Mixture

Definition

A mixture is a combination of two or more substances that are physically mixed but not chemically combined.

• Unlike compounds, mixtures do not participate in the process of  chemical bonding.
• Substances held together by physical forces.
• Each item retains its properties in the mixture.
• Mixtures can exist in various forms: solutions, suspensions, or colloids.

Examples of Mixtures:

• Crude oil: A mixture of organic compounds (mainly hydrocarbons).
• Seawater: A mixture of various salts and water.
• Air: A mixture of gases like oxygen, carbon dioxide, nitrogen, and argon.
• Ink: A mixture of coloured dyes.
• Gunpowder: A mixture of sulphur, potassium nitrate, and carbon.

Type of Mixture

Homogeneous Mixtures: Homogeneous mixtures have a uniform composition. In these mixtures, the individual components are evenly distributed at the molecular level and cannot be visually distinguished.

• Example: saltwater (a solution of salt dissolved in water), air (a mixture of gases), and sugar dissolved in coffee.

Heterogeneous Mixtures: Heterogeneous mixtures have non-uniform compositions, with distinct regions containing different concentrations of the components.

• The individual substances in a heterogeneous mixture are typically visible to the naked eye and can be physically separated through methods such as filtration, decantation, or chromatography.
• Examples: sand and water, a salad containing various vegetables, and a mixture of oil and water.

In addition to being characterized as heterogeneous or homogeneous, mixture can also be described based on the particle size of the components:

• Solution: Solution contains very small particle sizes less than 1 nanometer in diameter. It is physically stable.
• Colloid: Colloid is a mixture of substances from 1 nanometer to 1 micrometer in diameter, dispersed in a solvent. 
• Suspension: Are heterogeneous mixtures consisting particles that are visible to the naked eye. Substances will settle over time.

Compound

Definition

A compound is a substance that is composed of two or more different elements chemically bonded together.

• The elements of a compound are combined in specific ratio.
• It have unique properties distinct from the elements they are composed of.
• Compounds are represented with its chemical formula. Chemical formula indicates the elements present in the compound and the relative number of atoms or ions of each elements.

Here’s an explanation with examples:

• Water (H2O): One of the most common compounds, water is composed of two hydrogen atoms bonded to one oxygen atom. 
• Sodium Chloride (NaCl): Commonly known as table salt, sodium chloride is composed of one sodium atom bonded to one chlorine atom

Types of Compounds

Compounds can be classified into two types:

• Molecular Compounds: Atoms bind through covalent bonds. Examples include: Carbon dioxide (CO2) Methane (CH4)
• Ionic Compound: Held together by ionic bonds. Example: Table salt (NaCl): Comprising one sodium atom and one chlorine atom.

Element

Definition

An chemical element is a substance made up of identical atoms (atoms that have the same number of protons in their nucleus). For example, a hydrogen atom has 1 proton, while a carbon atom has 6 protons.

• The number of electron and neutron in an element can be different.
• It cannot be broken down into another substance.
• Each element is represented by a unique chemical symbol, such as “H” for hydrogen and “O” for oxygen.
• Change in the number of electrons in an atom of an element produces ions.
• Change in the number of neutrons produces isotopes.
• The term “element” was first introduced by the English scientist Robert Boyle.

Atomic Structure: Atoms of an element have a specific structure consisting of a nucleus composed of protons (positively charged particles) and neutrons (particles with no charge), surrounded by a cloud of electrons (negatively charged particles) in various energy levels or orbitals.

Chemical Symbols: Each element is represented by a chemical symbol, typically one or two letters derived from its name. For example, “H” represents hydrogen.

Periodic Table: Elements are organized systematically in the periodic table based on their atomic number, electron configuration, and chemical properties.

