JJ Thomson and Rutherfords Experiments

🔬 Life and Experiments of J.J. Thomson and Ernest Rutherford

🌟 Introduction (परिचय)

Atomic science started to shine in the early 20th century with the contributions of two great physicists: J.J. Thomson and Ernest Rutherford. Their groundbreaking experiments led to models that changed how we understand atoms today. These foundations are now seen in technologies like television displays, mobile processors, and even nuclear medicine.

👨‍🔬 J.J. Thomson: Life & Achievements (जीवन और उपलब्धियाँ)

Sir Joseph John Thomson was born in 1856 in Manchester, UK. He studied engineering but became more inclined towards physics. He later became the director of the famous Cavendish Laboratory. In 1897, he discovered the electron — a fundamental negatively charged particle — for which he received the Nobel Prize in Physics (1906).

📎 Thomson's Cathode Ray Experiment

Thomson used a cathode ray tube — a glass tube from which air is removed — and observed a stream of rays coming from the cathode. These rays:

• Were negatively charged
• Could be deflected by electric and magnetic fields
• Had mass-to-charge ratio (e/m) lower than hydrogen ion

Conclusion: He proposed the "Plum Pudding Model" — atom is a positively charged sphere with electrons embedded in it like raisins in pudding.

⚠️ Drawbacks of Thomson Model

- Could not explain the nuclear structure
- No explanation for alpha scattering or spectral lines

🔬 Rutherford: Life & Achievements (जीवन और प्रयोग)

Ernest Rutherford, born in New Zealand in 1871, is known as the "Father of Nuclear Physics." He studied at Cambridge and worked with J.J. Thomson. In 1909, he and his assistants Geiger and Marsden performed the famous Gold Foil Experiment. He received the Nobel Prize in Chemistry (1908).

💡 Rutherford's Alpha Scattering Experiment

A thin gold foil was bombarded with alpha particles (positively charged helium nuclei). Observations:

• Most passed through undeflected
• Some deflected at small angles
• Few bounced back

Conclusion: Atom has a small, dense, positively charged center called the nucleus. Electrons revolve around it like planets around the sun — called the Rutherford Model.

❌ Limitations of Rutherford's Model

- Failed to explain electron stability in orbits
- Did not justify atomic spectra
- Classical physics predicted electrons would spiral into the nucleus

📱 Applications in Modern Devices (आज के उपकरणों में प्रयोग)

The experiments of Thomson and Rutherford have deep influence in:

• Television Screens: CRTs use cathode rays (based on Thomson’s work)
• Mobile Phones: Atomic models help in transistor-level design
• Nuclear Reactors: Based on understanding of nuclear structure
• X-rays and Imaging: Use behavior of atomic particles

These devices wouldn't have been possible without the pioneering work on atomic structure.

🔗 Internal Blog Links

• 🧠 Play Colligative Memory Game
• 🎯 Limiting Reagent Game
• 📚 Atomic Structure Related Articles

📌 Hindi Summary (हिन्दी में सारांश)

जे.जे. थॉमसन ने इलेक्ट्रॉन की खोज की और 'प्लम पुडिंग मॉडल' दिया।
अर्नेस्ट रदरफोर्ड ने न्यूक्लियस की खोज की और परमाणु का सौरमंडल जैसा मॉडल प्रस्तुत किया।
इनके मॉडल की सीमाएं भी थीं, लेकिन आधुनिक विज्ञान और तकनीक — जैसे मोबाइल, टीवी, रिएक्टर — इन्हीं खोजों पर आधारित हैं।

🔍 Final Thoughts

The evolution from Thomson’s cathode rays to Rutherford’s nucleus marked a revolution in atomic theory. While modern quantum models like Bohr’s or Schrödinger’s have advanced further, the legacy of these early experiments remains strong. Every time you unlock your smartphone, watch TV, or undergo an MRI, you’re witnessing the legacy of these brilliant scientists.

