Matter surrounds us in every aspect of our environment. Everything from the book in your hand to the distant stars in the sky, including your own body, is composed of different forms of matter in various states. To comprehend processes such as the sun's heat emission and the origin of life on Earth, studying matter and its states is crucial.

Matter is anything that has mass and occupies space. A material is a specific kind of matter with distinct properties and composition. Matter exhibits the following distinct properties:

• Matter is made up of particles that are always in motion.
• There is space between the particles, known as interparticle or intermolecular space.
• The particles of matter attract each other.

States of Matter

Previously, three states of matter were recognized, but scientists have now confirmed the existence of five states of matter. The three common states of matter in our surroundings are solid, liquid, and gas, while the two new states, plasma and Bose-Einstein Condensate (BEC), are produced under extreme conditions. The three main states of matter are:

Solid State

• Solids have a definite shape and fixed volume, exhibiting negligible compressibility.
• They tend to retain their shape when subjected to external forces.
• Solids are rigid; they may break when force is applied, but their shape is difficult to change.

Liquid State

• Liquids have a definite volume but no definite shape.
• They flow and change shape, making them non-rigid and classified as fluids.

Gaseous State

• Gases neither have a definite volume nor a definite shape.
• Gaseous particles are in random motion, constantly colliding with each other and the walls of their container.
• The pressure exerted by a gas is due to the force exerted by gas particles per unit area on the container walls.
• The rate of diffusion is higher in gases compared to solids.

Interconversion of States of Matter

The states of matter can be interchanged by altering temperature and pressure.

Effect of Temperature

• As the temperature of a solid increases, the kinetic energy of its particles rises, causing them to vibrate more vigorously.
• The heat energy supplied overcomes the attractive forces between particles, allowing them to move freely and change state.
• The melting point is the temperature at which a solid turns into a liquid at atmospheric pressure. It decreases in the presence of impurities. The melting point of ice is 273.16 K.
• Kelvin (K) is the SI unit of temperature: 0°C = 273.16 K, for convenience, 0°C = 273 K.
• The process of a solid turning into a liquid at its melting point is called fusion.
• The latent heat of fusion is the amount of heat required to change 1 kg of solid into liquid at atmospheric pressure at its melting point.
• The boiling point is the temperature at which a liquid begins to boil at atmospheric pressure. It increases in the presence of impurities. The boiling point of water is 373 K (100°C).
• The conversion of a liquid into its vapor (gaseous state) is called vaporization or boiling.
• The rate of evaporation increases with an increase in surface area, temperature, and wind speed, and decreases with higher humidity.
• The latent heat of vaporization is the amount of energy required to change 1 L of water into vapor at its boiling point.
• Sublimation is the process of a solid converting directly into a gaseous state and vice-versa. Substances that undergo sublimation are called sublimates (e.g., naphthalene, benzoic acid, iodine).
• Steam has more energy than water at the same temperature because particles in steam have absorbed extra energy as the latent heat of vaporization.

Effect of Pressure

• Applying pressure compresses a gas as the particles come closer, decreasing the interparticle distance, thus converting the gas into a liquid.
• A gas can be liquefied by applying pressure and reducing temperature.
• Higher external pressure increases the boiling point of a liquid because more heat is needed to make the vapor pressure equal to the external pressure.
• Solid carbon dioxide (dry ice) converts directly to the gaseous state at a pressure of 1 atmosphere without becoming liquid.

Classification of Matter

At the macroscopic or bulk level, matter can be classified as mixtures or pure substances, further categorized based on chemical composition.

Pure Substance

A pure substance has a fixed chemical composition and distinct properties. Pure substances are mainly of two types: elements and compounds.

Element

• An element consists of only one type of particle called atoms and cannot be divided into simpler substances by chemical or physical methods.
• There are 118 known elements, with 92 occurring naturally and the rest synthesized artificially.
• Elements can be metals (e.g., copper, aluminum, iron), non-metals (e.g., sulfur, oxygen, nitrogen), or metalloids (e.g., antimony, germanium).
• Elements in the Earth's crust: oxygen > silicon > aluminum > iron > calcium > sodium ≈ potassium > magnesium > others.
• Elements in the human body: oxygen > carbon > hydrogen > nitrogen > phosphorus > sodium.

