Metals and Non-metals

Physical Properties

Introduction

  • In Class IX, you learned about metals and non-metals.
  • Elements are classified based on their properties.

Uses in Daily Life

  • Think about how metals and non-metals are used around you.
  • Consider properties when categorizing elements as metals or non-metals.
  • Properties influence their uses.

Physical Properties

Metals

Activity 3.1

  • Collect samples: iron, copper, aluminium, magnesium.
  • Note their appearance.
  • Clean with sandpaper and observe again.
  • Metallic Lustre: Metals shine in their pure state.

Activity 3.2

  • Try cutting small pieces of iron, copper, aluminium, magnesium.
  • Handle sodium metal with care and try cutting it.
  • Hardness: Metals are generally hard, but hardness varies.

Activity 3.3

  • Take pieces of iron, zinc, lead, copper.
  • Strike each with a hammer on an iron block.
  • Malleability: Metals can be beaten into thin sheets. Gold and silver are very malleable.

Activity 3.4

  • List metals whose wires you’ve seen.
  • Ductility: Metals can be drawn into thin wires. Gold is the most ductile.

Activity 3.5

  • Clamp an aluminium or copper wire.
  • Fix a pin with wax.
  • Heat the wire and observe.
  • Conductivity: Metals are good conductors of heat and have high melting points. Best conductors: silver and copper. Poor conductors: lead and mercury.

Activity 3.6

  • Set up an electric circuit.
  • Place metal between terminals A and B.
  • Electrical Conductivity: Metals conduct electricity, making the bulb glow.

Additional Observations

  • Electric wires are coated with PVC or rubber for safety.
  • Sonorous: Metals make a sound when struck. School bells are made of metals for this reason.

Non-metals

In the previous class, you learned that there are fewer non-metals than metals. Examples of non-metals include:

  • Carbon
  • Sulphur
  • Iodine
  • Oxygen
  • Hydrogen

Non-metals can be solids or gases. Bromine is an exception as it is a liquid.

Physical Properties: Metals vs. Non-metals

Let’s see if non-metals have similar physical properties to metals.

Activity 3.7: Observing Non-metals

  • Collect samples of carbon (coal or graphite), sulphur, and iodine.
  • Perform the same tests as in Activities 3.1 to 3.4 and 3.6.
  • Record your observations.

Discussion Points:

  • Metals vs. Non-metals:
    • Metals: Mostly solid at room temperature (except mercury).
    • Non-metals: Can be solid, liquid, or gas.
  • Melting Points:
    • Metals usually have high melting points (except gallium and caesium).
    • Non-metals vary; for example, iodine is lustrous but carbon can exist as diamond (very hard) or graphite (conducts electricity).

Examples of Exceptions:

  1. Mercury: A liquid metal at room temperature.
  2. Gallium and Caesium: Metals with low melting points that can melt in your hand.
  3. Iodine: A lustrous non-metal.
  4. Carbon: Exists as different allotropes like diamond (very hard) and graphite (conducts electricity).
  5. Alkali Metals: Soft, low-density metals that can be cut with a knife.

Chemical Properties: Metals vs. Non-metals

Activity 3.8: Testing Chemical Properties

Activity 3.8: Testing Chemical Properties

  • Materials Needed: Magnesium ribbon, sulphur powder.
  • Steps:
    1. Burn the magnesium ribbon, dissolve the ashes in water, and test with litmus paper.
    2. Burn sulphur powder, collect the fumes in a test tube, add water, and test with litmus paper.

Observations:

  • Magnesium: Forms a basic solution.
  • Sulphur: Forms an acidic solution.

Conclusion:

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

In the next section, you will learn more about metal oxides.

Chemical Properties of Metals

We’ll explore the chemical properties of metals in Sections 3.2.1 to 3.2.4. For this, we’ll use the following metals:

  • Aluminium
  • Copper
  • Iron
  • Lead
  • Magnesium
  • Zinc
  • Sodium

What Happens When Metals are Burnt in Air?

Magnesium burns in air with a dazzling white flame, but do all metals react the same way? Let’s find out.

Activity 3.9: Burning Metals

Activity 3.9: Burning Metals

  • Caution: Perform with teacher’s assistance and wear eye protection.
  • Steps:
    • Hold a metal sample with tongs and burn it over a flame.
    • Collect the product formed.
    • Observe and record:
      • Which metals burn easily?
      • Flame colors.
      • Appearance of the metal surface after burning.
    • Arrange metals in order of reactivity towards oxygen.
    • Check if the products are soluble in water.