🧪 Rutherford & J.J. Thomson Quiz

📚 Empirical formula

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Atomic Structure Puzzle

Structure of Atom - Drag & Drop Matching

Score: 0 Time: 60s

Electron

Proton

Neutron

Atomic Number

Mass Number

Isotope

Isobar

Bohr Model

Orbital Shape

Quantum Numbers

Negative charge particle

Positive charge particle

Neutral particle

Number of protons

Protons + Neutrons

Same Z, different A

Same A, different Z

Electrons in shells

s, p, d, f

n, l, m, s

📘 Structure of Atom – Quick Definitions

• Electron: A negatively charged particle orbiting the nucleus in specific energy levels.
• Proton: A positively charged particle found inside the nucleus; defines the atomic number.
• Neutron: A neutral particle in the nucleus that adds mass but no charge.
• Atomic Number (Z): Number of protons in the nucleus; determines the element.
• Mass Number (A): Total number of protons and neutrons in an atom.
• Isotopes: Atoms of the same element with different numbers of neutrons.
• Bohr’s Model: Electrons revolve in fixed orbits with quantized energy.
• Quantum Numbers: Numbers that describe the position and energy of an electron.

📘 Quick Definitions — Structure of Atom

Review the key terms used in the game. Tap/click to expand.

Electron

A negatively charged subatomic particle (charge = −1e) that occupies orbitals around the nucleus. Mass ≈ 1/1836 of a proton.

Proton

A positively charged particle (charge = +1e) found in the nucleus. Its count defines the atomic number (Z).

Neutron

An electrically neutral particle in the nucleus. Together with protons it contributes to the atom’s mass number.

Atomic Number (Z)

Number of protons in the nucleus. It uniquely identifies an element (e.g., Z=6 → Carbon).

Mass Number (A)

Total number of nucleons: A = Z + N (protons + neutrons). Determines the isotope of an element.

Isotope

Atoms of the same element (same Z) with different mass numbers (different neutrons). Example: ¹²C and ¹⁴C.

Isobar

Different elements with the same mass number (A) but different atomic numbers (Z). Example: ⁴⁰Ar and ⁴⁰Ca.

Bohr Model

Electrons move in discrete energy shells (n = 1,2,3…). Energy is quantized; electrons jump between shells by absorbing/emitting quanta.

Orbital Shapes (s, p, d, f)

Probability regions for electrons: s (spherical), p (dumbbell), d (clover-like), f (complex). Shape depends on the quantum number ℓ.

Quantum Numbers (n, ℓ, m_ℓ, m_s)

• n: Principal (shell, size/energy)
• ℓ: Azimuthal (subshell/orbital shape)
• m_ℓ: Magnetic (orbital orientation)
• m_s: Spin (+½ or −½)

Valence Electrons & Shell

Electrons in the outermost shell; they control chemical bonding and reactivity (octet/duet rules).

Pauli Exclusion & Hund’s Rule

Pauli: No two electrons in an atom have the same four quantum numbers.
Hund: Electrons occupy degenerate orbitals singly with parallel spins before pairing.

Aufbau Principle

Electrons fill lower-energy orbitals first (approx. order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → …).

Different Bond Angles and Molecular structures

Different Bond Angles and Molecular Structures by MOT

Understanding how Molecular Orbital Theory (MOT) explains bond angles and molecular geometry

Introduction

Molecular Orbital Theory (MOT) is a powerful concept that explains how atomic orbitals combine to form molecular orbitals, determining the stability, shape, and bond angles of molecules. While VSEPR theory predicts geometry based on electron repulsion, MOT gives a deeper quantum-level insight into bond formation, hybridization, and bond order. Let us explore how MOT helps us understand different bond angles and structures of various compounds.

