Matter, Atoms & Elements (Metals & Non-Metals)
matter, which is defined as anything that has mass and occupies space. Everything around us – air, water, food, stones, stars, plants – is considered matter. Modern science, however, classifies matter based on physical and chemical properties. Key Physical Properties:
- Matter is made up of particles, these particles are extremely small and there is space between particles.
- Particles are continuously moving (kinetic energy) and attract each other.
States of Matter – Matter exists in three physical states (solid ↔ liquid ↔ gas)
- Matter changes state on changing temperature or pressure.
- Gases spread aroma over large distances due to rapid particle motion.
Melting and Latent Heat
- The temperature at which a solid turns into a liquid is its melting point (e.g., ice melts at 273.15 K).
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Even when melting, temperature remains constant. The heat energy is used to overcome the attraction between particles, called latent heat of fusion.
Boiling and Latent Heat of Vaporization
- On further heating, liquids reach their boiling point (e.g., water at 373 K or 100°C), where particles gain enough energy to overcome attraction and become vapour.
- This extra energy is known as the latent heat of vaporization.
Sublimation and Deposition
- Some substances like camphor or ammonium chloride change directly from solid to gas, this process called sublimation.
- Deposition is the reverse where gas directly converting to solid.
Effect of Pressure
- Increasing pressure and reducing temperature can convert gases into liquids.
- Example: Solid CO₂ (dry ice) sublimates directly into gas at 1 atmosphere pressure.
Factors affecting evaporevaporation:
- Humidity: More humidity → slower evaporation
- Wind speed: Higher wind → faster evaporatio
| Quantity | Unit | Symbol |
|---|---|---|
| Temperature | Kelvin | K |
| Length | Metre | m |
| Mass | Kilogram | kg |
| Weight | Newton | N |
| Volume | Cubic metre | m³ |
| Density | kg/m³ | kg m⁻³ |
| Pressure | Pascal | Pa |
Pure, Mixture and Componats
For scientists, a pure substance has only one kind of particle and a uniform composition. A mixture is made up of two or more substances that are not chemically combined. Examples air, seawater soil etc.
Mixtures: Physical combination of substances; properties of constituents are retained.
- Homogeneous Mixtures – Uniform composition (e.g., salt water, sugar solution)
- Heterogeneous Mixtures – Non-uniform composition (e.g., oil and water, salt and sulphur)
Compounds: Chemical combination of two or more elements; fixed composition and new properties.
Solutions involve a solute, solvent, and are classified as saturated, unsaturated, or supersaturated.
Solution is a homogeneous mixture involve a solute, solvent, and are classified as saturated, unsaturated, or supersaturated.
- Solvent: Substance present in a larger amount.
- Solute: Substance present in a lesser amount.
Properties of Solutions: Particles are very small. They don’t scatter light. They are stable and can’t be separated by filtration. Types of Solutions:
- Solid in liquid (e.g., sugar in water)
- Gas in liquid (e.g., CO₂ in soda)
- Gas in gas (e.g., air)
Concentration of a Solution: Dilute solution: Less solute Concentrated solution: More solute Saturated solution: No more solute can dissolve Unsaturated solution: More solute can still dissolve
Mass by mass percentage = (Mass of solute/Mass of solution) ×100
Suspensions:
- Non-homogeneous mixtures where particles remain suspended and visible.
- Particles can be separated by filtration.
- Scatter light and settle down when undisturbed.
Colloids:
- Heterogeneous mixtures with very small particles, Appear homogeneous but are not.
- Show Tyndall effect (scattering of light).
- Stable and don’t settle easily, Cannot be filtered, but can be separated using centrifugation.
- Examples include milk, fog, shaving cream.
Tyndall Effect:
- Scattering of light by colloidal particles, Can be observed in fog, forest mist, and milk.
Physical and Chemical Changes:
- Physical change: No change in chemical composition (e.g., melting, boiling).
- Chemical change: Change in composition; new substances formed (e.g., burning).
Burning of a candle is both physical (melting) and chemical (burning).
There are more than 100 known elements, with 92 naturally occurring.
Maharishi Kanad and Pakudha Katayama proposed that matter consists of indivisible particles called “Parmanu.” By the 18th century, scientific understanding advanced, leading to the discovery of laws of chemical combination.
Laws of Chemical Combination
- Law of Conservation of Mass Proposed by Lavoisier, this law states: “Mass can neither be created nor destroyed in a chemical reaction.”
