• Allows customization of metal properties.

2. Examples:

• Pure iron is soft and stretches easily when hot. Alloying with carbon (about 0.05%) makes it hard and strong.
• Stainless steel, an alloy of iron with nickel and chromium, is hard and does not rust.
• Alloys can be made from metals and non-metals.

3. Definition:

• An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal.
• Prepared by melting the primary metal and dissolving other elements in definite proportions, then cooling to room temperature.

4. Special Types:

• Amalgams: Alloys containing mercury.

5. Properties of Alloys:

• Lower electrical conductivity and melting points compared to pure metals.
• Examples:
  • Brass (copper and zinc) and bronze (copper and tin) are not good conductors of electricity.
  • Solder (lead and tin) has a low melting point, used for welding electrical wires.

6. Gold Alloying:

• Pure gold (24 carat) is very soft and unsuitable for jewelry.
• 22 carat gold, used in India for ornaments, is alloyed with silver or copper to increase hardness.

Ancient Indian Metallurgy

Iron Pillar of Delhi:

• Built over 1600 years ago by Indian iron workers.
• Known for its rust resistance, attracting scientific examination worldwide.
• The pillar is 8 meters high and weighs 6 tonnes (6000 kg).