1. Linear Molecular Structure (180° Bond Angle)

Example: BeCl₂, CO₂, N₂, C₂H₂ (Acetylene)

In molecules such as BeCl₂ or CO₂, the atoms align in a straight line. According to MOT, the bonding involves the overlap of sp hybrid orbitals (for Be or C) with p orbitals (for O or Cl). Because the orbitals are oriented directly opposite each other, the resulting molecular geometry is linear with a bond angle of 180°.

In CO₂, the bonding consists of two sigma (σ) and two pi (π) bonds between carbon and oxygen. The central carbon atom is sp hybridized, producing a linear shape that minimizes electron repulsion.

2. Trigonal Planar Structure (120° Bond Angle)

Example: BF₃, AlCl₃, SO₃

In BF₃, boron undergoes sp² hybridization where one 2s and two 2p orbitals mix to form three equivalent sp² orbitals lying in one plane. These orbitals overlap with the p orbitals of fluorine to form σ bonds. The result is a trigonal planar geometry with bond angles of 120°. MOT explains that delocalized π molecular orbitals in such systems further stabilize the planar structure, as seen in SO₃.

3. Tetrahedral Structure (109.5° Bond Angle)

Example: CH₄ (Methane), SiH₄, CCl₄

In methane, carbon undergoes sp³ hybridization forming four equivalent orbitals that point toward the corners of a tetrahedron. Each hydrogen 1s orbital overlaps with one sp³ orbital to form a σ bond. According to MOT, the molecular orbitals result in uniform electron density distribution, giving a perfect 109.5° angle between bonds.

This tetrahedral structure minimizes repulsion between bonding electron pairs, leading to high stability.

4. Trigonal Pyramidal Structure (107° Bond Angle)

Example: NH₃ (Ammonia), PCl₃

Ammonia has three bonding pairs and one lone pair on nitrogen. Due to sp³ hybridization, nitrogen forms four orbitals — one containing a lone pair. The lone pair-bond pair repulsion is greater than bond pair-bond pair repulsion, which slightly compresses the bond angle to 107°.

MOT explains that the nitrogen lone pair occupies a higher-energy nonbonding molecular orbital, altering the effective geometry from tetrahedral to trigonal pyramidal.

5. Bent or V-Shaped Structure (104.5° Bond Angle)

Example: H₂O (Water), H₂S

Water is a classic example of a bent molecule. The oxygen atom has two lone pairs and two bond pairs, forming an sp³ hybridized system. However, because lone pair-lone pair repulsion is maximum, the H–O–H angle is reduced to 104.5°.

MOT interprets this by showing that the molecular orbitals of O–H bonds are formed from overlap between O(2p) and H(1s), with lone pairs occupying higher-energy nonbonding orbitals that distort the ideal tetrahedral angle.

6. Trigonal Bipyramidal Structure (120° and 90° Bond Angles)

Example: PCl₅, PF₅

Phosphorus in PCl₅ uses sp³d hybridization where one 3d, one 3s, and three 3p orbitals combine to form five sp³d hybrid orbitals. Three of them form an equatorial plane (120° apart), and two are placed axially (90° to the plane). Thus, MOT helps visualize two types of bonding interactions — axial and equatorial — giving rise to distinct bond lengths and angles.

7. Octahedral Structure (90° Bond Angle)

Example: SF₆, [Co(NH₃)₆]³⁺

In SF₆, the sulfur atom undergoes sp³d² hybridization, forming six equivalent orbitals oriented towards the corners of an octahedron. All S–F bonds are equivalent with angles of 90°.

MOT further clarifies that bonding occurs through σ-type interactions, and π interactions are negligible due to the high electronegativity of fluorine.