Geochemical Classification of Elements

Elements can be classified based on their properties and affinity for other elements. Goldschmidt’s classify the chemical elements into four groups-

• Siderophile elements (iron loving): Elements that concentrate in metallic iron such as Ni, Co, Os, Ir
• Chalcophile elements: Elements that concentrate in sulfides, are chalcophile elements such as Se, As, Zn, Cd
• Lithophile elements (Rock Loving): Elements that concentrate in silicate phases are grouped in this category, such as Rb, Sr, Ba, Nb, Ta, Th, U, REE
• Atmophile elements: These are naturally occurring gaseous elements such as N and rare gases

Types of Elements

Elements can be categorised into several groups based on their properties:

• Metals: Elements that are typically shiny, malleable, ductile, and good conductors of heat and electricity (e.g., iron, copper, gold).
• Non-metals: Elements that are generally brittle, poor conductors of heat and electricity, and often exist in gaseous form (e.g., oxygen, carbon, sulphur).
• Metalloids: Elements that possess properties intermediate between metals and nonmetals (e.g., silicon, germanium, arsenic).
• Noble Gases: Elements that are inert, colourless, odourless gases with very low reactivity (e.g., helium, neon, argon).
• Transition Metals: Elements found in the d-block of the periodic table, characterised by their variable oxidation states and ability to form coloured compounds (e.g., iron, copper, zinc).
• Rare Earth Elements: A group of elements found in the lanthanide and actinide series of the periodic table, known for their unique magnetic, electrical, and optical properties.

Periodic Classification of Element

The periodic classification of elements refers to the systematic organisation of chemical elements according to their atomic number, electron configuration, and chemical properties.

In 1864, J.A.R. Newlands proposed classifying elements based on increasing atomic weights. The modern periodic table is based on the Periodic Law proposed by Dmitri Mendeleev. Mendeleev organized elements in ascending order of their atomic masses and classified them according to their atomic masses.

• The periodic table arranges elements in periods (rows) and groups (columns).
• Elements within the same group share similar chemical properties.

Structure of Modern Periodic Table

The modern periodic table is structured in a way that organises elements based on their atomic number, electron configuration, and chemical properties. It consists of:

• Periods (horizontal rows)
• Groups (vertical columns)

Each square in periodic table represents an element, displaying its chemical symbol and atomic number.

Period (Rows)

There are 7 periods in the periodic table. Elements within the same period have the same number of electron shells. For example, the first period contains only hydrogen and helium.

Groups (Columns)

There are 18 groups in the periodic table. Elements within the same group share similar chemical properties due to the similarities in their outer electron configuration. Some notable groups include:

• Group 1 (Alkali Metals): Reactive metals that form strong alkalis with water.
• Group 2 (Alkaline Earth Metals): Metals that also form alkalis but are weaker than Group 1 elements.
• Group 17 (Halogens): Elements that form salts.
• Group 18 (Noble Gases): Inert gases under normal conditions.

s,p,d and f block elements

Depending on which subshell the final electron enters, the elements in the periodic table are divided into four major groups. These blocks are the:

• s-block
• p-block
• d-block
• f-block.

s-block Elements

These elements have their outermost electron(s) in the s orbital. They are found in groups 1 and 2 of the periodic table.

Characteristics

• Typically, they are (except hydrogen) highly reactive metals .
• They  lose electrons to form positively charged ions (cations).
• They have low electronegativity and low ionisation energies. When burned, they produces a characteristic colour flame.
• Most s-block elements are solids at room temperature (except caesium, which is liquid around 35°C).
• Examples: Lithium (Li), Sodium (Na), Potassium (K), etc.

p-block Elements

These elements have their outermost electron(s) in the p orbital. They are found in groups 13 to 18 of the periodic table.

Characteristics

• They exhibit a wide range of properties, from metals to metalloids and nonmetals.
• Various States: They exist as solids, liquids, or gases at room temperature (e.g bromine is a liquid).
• Their electronegativity varies widely, with some being highly electronegative (e.g., oxygen) and others being less so (e.g., boron).
• They often form covalent bonds by sharing electrons.
• They do not impart any characteristic colour to flames.

d-block Elements (also known as transition metals)

These elements have their outermost electron(s) in the d orbital. They are found in groups 3 to 12 of the periodic table.