Observations:

  • Metals + Oxygen → Metal Oxides
    • Example: Copper + Oxygen → Copper(II) oxide (black oxide)
      • 2Cu + O₂ → 2CuO
    • Example: Aluminium + Oxygen → Aluminium oxide
      • 4Al + 3O₂ → 2Al₂O₃
  • Metal Oxides:
    • Usually basic.
    • Some, like aluminium oxide and zinc oxide, are amphoteric (react with both acids and bases).
      • Aluminium oxide reactions:
        • Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
        • Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O
  • Solubility:
    • Most metal oxides are insoluble in water.
    • Some, like sodium oxide and potassium oxide, dissolve to form alkalis:
      • Na₂O + H₂O → 2NaOH
      • K₂O + H₂O → 2KOH

Reactivity:

  • Highly Reactive: Sodium and potassium (kept in kerosene to prevent fires).
  • Moderately Reactive: Magnesium, aluminium, zinc, lead (form protective oxide layers).
  • Less Reactive: Iron (filings burn, solid iron does not).
  • Least Reactive: Copper (forms a black layer of copper(II) oxide), silver, and gold (no reaction with oxygen even at high temperatures).

Additional Information:

  • Anodising: A process to thicken the oxide layer on aluminium, making it more resistant to corrosion. This is done by electrolysing aluminium in dilute sulphuric acid, resulting in a thicker oxide layer that can be dyed for an attractive finish.

What Happens When Metals React with Water?

Activity 3.10: Reacting Metals with Water

  • Caution: Perform with teacher’s assistance.
  • Steps:
    • Collect metal samples (aluminium, copper, iron, lead, magnesium, zinc, sodium).
    • Put small pieces of metals in beakers with cold water.
    • Observe:
      • Which metals reacted with cold water?
      • Did any metal produce fire?
      • Does any metal float after some time?
    • Test metals that didn’t react with cold water in hot water.
    • For those that didn’t react with hot water, observe their reaction with steam.
    • Arrange metals by reactivity with water.

Reactions:

  • Metals react with water to form metal oxides and hydrogen gas.
    • Formula: Metal + Water → Metal oxide + Hydrogen
    • Example: Sodium
      • 2Na + 2H₂O → 2NaOH + H₂ + heat
    • Example: Potassium
      • 2K + 2H₂O → 2KOH + H₂ + heat
  • Reactivity Observations:
    • Sodium and Potassium: React violently with cold water, producing fire.
    • Calcium: Reacts with cold water less violently, starts floating due to hydrogen bubbles.
    • Magnesium: Reacts with hot water, not cold water, starts floating due to hydrogen bubbles.
    • Aluminium, Iron, Zinc: Do not react with cold or hot water but react with steam to form oxides and hydrogen.
      • Examples:
        • 2Al + 3H₂O → Al₂O₃ + 3H₂
        • 3Fe + 4H₂O → Fe₃O₄ + 4H₂
    • Lead, Copper, Silver, Gold: Do not react with water at all.

Reactivity Order: Sodium > Potassium > Calcium > Magnesium > Aluminium > Iron > Zinc > Lead > Copper > Silver > Gold

What Happens When Metals React with Acids?

Activity 3.11: Reacting Metals with Acids

  • Caution: Do not use sodium and potassium due to vigorous reactions.
  • Steps:
    • Collect metal samples (except sodium and potassium).
    • Clean tarnished samples with sandpaper.
    • Place samples in test tubes with dilute hydrochloric acid.
    • Suspend thermometers in test tubes, observe bubble formation.
    • Observe which metals react vigorously and which produced the highest temperature.
    • Arrange metals by reactivity with dilute acids.

Reactions:

  • Metals react with acids to form salt and hydrogen gas.
    • Formula: Metal + Dilute acid → Salt + Hydrogen
    • Example Reactions:
      • Magnesium: Mg + 2HCl → MgCl₂ + H₂
      • Aluminium: 2Al + 6HCl → 2AlCl₃ + 3H₂
      • Zinc: Zn + 2HCl → ZnCl₂ + H₂
      • Iron: Fe + 2HCl → FeCl₂ + H₂
  • Observations:
    • Magnesium: Fastest bubble formation, most exothermic reaction.
    • Reactivity Order: Mg > Al > Zn > Fe
    • Copper: No reaction with dilute HCl, no bubbles, no temperature change.

Note:

  • Hydrogen gas is not evolved when metals react with nitric acid (HNO₃) because it oxidizes hydrogen to water.
  • Exception: Magnesium and manganese react with very dilute HNO₃ to produce hydrogen.

Interesting Fact:

  • Aqua regia: A mixture of concentrated hydrochloric acid and nitric acid (3:1 ratio) that can dissolve gold and platinum.
Summary of Key Points
  • Metals react differently with water and acids.
  • Sodium and potassium react violently with cold water.
  • Magnesium reacts with hot water.
  • Aluminium, iron, and zinc react with steam.
  • Reactivity with acids varies, with magnesium being the most reactive among common metals.
  • Aqua regia can dissolve gold and platinum, highlighting unique chemical properties.