Comparison Table of Bond Angles and Hybridization

Compound	Shape	Hybridization	Bond Angle (°)	Explained by MOT
CO₂	Linear	sp	180	σ and π bonding via sp–p overlap
BF₃	Trigonal Planar	sp²	120	Delocalized π orbitals stabilize planar shape
CH₄	Tetrahedral	sp³	109.5	Equal overlap of sp³ and H 1s orbitals
NH₃	Trigonal Pyramidal	sp³	107	Nonbonding MO distorts ideal angle
H₂O	Bent	sp³	104.5	Lone pairs occupy higher MOs
PCl₅	Trigonal Bipyramidal	sp³d	90,120	Axial and equatorial bond differences
SF₆	Octahedral	sp³d²	90	6 σ-bonds equally distributed

Conclusion

Molecular Orbital Theory gives a quantum-level picture of how atomic orbitals combine to form molecular orbitals, helping us interpret not only bond order and magnetic properties but also bond angles and geometries. While hybridization and VSEPR are quick prediction tools, MOT explains the why — the true reason behind molecular shapes and angle variations.

Hence, by analyzing bonding and antibonding orbitals, MOT provides a deeper understanding of how electron distributions govern molecular structures in chemistry.

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The Periodic Table: A Gateway to Understanding Chemistry

 The Periodic Table: A Gateway to Understanding Chemistry

Exclusive Summary

The periodic table is one of the most important tools in chemistry, serving as a roadmap of all known elements. It organizes elements in a structured manner, highlighting their atomic properties, periodic trends, and chemical reactivity. Beyond being a list, it is a predictive model that guides scientific research, education, and practical applications in fields like medicine, materials science, and energy. This article explores the history, structure, classifications, trends, and modern applications of the periodic table, providing both students and educators with a comprehensive understanding.

1. Historical Development of the Periodic Table

1.1 Early Attempts at Classification

In the early 19th century, scientists like Johann Döbereiner introduced the Law of Triads, grouping elements based on similarities.

John Newlands proposed the Law of Octaves, observing that every eighth element had similar properties.

1.2 Mendeleev’s Breakthrough

Dmitri Mendeleev (1869) arranged elements by increasing atomic mass and left gaps for undiscovered ones.

His predictions about new elements (like gallium and germanium) confirmed the validity of his table.

1.3 Moseley’s Contribution

Henry Moseley (1913) redefined the table based on atomic number instead of mass, leading to the modern periodic law.

2. Structure and Layout of the Periodic Table

2.1 Periods and Groups

The table has 7 periods (rows) and 18 groups (columns).

Groups represent elements with similar valence electron configurations, hence similar chemical properties.

2.2 Blocks of Elements

s-block: Alkali metals and alkaline earth metals.

p-block: Includes halogens, noble gases, and metalloids.

d-block: Transition metals with variable oxidation states.

f-block: Lanthanides and actinides, usually placed below.

2.3 Special Categories

Metals, nonmetals, and metalloids.

Noble gases: stable and inert.

Hydrogen: unique placement, resembling both Group 1 and Group 17.

3. Periodic Trends and Their Importance

3.1 Atomic Radius

Decreases across a period (due to stronger nuclear pull).

Increases down a group (due to additional electron shells).

3.2 Ionization Energy

Increases across a period (harder to remove electrons).

Decreases down a group.

3.3 Electronegativity and Reactivity

Nonmetals have high electronegativity, while metals have low.

Alkali metals are highly reactive; noble gases are least reactive.

3.4 Electron Affinity

Tendency to accept electrons increases across a period, with halogens being highest.

4. Classifications and Chemical Families

4.1 Alkali and Alkaline Earth Metals

Highly reactive, especially with water.

Important in industries like batteries and construction.

4.2 Transition Metals

Conductive, malleable, and form colorful compounds.

Used in catalysis, alloys, and electronics.

4.3 Halogens and Noble Gases

Halogens: toxic and reactive, forming salts.

Noble gases: inert, used in lighting, lasers, and cooling.

4.4 Lanthanides and Actinides

Lanthanides: “rare earths” used in magnets and screens.

Actinides: radioactive, with uranium and plutonium vital for nuclear energy.

