General Chemistry — UPSC Science & Tech Notes

1. What is Matter?

Matter is anything that has mass and occupies volume. In simple terms, if it is 'stuff' you can touch or weigh, it is matter. Matter is made of tiny particles called atoms.

Molecules: Atoms rarely stay alone. They join together in groups called molecules. For example, two Hydrogen atoms and one Oxygen atom join to form a Water molecule.

2. The Evolution of Atomic Theory

Scientists didn't always know what an atom looked like. Our current understanding is the result of centuries of discovery:

1. Dalton’s Atomic Theory: He proposed that atoms are like solid, indivisible billiard balls. He was the first to say each element has its own type of atom.

2. J.J. Thomson (Plum Pudding Model): He discovered the Electron. He thought an atom was a positive sphere with negative electrons stuck in it like raisins in a pudding.

3. Ernest Rutherford (Gold Foil Experiment): He discovered the Nucleus. He realized the atom is mostly empty space with a tiny, dense, positive center.

4. Niels Bohr: He suggested that electrons move in fixed circular paths called shells or orbits, similar to planets orbiting the Sun.

5. Schrödinger (Quantum Model): The modern view. Electrons don't follow neat paths; they exist in orbitals (clouds) where they are most likely to be found.

3. Subatomic Particles: The Building Blocks

An atom is composed of three main particles. Their balance determines how an element behaves.

A. Protons (p+)

Identity: The number of protons is the Atomic Number. It acts like a fingerprint—if you change the number of protons, you change the element itself.

Location: Inside the dense central Nucleus.

B. Neutrons (n0)

The Nuclear Glue: Since protons are all positive, they want to repel each other. Neutrons provide the Strong Nuclear Force that holds the nucleus together.

Isotopes: Atoms of the same element can have different numbers of neutrons. These are called Isotopes (e.g., Carbon-12 and Carbon-14).

C. Electrons (e-)

Chemical Reactivity: Electrons are found in the Electron Cloud outside the nucleus. They are responsible for chemical bonding and electricity.

Mass: They are so light (1/1840th of a proton) that their mass is usually ignored when calculating atomic weight.

4. Electron Configuration: Organizing the Shells

Electrons are not just floating randomly; they occupy specific energy levels or Shells.

Shell Capacity (2n²): The 1st shell (K) holds 2, the 2nd (L) holds 8, the 3rd (M) holds 18.

The Octet Rule: Atoms are most stable when their outermost shell is full (usually with 8 electrons). This is why atoms react—to find or share electrons to reach that 'magic 8'.

Aufbau Principle: This rule states that electrons fill the lowest energy levels first (like filling the ground floor of a building before moving up).

Example: Sodium (Atomic No. 11) has 11 electrons. They fill up as 2 in the first shell, 8 in the second, and 1 in the last (2-8-1).

1. Atomic Number (Z): The Chemical Fingerprint

The Atomic Number is the most fundamental property of an atom. It is the number of protons found in the nucleus.

Why is it important?

Identity: It determines which element the atom is. If an atom has 6 protons, it is Carbon; if it has 92, it is Uranium. No two different elements have the same atomic number.

Neutrality: In a neutral atom, the number of protons equals the number of electrons. Since electrons govern chemical reactions, the atomic number indirectly dictates an element's chemical behavior.

Symbol: Represented by the letter Z.

2. Atomic Mass and AMU

The mass of an atom is concentrated in its nucleus (protons + neutrons). Electrons are so light that their mass is negligible.

Atomic Mass Unit (amu): Because atoms are so small, we use a relative scale. 1 amu is defined as exactly 1/12th the mass of a Carbon-12 atom.

Dalton (Da): Another name for the atomic mass unit.

Standard Reference: Carbon-12 (6 protons and 6 neutrons) is the international standard for measuring the masses of all other elements.

3. Elements and Symbols

An Element is a pure substance consisting of atoms that all have the same atomic number.

Rules for Symbols:

Symbols are shorthand for the element's name (English or Latin).

The first letter is always Uppercase, and the second letter (if present) is always lowercase.

Examples:

Hydrogen: H

Aluminum: Al (not AL)

Cobalt: Co (not CO, which is Carbon Monoxide)

Gold (from Latin Aurum): Au

4. Isotopes: Same Identity, Different Weight

Isotopes are atoms of the same element (same Atomic Number) that have different numbers of neutrons (different Mass Numbers).