How Do Metals React with Solutions of Other Metal Salts?

Activity 3.12: Displacement Reactions

  • Steps:
    • Take a clean copper wire and an iron nail.
    • Place the copper wire in a solution of iron sulphate.
    • Place the iron nail in a solution of copper sulphate.
    • Observe after 20 minutes.
    • Record which test tube shows a reaction.
    • Determine how you know a reaction occurred.
    • Correlate your observations with previous activities.

Observation:

  • Reaction Evidence: If a metal displaces another metal from its salt solution, a reaction has taken place.
  • Example Reaction: Iron nail in copper sulphate solution
    • Copper gets deposited on the nail.
    • Chemical Equation: Fe + CuSO₄ → FeSO₄ + Cu

Conclusion:

  • Reactivity: More reactive metals displace less reactive metals from their compounds.
  • Reactivity Comparison: Iron is more reactive than copper.

The Reactivity Series

  • Definition: A list of metals arranged in order of decreasing reactivity.
  • Importance: Helps predict how metals will react with other substances.

How Do Metals and Non-Metals React?

  • Why Metals React: Metals react to achieve a stable electron configuration, similar to noble gases.
  • Example: Sodium (Na) and Chlorine (Cl)
    • Sodium loses one electron to achieve a stable octet, forming Na⁺.
    • Chlorine gains one electron to achieve a stable octet, forming Cl⁻.
    • Reaction: Na + Cl → NaCl
    • Result: Na⁺ and Cl⁻ ions are held together by strong electrostatic forces.

Ionic Compounds:

  • Formation: By transfer of electrons from metals to non-metals.
  • Examples:
    • Sodium Chloride (NaCl)
    • Magnesium Chloride (MgCl₂)
  • Ions:
    • In MgCl₂: Mg²⁺ (cation) and Cl⁻ (anion).

Key Points:

  • Reactive Metals: Can displace less reactive metals from solutions.
  • Displacement Reaction: Used to determine metal reactivity.
  • Ionic Compounds: Formed by transfer of electrons between metals and non-metals.

Properties of Ionic Compounds

Activity 3.13: Observing Ionic Compounds

Steps:

  • Take samples like sodium chloride, potassium iodide, and barium chloride.
  • Observe their physical state.
  • Heat samples on a metal spatula and note any color change or melting.
  • Try dissolving samples in water, petrol, and kerosene.
  • Test their conductivity in water using a simple circuit.

Observations:

  • Physical Nature:
    • Ionic compounds are solid and hard.
    • They are brittle and break easily under pressure.
  • Melting and Boiling Points:
    • High melting and boiling points due to strong inter-ionic attractions.
  • Solubility:
    • Generally soluble in water.
    • Insoluble in solvents like kerosene and petrol.
  • Conduction of Electricity:
    • Conduct electricity in water due to free-moving ions.
    • Do not conduct electricity in solid form.
    • Conduct electricity when molten because the ions can move freely.

Occurrence of Metals

  • Source: Metals are found in the earth’s crust and seawater.
  • Minerals: Naturally occurring elements or compounds in the earth’s crust.
  • Ores: Minerals with a high percentage of metal that can be profitably extracted.

Extraction of Metals

  • Reactivity Series: Metals are extracted based on their reactivity.
    • Low Reactivity Metals: Found in a free state (e.g., gold, silver, platinum, copper).
    • Medium Reactivity Metals: Found as oxides, sulphides, or carbonates (e.g., zinc, iron, lead).
    • High Reactivity Metals: Never found free in nature, always in compounds (e.g., potassium, sodium, calcium, magnesium, aluminium).
  • Extraction Process:
    • Steps:
      • Mining the ore.
      • Concentrating the ore.
      • Reducing the ore to get the metal.
      • Refining the metal.
    • Different methods are used depending on the metal’s reactivity.

Summary:

  • Properties of Ionic Compounds:
    • Solid, hard, brittle.
    • High melting and boiling points.
    • Soluble in water, not in kerosene or petrol.
    • Conduct electricity in molten state or water solution, not as solids.
  • Occurrence and Extraction of Metals:
    • Metals are found as minerals and ores in the earth’s crust.
    • Extracted based on their reactivity using various methods.

Enrichment of Ores

  • Gangue: Ores mined from the earth contain impurities like soil and sand.
  • Removal: Gangue is removed by using differences in physical or chemical properties.

Extracting Metals Low in the Activity Series

  • Unreactive Metals: These metals can be reduced by heating alone.
  • Example:
    • Mercury: Cinnabar (HgS) is heated in air to form HgO, which is then reduced to mercury.
    • Copper: Copper (Cu) can be extracted by heating Cu₂S in air.