5. Applications and Modern Relevance

5.1 Education and Research

A foundational tool in classrooms and scientific laboratories.

5.2 Industry and Technology

Elements like silicon (semiconductors) revolutionize electronics.

Platinum group metals used in catalytic converters.

5.3 Environmental and Medical Uses

Heavy metals raise environmental concerns (toxicity, pollution).

Radioisotopes aid in cancer treatment and imaging.

5.4 Future of the Periodic Table

Discovery of superheavy elements (up to atomic number 118 confirmed).

Research continues on stability and potential new materials.

Conclusion and Actionable Takeaways

The periodic table is not just a chart—it is a scientific map of matter. Its predictive power allows chemists to understand trends, anticipate reactions, and develop new materials. From classrooms to cutting-edge laboratories, the periodic table remains a central tool in exploring the universe of chemistry.

Actionable Takeaways:

Students: Learn periodic trends thoroughly—they are the key to solving chemistry problems.

Teachers: Use real-life examples (e.g., noble gases in neon lights) to connect concepts.

Researchers: Explore periodic predictions for discovering new materials.

General readers: Recognize how elements from the table impact daily life (from batteries to medicines).

Aufbau, Pauli's , Hund's Rule

 Aufbau Principle, Hund’s Rule, and Pauli’s Exclusion Principle: Foundations of Atomic Structure

Introduction

The structure of the atom is one of the central themes in chemistry and physics. Atoms contain electrons, protons, and neutrons, but the way electrons are arranged inside the atom determines nearly all of chemistry — bonding, magnetism, reactivity, and even color of compounds. To understand this arrangement, scientists rely on three essential rules:

1. The Aufbau Principle – tells us the order in which orbitals are filled.

2. Hund’s Rule of Maximum Multiplicity – explains how electrons distribute themselves in orbitals of the same energy.

3. Pauli’s Exclusion Principle – establishes the fundamental restriction on how many electrons can occupy an orbital.

Together, these principles form the foundation of electronic configuration, which is critical in explaining the periodic table, chemical bonding, and spectroscopy. Let’s explore each principle deeply.

The Aufbau Principle

Meaning

The term “Aufbau” comes from the German word aufbauen, which means “to build up.” The Aufbau Principle states that electrons occupy orbitals in the order of increasing energy, filling the lowest energy orbital first before moving to higher energy levels.

This is sometimes called the building-up principle because it describes how electron configurations are constructed.

The Order of Filling

The general filling order is determined by the (n + l) rule, where:

n = principal quantum number (shell number: 1, 2, 3, …)

l = azimuthal quantum number (s = 0, p = 1, d = 2, f = 3)

The orbital with the lower (n + l) value has lower energy and is filled first. If two orbitals have the same (n + l), the orbital with lower n is filled first.

Order of Orbitals by Energy:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p

This sequence explains why, for example, the 4s orbital fills before the 3d orbital.

Example

Hydrogen (Z = 1): 1s¹

Carbon (Z = 6): 1s² 2s² 2p²

Calcium (Z = 20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Exceptions

Some elements like copper (Cu, Z = 29) and chromium (Cr, Z = 24) show exceptions:

Chromium: [Ar] 3d⁵ 4s¹ (instead of [Ar] 3d⁴ 4s²)

Copper: [Ar] 3d¹⁰ 4s¹ (instead of [Ar] 3d⁹ 4s²)

This happens because half-filled (d⁵) and fully-filled (d¹⁰) configurations are particularly stable due to symmetry and exchange energy.

Hund’s Rule of Maximum Multiplicity

Statement

Hund’s Rule says: When electrons occupy orbitals of the same energy (degenerate orbitals), electrons fill them singly with parallel spins before pairing begins.

In simple words:

Every orbital in a subshell (like 2p, 3d, 4f) gets one electron each before any orbital gets two.

The electrons prefer to remain unpaired and with parallel spins as long as possible.

