Metals and Non-Metals

Metals and Non-Metals: Complete Guide for CBSE Class 8 Science

Introduction to Metals and Non-Metals

This chapter covers the physical and chemical properties of metals and non-metals, including reactions and uses. Refer to NCERT Solutions for Class 8 Science for detailed answers and Class 8 Notes for short summaries. Join Home Tuition for Class 8 for practical understanding through experiments and daily examples like rusting and conductivity. This lesson helps in understanding elements used in daily life.

Understanding the classification of elements into metals, non-metals, and metalloids is fundamental to chemistry. With over 114 known chemical elements, scientists have organized them based on their physical and chemical properties to make studying them more manageable.

Elements are pure substances that cannot be broken down into simpler substances by any chemical means. Examples include silver, iron, oxygen, and nitrogen. These elements are classified into three main categories:

  • Metals (e.g., Sodium, Copper, Gold)
  • Non-metals (e.g., Carbon, Nitrogen, Oxygen)
  • Metalloids (e.g., Germanium, Arsenic, Silicon)

This classification helps us understand how different elements behave, react, and can be used in our daily lives.

What Are Elements?

An element is a pure substance that cannot be subdivided into two or more new substances by any means. Chemists have identified 114 elements that exist in nature, while some have been created synthetically in laboratories.

Classification of Elements

Elements exhibit wide variations in properties, which led to their classification into:

  1. Metals - Located on the left side of the periodic table
  2. Non-metals - Located on the right side of the periodic table
  3. Metalloids - Located along the "staircase" line separating metals and non-metals

Understanding Metals

What Are Metals?

Metals are elements that typically have 1, 2, or 3 electrons in their outermost shell. They tend to lose these electrons easily to form positively charged ions (cations), making them electropositive elements.

Physical Properties of Metals

1. Physical State

  • Most metals are solid at room temperature
  • Exceptions: Mercury and gallium are liquid; mercury is the only common liquid metal at room temperature

2. Malleability

  • Metals can be beaten into thin sheets without breaking
  • Gold and silver are among the most malleable metals
  • Aluminum and copper are also highly malleable
  • Exceptions: Sodium, potassium, and calcium are not malleable

3. Ductility

  • Metals can be drawn into thin wires
  • Gold and silver are the most ductile
  • Copper, aluminum, and tungsten are very ductile and used in electrical wiring
  • Exceptions: Sodium, potassium, calcium, tin, and lead have low ductility

4. Conductivity

  • Metals are excellent conductors of heat and electricity
  • Silver is the best conductor, followed by copper and aluminum
  • This property is due to the presence of free electrons
  • Lead and mercury are relatively poor conductors among metals

5. Lustre

  • Metals have a characteristic shine called metallic lustre
  • They can be polished to enhance their shine
  • Examples: Gold, silver, and copper have excellent lustre

6. Density

  • Most metals have high densities
  • Mercury has a very high density (13.6 g/cm³)
  • Exceptions: Sodium, potassium, magnesium, and aluminum have low densities
  • Osmium has the maximum density (22 g/cm³) among all elements

7. Hardness

  • Most metals are hard and cannot be cut with a knife
  • Exceptions: Sodium and potassium are soft and can be cut with a knife

8. Melting and Boiling Points

  • Most metals have high melting and boiling points
  • Exceptions: Sodium, potassium, and gallium have low melting points

9. Sonority

  • Metals produce sound when struck (they are sonorous)
  • This property is utilized in making bells and musical instruments

10. Color

  • Most metals are silvery-white or grey
  • Exceptions: Gold is yellow, copper is reddish-brown

Chemical Properties of Metals

1. Reaction with Oxygen (Air)

Metals react with oxygen to form metal oxides, which are generally basic in nature:

Metal + Oxygen → Metal Oxide

Examples:

  • 4Na + O₂ → 2Na₂O (Sodium oxide)
  • 2Mg + O₂ → 2MgO (Magnesium oxide)
  • 4Al + 3O₂ → 2Al₂O₃ (Aluminum oxide)

When dissolved in water, these oxides form alkaline solutions:

  • Na₂O + H₂O → 2NaOH (Sodium hydroxide)

This alkaline solution turns red litmus paper blue.