Compound

• A compound forms when two or more elements combine in a fixed proportion, resulting in a homogeneous substance with properties different from its constituent elements.
• Examples of compounds include water, ammonia, sugar, and carbon dioxide.
• Constituents of a compound cannot be separated into simpler substances by physical methods but can be separated by chemical processes.

Mixtures

A mixture contains two or more elements or compounds in any proportion. Most of the matter around us exists as mixtures of two or more pure components, such as sea water, soil, and minerals.

Key characteristics of mixtures:

• They can be separated into their constituents by physical methods.
• They do not have a fixed melting and boiling point.
• They do not have a definite chemical formula.

Types of Mixtures

Homogeneous Mixtures

• Homogeneous mixtures have a uniform composition throughout and exhibit the same properties as their component elements.
• Examples include air, solutions like sugar dissolved in water, salt dissolved in water, and alloys (mixtures of two or more metals).
• Homogeneous mixtures can have variable compositions as long as they contain the same components.

Heterogeneous Mixtures

• Heterogeneous mixtures do not have a uniform composition, and different parts of the mixture may have different properties.
• Examples include suspensions (mixture of salt and sugar), and colloids.

Purification & Characterization of Substances

Heterogeneous mixtures can be separated into their respective components by simple physical methods:

1. Handpicking: Manual separation of undesired impurities, e.g., removing stones from pulses.
2. Threshing: Separating grains from husks by beating the stems with machines or a wooden stick.
3. Winnowing: Separating lighter components of the mixture from heavier ones by allowing the mixture to fall from a height, with lighter ones being blown away by wind.
4. Sieving: Separating particles of different sizes through a sieve.
5. Sedimentation and Decantation: Allowing a mixture of insoluble components to settle, then pouring off the upper lighter liquid layer.
6. Filtration: Allowing solid ingredients in a liquid to pass through a filter paper with fine pores, leaving the solid on the paper and collecting the liquid at the bottom of the container.

Homogeneous mixtures can be separated by special techniques:

1. Evaporation: Used to separate a mixture of volatile and non-volatile substances. On heating, the volatile substance evaporates, leaving behind the non-volatile substance.
2. Centrifugation: Based on the density of particles. Used in diagnostic laboratories, dairies, and washing machines to separate substances like blood and urine, butter from cream, and to squeeze out water from wet clothes.
3. Sublimation: Used for sublimates, i.e., solid substances which convert directly into vapors without passing through the liquid state, such as ammonium chloride, camphor, naphthalene, benzoic acid, and anthracene.
4. Crystallization: The process of separating a pure solid in the form of its crystals, for example, purifying salt from sea water.
5. Distillation: Used to separate two miscible liquids. The mixture is boiled in a flask, and the more volatile liquid boils first, with its vapors passed through a condenser and collected in another container. Steam distillation is used for purifying steam-volatile compounds that are immiscible with water.
6. Chromatography: A modern technique for separation and purification of organic compounds, based on the distribution of components between a stationary phase (solid or liquid supported over a solid) and a mobile phase (liquid or gas).

Solution

A solution is a homogeneous mixture of two or more substances.

Key characteristics of solutions:

• The particle size in a solution is smaller than 1 nm in diameter.
• The component present in larger amount is called the solvent, while the component present in smaller amount is called the solute.
• Solutions are stable and their particles do not settle down when left undisturbed.
• Particles in solutions are too small to be seen and cannot be separated from the mixture by filtration.
• An example of a solution is tincture of iodine, which is a solution of iodine in alcohol.
• Solutions do not scatter a beam of light passing through them, making the path of light invisible.

Suspension

A suspension is a heterogeneous mixture in which the solute particles do not dissolve but remain suspended throughout the bulk of the medium.

Key characteristics of suspensions:

• The particle size is more than 100 nm, making them visible to the naked eye.
• Suspensions can be separated by filtration as the particles cannot pass through filter paper.
• Examples of suspensions include mixtures like dust in air and muddy water.

Colloids or Colloidal Solution

A colloidal solution is a heterogeneous solution containing two phases: dispersed phase and dispersion medium.

Key characteristics of colloids:

• The particle size is in between 1 nm to 100 nm, making them too small to be seen by the naked eye.
• Colloids do not settle down when left undisturbed due to Brownian motion.
• Colloids scatter light passing through them, making the path of light visible (Tyndall effect).
• They cannot be separated by filtration, but can be separated by techniques like centrifugation.
• Examples of colloids include foams, emulsions like milk, and aerosols like smoke.