Example: Hydrogen's Three Species

1. Protium (1H1): 1 proton, 0 neutrons (Most common).

2. Deuterium (1H2): 1 proton, 1 neutron (Used in heavy water).

3. Tritium (1H3): 1 proton, 2 neutrons (Radioactive).

Properties of Isotopes:

Chemical Properties: They are identical because they have the same number of valence electrons.

Physical Properties: They are different. Their densities, boiling points, and melting points vary slightly due to the difference in mass.

5. Representation of an Atom

In chemical notation, we write the symbol with the Mass Number (A) as a superscript on the top left and the Atomic Number (Z) as a subscript on the bottom left.

Formula: Mass Number (A) = Protons (Z) + Neutrons (n)

1. Synthetic Elements: Human-Made Wonders

There are 24 known chemical elements that do not occur naturally on Earth. These are called Synthetic Elements.

How they are made: They are created by scientists in labs by manipulating fundamental particles or smashing atoms together in particle accelerators.

Stability: Most synthetic elements are unstable and decay into other elements very quickly (radioactive).

2. Molecules and Atomicity

A Molecule is a group of two or more atoms chemically bonded together. It is the smallest particle of a substance that can exist independently.

A. Molecules of Elements

These contain only one type of atom. The number of atoms in a molecule is its Atomicity.

Monoatomic: Made of 1 atom (e.g., Helium - He, Argon - Ar).

Diatomic: Made of 2 atoms (e.g., Oxygen - O2, Hydrogen - H2).

Triatomic: Made of 3 atoms (e.g., Ozone - O3).

B. Molecules of Compounds

Atoms of different elements join in fixed proportions. For example, Water (H2O) always has 2 Hydrogen atoms for every 1 Oxygen atom.

3. Ions: The Charged Particles

When atoms lose or gain electrons, they become Ions. These carry an electric charge.

Cation: A positively charged ion (formed when an atom loses electrons). Example: Sodium ion (Na+).

Anion: A negatively charged ion (formed when an atom gains electrons). Example: Chloride ion (Cl-).

Polyatomic Ions: A group of atoms that carry a net charge (e.g., Ammonium - NH4+).

4. Chemical Bonding: The Atomic Glue

Chemical bonds are the forces that hold atoms together. Bonding happens because atoms want to achieve a stable electron arrangement (usually 8 electrons in the outer shell).

A. Ionic Bond (Electrovalent Bond)

This bond is formed by the complete transfer of electrons from one atom to another.

Formation: Usually occurs between a Metal (which loses electrons) and a Non-metal (which gains them).

Properties of Ionic Compounds:

1. State: Mostly crystalline solids.

2. Melting/Boiling Points: Very high because of strong electrostatic attraction between ions.

3. Conductivity: They conduct electricity when dissolved in water or melted.

B. Covalent Bond

This bond is formed by the equal sharing of electrons between atoms. This usually happens between two non-metals (e.g., H2O, CH4).

C. Metallic Bond

Found in metals, where electrons are not bound to one nucleus but float in a 'sea' of electrons, allowing metals to conduct electricity and heat easily.

1. Covalent Bond: The Power of Sharing

A Covalent Bond is formed when two atoms share electrons equally to achieve stability. This typically happens between two non-metals.

Key Characteristics:

1. Physical State: Unlike ionic compounds, these can be found as solids, liquids, or gases (e.g., Water is liquid, Oxygen is gas, Diamond is solid).

2. Boiling and Melting Points: Generally low compared to ionic bonds because the forces holding individual molecules together are weaker.

3. Thermal Stability: Lower than ionic bonds, meaning they break down more easily under heat.

2. Metallic Bond: The 'Electron Sea'

Metallic bonding occurs in metals like Iron (Fe), Gold, and Copper. In these structures, the outer-shell electrons are not tied to any single atom.

How it works:

Free Electrons: The valence electrons detach from their atoms and form a 'cloud' or 'sea' of electrons.

Attraction: The bond is the strong attraction between these floating negative electrons and the positive metal ions.

Properties of Metals:

High Conductivity: Because the electrons are free to move, they can carry electricity and heat very efficiently.

Strength: They form 'giant structures' that are held together tightly by these forces.

3. Molecular Bonds and Intermolecular Forces

While chemical bonds hold atoms together inside a molecule, Intermolecular Forces are the attractions that exist between different molecules. These forces decide if a substance is a liquid or a gas.

Types of Intermolecular Forces:

1. Hydrogen Bonding: A very strong attraction found in molecules like water. It's why water has such a high boiling point compared to other liquids.