Reason

This happens because electrons are negatively charged and repel each other. By occupying separate orbitals, they reduce electron-electron repulsion, increasing stability. The parallel spins also provide extra stabilization known as exchange energy.

Example

For oxygen (Z = 8): configuration is 1s² 2s² 2p⁴

The 2p orbitals are three: px, py, pz

Hund’s Rule ensures electrons fill like this: ↑↓, ↑, ↑ (not ↑↓, ↑↓, empty)

This gives two unpaired electrons.

For nitrogen (Z = 7): configuration is 1s² 2s² 2p³

The 2p orbitals fill as ↑, ↑, ↑ — each orbital singly filled with parallel spins.

Applications

Explains magnetism: Oxygen has two unpaired electrons → paramagnetic.

Explains why half-filled and fully-filled subshells (like d⁵ and d¹⁰) are particularly stable.

Pauli’s Exclusion Principle

Statement

Proposed by Wolfgang Pauli in 1925, the principle states:

“No two electrons in an atom can have the same set of four quantum numbers.”

The four quantum numbers are:

1. Principal quantum number (n) – shell

2. Azimuthal quantum number (l) – subshell

3. Magnetic quantum number (mₗ) – orbital orientation

4. Spin quantum number (mₛ) – spin (+½ or −½)

Since only two possible spin values exist, each orbital can contain a maximum of two electrons with opposite spins.

Example

In the 1s orbital:

First electron: (n=1, l=0, mₗ=0, mₛ=+½)

Second electron: (n=1, l=0, mₗ=0, mₛ=−½)

Thus, the Pauli Exclusion Principle is the reason why electrons pair up in orbitals with opposite spins.

Relationship Between the Three Principles

1. Aufbau Principle tells us the sequence in which orbitals are filled.

2. Pauli Exclusion Principle restricts each orbital to maximum two electrons with opposite spins.

3. Hund’s Rule explains how electrons distribute themselves within degenerate orbitals before pairing.

Together, they allow chemists to write electron configurations correctly and predict the chemical behavior of elements.

Applications in Modern Chemistry and Physics

1. Periodic Table Structure – These principles explain the structure of the periodic table and periodic trends like ionization energy and atomic radius.

2. Magnetism – Hund’s Rule explains paramagnetism (unpaired electrons) and diamagnetism (all paired).

3. Spectroscopy – Electron transitions between orbitals explain emission and absorption spectra.

4. Chemical Bonding – Molecular orbital theory and valence bond theory both rely on these principles.

5. Quantum Mechanics – Pauli’s principle applies not only to electrons but to all fermions (particles with half-integer spin).

6. Computational Chemistry – Aufbau filling order is programmed into software that predicts molecular structures.

Historical Background

Aufbau Principle developed gradually in the early 20th century from spectroscopic observations and Bohr’s model.

Hund’s Rule was proposed by Friedrich Hund in 1925 while studying molecular spectra.

Pauli’s Exclusion Principle was formulated in 1925 by Wolfgang Pauli and later became a cornerstone of quantum mechanics.

All three emerged around the same time quantum mechanics was being developed, showing their deep connection to the new physics of the 1920s.

Modern Connections

These principles are not just theoretical; they connect to real-world technology:

Lasers depend on controlled electron transitions.

Magnetic resonance imaging (MRI) uses principles of electron spin.

Semiconductors (transistors, computer chips) rely on quantum mechanical behavior of electrons.

Solar cells and LEDs depend on electron orbital transitions explained by these principles.

Examples of Electronic Configurations

1. Hydrogen (Z=1): 1s¹

2. Helium (Z=2): 1s² → Pauli principle ensures two electrons are paired.

3. Carbon (Z=6): 1s² 2s² 2p² → Hund’s rule → two unpaired electrons in 2p.

4. Neon (Z=10): 1s² 2s² 2p⁶ → complete octet, stable noble gas.

5. Iron (Z=26): [Ar] 3d⁶ 4s² → Hund’s Rule → 3d has four unpaired electrons.

These configurations help explain chemical reactivity, oxidation states, and magnetism.