Types of Colloids

Colloids are classified based on the state of the dispersed phase and dispersion medium:

Physical Changes

A physical change is a change in which the composition of a substance remains unchanged, meaning no new substance is formed.

Examples of physical changes:

• Interconversion of states of matter (solid, liquid, gas, plasma, BEC).
• Magnetisation of iron in the presence of a magnetic field.
• Evaporation, distillation, sublimation, condensation, and crystallisation.

Example: Blue Copper Sulphate Crystal

The blue color of copper sulphate crystal disappears when heated, due to the loss of water of crystallisation. It regains its blue color when moistened with water.

CuSO4 • 5H2O → CuSO4 • 2H2O + 3H2O

This is an example of a physical change.

Characteristics of Physical Changes

• Physical changes are temporary and reversible.
• They are due to the change in physical properties like density, volume, and state temporarily.
• The composition of the substance remains the same during physical changes.

Chemical Changes

Chemical changes result in the formation of new substances with different chemical properties.

Examples of chemical changes:

• Rusting of iron when exposed to moist air.
• Digestion of food in the stomach where complex food molecules are converted into simple molecules of glucose.
• Burning of substances, accompanied by the production of heat.

Example: Blue Copper Sulphate Crystal

Blue copper sulphate crystals, when heated, lose their water molecules and turn into white anhydrous salt, which further decomposes into black cupric oxide and sulphur trioxide on strong heating.

CuSO4 • 5H2O → CuSO4 + 5H2O → CuO + SO3

This is an example of a chemical change.

Characteristics of Chemical Changes

• During chemical changes, the composition of the substance changes due to a change in constituent particles.
• The identity of the substance is lost during a chemical change.
• Energy is absorbed or released during a chemical change.
• Chemical changes are permanent and irreversible.

Non-metals

• Non-metals are elements that tend to form negative ions by gaining electrons.
• Examples include hydrogen, oxygen, sulphur, etc.
• Non-metals are usually in solid or gaseous form under normal conditions.

Physical Properties of Non-metals

• Physical State: Most non-metals are soft if solid. Diamond (an allotrope of carbon) is the hardest known substance. Bromine is the only non-metal that is a liquid.
• Lustre: Non-metals do not have lustre, except for diamond, graphite, and iodine which exhibit lustre.
• Malleability and Ductility: Non-metals cannot be beaten into sheets or drawn into wires due to weak attractive forces between their atoms.
• Electrical and Thermal Conductivity: Non-metals are generally poor conductors of heat and electricity.
• Brittleness: Non-metals are brittle in nature.
• Melting and Boiling Points: Non-metals generally have low melting and boiling points. Solid non-metals like boron, silicon, and carbon have comparatively higher boiling points.

Chemical Properties of Non-metals

• Electronegativity: Non-metals are electronegative and tend to form anions by gaining electrons.
• Reaction with Oxygen: Non-metals combine with oxygen to form oxides. For example, sulphur burns in air to form sulphur dioxide.
• Formation of Oxides: Non-metals form acidic oxides (CO2, SO2, SO3, NO2, P2O5) as well as neutral oxides (CO, NO, N2O, H2O2).
• Reaction with Water: Non-metals generally do not react with water. Phosphorus is a notable exception, reacting vigorously with air.
• Reaction with Bases: Reactions of non-metals with bases are complex.

Metalloids

Metalloids possess properties that are intermediate between metals and non-metals. They are positioned in the periodic table between metals and non-metals.

• Examples include Ge (Germanium), Ga (Gallium), Si (Silicon), etc.

Metallurgy

The process of obtaining a pure metal from its ore is called metallurgy. Various steps involved in the extraction of metals are termed as metallurgical operations.

Principal Ores of Some Important Metals

Methods used in the Extraction of Metals

Metallurgical operations depend upon the reactivity of metals:

• A more reactive metal (like Na, K, Mg, Ca, Al) is obtained by the electrolysis of its salt in molten state. For example, sodium is obtained by the electrolysis of fused sodium chloride.