2. Dipole-Dipole: Occurs between molecules that have 'poles' (one side slightly positive, one side slightly negative).

3. London Dispersion Forces: Weak, temporary attractions that happen in all molecules (also called Ion-Dipole forces).

4. Ion-Ion Interaction: The strongest attraction, typically seen in ionic crystals.

1. Polarity: The Tug-of-War for Electrons

Polarity describes how electrons are shared between atoms in a covalent bond. It depends on Electronegativity (the 'hunger' an atom has for electrons).

A. Non-Polar Covalent Bonds

Definition: Formed between atoms with the same electronegativity. Electrons are shared equally.

Dipole Moment: It is zero because there is no partial charge.

Example: H2, O2, Cl2 (identical atoms pull equally).

B. Polar Covalent Bonds

Definition: Formed between atoms with different electronegativities. Electrons are shared unequally, spending more time near the stronger atom.

Dipole Moment: Created because one end of the molecule becomes slightly negative (delta -) and the other slightly positive (delta +).

Example: Water (H2O) and Hydrogen Chloride (HCl).

2. Chemical Compounds

A Compound is a pure substance formed when atoms of two or more different elements combine chemically in a fixed ratio.

Properties: The properties of a compound are entirely different from the elements that make it. For example, Sodium (a reactive metal) and Chlorine (a toxic gas) combine to form Table Salt (NaCl), which is essential for life.

Diversity: Everything from our DNA to the water we drink is a compound. Even changing one single atom in a large compound can turn a medicine into a poison.

3. Organic Compounds: The Chemistry of Life

Organic compounds are substances that primarily contain Carbon (C). Carbon is unique because it can form stable bonds with itself and other elements, creating complex structures.

Biological Importance: Fats, proteins, carbohydrates, enzymes, and hormones are all organic compounds.

Structural Framework: Carbon atoms link together to form chains or rings, providing the 'skeleton' of the molecule.

4. Carbon-Carbon Bonds

Carbon can bond with other carbon atoms in three different ways, leading to various types of hydrocarbons:

1. Single Bonds (Alkanes): Known as 'Sigma' bonds. These are saturated hydrocarbons (e.g., Methane, Ethane).

2. Double Bonds (Alkenes): Involves four shared electrons. These are unsaturated (e.g., Ethene).

3. Triple Bonds (Alkynes): Involves six shared electrons. These are highly reactive and unsaturated (e.g., Ethyne/Acetylene).

1. Polarity: The Tug-of-War for Electrons

Polarity describes how electrons are shared between atoms in a covalent bond. It depends on Electronegativity, which is the 'hunger' an atom has for electrons. Think of it as a game of tug-of-war where the stronger atom pulls the electrons closer to itself.

A. Non-Polar Covalent Bonds

Definition: Formed between atoms with the same electronegativity (identical pull). Because both atoms are equally strong, the electrons stay right in the middle.

Dipole Moment: It is zero because there is no partial charge on either side.

Example: H2, O2, Cl2. Since both atoms are the same, they share equally.

B. Polar Covalent Bonds

Definition: Formed between atoms with different electronegativities. One atom is a 'bully' and pulls the shared electrons closer to its side.

Dipole Moment: This creates a Dipole. The side with more electrons becomes slightly negative (delta negative), and the side with fewer becomes slightly positive (delta positive).

Example: Water (H2O) and Hydrogen Chloride (HCl). In water, Oxygen pulls electrons away from Hydrogen.

2. Chemical Compounds

A Compound is a pure substance formed when atoms of two or more different elements combine chemically in a fixed ratio (like a specific recipe).

Transformation of Properties: The properties of a compound are entirely different from the elements that make it. For example, Sodium (an explosive metal) and Chlorine (a poisonous gas) combine to make Table Salt (NaCl), which we eat every day!

Complexity: Compounds make up everything from the DNA in your cells to the CO2 we exhale. Even changing one single atom in a large molecule can change a life-saving medicine into a deadly toxin.

3. Organic Compounds: The Chemistry of Life

Organic compounds are substances that primarily contain Carbon (C). Carbon is the 'LEGO brick' of life because it can bond with almost anything, including itself, to form long chains and complex rings.

Biological Importance: Your body is essentially a collection of organic compounds like fats, proteins, carbohydrates, enzymes, and hormones.

Versatility: Carbon's unique ability to form four bonds allows for a staggering variety of millions of different structures.

4. Carbon-Carbon Bonds

The way carbon atoms link together determines what kind of material is formed. These are categorized by the number of shared electron pairs:

1. Single Bonds (Alkanes): These share one pair of electrons. They are called Saturated because they have as much hydrogen as they can possibly hold (e.g., Methane - the gas in your stove).