2. Reaction with Water

The reactivity of metals with water varies:

Highly Reactive Metals (Sodium, Potassium):

  • React vigorously with cold water
  • 2Na + 2H₂O → 2NaOH + H₂↑
  • 2K + 2H₂O → 2KOH + H₂↑

Moderately Reactive (Calcium):

  • Ca + 2H₂O → Ca(OH)₂ + H₂↑

Less Reactive (Magnesium):

  • Reacts mildly with cold water but vigorously with boiling water
  • Mg + 2H₂O → Mg(OH)₂ + H₂↑

Even Less Reactive (Zinc, Iron):

  • React with steam to form oxides
  • Zn + H₂O → ZnO + H₂↑
  • 3Fe + 4H₂O → Fe₃O₄ + 4H₂↑

Least Reactive (Copper, Silver, Gold):

  • Do not react with water or steam

3. Reaction with Dilute Acids

Metals more reactive than hydrogen displace hydrogen from dilute acids:

Metal + Dilute Acid → Salt + Hydrogen gas

Examples:

  • Zn + 2HCl → ZnCl₂ + H₂↑
  • Zn + H₂SO₄ → ZnSO₄ + H₂↑
  • Mg + 2HCl → MgCl₂ + H₂↑
  • Fe + 2HCl → FeCl₂ + H₂↑

Note: Copper, silver, and gold do not react with dilute acids as they are less reactive than hydrogen.

4. Reaction with Salt Solutions

A more reactive metal can displace a less reactive metal from its salt solution:

Example:

  • Zn + CuSO₄ → ZnSO₄ + Cu

Zinc displaces copper from copper sulfate solution because zinc is more reactive than copper.

Examples of Common Metals and Their Uses

1. Copper

  • Properties: Reddish-brown, highly ductile, excellent conductor
  • Uses:
    • Electric wires and cables
    • Utensils and ornaments
    • Electronic devices
    • Coins and statues (as brass and bronze alloys)

2. Iron

  • Properties: Strong, magnetic, abundant
  • Uses:
    • Construction (buildings, bridges)
    • Automobiles and machinery
    • Railings and pipes
    • Utensils (as steel)

3. Aluminum

  • Properties: Silvery-white, light, malleable, corrosion-resistant
  • Uses:
    • Aircraft bodies
    • Utensils and drink cans
    • Electric wires
    • Aluminum foil for packaging
    • Window frames

4. Zinc

  • Properties: Bluish-white, brittle
  • Uses:
    • Galvanizing iron (preventing rust)
    • Dry cells and batteries
    • Brass and bronze alloys
    • Extracting silver and gold

5. Gold

  • Properties: Yellow, extremely malleable, corrosion-resistant
  • Uses:
    • Jewelry and coins
    • Electronic devices
    • Dental fillings (amalgam)

6. Silver

  • Properties: White, lustrous, best electrical conductor
  • Uses:
    • Jewelry
    • Photography
    • Water purification
    • Electronics

Understanding Non-Metals

What Are Non-Metals?

Non-metals are elements that typically have 4, 5, 6, or 7 electrons in their valence shell. They tend to gain electrons to form negatively charged ions (anions), making them electronegative elements.

Physical Properties of Non-Metals

1. Physical State

  • Exist as solids, liquids, or gases at room temperature
  • Solids: Carbon, sulfur, phosphorus, iodine
  • Liquids: Bromine (the only liquid non-metal at room temperature)
  • Gases: Oxygen, hydrogen, nitrogen, chlorine

2. Brittleness

  • Solid non-metals are brittle and break easily
  • Exception: Diamond (carbon) is the hardest naturally occurring substance

3. Non-Ductility

  • Non-metals cannot be drawn into wires
  • Exception: Carbon fibers are ductile

4. Conductivity

  • Poor conductors of heat and electricity
  • Exception: Graphite (carbon) conducts electricity due to free electrons

5. Lustre

  • Generally dull and non-lustrous
  • Exceptions: Iodine and graphite have lustre

6. Malleability

  • Non-metals are not malleable

7. Density

  • Low density compared to metals
  • Exceptions: Iodine and diamond have relatively high densities

8. Melting and Boiling Points

  • Generally low
  • Exceptions: Carbon (graphite: 3730°C), boron, and silicon have high melting points

9. Sonority

  • Non-sonorous (do not produce sound when struck)

10. Color

  • Varied colors: Chlorine (greenish-yellow), bromine (brown), iodine (violet)