Ore Concentration

The removal of impurities from the ore is called its concentration. It is carried out by the following methods:

1. Gravity Separation or Hydraulic Washing: This method is generally used for the concentration of oxide ores. Powdered ore is agitated with a running stream of water. The lighter gangue particles are taken away by water while heavier ore particles settle down.
2. Froth Floatation Process: This method is generally used for the concentration of sulphide ores. In this process, powdered ore is mixed with water and pine oil in a tank. Compressed air is blown into the mixture, causing froth to form. The ore particles stick to the froth which floats on the top and can be separated easily, leaving the gangue particles behind.
3. Magnetic Separation: This method is used when one component, either the ore or impurity, is magnetic in nature. For example, wolframite (magnetic) is separated from non-magnetic ore cassiterite (SnO2) by this method.
4. Leaching: In this method, concentration is done by chemical methods using chemical reagents. Gold, silver, and aluminium are concentrated by this method.

Conversion of Ore to Metal Oxide

1. Roasting: In this method, concentrated ore is heated in excess air. It converts sulphide ores into oxides and oxidizes impurities.
2. Calcination: The concentrated ore is heated in a reverberatory furnace below its melting point, in the absence of air and without the addition of any external substance. Calcination is mainly done for carbonate or hydroxide ores.

Reduction of Metal Oxide

The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon.

• For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc:
  ZnO (s) + C (s) → Zn (s) + CO (g)

Refining of Metals

The metals produced by various reduction processes described above are not very pure. They contain impurities that must be removed to obtain pure metals. The most widely used method for refining impure metals is electrolytic refining. 

Corrosion

Corrosion is the slow process of decaying away of metals by reaction with atmospheric air, moisture, acids, bases, etc. Examples include rusting of iron, tarnishing of silver, and formation of green coating over copper. It is an electrochemical process where metals oxidize and form oxides and other salts.

Prevention of Corrosion

1. Galvanisation: Coating iron and steel objects with a thin layer of zinc by dipping them in molten zinc. The zinc coating protects against rusting even if it is scratched.
2. Alloying: Improving the properties of a metal by mixing it with other metals.
3. Painting: Coating iron surfaces with paint to protect them from air and moisture.
4. Greasing or Oiling: Applying grease or oil to the surface of iron objects to prevent contact with air and moisture, thus preventing rust.
5. Tin-Plating and Chromium Plating: Depositing a thin layer of tin or chromium on iron objects through electroplating to protect them from rusting.

Alloys

An alloy is a homogeneous mixture of two or more metals. Alloys have lesser chemical activity and lower melting points than their pure constituents, and they are usually harder.

Some Important Alloys and Their Uses

Atomic Structure

Atoms are composed of smaller particles known as sub-atomic particles, including electrons, protons, and neutrons. Protons and neutrons are located in the nucleus, while electrons revolve around the nucleus. The arrangement of electrons around the nucleus determines the structure of an atom.

Atom

An atom is the smallest particle of an element that takes part in a chemical reaction. It consists of three fundamental particles: electrons, protons, and neutrons. Protons and neutrons are located inside the nucleus, while electrons revolve around the nucleus in closed orbits.

Properties of Atom

• Atomic Number: The atomic number (Z) is the number of protons present inside the nucleus of an atom. In an isolated atom, atomic number = Number of protons = Number of electrons.
• Mass Number: The mass number (A) is equal to the sum of the number of protons and neutrons present in the nucleus of an atom. Mass number = Number of protons + Number of neutrons = Atomic number (Z) + Number of neutrons (N).
• An atom with atomic number Z and mass number A is represented as Z^AX.

Molecules

A molecule is the smallest particle that can remain in a free state. It can consist of one or more different or same atoms. Noble gases are examples of monoatomic molecules.

Different Atomic Species

• Isotopes: Atoms of an element with the same atomic number but different mass numbers are called isotopes. For example, ¹H (protium), ²H (deuterium), and ³H (tritium) are isotopes of hydrogen.
• Isobars: Atoms of different elements that have the same mass number but different atomic numbers are called isobars. For example, ⁴⁰Ar and ⁴⁰Ca.
• Isotones: Atoms of different elements that have the same number of neutrons but different atomic numbers and mass numbers are called isotones. For example, ¹³C and ¹⁴N.