Characteristics

• Most d-block elements are metals.
• Many of their compounds are coloured.
• Many transition metals are good conductors of heat and electricity.
• They often act as catalysts in chemical reactions.
• Examples: Iron (Fe), Copper (Cu), Zinc (Zn), etc.

f-block Elements (also known as inner transition metals)

These elements have their outermost electron(s) in the f orbital. They are placed at the bottom of the periodic table, separated to keep it compact. There are two series of f-block elements:

• The lanthanides (elements 57-71):  Lanthanide (58Ce-71Lu) closely resemble lanthanum, have similar chemical properties, and are widely used in lasers.
• The actinides (elements 89-103): Actinides (90Th-103Lr) are radioactive, highly electropositive, and exhibit various oxidation states.

Characteristics

• Many are radioactive.
• They are typically soft metals. Some have important uses in nuclear reactors, medicine (e.g., radioisotopes for imaging), and other specialized applications.
• Examples: Cerium (Ce), Uranium (U), Neptunium (Np), etc.

Fact Sheet

• In periodic table, elements are arranged according to their increasing atomic number.
• Horizontal rows on the periodic table is termed as periods.
• The number of horizontal rows in periodic table are seven and have been numbered from 1 to 7.
• The 1st period is the shortest period of all and contains only 2 elements, H and He.
• The 2nd and 3rd periods are termed as short periods and contain 8 elements each.
• 4th and 5th periods are long periods and contain 18 elements each.
• 6th and 7th period is very long period containing 32 elements.
• Vertical columns are called groups. There are 18 groups in the periodic table.
• Group 1 on extreme left position contains alkali metals (Li, Na, K, Rb, Cs and Fr).
• Group 18 on extreme right side position contains noble gases (He, Ne, Ar, Kr, Xe and Rn).
• In the middle of periodic table, we have semi-metals or metalloid because they exhibit some properties of metals and non-metals.

Molecule

Definition

A molecule is defined as the fundamental units of chemical compounds formed when two or more atoms chemically bond together.

• It forms the smallest  unit of the substance.
• The atoms of the molecule can be of the same element like oxygen (O2) and hydrogen (H2) (diatomic molecule), or they can be different elements, such as water (H2O) or carbon dioxide (CO2).
• Molecule retains the physical and chemical properties of that substance.
• Molecules are held together by covalent bonds or ionic bonds.

Chemical Bond

A chemical bond is a force that holds two or more atoms together in a molecule. The interactions between atoms is due to the electrostatic forces (as in ionic bonds) or the sharing of electrons (as in covalent bonds). These bonds determine the stability, properties, and behaviour of substances.

Type of Chemical Bond

There are three main types of chemical bonds:

• Covalent Bond
• Ionic Bond
• Hydrogen Bond

Covalent Bond

Covalent bonds also referred to as molecular bonds involve the equal sharing of electrons between atoms. It is common between two nonmetals. There are two main types of covalent bonds:

• Single Covalent Bonds: In a single covalent bond, atoms share one pair of electrons. Examples include hydrogen gas (H2) and chlorine gas (Cl2).
• Multiple Covalent Bonds: In multiple covalent bonds, atoms share two or more pairs of electrons. Examples include oxygen gas (O2) and nitrogen gas (N2).

Ionic Bond

Ionic bonds form between atoms when one atom transfers one or more electrons to another atom. This transfer results in the formation of positively charged ions (cations) and negatively charged ions (anions). The oppositely charged ions are attracted to each other by electrostatic forces, forming an ionic bond. Ionic bonds are typically formed between a metal atom (which tends to lose electrons) and a nonmetal atom (which tends to gain electrons).

Examples of compounds with ionic bonds are sodium chloride (NaCl), potassium iodide (KI), and magnesium oxide (MgO).

Hydrogen Bond

It is not ionic or covalent; it’s a dipole-dipole attraction that forms between a hydrogen atom and an electronegative a tom. Important in biological molecules like DNA and proteins.

Isomer

Isomers are those molecules which have same molecular formula but different arrangements of atoms. Isomers can exhibit different physical and chemical properties due to their distinct structures.

There are several types of isomerism, including structural isomerism, geometric isomerism, and optical isomerism.

Electron Configuration

Definition

Electron configuration is the arrangement of negatively charged sub-atomic electron particle in orbital shell and subshell around an atomic nucleus. Electron configuration follows specific rules based on the energy levels and sublevels of electrons. 

