Bohr's Model and Spectroscopy

• Bohr's Model of the Atom continues the evolution of understanding and models of atoms.

Problems with the Rutherford Model

• Rutherford's atomic model proposed an orbit system for electrons, similar to planets orbiting the Sun.
• This model was reasonable because the Sun's gravity is counteracted by the planet's movement.
• It seemed electrons orbiting an atomic nucleus would act the same way.
• But there was a problem with this idea.
• Moving charged particles produce EM energy.
• An electron traveling in an orbit emits energy as photons and loses energy.
• If an electron loses energy as it orbits, it should spiral in toward the positively charged nucleus and the atom would collapse.
• This collapse doesn't happen.
• Electrons aren't stationary.

Spectroscopy

• Robert Bunsen and Gustav Kirchhoff invented the spectroscope in 1859.
• Spectroscopy studies analyzing spectra of visible light, UV, and X-rays.
• When light goes through a spectroscope, an emission spectrum forms.
• A bright-line spectra is a pattern of light.
• A continuous spectrum has every wavelength in a particular region of the EM spectrum, like when white light goes through a prism.
• An emission spectrum of gas occurs when light from a heated gas goes through a prism to separate the emitted light into its component colors.
• Every element has its own unique line spectrum, which is like a fingerprint.
• A spectrometer identifies unknown elements.

Line Spectrum of Hydrogen

• Niels Bohr used these techniques and the emission spectrum of hydrogen to develop a quantum model.
• Unique line spectrum of hydrogen atoms shows electrons exist only at discrete energy levels, meaning energy is quantized.
• This aligns with Planck's quantum theory.
• Energy is required to move electrons from various "states."
• Electrons can only attain particular discrete energy levels.
• Bohr's model has two main postulates.
• First: Electrons don't radiate energy as they orbit the nucleus.
• Orbits have constant energy states called a ground state, which is the lowest energy state.
• Second: Electrons can only change energy by undergoing a transition from one stationary state to another (excited state).
• Electrons need to absorb sufficient energy to do this.
• Gaining a photon moves e- to a higher energy level.
• Losing a photon moves e- from a higher to a lower energy level.
• A quantum of energy, equal to the absorbed quantity, releases a "photon" (EMR) when the electron returns to its ground state.
• The higher the energy levels become more closely packed together.
• Electrons make jumps of varying distances, leading to lines on the emission spectrum that are not evenly spaced.
• Not all EM radiation is visible.
• Many electron transitions where the released energy returns to the ground state, have wavelengths beyond the visible part of the spectrum.
• That is why we only see certain lines and colors on the emission spectrum.

Successes and Failures of Bohr's Model

• Bohr's model successfully showed the filling of successive orbitals and predicted the number of electrons in each energy level well.
• It explained that higher energy levels contain more energy.
• But it didn't account for electron motion.
• It only explained spectra for hydrogen.
• The energy levels Bohr calculated for the e- in the H atom were similar to values obtained from the emission spectrum via spectroscopy.
• Data observed for atoms with more than one e did not match his predicted models.

Explore Bohr's atomic model, addressing the limitations of Rutherford's model. Learn how spectroscopy, invented by Bunsen and Kirchhoff, aids in analyzing spectra of visible light, UV, and X-rays, providing insights into atomic structure and behavior.