Distribution of Electrons in Various Orbitals

The distribution of electrons in various orbitals follows specific rules:

Bohr Bury Scheme (1921)

• The maximum number of electrons in an orbit is 2²ⁿ, where n is the number of the orbit.
• The maximum number of electrons in the outermost orbit is 8, and the penultimate orbit cannot have more than 18 electrons.
• The outer orbit cannot have more than 2 electrons, and the penultimate shell cannot have more than 9 electrons until the inner penultimate shell has 2n² electrons.

Chemical Reactions, Organic Chemistry, and Hydrocarbons

Rules for Filling Electrons in the Orbitals

1. Aufbau Principle: Orbitals are filled up according to the increase in their energy. The order of filling electrons is:

   Orbital Type	Order of Filling
   s	1s < 2s < 3s < 4s < 5s < 6s < 7s
   p	2p < 3p < 4p < 5p
   d	3d < 4d < 5d < 6d
   f	4f < 5f < 6f < 7f

2. Pauli’s Exclusion Principle: Each orbital can have a maximum of two electrons with opposite spins.
3. Hund’s Rule: Electrons fill each orbital singly before pairing up.

Chemical Reaction

A chemical reaction is a change in which one or more substances react to form new substances with different properties. Various types of chemical reactions include:

1. Combination Reactions: Two or more materials combine to form one material. Example: reaction of calcium oxide with water:
   CaO + H₂O → Ca(OH)₂
2. Decomposition Reactions: One material breaks down to form two or more materials. Examples:
   • Ca(OH)₂ → CaO + H₂O
   • HgO → Hg + 1/2O₂
   • 2H₂O → 2H₂ + O₂
   • Na₂CO₃ → Na₂O + CO₂
3. Displacement Reactions: More reactive metal displaces less reactive element from its salt. Examples:
   • Mg + 2HCl → MgCl₂ + H₂
   • Pb + CuCl → PbCl₂ + Cu
   • Zn + CuSO₄ → ZnSO₄ + Cu
4. Double Displacement Reactions: Metals of two compounds interchange to form two new salts. Examples:
   • FeS + 2HCl → FeCl₂ + H₂S
   • 2KOH + H₂SO₄ → K₂SO₄ + 2H₂O
   • AgNO₃ + NaCl → AgCl + NaNO₃
5. Oxidation and Reduction Reactions: Oxidation is the addition of oxygen or removal of hydrogen, while reduction is the removal of oxygen or addition of hydrogen. When both occur in a reaction, it is called a redox reaction. Example: rusting of iron.

Organic Chemistry

Organic chemistry involves the study of compounds of carbon except its oxides, carbonates, and hydrogen carbonate salts. These compounds were initially thought to be extracted only from natural substances, leading to the term organic compounds, but now it is known that these compounds can be synthesized in laboratories.

Hydrocarbons

Hydrocarbons are organic compounds containing only carbon and hydrogen.

Saturated Hydrocarbons (Alkanes)

These hydrocarbons have all carbon atoms linked by single bonds. They are chemically inert and burn with a blue, non-smoky flame due to complete combustion.

General formula: C_nH_2n+2

Unsaturated Hydrocarbons

These compounds have at least one double or triple bond along with single bonds. They burn with a sooty or smoky flame due to incomplete combustion.

Alkenes (Olefins)

Contain at least one double bond.

General formula: C_nH_2n

Alkynes

Contain at least one triple bond.

General formula: C_nH_2n-2

The minimum number of carbon atoms in an unsaturated compound is 2 because double or triple bonds are only possible between 2 carbon atoms.

Homologous Series

A homologous series is a series of compounds with similar chemical properties, where each member contains the same functional group and differs from the next by a \(\text{—CH}_2\) unit in its molecular formula.

Examples

For example, CH₄, C₂H₆, C₃H₈, C₄H₁₀ are members of the alkane family.

Characteristics of a Homologous Series

• All members of a homologous series can be represented by the same general formula.
• Successive members in the series differ by 1 carbon atom and 2 hydrogen atoms in their molecular formula.
• All compounds in a homologous series exhibit similar chemical properties.
• Physical properties gradually change with increasing molecular mass; for instance, melting and boiling points increase as molecular mass increases.
• The difference in molecular mass between any two adjacent members of a homologous series is approximately 14 atomic mass units (u).

Examples of Homologous Series