Example: Based on Aufbau Principle electron configuration of Helium and Sodium are:

• Helium has an atomic number of 2. Its electron configuration can be represented as 1s².
• Sodium has an atomic number of 11. Its electron configuration can be 1s² 2s² 2p⁶ 3s¹.

Electron Shells

The electron shell is a grouping of electrons in an atom according to their energy level, which incircles the atomic nucleus. The electron shell of an atom can hold 2n² electrons, where n is the energy level. For example, the first shell (closest to the nucleus) can hold 2x(1)²or 2 electrons, the second shell can accommodate 2x(2)² or 8 electrons.

• Electrons that are in the first energy level are closest to the nucleus and will have the lowest energy on the other hand electrons further away from the nucleus will have higher energy.

Subshell

Within the shells, electrons are further grouped into subshells. There are four different types of sub-shell named as s, p, d, and f (in increasing energy order).

• The first shell has only one subshell s subshell
• The second shell has s and  p subshell
• The third shell has s, p, and d subshells
• The fourth shell has s, p, d and f subshells

Atomic Orbital

Within each subshell, electrons are grouped into orbitals. The area surrounding the nucleus where electrons in their respective sub-shell are estimated to be present is known as the orbitals.

Key principles governing Electronic Configuration

The arrangement of electrons within electron shell is based on certain rules and principles. These principles/Rules are:

• Aufbau principle
• Hund’s rule
• Pauli exclusion principle

The above three principles explain the filling order of atomic orbitals and the distribution of electrons among them.

Aufbau Principle

The German word “Aufbau” means “building up.” The Aufbau principle defines a set of rules. These rules are:

Rule 1: Electrons 1st occupy the lowest energy orbitals. It means electrons of an atom occupy their position in atomic orbitals in a specific order, starting from the lowest energy level to higher energy levels. 

Rule 2: Orbitals will be filled in of Increasing energy. It means the order in which orbitals are filled follows a specific sequence such as 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, and so on, is based on their energy levels.

Rule 2: Each orbital can hold only two electrons (must be of opposite spin).

Rule 3: Two or more orbitals with the same energy are each half-filled by one electron before any one orbital is completely filled by addition of the second electron.

Hund’s Rule

This rule states that:

• Electrons occupy orbitals of the same energy singly before they start pairing up.
• All of the electrons in singly occupied orbitals have the same spin.

This rule helps in determining the arrangement of electrons within subshells.

Pauli Exclusion Principle

Proposed by Wolfgang Pauli. This Principle states that, Each electron can be described by a unique set of four quantum numbers; it means no two electrons in an atom can have the same set of quantum numbers.

Valence Electrons

Valence electrons are the electrons in the outermost shell (highest occupied energy level) of the atom. The shell in which valence electron are present is termed as valence shell.

• They participate in the formation of chemical bond.
• They are held loosely
• The number of valence electron determine the properties of the atom and the way atom will bond chemically (reactivity).
• More reactive element have fewer electron in the valence cell.
• Stable atom has 8 electron in their valence cell.

Example: Boron

• Boron has 5 electrons and the electron configuration of boron is 1s² 2s² 2p¹, indicating that it has 2 electrons in the 1s orbital and 2 electrons in the 2s orbital, leaving only 1 electron in the 2p orbital, which is the outermost electron shell. Therefore, boron has 3 valence electrons. These valence electrons are crucial in defining boron’s chemical behavior.
• To achieve a stable electron configuration Boron tends to lose these 3 valence electrons.

Ionisation

Ionization is the process by which an atom or a molecule gains or loses one or more electrons, resulting in the formation of an ion with a net positive charge (cation) or negative charge (anion).

• Ion Formation: Anion is formed when an atom or molecule gains one or more electrons and becomes negatively charged. Conversely, when an atom or molecule loses one or more electrons, it becomes positively charged and termed as cation.

Atomic Number, Mass Number, Isotope, Isobar, Isotone, Atomic Mass

Atomic Number (Z)

The atomic number of an element is the number of protons in the nucleus of its atoms. It determines the identity of the element and its position in the periodic table.  For example: Helium The atomic number of helium is 2 because the number of protons in helium nucleus is 2.