Common Misconceptions

1. Aufbau is absolute – No, exceptions exist (like Cu, Cr).

2. Hund’s Rule is about pairing only – Actually, it’s about maximizing stability through parallel spins.

3. Pauli’s principle is only for chemistry – In reality, it applies to all fermions in physics.

Conclusion

The Aufbau Principle, Hund’s Rule, and Pauli’s Exclusion Principle together provide the fundamental framework for understanding how electrons are arranged inside atoms. These rules explain the periodic table, magnetism, bonding, and even modern devic

es like semiconductors and lasers.

Without them, the entire structure of chemistry would be impossible to rationalize. They remain among the most elegant and powerful ideas in science — simple rules that explain the extraordinary complexity of the material world.

Hybridisation of Orbitals: Concept, Types, and Applications

 🌐 Hybridisation of Orbitals: Concept, Types, and Applications

🔹 Introduction

In chemistry, understanding how atoms bond with each other is essential to explain the structure and properties of molecules. One of the most important concepts that helps us in this explanation is Hybridisation of Orbitals.

The concept was introduced by Linus Pauling in 1931 to explain the equivalent bond formation in molecules like methane (CH₄). Hybridisation is not just a theoretical idea—it is widely used in explaining molecular geometry, bond strength, and bond angles.

In this article, we will explore:

What hybridisation means.

The theory behind orbital mixing.

Different types of hybridisation (sp, sp², sp³, sp³d, sp³d², sp³d³).

Examples and geometries.

Applications in daily life and advanced chemistry

🔹 What is Hybridisation?

Definition:

Hybridisation is the process of mixing atomic orbitals of similar energy levels to form new orbitals, called hybrid orbitals, which are equivalent in shape and energy.

These hybrid orbitals then overlap with orbitals of other atoms to form stable chemical bonds.

Key Features of Hybridisation:

1. Only orbitals of similar energy mix together.

2. The number of hybrid orbitals formed = number of atomic orbitals mixed.

3. Hybrid orbitals have a definite geometry, which determines the shape of molecules.

4. Hybridisation explains why equivalent bonds (like four C–H bonds in CH₄) exist.

🔹 Conditions for Hybridisation

Orbitals should have comparable energy (e.g., 2s and 2p).

Orbitals must belong to the same atom.

Hybridisation occurs during bond formation, not in isolated atoms.

Both half-filled and fully filled orbitals may participate.

🔹 Types of Hybridisation

Hybridisation depends on the number of orbitals involved. Below are the main types with geometries and examples:

1. sp Hybridisation

Mixing: 1 s + 1 p orbital → 2 sp orbitals.

Geometry: Linear (180° bond angle).

Examples: BeCl₂, CO₂, C₂H₂ (acetylene).

👉 Explanation: In BeCl₂, the central atom beryllium has two sp hybrid orbitals oriented linearly, forming σ bonds with chlorine atoms.

2. sp² Hybridisation

Mixing: 1 s + 2 p orbitals → 3 sp² orbitals.

Geometry: Trigonal Planar (120° bond angle).

Examples: BF₃, C₂H₄ (ethene), SO₃.

👉 Explanation: In BF₃, boron forms three equivalent sp² orbitals arranged in a plane at 120°, bonding with fluorine.

3. sp³ Hybridisation

Mixing: 1 s + 3 p orbitals → 4 sp³ orbitals.

Geometry: Tetrahedral (109.5° bond angle).

Examples: CH₄ (methane), NH₃ (trigonal pyramidal), H₂O (bent shape).

👉 Explanation: In methane, carbon undergoes sp³ hybridisation, giving rise to four identical tetrahedral bonds with hydrogen.