Chemical Properties of Non-Metals

1. Reaction with Oxygen

Non-metals react with oxygen to form acidic or neutral oxides:

Non-metal + Oxygen → Non-metallic Oxide

Examples:

  • C + O₂ → CO₂ (Carbon dioxide - acidic)
  • S + O₂ → SO₂ (Sulfur dioxide - acidic)
  • 4P + 5O₂ → 2P₂O₅ (Phosphorus pentoxide - acidic)
  • 2H₂ + O₂ → 2H₂O (Water - neutral)

Acidic oxides react with bases to form salts:

  • CO₂ + 2NaOH → Na₂CO₃ + H₂O

2. Reaction with Water

Examples:

  • CO₂ + H₂O → H₂CO₃ (Carbonic acid)
  • SO₂ + H₂O → H₂SO₃ (Sulfurous acid)

3. Reaction with Acids

Most non-metals do not react with dilute acids. However, they react with concentrated acids:

  • C + 2H₂SO₄ (conc.) → CO₂ + 2SO₂ + 2H₂O
  • S + 2H₂SO₄ (conc.) → 3SO₂ + 2H₂O

4. Reaction with Chlorine

Non-metals form covalent compounds with chlorine:

  • H₂ + Cl₂ → 2HCl
  • P₄ + 10Cl₂ → 4PCl₅

Why Do Metals Conduct Electricity and Non-Metals Do Not?

Metals conduct electricity because:

  • They have free electrons in their structure
  • These electrons are loosely bound and can move freely throughout the metallic structure
  • When voltage is applied, these free electrons flow, creating an electric current
  • This is why metals like silver, copper, and aluminum are excellent conductors

Non-metals do not conduct electricity because:

  • Their electrons are tightly bound to atoms
  • No free electrons are available to carry electric charge
  • Their covalent bonds hold electrons firmly in place
  • Exception: Graphite conducts electricity because it has delocalized electrons in its layered structure

Metalloids: The Bridge Between Metals and Non-Metals

Which Elements Are Metalloids and What Are Their Properties?

Metalloids (or semi-metals) are elements that exhibit properties of both metals and non-metals. They are located along the "staircase" line in the periodic table.

Common Metalloids:

  1. Boron (B)
  2. Silicon (Si)
  3. Germanium (Ge)
  4. Arsenic (As)
  5. Antimony (Sb)
  6. Tellurium (Te)

Properties of Metalloids

Element Metallic Properties Non-Metallic Properties
Germanium Semiconductor; conductivity increases with temperature GeO₂ is acidic
Arsenic Grey form conducts electricity Forms hydride (AsH₃), a weak base
Antimony Grey form conducts electricity Forms stibine (SbH₃), a weak base
Silicon Good conductor at high temperatures Forms acidic oxide (SiO₂)
Boron Good conductor at high temperatures Poor conductor at room temperature

Key Characteristics:

  • Appearance: May have metallic lustre (e.g., germanium)
  • Conductivity: Semiconductors - conduct electricity better than non-metals but not as well as metals
  • Chemical behavior: Can behave as metals or non-metals depending on conditions
  • Uses: Extensively used in electronics and semiconductor industry

Compare Physical Properties of Metals, Non-Metals, and Metalloids

Property Metals Non-Metals Metalloids
State at Room Temperature Mostly solid (except Hg, Ga) Solid, liquid (Br), or gas Solid
Malleability Malleable Brittle (if solid) Varies
Ductility Ductile Non-ductile Limited ductility
Conductivity Good conductors Poor conductors (except graphite) Semiconductors
Lustre Lustrous Dull (except I₂, graphite) May have lustre
Density High (except Li, Na, K) Low (except diamond) Moderate
Melting/Boiling Points High (except Na, K) Low (except C, Si, B) Moderate to high
Sonority Sonorous Non-sonorous Varies
Hardness Hard (except Na, K) Soft (except diamond) Varies
Electron Tendency Lose electrons (1-3 valence e⁻) Gain electrons (5-7 valence e⁻) 4 valence electrons

How Metals React with Oxygen, Water, and Acids Differently

Reactivity Series of Metals

The reactivity series arranges metals in order of decreasing reactivity:

Most Reactive:

  • Potassium (K)
  • Sodium (Na)
  • Calcium (Ca)
  • Magnesium (Mg)
  • Aluminum (Al)
  • Zinc (Zn)
  • Iron (Fe)
  • Lead (Pb)
  • Hydrogen (H) - reference point
  • Copper (Cu)
  • Mercury (Hg)
  • Silver (Ag)
  • Gold (Au)
  • Platinum (Pt) Least Reactive

Reaction Patterns Based on Reactivity

With Oxygen:

  • Highly reactive (K, Na): React vigorously, may catch fire
  • Moderately reactive (Mg, Al, Zn): React on heating
  • Less reactive (Fe, Cu): React slowly or require high temperature
  • Least reactive (Au, Pt): Do not react with oxygen

With Water:

  • K, Na: React violently with cold water
  • Ca: Reacts vigorously with cold water
  • Mg: Reacts with hot water/steam
  • Al, Zn, Fe: React with steam
  • Cu, Ag, Au: Do not react with water

With Acids:

  • Above H in series: React with dilute acids, release H₂ gas
  • Below H in series: Do not react with dilute acids
  • Reactivity decreases down the series

Practical Implications:

  1. Storage: Highly reactive metals (Na, K) stored in kerosene
  2. Extraction: More reactive metals harder to extract from ores
  3. Corrosion: More reactive metals corrode faster
  4. Displacement reactions: More reactive metals displace less reactive ones from solutions

Chemical Bonding in Metals and Non-Metals

Ionic Bonds (Electrovalent Bonds)

Formed by transfer of electrons from metals to non-metals:

Example: Sodium Chloride (NaCl)

  • Na (2,8,1) loses 1 electron → Na⁺ (2,8)
  • Cl (2,8,7) gains 1 electron → Cl⁻ (2,8,8)
  • Electrostatic attraction forms NaCl

Covalent Bonds

Formed by sharing of electrons between non-metal atoms:

Example: Water (H₂O)

  • Each H atom shares 1 electron with O
  • O atom shares 2 electrons (one with each H)
  • Both atoms achieve stable configurations

Types of Covalent Bonds:

  • Single bond: One shared pair (H₂, Cl₂)
  • Double bond: Two shared pairs (O₂)
  • Triple bond: Three shared pairs (N₂)

Occurrence of Metals and Non-Metals

Metals in Nature

Minerals

Naturally occurring substances with definite composition and crystalline structure containing metals.

Ores

Minerals from which metals can be extracted profitably.

Gangue

Unwanted impurities present in ores.

Types of Metal Ores:

Ore Type Examples
Oxides Bauxite (Al₂O₃·2H₂O), Hematite (Fe₂O₃)
Sulfides Copper Pyrites (CuFeS₂), Zinc Blende (ZnS)
Carbonates Limestone (CaCO₃), Calamine (ZnCO₃)
Halides Rock Salt (NaCl), Horn Silver (AgCl)

Non-Metals in Nature

Non-Metal Free State Combined State
Oxygen Air (21%) Water, oxides, carbonates
Nitrogen Air (78%) Proteins, nitre (KNO₃)
Carbon Diamond, graphite, coal CO₂, carbonates, petroleum
Sulfur Volcanic regions Metal sulfides, H₂S
Hydrogen Very rare Water, hydrocarbons

Metallurgy: Extraction of Metals

Metallurgy is the science of extracting metals from their ores and refining them for use.

Steps in Metallurgy:

1. Crushing and Grinding

  • Breaking large ore pieces into smaller fragments
  • Grinding into fine powder using ball mills

2. Concentration (Ore Enrichment)

Methods:

a) Hydraulic Washing

  • Based on density differences
  • Ore particles heavier than gangue
  • Used for oxide ores (Sn, Pb)

b) Froth Flotation

  • For sulfide ores
  • Uses oil (pine oil) and water
  • Ore particles attach to froth, gangue sinks

c) Magnetic Separation

  • For magnetic ores (Fe₃O₄)
  • Uses electromagnetic separation

d) Chemical Separation

  • Example: Bayer's process for purifying bauxite
  • Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

3. Conversion to Metal Oxide

a) Calcination (for carbonates):

  • ZnCO₃ → ZnO + CO₂↑

b) Roasting (for sulfides):

  • 2ZnS + 3O₂ → 2ZnO + 2SO₂↑

4. Reduction to Metal

a) Reduction with Carbon:

  • ZnO + C → Zn + CO↑
  • Fe₂O₃ + 3C → 2Fe + 3CO↑

b) Reduction by Heating:

  • 2HgO → 2Hg + O₂↑

c) Electrolytic Reduction (for highly reactive metals):

  • Used for Na, K, Ca, Al

5. Refining (Purification)

  • Electrolytic refining
  • Liquation
  • Distillation

Corrosion of Metals

Corrosion is the deterioration of metals due to reaction with air and moisture.