Mass Number (A)

The mass number of an atom is the sum of its protons and neutrons.  It represents the total number of nucleons (protons and neutrons) in the nucleus of an atom.

Mass Number (A) = Number of Protons (p) + Number of Neutron (n)

For example: The mass number of helium-4 (most common isotope of helium) refers to the total number of nucleons (protons + neutrons) present in the nucleus of a helium atom. Helium has an atomic number of 2, indicating it has two protons in its nucleus. As well as helium has two neutrons in its nucleus.

Thus the atomic mass number of helium 4 will be 4 (2 protons + 2 neutrons = 4 nucleons).

However, there are other isotopes of helium, such as helium-3, which have different numbers of neutrons and, therefore, different mass numbers.

Isotopes

Isotopes (iso means equal and topos means place) are atoms of the same element that have the same number of protons but different numbers of neutrons.

• Isotopes have the same atomic number but different mass number.
• Isotopes of an element have similar chemical properties but different physical properties due to differences in mass.
• Almost every chemical substance has an isotope. Approximately 275 isotopes of approximately 81 stable chemical components have been confirmed.
• Isotopes are of two types: Stable isotopes and Radioactive isotopes

Example: Isotopes of helium:  Helium has two stable isotopes: helium-3 and helium-4.

• Helium-3: Helium-3 has two protons like all helium atoms, but it contains only one neutron in its nucleus. It is less common on Earth than helium-4.
• Helium-4: Helium-4, also known as alpha particles, is the most common and stable isotope of helium. It has two protons and two neutrons in its nucleus. Commonly it is found in natural gas deposits.

Isobars

Isobar (iso means equal, and baros means weight) are the atoms of different chemical element have the same mass number but different atomic numbers. The term “isobars” was coined by Alfred Walter Stewart in 1918.

Example: Argon, potassium, and calcium

• Argon (Ar): Atomic number of Argon is 18 and Mass Number is 40. It has 18 protons and  22 neutrons in its nucleus. Chemically it is a noble gas.
• Potassium (K): Atomic number of Potassium is 19 and Mass Number is 40. It has 19 protons and 21 neutrons in its nucleus. Chemically it is an alkali metal.
• Calcium (Ca): Atomic number of calcium is 20 and Mass Number is 40. It has 20 protons and 20 neutrons in its nucleus. Chemically it is an alkaline earth meta.

Isotones

Isotones are two or more different chemical element that have the same number of neutrons but different number of protons. The value of (A – Z) for these elements is the same even when the atomic number Z and the mass number A are different.

Example: Carbon, nitrogen, and oxygen

• Carbon (C): Carbon has 6 protons and 8 Neutron in its nucleus
  • Value of A-Z = 14-6 = 8
• Nitrogen (N): Nitrogen has 7 protons and 8 Neutron in its nucleus
  • Value of A-Z = 15-7 = 8
• Oxygen (O): Carbon has 8 protons and 8 Neutron in its nucleus
  • Value of A-Z = 16-8 =8 

Isoelectronic

Isoelectronic are atoms, ions, or molecules that have the same number of electrons. Despite potentially having different numbers of protons, they have identical electron configurations and thus similar chemical properties.

Example: Sodium (Na⁺ ) and fluorine (F⁻) 

Sodium (Na): Atomic number of Sodium is 11.

• Electron configuration of sodium is 1s² 2s² 2p⁶ 3s¹.
• Sodium has 11 electrons and the outermost electron is in the 3s orbital.
• To form a stable Sodium Na⁺ ion sodium loses its outermost electron resulting in an electron configuration of 1s² 2s² 2p⁶

 Fluorine (F): Atomic number of Fluorine is 9.

• Electron configuration of fluorine is 1s² 2s² 2p⁵.
• Fluorine has 9 electrons and the outermost electron is in the 2p orbital.
• Fluorine gains one electron to form a stable F⁻ ion resulting in an electron configuration of  1s² 2s² 2p⁶

Atomic Mass

Atomic mass is the total mass of the subatomic particles (proton, neutron and electron) in an atom and is measured in atomic mass units (amu).

