Bohr Model of Atom
- Developed by Niels Bohr in 1913
- Based on Rutherford’s model of atom
- Explains stability of atoms
Consequence of Rutherford Model
- According to Rutherford’s model, electrons should continuously emit energy as they orbit around the nucleus
- This would cause electrons to spiral into the nucleus
- But, observations did not match this prediction
Instability of Rutherford Model
- Instability arises from continuous energy loss
- Electrons would emit energy in the form of electromagnetic radiation (radiation of energy)
- As a result, electrons would spiral inward, eventually causing the atom to collapse
Inspired by Planetary Motion
- Bohr applied Planck’s quantum theory to modify Rutherford’s model
- Proposed a new model where electrons occupy specific energy levels around the nucleus
- Electrons can only exist in these defined energy levels, called “shells”
Quantization of Energy Levels
- Electrons can only occupy specific energy levels or shells
- The energy of an electron is quantized, meaning it can only have certain discrete values
- Energy increases as we move further away from the nucleus
Absorption and Emission of Energy
- Electrons can absorb or emit energy when transitioning between energy levels
- Absorption: Electrons absorb energy and move to higher energy levels
- Emission: Electrons release energy and move to lower energy levels
Energy-Level Diagram
- Energy levels are represented by whole numbers (n = 1, 2, 3, …)
- The closest energy level to the nucleus has the lowest energy (n=1)
- As n increases, energy levels move further away from the nucleus and have higher energy
Ground State and Excited State
- Ground State: Electrons occupy the lowest energy level (n=1)
- Excited State: Electrons occupy higher energy levels (n>1) after absorbing energy
Emission of Light
- When electrons return to lower energy levels (from higher energy levels), they emit energy
- This emitted energy corresponds to different wavelengths of light
- The emission spectrum consists of discrete lines specific to the element
Equation for Energy Levels
- Bohr proposed an equation to calculate the energy levels of hydrogen atom
- E = -13.6 eV / n^2
- E is the energy of a specific energy level
- n is the principal quantum number representing the energy level
Consequence of Rutherford Model
- According to Rutherford’s model, electrons should continuously emit energy as they orbit around the nucleus
- This would cause electrons to spiral into the nucleus
- But, observations did not match this prediction
Instability of Rutherford Model
- Instability arises from continuous energy loss
- Electrons would emit energy in the form of electromagnetic radiation (radiation of energy)
- As a result, electrons would spiral inward, eventually causing the atom to collapse
Inspired by Planetary Motion
- Bohr applied Planck’s quantum theory to modify Rutherford’s model
- Proposed a new model where electrons occupy specific energy levels around the nucleus
- Electrons can only exist in these defined energy levels, called “shells”
Quantization of Energy Levels
- Electrons can only occupy specific energy levels or shells
- The energy of an electron is quantized, meaning it can only have certain discrete values
- Energy increases as we move further away from the nucleus
Absorption and Emission of Energy
- Electrons can absorb or emit energy when transitioning between energy levels
- Absorption: Electrons absorb energy and move to higher energy levels
- Emission: Electrons release energy and move to lower energy levels
Energy-Level Diagram
- Energy levels are represented by whole numbers (n = 1, 2, 3, …)
- The closest energy level to the nucleus has the lowest energy (n=1)
- As n increases, energy levels move further away from the nucleus and have higher energy
Ground State and Excited State
- Ground State: Electrons occupy the lowest energy level (n=1)
- Excited State: Electrons occupy higher energy levels (n>1) after absorbing energy
Emission of Light
- When electrons return to lower energy levels (from higher energy levels), they emit energy
- This emitted energy corresponds to different wavelengths of light
- The emission spectrum consists of discrete lines specific to the element
Equation for Energy Levels
- Bohr proposed an equation to calculate the energy levels of hydrogen atom
- E = -13.6 eV / n^2
- E is the energy of a specific energy level
- n is the principal quantum number representing the energy level
Absorption Spectrum
- When an electron absorbs energy and moves to a higher energy level, it creates absorption lines in the spectrum
- Absorption lines are dark lines in the spectrum where specific wavelengths of light are absorbed by the atoms
- Each element has a unique pattern of absorption lines
- Absorption spectrum can be used to identify the presence of elements in a sample
Planetary Model Comparison
- Bohr’s model is often compared to the planetary model of our solar system
- Just as planets orbit around the sun due to gravitational attraction, electrons orbit around the nucleus due to electrostatic attraction
- However, the comparison is not entirely accurate as electrons do not move in a well-defined circular path as planets do
Limitations of Bohr Model
- Bohr’s model successfully explains the stability of atoms and the discrete nature of energy levels
- However, it fails to explain the behavior of atoms with more than one electron
- It cannot account for the complex spectra observed in multi-electron atoms
- Bohr’s model is considered a limited explanation but was a significant step in the development of quantum mechanics
Quantum Mechanics
- Quantum mechanics is the branch of physics that deals with the behavior of subatomic particles
- It provides a more accurate description of the atomic structure and behavior
- Quantum mechanics uses mathematical equations and probability theory to describe the behavior of particles on a microscopic scale
- It is a fundamental theory in modern physics and has numerous applications in technology and research
Schrödinger Equation
- The Schrödinger equation is the fundamental equation of quantum mechanics
- It describes the behavior of particles as waves
- The equation uses wave functions to mathematically represent the probability of finding a particle in a particular state
- Solving the Schrödinger equation provides information about the energy levels and wave functions of electrons in an atom
Wave-Particle Duality
- Electrons and other subatomic particles exhibit both wave-like and particle-like properties
- This is known as wave-particle duality
- It is a fundamental principle of quantum mechanics
- The wave nature of electrons explains phenomena such as interference and diffraction
Heisenberg Uncertainty Principle
- The Heisenberg uncertainty principle states that it is impossible to simultaneously determine the exact position and momentum of a particle with precision
- The more accurately we try to measure one property, the less accurately we can measure the other
- This principle is a fundamental limitation of measuring subatomic particles
Electron Cloud Model
- The electron cloud model of an atom describes the probable locations of electrons around the nucleus
- It represents the three-dimensional probability distribution of finding an electron in a specific region
- The cloud represents the likelihood of finding an electron at a particular point in space
- Electron density is higher in regions of higher probability
Quantum Numbers
- Quantum numbers are used to describe the properties and characteristics of electrons in an atom
- Principal quantum number (n) describes the energy level or shell of the electron
- Azimuthal quantum number (l) defines the shape of the orbital
- Magnetic quantum number (m) determines the orientation of the orbital in space
- Spin quantum number (s) specifies the spin direction of the electron
Orbital Shapes
- Orbitals are regions in space where there is a high probability of finding an electron
- Orbitals have different shapes depending on the values of the quantum numbers
- S orbitals are spherical and have the lowest energy
- P orbitals are dumbbell-shaped and have higher energy
- D and F orbitals have even more complex shapes and higher energies