4. sp³d Hybridisation

Mixing: 1 s + 3 p + 1 d orbital → 5 sp³d orbitals.

Geometry: Trigonal Bipyramidal.

Bond Angles: 120° (equatorial), 90° (axial).

Examples: PCl₅, SF₄.

5. sp³d² Hybridisation

Mixing: 1 s + 3 p + 2 d orbitals → 6 sp³d² orbitals.

Geometry: Octahedral (90° bond angle).

Examples: SF₆, [Co(NH₃)₆]³⁺.

6. sp³d³ Hybridisation

Mixing: 1 s + 3 p + 3 d orbitals → 7 sp³d³ orbitals.

Geometry: Pentagonal Bipyramidal.

Examples: IF₇.

🔹 Table of Hybridisation, Geometry, and Examples

Type of Hybridisation Orbitals Involved Geometry Bond Angle Examples

sp 1s + 1p Linear 180° BeCl₂, CO₂

sp² 1s + 2p Trigonal Planar 120° BF₃, C₂H₄

sp³ 1s + 3p Tetrahedral 109.5° CH₄, NH₃, H₂O

sp³d 1s + 3p + 1d Trigonal Bipyramidal 120°/90° PCl₅, SF₄

sp³d² 1s + 3p + 2d Octahedral 90° SF₆, XeF₄

sp³d³ 1s + 3p + 3d Pentagonal Bipyramidal 72°/90° IF₇

🔹 Difference Between Hybrid Orbitals and Atomic Orbitals

Property Atomic Orbitals Hybrid Orbitals

Shape s – spherical, p – dumbbell Same shape in a set

Energy Different for s, p, d Equal energy (degenerate)

Orientation Random Definite geometry

Formation Natural Formed during bond making

🔹 Applications of Hybridisation

1. Explains Molecular Geometry – Why methane is tetrahedral, ethene is planar, and acetylene is linear.

2. Bond Strength & Stability – Hybrid orbitals give stronger σ bonds than pure orbitals.

3. Valence Bond Theory (VBT) – Hybridisation supports VBT in explaining chemical bonding.

4. Transition Metal Complexes – Explains geometry in coordination compounds (octahedral, square planar).

5. Organic Chemistry – Essential in explaining structures of alkanes, alkenes, alkynes, and aromatic compounds.

6. Material Science – Helps understand bonding in diamond (sp³) and graphite (sp²).

🔹 Hybridisation in Real Life Examples

Diamond → Each carbon is sp³ hybridised (tetrahedral, hardest structure).

Graphite → Each carbon is sp² hybridised, with delocalised electrons (good conductor).

Acetylene (C₂H₂) → sp hybridisation, explaining linear shape and triple bonds.

Ammonia (NH₃) → sp³ hybridisation with lone pair, giving trigonal pyramidal shape.

🔹 Hybridisation and Bond Character

The percentage of s-character in hybrid orbitals influences bond properties:

sp (50% s + 50% p) → Shortest and strongest bonds.

sp² (33% s + 67% p) → Intermediate bond strength.

sp³ (25% s + 75% p) → Longest and weakest bonds among these three.

👉 This explains why C≡C bonds are stronger than C=C or C–C bonds.

🔹 Common Misconceptions in Hybridisation

1. Hybridisation is not a physical mixing of orbitals; it’s a mathematical concept.

2. Not all molecules follow hybridisation rules perfectly (e.g., some d-orbital participation is debated).

3. Lone pairs also occupy hybrid orbitals, influencing shape.

🔹 Conclusion

Hybridisation is a cornerstone in understanding chemical bonding and molecular structures. It explains why bonds are equivalent, why molecules adopt specific geometries, and how bond strength varies. From the hardness of diamond to the conductivity of graphite, hybridisation provides a simple yet powerful explanation of chemical behavior.

As chemistry advances into nanotechnology, materials science, and drug design, hybridisation remains a fundamental concept bridging atomic theory and molecular properties.