Atomic Structure and Spectra

Based on the alpha-particle scattering experiment, Rutherford proposed that an atom consists of a tiny, dense, positively charged nucleus at its centre, with electrons orbiting it in a vast empty space. This is often referred to as the 'nuclear model' or 'planetary model'. Limitations: 1. Atomic Stability: Classical physics predicts that orbiting electrons, being accelerating charges, should continuously radiate energy and spiral into the nucleus, causing the atom to collapse. This contradicts the observed stability of atoms. 2. Atomic Spectra: The model predicts a continuous spectrum of light emitted by atoms, whereas experiments show discrete line spectra (specific, distinct wavelengths of light).

Niels Bohr proposed a model for the hydrogen atom that addressed the limitations of Rutherford's model. His model introduced the concept of quantised energy levels. Postulates of the Bohr Model: 1. Electrons orbit the nucleus in certain stable, circular orbits (called stationary states) without radiating energy. These orbits correspond to discrete, quantised energy levels. 2. Electrons in these stable orbits have discrete, quantised angular momentum, L = n(h/2π), where n is an integer (principal quantum number) and h is Planck's constant. 3. An electron can only jump from one stationary state to another by absorbing or emitting a photon of specific energy. The energy of the photon (hf) must be exactly equal to the energy difference between the two states: hf = E₂ - E₁. How it addresses Rutherford's limitations: 1. Atomic Stability: Postulate 1 states that electrons do not radiate energy when in stable orbits, thus preventing the atom from collapsing. 2. Atomic Spectra: Postulate 3 explains the discrete line spectra, as only specific energy differences (and thus specific photon frequencies/wavelengths) are possible when electrons transition between quantised energy levels.

Electrons in an atom can only occupy specific, discrete energy states or levels. These energy levels are quantised, meaning electrons cannot exist at energies between these levels. * Ground State: The lowest possible energy level an electron can occupy (n=1). * Excited States: Higher energy levels (n=2, 3, 4, etc.) that an electron can occupy if it gains sufficient energy. * Ionisation: If an electron gains enough energy to completely escape the atom, the atom becomes an ion. The ionisation energy is the minimum energy required to remove an electron from the ground state to infinity (where its energy is considered 0 eV).

An emission spectrum is produced when electrons in excited atoms (e.g., in a gas discharge tube) fall from higher energy levels to lower energy levels. As they transition, they emit photons. Since only specific energy differences are allowed, only photons of specific frequencies (and thus specific wavelengths) are emitted. This results in a spectrum of bright lines on a dark background. Each element has a unique emission spectrum, acting like a 'fingerprint'.

An absorption spectrum is produced when white light (containing a continuous range of wavelengths) passes through a cool gas. Electrons in the gas atoms absorb photons that have energies exactly matching the energy differences between their allowed energy levels. These absorbed photons cause electrons to jump to higher energy levels. The wavelengths corresponding to these absorbed photons are removed from the continuous spectrum, appearing as dark lines on a bright background. The dark lines in an absorption spectrum occur at the same wavelengths as the bright lines in the emission spectrum of the same element.