2. Double Bonds (Alkenes): These share two pairs of electrons. They are Unsaturated and more reactive (e.g., Ethene, used to ripen fruit).

3. Triple Bonds (Alkynes): These share three pairs of electrons. They are highly reactive and contain a lot of energy (e.g., Acetylene, used in welding torches).

1. What are Polymers?

A Polymer is a giant molecule made by linking hundreds or thousands of smaller units called Monomers.

Polymerization: The chemical process of joining these monomers into a long, chain-like network.

The Hydrocarbon Backbone: Most synthetic polymers (like plastics) have a central chain of Carbon and Hydrogen atoms. This is possible because Carbon is tetravalent (it can form four bonds), allowing it to create incredibly long, stable chains.

Diversity of Polymers:

Flexible: Rubber, polyester, and polythene.

Rigid: Epoxy resins, Bakelite (used in electrical switches), and toughened glass.

2. The Chemistry of Cleaning: Soaps vs. Detergents

How does soap actually remove grease? It uses a special molecular structure that bridges the gap between water and oil.

Micelles: Soap molecules have a 'head' that loves water and a 'tail' that loves oil. They surround oil droplets, lifting them off the skin and suspending them in water so they can be rinsed away.

Soaps: They are sodium or potassium salts of long-chain fatty acids. They are mild but don't work well in 'hard water' (water with high mineral content).

Detergents: These are synthetic derivatives with a much stronger cleaning action. They work effectively even in hard water.

3. Natural Chemistry: Onions and Fruit Ripening

Why do Onions make us cry?

When you cut an onion, you break its cells, releasing enzymes and sulfenic acid. These react to form a gas called Syn-propanethial-S-oxide. When this gas touches the moisture in your eyes, it turns into a mild sulfuric acid, causing your tear glands to react to flush it out.

How do fruits ripen?

The secret is a gaseous plant hormone called Ethylene.

The Signal: As a fruit matures, it begins producing more ethylene. This gas acts as a signal to the fruit to change its color, soften its texture, and convert starches into sugars.

Commercial Use: Farmers often use ethylene gas to ripen green-harvested fruits (like bananas) before they reach the market.

4. Everyday Chemistry Questions for UPSC

Why does Ice float? Because of Hydrogen Bonding, ice forms a crystal lattice that is less dense than liquid water.

What is an Antacid? It is a Base (like Magnesium Hydroxide) used to neutralize excess Hydrochloric acid in the stomach.

Do bottled water expire? The water doesn't, but the Plastic (Polymer) bottle can eventually leak chemicals into the water over time.

How do Sunscreens work? They contain chemicals like Zinc Oxide that either reflect or absorb harmful UV radiation.

1. Thermoplastics vs. Thermosetting Plastics

Not all plastics are created equal. Their molecular structure determines whether they can be melted and reshaped or if they stay rigid forever.

A. Thermoplastics

Behavior: These soften on heating and can be reshaped multiple times. They have linear or slightly branched chains.

Recyclability: Highly recyclable because they can be melted down and molded into new products.

Examples: Polythene (bags), PVC (pipes), and Polystyrene (containers).

B. Thermosetting Plastics

Behavior: Once molded and set, they cannot be softened or reshaped by heating. This is due to 'cross-links' between polymer chains that create a permanent 3D network.

Properties: They are heat-resistant and durable.

Examples: Bakelite (used in electrical switches and saucepan handles) and Melamine (used in fire-resistant fabrics and floor tiles).

2. Chemistry of Food Preservatives

Preservatives are substances added to food to prevent spoilage by bacteria, molds, and oxidation.

A. Natural Preservatives

Salt and Sugar: These work through Osmosis. They draw moisture out of food and bacterial cells, preventing microbes from growing.

Oil and Vinegar: Create an environment where most bacteria cannot survive due to high acidity or lack of air.

B. Chemical Preservatives

Sodium Benzoate: Commonly used in fruit juices and squashes to inhibit yeast and mold.

Sodium Metabisulphite: Used in jams and pickles to preserve color and prevent bacterial growth.

Nitrates/Nitrites: Used in processed meats to prevent the growth of deadly botulism bacteria and maintain red color.

3. Other Food Chemicals

Antioxidants: These prevent the Rancidity of fats and oils (the bad smell and taste in stale food) by stopping oxygen from reacting with the food. Example: BHT (Butylated Hydroxytoluene).

Artificial Sweeteners: Compounds like Aspartame or Saccharin provide sweetness without the calories, as they are not metabolized by the body like sugar.

Emulsifiers: Help mix substances that usually don't mix, like oil and water (found in mayonnaise and ice cream).