- Law of Constant Proportions Stated by Proust: “In a chemical substance, the elements are always present in definite proportions by mass.”
John Dalton’s Atomic Theory
- All matter is made of tiny particles called atoms.
- Atoms are indivisible and indestructible.
- Atoms of the same element are identical in mass and properties.
- Atoms of different elements differ in mass and properties.
- Atoms combine in whole number ratios to form compounds.
- The relative number and kinds of atoms in a compound are constant.
Atoms
- Atomic radius is measured in nanometres (nm).
Dalton introduced the concept of atomic mass, and the modern standard is based on the carbon-12 isotope, with mass units in unified mass unit (u). (Hydrogen’s atomic mass = 1 u, oxygen = 16 u, etc.)
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Atomic mass unit (u) = 1/12th the mass of one atom of carbon-12.
Atoms are incredibly small and can’t exist independently; they usually combine to form molecules or ions. Molecules consist of one or more atoms held together tightly, while ions are charged species formed by the gain or loss of electrons (e.g., Na⁺, Cl⁻).
The atomicity of elements describes how many atoms are present in their molecule form (e.g., O₂ = diatomic).
- Molecules can be monoatomic (Ar), diatomic (O₂), polyatomic (S₈), etc.
- Atoms of different elements combine in fixed ratios to form compounds.
- Valency is the combining capacity of an element.
- chemical formula represents the composition of a compound.
- Criss-cross method helps balance charges of ions to form a compound.
- Polyatomic ions require brackets if more than one is used in a formula, Examples include formulas of compounds like CaCl₂, Al₂O₃, NaNO₃, and (NH₄)₂SO₄.
- Molecular mass = Sum of atomic masses in a molecule.
- Formula unit mass applies to ionic compounds and is calculated similarly to molecular mass.
in the late 19th and early 20th centuries proved that atoms are made up of sub-atomic particles — electrons, protons, and later neutrons.
- J.J. Thomson discovered the electron and proposed the “plum pudding” model where electrons were embedded in a positively charged sphere.
- E. Goldstein discovered canal rays, leading to the identification of positively charged particles called protons.
- James Chadwick discovered neutrons—neutral subatomic particles found in the nucleus along with protons.
Rutherford’s gold foil (alpha-particle scattering) experiment revolutionized the understanding of atomic structure. The Bohr model of the atom was introduced to address the limitations of Rutherford’s atomic model.
- Bohr proposed that electrons revolve in fixed orbits called energy levels or shells without radiating energy. These shells are designated as K, L, M, N… or n = 1, 2, 3, etc.
- Bohr and Bury proposed : a shell can hold a maximum of 2n² electrons, and electrons fill the inner shells first. The outermost shell can hold a maximum of 8 electrons.
- Electrons in the outermost shell are called valence electrons, and their number determines the valency of an element. If the shell is full, the atom is chemically inactive (like noble gases).
The atomic number of an element is defined as the number of protons in its nucleus, denoted by ‘Z’. The mass number is the total number of protons and neutrons in an atom’s nucleus, denoted by ‘A’.
isotopes are atoms of the same element having the same atomic number but different mass numbers. For example, hydrogen has three isotopes: protium, deuterium, and tritium. (e.g., ¹H, ²H, ³H).
- Isotopes have similar chemical properties but different physical properties.
- Applications of isotopes include their use in nuclear reactors, cancer treatment, and goitre treatment.
Isobars are atoms of different elements with different atomic numbers but the same mass number, like calcium (Z = 20) and argon (Z = 18) both having a mass number of 40.
Metals and Non-metals
Physical Properties of Metals
- Metallic Lustre: Metals in pure state have a shiny appearance.
- Hardness: Most metals are hard, though hardness varies.
- Malleability: Metals can be beaten into thin sheets.
- Ductility: Metals can be drawn into thin wires (e.g., gold is most ductile).
- Conductivity:
- Heat: Silver and copper are excellent heat conductors.
- Electricity: Metals like copper and aluminium conduct electricity.
- Sonority: Metals produce a ringing sound when struck.
- Melting Point: Generally high in metals, except for metals like gallium and caesium.
Physical Properties of Non-Metals
- Include carbon, sulphur, iodine, etc.
- Can be solids or gases, except bromine (liquid).
- Not lustrous, poor conductors, not malleable or ductile.
- Produce acidic oxides when burnt.
- Some non-metals like iodine are exceptions (lustrous).
Important Observations
- Mercury is the only metal liquid at room temperature.
