Absorption & Emission Spectra

Emission Spectrum Hot, glowing objects emit a continuous spectrum of light  temperature. Fill a tube a glass tube with a low- pressure gas and heat it. It will emit discrete wavelengths of light. This is called the emission spectrum.

Absorption Lines When bright light passes through a cold gas and then a prism, not all of the light from the distance source gets through, i.e., some is absorbed. When light passes through a prism, it produces a continuous spectrum. When light from a hot gas passes through a prism, it produces a emission lines.

Comparison Every wavelength absorbed by the gas is also emitted NOT every emitted wavelength is absorbed. Absorbed Emitted

Some big questions… Beginning of 20 th century, physics could not explain … structure of matter the stability of matter the discrete spectra the origins of X-rays or radioactivity.

Bohr Model Danish physicist Niels Bohr introduced a model that pointed a way forward. He proposed the quantization of electrons’ orbits. Electrons are allowed here and here but NOT here

Bohr Model Bohr called stable electron orbits “stationary states”. This is one stationary state This is another

In this model, electrons in an atom can possess discrete amounts of energy, i.e., only E 1, E 2, E 3, etc., and never E 1.5, or E At E 1 or ‘ground state’, the electrons have the least energy and are indefinitely stable.

When electrons drop from a higher energy level to a lower energy level, they emit a photon. E.g., from n=5 to n=4 E.g., from n=3 to n=1 Photon emission

Electrons only absorb those photons that carry them to a discrete, more excited state. Photon absorption

At some point, photons have enough energy to rip the electron free. This is called ‘ionization energy’. For the electrons held in a hydrogen atom, the ionization energy is 13.6 electron-volts. Ionization energy

Application Electrons can absorb discrete amounts of energy from collisions with other particles. The electrons jump a more excited state and, a couple of nanoseconds later, emit a photon. e.g., fluorescent lights

Think about this… Atoms in a gas are struck by 3.0 eV photons. As a result, the electrons move from their stable ground state to the n=3 excited state. Shortly afterwards, the atoms emit 2.0 eV photons. What other emitted photons might be observed? A)3.0 eV and 5.0 eV photons B)3.0 eV and 1.0 eV photons C)1.0 eV and 5.0 eV photons D)3.0 eV photons only E)1.0 eV photons only Electrons gained 3.0 eV so have to lose 3.0 eV to return to ground state. They could do it in one big leap (n=3 to 3.0 eV) or in two small jumps (n=3 to 1 eV and then n=2 to 2 eV)

Implications of Bohr’s model 1)Matter is stable Once an atom is in its ground state, it will remain there indefinitely 2)Atoms emit and absorb a discrete spectrum. Energy of photon (E=hf) must match energy level available to atom Otherwise, energy from light is not absorbed.

Implications of Bohr’s model 3)Absorption wavelengths are a subset of emission wavelengths Emission: Electrons can drop from n=3 to n=1 or from n=2 to n=1 any time and emit that particular wavelength of light. Absorption: Absorption lines: n=1 to n=2 or to n=3 or to n=4, etc. However, if there are no electrons in an excited state when a photon hits the atom, the electron cannot jump from n=2 to n=3, or from n=3 to n=6, etc., i.e., no absorption 4)Each element has a unique spectrum different elements with different number of electrons have different stable stationary states and therefore emit and absorb photons of different wavelengths.

Success & limitations of Bohr model Success: Matter is stable Hydrogen emission and absorption modeled Limitations Does not work with any other elements

Quantum mechanical hydrogen Electron energy is quantized; what about other properties? In 1925, Edwin Schrodinger (of cat fame) introduced theory that described spectra of all atoms by quantizing the following: 1)Energy: n, principal quantum number 2)Angular momentum: L, orbital quantum number 3)Plane of electron’s revolution or tilt: m, magnetic quantum number 4)Spin: m s, spin quantum number

But… Shouldn’t most electrons be in ground state? They aren’t. In 1925, Austrian Wolfgang Pauli suggested that no two electrons in the same quantum mechanical system could have the exact same set of quantum numbers (n, l, m, m s ). The so-called Pauil Exclusion Principle saved the periodic table!

Implications

Emission spectra revisited

Sodium, for example