Bohr Model Of Atom - Bohr Model Of Atom-I

Bohr Model of Atom-II – Assumptions

Electrons revolve in circular orbits around the nucleus
Electrons can only exist in certain discrete energy levels
Electrons can jump to higher or lower energy levels by absorbing or emitting energy respectively
Electrons in stable orbits do not radiate energy "

Electrons exhibit stable orbits at certain energy levels
When an electron transitions between energy levels, energy is absorbed or emitted
Emission: Electron jumps to lower energy level and emits energy in the form of photon
Absorption: Electron jumps to higher energy level by absorbing energy "

Line spectrum: Emission or absorption of energy results in the appearance of a spectrum with discrete lines
Each line represents a specific energy transition
Spectrum unique for each atom and helps identify elements
Lyman, Balmer, and Paschen series prominent for hydrogen atom "

Energy of an electron in nth orbit given by E = -13.6/n^2 eV
Negative sign indicates the energy is bound and stable
Energy difference between levels determines the wavelength/frequency of emitted/absorbed radiation "

Heisenberg uncertainty principle: It is not possible to precisely measure the position and momentum of a particle simultaneously
Limitations in knowledge of electron path in Bohr’s model resolved by the uncertainty principle
Describes the fundamental limitations of measurement in quantum mechanics
Δx Δp ≥ h/4π
- Δx is the uncertainty in position
- Δp is the uncertainty in momentum
- h is Planck’s constant "

Each element has a unique atomic spectrum
Atomic spectra result from the arrangement of electrons in energy levels
Ground state: Lowest energy state where electrons occupy the lowest energy levels
Excited state: Higher energy state where electrons occupy higher energy levels
Emission and absorption spectra provide information about the energy levels in an atom "

Quantum numbers describe the properties and characteristics of electrons in an atom
Principal quantum number (n): Determines the energy level and size of the orbital (n = 1, 2, 3, …)
Angular momentum quantum number (l): Determines the shape of the orbital (l = 0 to n-1)
Magnetic quantum number (ml): Determines the orientation of the orbital (-l to +l)
Spin quantum number (ms): Describes the spin direction of the electron (+1/2 or -1/2) "

Electron configuration: Distribution of electrons among the energy levels and orbitals in an atom
Follows the order of filling: 1s, 2s, 2p, 3s, 3p, …
Aufbau principle: Electrons occupy the lowest energy level available
Pauli exclusion principle: No two electrons in an atom can have the same set of quantum numbers
Hund’s rule: Electrons occupy orbitals of the same energy level singly, with parallel spin, before pairing up "

Orbital diagrams represent the distribution of electrons in energy levels and orbitals
Each box or line represents an orbital, with up and down arrows for electron spin
Example: Oxygen (O) has atomic number 8
- Electron configuration: 1s^2 2s^2 2p^4
- Orbital diagram: 1s↓↑ 2s↓↑ 2p↑ ↑ ↑ ↑ "

Valence electrons: Electrons in the outermost energy level of an atom
Determine the chemical properties and reactivity of an element
Elements in the same group have the same number of valence electrons
Example: Carbon (C) has atomic number 6 and electron configuration of 1s^2 2s^2 2p^2
- Carbon has 4 valence electrons "

Electron dot structures (Lewis structures) represent valence electrons as dots around the chemical symbol
Each dot represents one valence electron
Example: Nitrogen (N) has atomic number 7 and electron configuration of 1s^2 2s^2 2p^3
- Electron dot structure: N: . . "

Ionization energy: Energy required to remove an electron from an atom or ion in the gas phase
Trend: Ionization energy generally increases across a period and decreases down a group
Atoms with low ionization energy are more likely to form cations (lose electrons)
Example: Ionization energy of lithium (Li) is 520 kJ/mol
- Li(g) → Li^+(g) + e^- ΔH = 520 kJ/mol "

Electron affinity: Energy change when an electron is added to an atom or ion in the gas phase
Trend: Electron affinity generally increases across a period and decreases down a group
Atoms with high electron affinity are more likely to form anions (gain electrons)
Example: Electron affinity of chlorine (Cl) is -349 kJ/mol
- Cl(g) + e^- → Cl^-(g) ΔH = -349 kJ/mol "

Bohr’s model of the atom explained the stability and line spectra of hydrogen
Energy levels in the Bohr model are quantized and determined by the principal quantum number (n)
Electrons transition between energy levels by absorbing or emitting energy
The Bohr model has limitations, but it laid the foundation for further developments in quantum mechanics
Quantum numbers, electron configurations, and orbital diagrams are used to describe electron arrangement
Valence electrons and electron dot structures are important for chemical bonding
Ionization energy and electron affinity provide insights into the reactivity of elements