Rusting of Iron

Process:

  • 4Fe + 3O₂ → 2Fe₂O₃
  • Fe₂O₃ + xH₂O → Fe₂O₃·xH₂O (rust)

Conditions Required:

  • Presence of oxygen
  • Presence of moisture

Prevention of Corrosion:

  1. Painting or Oiling: Creates protective barrier
  2. Galvanizing: Coating with zinc
  3. Electroplating: Depositing protective metal layer (tin, chromium)
  4. Anodizing: For aluminum - forms uniform oxide layer
  5. Alloying: Making stainless steel (Fe + Cr + Ni)

Alloys

An alloy is a homogeneous mixture of two or more metals (or a metal and non-metal) melted together.

Objectives of Alloy Making:

  • Increase hardness
  • Increase tensile strength
  • Improve corrosion resistance
  • Lower melting point
  • Modify chemical reactivity

Important Alloys:

Alloy Composition Properties Uses
Steel Fe (98%) + C (1%) Strong, hard Construction, machinery
Stainless Steel Fe (83%) + Cr (15%) + Ni (1%) + C (1%) Corrosion-resistant Utensils, surgical instruments
Brass Cu (60-80%) + Zn (20-40%) Malleable, lustrous Electric switches, utensils
Bronze Cu (80%) + Sn (10%) + Zn (10%) Hard, corrosion-resistant Statues, coins
Duralumin Al (95%) + Cu (4%) + Mn (0.5%) + Mg (0.5%) Light, strong as steel Aircraft, rockets
Solder Pb + Sn Low melting point Joining metals

Key Chemical Formulas

Formula Name Chemical Equation Explanation
Sodium oxide formation 4Na + O₂ → 2Na₂O Sodium reacts with oxygen
Sodium hydroxide formation Na₂O + H₂O → 2NaOH Basic oxide dissolves in water
Magnesium combustion 2Mg + O₂ → 2MgO Magnesium burns in air
Sodium-water reaction 2Na + 2H₂O → 2NaOH + H₂↑ Vigorous reaction with hydrogen gas evolution
Zinc-acid reaction Zn + 2HCl → ZnCl₂ + H₂↑ Displacement of hydrogen from acid
Zinc-copper displacement Zn + CuSO₄ → ZnSO₄ + Cu More reactive metal displaces less reactive
Calcination of zinc carbonate ZnCO₃ → ZnO + CO₂↑ Thermal decomposition
Roasting of zinc sulfide 2ZnS + 3O₂ → 2ZnO + 2SO₂↑ Oxidation of sulfide ore
Reduction of zinc oxide ZnO + C → Zn + CO↑ Carbon reduction
Rusting of iron 4Fe + 3O₂ → 2Fe₂O₃ then Fe₂O₃ + xH₂O → Fe₂O₃·xH₂O Hydrated iron oxide formation
Carbon dioxide in water CO₂ + H₂O → H₂CO₃ Acidic oxide forms acid
Hydrogen combustion 2H₂ + O₂ → 2H₂O Neutral oxide formation

Conclusion

Understanding metals and non-metals is fundamental to chemistry and has practical applications in everyday life. From the electrical wires in our homes (copper) to the vehicles we drive (steel), from the packaging of our food (aluminum) to the electronics we use (silicon), these elements and their compounds are integral to modern civilization.

The classification of elements based on their properties helps us predict their behavior, understand their reactions, and utilize them effectively. Whether it's the highly reactive sodium stored in kerosene, the noble gold used in jewelry, or the semiconductor silicon powering our computers, each element has unique characteristics that make it suitable for specific applications.

As you continue your study of chemistry, remember that these foundational concepts about metals, non-metals, and metalloids will help you understand more complex chemical phenomena. The reactivity series, types of chemical bonding, and metallurgical processes are not just academic concepts but principles that have shaped human technological