Extracting Metals in the Middle of the Activity Series

  • Moderately Reactive Metals: Usually found as sulphides or carbonates.
  • Conversion to Oxides:
    • Roasting: Heating sulphides in excess air (e.g., ZnS to ZnO).
    • Calcination: Heating carbonates in limited air (e.g., ZnCO₃ to ZnO).
  • Reduction: Oxides are reduced using carbon (e.g., ZnO + C → Zn + CO).
  • Displacement Reactions: Highly reactive metals like aluminium can reduce other metals from their compounds.
    • Example: Thermit reaction (Fe₂O₃ + Al → Fe + Al₂O₃).

Extracting Metals towards the Top of the Activity Series

  • Highly Reactive Metals: Cannot be reduced by carbon.
  • Electrolytic Reduction: Metals are obtained by electrolysis of their molten chlorides.
    • Example: Sodium, magnesium, and calcium are obtained by electrolysis.
    • At Cathode: Na⁺ + e⁻ → Na
    • At Anode: 2Cl⁻ → Cl₂ + 2e⁻

Refining of Metals

  • Impurities: Metals from reduction processes contain impurities.
  • Electrolytic Refining: Used for metals like copper, zinc, and gold.
    • Process:
      • Anode: Impure metal.
      • Cathode: Pure metal strip.
      • Electrolyte: Solution of metal salt.
    • Outcome: Pure metal deposits on the cathode; impurities settle as anode mud.

Corrosion

What is Corrosion?

  • Silver: Turns black due to silver sulphide formation.
  • Copper: Turns green because of basic copper carbonate.
  • Iron: Rusts, forming a brown flaky substance.

Activity 3.14: Observing Rusting

  • Setup:
    • Test Tube A: Iron nails in water, corked.
    • Test Tube B: Iron nails in boiled distilled water with oil, corked.
    • Test Tube C: Iron nails with anhydrous calcium chloride, corked.
  • Observation:
    • Nails rust in Test Tube A (air and water present).
    • Nails do not rust in Test Tube B (only water, no air).
    • Nails do not rust in Test Tube C (only dry air).
  • Conclusion: Rusting requires both air and water.

Concept: Rusting requires both air and water.

Prevention of Corrosion

  • Methods:
    • Painting, Oiling, Greasing: Coats iron to prevent exposure to air and water.
    • Galvanizing: Coating iron with zinc; zinc protects even if scratched.
    • Chrome Plating, Anodising: Adding protective layers.
    • Alloying: Mixing iron with other elements to improve properties (e.g., stainless steel).

Alloys:

  • Definition: Homogeneous mixture of metals or metals with non-metals.
  • Preparation: Melting primary metal, adding other elements, and cooling.
  • Examples:
    • Brass: Copper and zinc.
    • Bronze: Copper and tin.
    • Solder: Lead and tin (used for welding due to low melting point).

Interesting Facts:

  • Pure Gold: Too soft, so it’s alloyed with silver or copper for jewelry (e.g., 22 carat gold).
  • Iron Pillar of Delhi: Built 1600 years ago, it resists rusting due to ancient Indian metallurgy techniques. It is 8 meters high and weighs 6 tonnes.

Key Takeaways

  • Rusting needs air and water.
  • Preventing rust can be done by coating or alloying metals.
  • Alloys improve the properties of metals, making them more useful.

Chapter Summary:

  • Elements can be classified as metals and non-metals.
  • Metals are lustrous, malleable, ductile, and good conductors of heat and electricity.
  • Metals are solids at room temperature, except mercury which is a liquid.
  • Metals form positive ions by losing electrons to non-metals.
  • Metals combine with oxygen to form basic oxides.
  • Aluminium oxide and zinc oxide are amphoteric oxides, showing both basic and acidic properties.
  • Different metals have different reactivities with water and dilute acids.
  • The activity series is a list of metals arranged in order of decreasing reactivity.
  • Metals above hydrogen in the activity series can displace hydrogen from dilute acids.
  • A more reactive metal displaces a less reactive metal from its salt solution.
  • Metals occur in nature as free elements or as compounds.
  • Metallurgy is the extraction of metals from their ores and refining them for use.
  • An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal.
  • Corrosion is when the surface of metals like iron gets corroded when exposed to moist air for a long time.
  • Non-metals have properties opposite to metals; they are neither malleable nor ductile and are bad conductors of heat and electricity (except graphite).
  • Non-metals form negatively charged ions by gaining electrons when reacting with metals.
  • Non-metals form oxides that are either acidic or neutral.
  • Non-metals do not displace hydrogen from dilute acids but react with hydrogen to form hydrides.
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