- Alkali metals like sodium and potassium can be cut with a knife.
Reaction with Air
- Metals combine with oxygen to form metal oxides.
- Copper → forms black Copper(II) oxide.
- Aluminium → forms Aluminium oxide.
- Metal oxides are usually basic, but some like Al₂O₃ and ZnO are amphoteric (react acids & bases both).
- Highly reactive metals like sodium & potassium catch fire in air and are stored in kerosene.
- Less reactive metals form a thin oxide layer which protects them from further oxidation.
- Anodising thickens this oxide layer on aluminium.
Reaction with Water
- Metals + Water → Metal Oxide + Hydrogen
- Metal Oxide + Water → Metal Hydroxide
- Potassium, Sodium → React violently with cold water
- Calcium → Reacts less violently
- Magnesium → Reacts only with hot water
- Iron, Zinc, Aluminium → React with steam, not cold water
- Lead, Silver, Gold → Do not react with water
Reaction with Acids
- Metal + Dilute Acid → Salt + Hydrogen Gas
- More reactive metals like Mg react vigorously.
- Copper does not react with dilute HCl.
- Nitric acid does not evolve H₂ because it’s a strong oxidising agent.
- Reactivity order: Mg > Al > Zn > Fe > Cu
Reaction with Solutions of Other Metal Salts
- Displacement reaction: A more reactive metal displaces a less reactive one from its salt solution. E.g., Iron displaces Copper from copper sulphate.
Reactivity of Metals and the Reactivity Series
- Reaction occurred in the copper sulphate solution, Iron is more reactive than copper, so it displaces copper from its solution. Fe (s) + CuSO₄ (aq) → FeSO₄ (aq) + Cu (s) [Displacement reaction.]
Reactivity Series: K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au
Properties of Ionic Compounds
- Physical state: Solid.
- Melting: They melt on heating.
- Solubility: Soluble in water, not in petrol or kerosene.
- Conductivity: Conduct electricity when dissolved in water.
- Inference:
- Ionic compounds are crystalline solids with high melting and boiling points.
- They conduct electricity in molten or aqueous states due to free ions.
Properties of Ionic Compounds
- Physical Nature: Ionic compounds are solid and brittle due to strong electrostatic forces.
- Melting and Boiling Points: These compounds have high melting and boiling points due to strong inter-ionic attractions.
- Solubility: Soluble in water, but not in organic solvents like kerosene and petrol.
- Electrical Conductivity: Conduct electricity in molten or dissolved form, not in solid form.
Occurrence and Extraction of Metals
- Minerals and Ores: Naturally occurring elements/compounds are called minerals. When a mineral contains a high percentage of metal and is economically viable to extract, it is termed an ore.
- Extraction Based on Reactivity:
- Low Reactivity Metals: Extracted by heating alone (e.g., Mercury from cinnabar).
- Moderate Reactivity Metals: Extracted via roasting or calcination, followed by reduction (e.g., Zinc).
- High Reactivity Metals: Extracted by electrolysis (e.g., Sodium, Aluminium).
Refining of Metals
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Electrolytic Refining: Used to purify metals like copper, zinc, gold. The impure metal is made the anode, and the pure metal is deposited on the cathode.
Corrosion – The slow degradation of metals due to environmental exposure.
- Silver: Turns black due to silver sulphide.
- Iron: Forms rust in presence of moisture and air.
- Copper: Turns green due to basic copper carbonate.
Prevention of Corrosion
- Painting, oiling, greasing Galvanising (coating with zinc) Chrome plating, anodising, or making alloys
- Galvanisation: Coating iron/steel with zinc. Even if the zinc layer is scratched, iron doesn’t rust easily.
Alloying: Improving Metal Properties
- Alloying improves metal properties by mixing them with other metals or non-metals. Example: Iron + Carbon → Stronger iron Iron + Nickel + Chromium → Stainless Steel (hard, rust-free)
- Alloy: A homogeneous mixture of two or more metals or a metal and a non-metal. Formed by melting and mixing elements in definite proportions.
Interesting Facts
- 24-carat gold is soft and not used in jewellery. 22-carat gold is used in India, alloyed with copper or silver.
- Amalgam: Alloy of mercury.
- Examples of alloys: Brass: Copper + Zinc Bronze: Copper + Tin Solder: Lead + Tin (used for joining wires)
- Iron Pillar of Delhi: Built 1600+ years ago. Resistant to rust due to advanced ancient technique. 8 m tall and weighs 6000 kg (6 tonnes).
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