• Once the ideas of Planck and Einstein (re: quanta) had been accepted, the study of atomic and molecular structure began.
• With the invention of the spectroscope (Bunsen & Kirchhoff, Fig 2/175) the quantitative analysis of the atomic spectra of elements began.
• The range of frequencies of light emitted or absorbed by an atom is called its spectrum (pl. spectra).

Spectroscope

• When an element is excited, atoms absorb energy, then as e-s fall back to “ground state”, energy is released in the form of light
• Visible line spectra called emission spectra of discrete(distinct) coloured lines – specific wavelengths = specific energies, not a continuous spectrum
• Atoms absorbing energy – discrete black lines – absorption spectrum.

Emission Spectra:

• Atomic spectra not explained – b/c if R’s model was correct, continual acceleration means continuous spectrum of light – spectral evidence shows this is not the case!

The Confusion About Bright Line Spectra

• Niels Bohr – Danish physicist, working for Thomson (raisin bun guy) suggested that line spectra could be explained using ideas from  the quantum theory!
• Bohr began to work with Rutherford (1913) – modified Rutherford’s atomic theory

Bohr’s Model of the Atom

• 54 years after line spectra were first described qualitatively by Bunsen and Kirchhoff.
• 28 years after line spectra were  described quantitatively by Jacob Balmer.

Bohr’s Postulates

• The atom has only specific allowable energy levels (stationary states) & each stat.state corresponds to the e-s occupying fixed circular orbits around nucleus.
• While in stat.state, atoms do not emit energy.
• An e- changes stat. state by emitting or absorbing a specific quantity of energy that is exactly equal to the difference in energy btw the two stat. states. (Fig 4/176)
• Higher energy level -> lower: loss of energy in the form of light
• Lower energy level -> higher: energy is absorbed
• Only certain quanta of light can be emitted or absorbed by an atom, so the energy of the electrons inside the atom is also quantized. In other words, an electron can only have certain allowable energies. (staircase analogy).
• H-spectra example
• If an e- that has been excited to the 3rd NRG level & then falls from 3rd to 2nd NRG level, emits a photon of red light with wavelength 656nm & red line is visible on emission spectrum.
• Each line on e. spectrum corresponds to a specific energy transition (I.e. 4th to 2nd = green, 5th to 2nd = blue)
• Hydrogen

Evaluating Bohr’s Model

• Was reasonably successful at explaining Mendeleev’s PT
• Periods result in filling an additional energy level (i.e. Na: 3 levels)
• A period ends when the maximum # of electrons in the valence level is reached (i.e. 2, 8, 8, 18)
• BUT – could only explain one-electron systems (so still more complex than he thought)

Bohr Energy-Level Diagrams

• Same rationale as orbit diagrams, but the focus is on the energy that the electrons have, rather than the position of the electrons, so orbits are not used.
• For a Bohr energy level diagram, start from the bottom and work your way up!

Example:

• 5e^– (3^rd energy level)                   (from group 15)
• 8e^– (2^nd energy level)                   (from 8 elements in period 2)
• 2e^– (1^st energy level)                   (from 2 elements in period 1)
• 15p⁺ (protons)                                (from atomic number)
• P                              (symbol)                                 (from PT)
• phosphorous   (name)                                         (lowercase)