Atomic mass of an element = Mass of protons + Mass of electrons + Mass of neutrons

• The calculated value for atomic mass is a whole number.
• Atomic Mass Unit: One atomic mass unit (u) is defined as one twelfth of the mass of a carbon-12 atom. (The International Union of Chemists adopted Carbon-12 as the standard for comparing the atomic and molecular weights of elements and compounds in 1961.)

Example: Carbon: the atomic mass of carbon is 12 u.

Atom and Atomic Theory

An atom is the fundamental unit of matter, consisting three sub-atomic particles proton, neutron and electron. Two of the sub-atomic particle (proton and electron) of the atom is electrically charge and one sub-atomic particle (neutron) has no charge.

Atom consist a nucleus made up of two sub-atomic particle protons and neutrons, which form a dense and positively charged core, while negatively charged third sub-atomic particle electrons form an electron cloud around the nucleus.

• The chemical properties of an atom are determined by the negatively charged electrons masking the nucleus.
• The physical properties of an atom are determined by the number of protons and neutrons in the nucleus of  an atom.

Electrical Charge on three sub-atomic particles: 

• Protons carry a positive (+1) electrical charge
• Neutrons have no electrical charge (0)
• Electrons carry a negative (-1) electrical charge

Atomic Nucleus/Nuclei

Atomic nuclei is a dense core of an atom comprised of electrically positive charged proton and electrically neutral neutron held together with a strong force. The nucleus accounts for less than 0.01% of the atom’s volume but more than 99.9% of its mass. 

The number of protons in nucleus of an atom is the atomic number, and the number of protons plus neutrons is the atomic mass.

In 1911, Ernest Rutherford discovered that at the core of every atom is a nucleus.Ernest Rutherford, the discovery of the proton is credited to Ernest Rutherford. Rutherford is also credited with the discovery of the atomic nucleus.

Sir James Chadwick, Sir James Chadwick (1891–1974) was a British physicist who is best known for his discovery of the neutron. For his revolutionary discovery of the neutron in 1932, he was granted the 1935 Nobel Prize in Physics.

J.J. Thomson, is a British physicist discovered electron in 1897. 

George Johnstone Stoney, an Irish physicist, is credited with introducing the term “electron”.

Atomic Models

Atomic models are conceptual frameworks used to understand the structure and behavior of atoms. From the early plum pudding model to modern quantum mechanical models, these representations have evolved, revealing the complex nature of atomic particles and their interactions and shaping our understanding of the universe. Here are some of these models:

John Dalton’s Atomic Theory (1803): Dalton suggested that:

• All matter is made up of atoms.
• All matter is composed of indivisible and indestructible atoms.
• Atoms of the same element are identical, but atoms of different elements vary in size and mass.
• Chemical reactions involve rearrangement of atoms to form products.

J.J. Thomson’s Model (1897): Thomson discovered the electron and proposed the “plum pudding” model.

• He proposed that atoms contain a positively charged sphere with embedded electrons.
• Electrons are distributed throughout the atom.

Ernest Rutherford’s Model (1911): Rutherford conducted the famous gold foil experiment:

• Discovered the nucleus.
• He proposed that most of an atom’s mass is concentrated in the nucleus. The nucleus is positively charged and electrons orbit around the nucleus.
• Protons and neutrons are collectively called nucleons.

Niels Bohr’s Model (1913): Bohr introduced the concept of quantized energy levels. He proposed that electrons move in specific orbits or energy shells and can jump between energy levels by absorbing or emitting energy. Energy is absorbed when an electron jumps from a lower orbit to a higher one and energy is emitted when an electron falls from a higher orbit to a lower orbit.

Quantum Model or Quantum-Wave Model or Quantum Mechanical Model (1926): Proposed by Erwin Schrodinger, in 1926.This model blends physics and mathematics. The Quantum Mechanical Model describes the behaviour of electrons in atoms as both particles and waves.

Definition

Chemistry is a branch of natural science which deals with the scientific study of the properties, composition, structure and behaviour of matter. Chemistry explains:

• How the substances interact with each other
• How substances behave under different conditions
• The changes they undergo and principles governing these changes

Antoine-Laurent Lavoisier (1743–1794) was a French chemist widely regarded as the “Father of Modern Chemistry”. Lavoisier transforms chemistry from a qualitative science to a quantitative one. He is most renowned for his discovery of the role of oxygen in combustion.

Branches of Chemistry

Due to the vast scope of chemistry it can be categorised into several branches. Traditionally, five major branches of chemistry are:

• Organic Chemistry
• Inorganic Chemistry
• Analytical Chemistry
• Physical Chemistry
• Biochemistry

Apart from these five branches, chemistry  can be divided into many sub branches that deals with cross-disciplinary study such as, environmental chemistry, forensic chemistry, etc.

Organic Chemistry

Definition: Organic chemistry is the branch of chemistry that deals with the study of carbon-containing compounds, which are often associated with living organisms. It deals with the study of structure, properties, reactions, synthesis, and applications of organic compounds (consist primarily of carbon and hydrogen).

Scope:

• Study the chemistry of life, including reactions occurring in living organisms.
• Study organic reaction, structure and property of organic molecules, polymers, hydrocarbons, drugs. 

Sub-branches of organic chemistry include:  Medical Chemistry, Physical Organic Chemistry, Organometallic Chemistry, Stereochemistry, and Polymer Chemistry.

Inorganic Chemistry

Definition: Inorganic chemistry is the branch of chemistry that deals with the study of inorganic compounds, which typically do not contain carbon-hydrogen (C-H) bonds.

Scope: 

• Inorganic compounds encompass a diverse type substances, including minerals, metals, salts, organometallic compounds, cluster compounds, and solid-state compounds.
• It explores the properties, structures, synthesis, reactions, and applications of these compounds.

Sub-branches of inorganic chemistry include: Nuclear Chemistry, Geochemistry, Bioinorganic Chemistry, Solid-State Chemistry, and Organometallic Chemistry.

Analytical Chemistry

Definition: Analytical chemistry is a branch of chemistry that focuses on the qualitative (measuring the amount or concentration of substances in a sample) and quantitative (determining the identity or presence of specific substances or components in a sample) determination of chemical component of the substances.

Scope: 

• Includes Quantitative and qualitative analysis of chemical substances
• Includes  separations, extractions, distillation, spectrometry and spectroscopy, chromatography, and electrophoresis of chemical substances. 
• Discovery and development of standards, chemical methods, and instrumental methods.

Sub-branches of analytical chemistry include: Environmental Chemistry, Forensic Chemistry, and Bioanalytical Chemistry.

Physical Chemistry

Definition: Physical chemistry applies the principle of physics to study the chemistry. It deals with the study of the physical properties and behaviour of matter at the molecular and atomic levels for example, thermodynamics and quantum mechanics. 

Physical chemists typically study :

• The rate of a chemical reaction
• The interaction of molecules with radiation
• The calculation of structures and properties

Scope: Utilises concepts from thermodynamics and quantum mechanics.

• Thermodynamics, deals with the study of energy changes in chemical and physical processes. 
• Quantum chemistry, applies quantum mechanics to understand the behaviour of atoms and molecules. 

Sub-branches of physical chemistry includes: Quantum Chemistry, Chemical Kinetics, Surface Chemistry.

Biochemistry

Definition: Biochemistry is a branch of science that combines principles and techniques from both biology and chemistry to study the chemical processes and substances that occur within living organisms. It spans molecular biology, genetics, biochemical pharmacology, clinical biochemistry, and agricultural biochemistry.

Scope: Studies key molecules such as proteins, nucleic acids, carbohydrates, lipids, drugs, and neurotransmitters. Closely related to molecular biology, cell biology, and genetics.

Sub-branches of biochemistry include: genetics, molecular biology, clinical biochemistry, pharmacology, toxicology, and agricultural biochemistry.

States of matter are various forms of matter that form under particular conditions, such as pressure, temperature, etc. Based on the molecular arrangements and energy levels, each state is characterized by its own unique physical properties and behaviors. There are three States of Matter:

• Solid (particles are closely packed together in a regular geometric arrangement)
• liquid (particles are loosely arranged)
• Gas (particles move freely because they are widely spaced)

Apart from above three states there are two more forms of states of matter exist, they are:

• Plasma (form at extremely high temperature)
• Bose-Einstein condensates (form at ultra-low temperature)