Bohr Atomic Theory, Section 3.2

• The Bohr Atomic Theory is in Section 3.2.

Main Achievement of Rutherford's Model

• Advanced the understanding of the atom.
• Established the existence of a positively charged nucleus around which electrons orbit.

Main Problem with Rutherford's Model

• Predicted that electrons orbit the nucleus like planets orbit the sun.
• Electrons are attracted to the positively charged nucleus but maintain their orbit because they are moving.
• Accelerating charges produce electromagnetic radiation continuously.
• Classical theory states that orbiting electrons should emit photons, lose energy, and spiral into the nucleus, causing the atom to collapse, which does not happen.

Bohr's Solution to Rutherford's Problem

• If light released/absorbed from an atom was quantized, then the energy of electrons inside the atom must also be quantized.
• Each orbit has a fixed distance and fixed energy, so electrons can only have certain allowed energies.
• Bohr stated that there were special energy states (stationary states) and existing rules did not apply inside an atom.
• Bohr used a new theory combined with experimental study to make his atomic theory.

Bohr's Postulates

• Electrons exist in circular orbits, like planetary orbits. The central force is the electrostatic force between the positive nucleus and the negative electrons.
• Electrons can exist only in a series of "allowed" orbits that can also be described as "energy levels." Because only certain orbits are allowed, only certain energy levels are allowed, which means that the energy in electrons is quantized.
• Electrons can jump between orbits by absorbing or emitting photons carrying an amount of energy that equals the difference in the energy levels.
• Emission spectra (bright line spectrum) are a series of bright lines of light produced or emitted by a gas excited by heat or electricity.
• When an electron transitions from a higher to a lower energy state, it loses energy, releasing a photon of light, resulting in a bright line in the light spectrum.
• When energy is absorbed, an electron transitions from a lower energy state to a higher one, explaining a dark line in the absorption spectrum.

Bohr's Explanations

• Bright and dark line spectra of elements show that only certain quanta of light can be emitted or absorbed by an atom.
• If light released/absorbed from an atom is quantized, then the energy of the electrons inside the atom must also be quantized.
• Electrons can only have certain energies, similar to the planetary model, where each electron has a fixed distance with fixed energy (only certain energies were allowed).

Energy Level Diagrams

• From bottom up, the energy levels are ordered as follows: 6th (3rd E level), 5th(2nd E level).

Other Points

• 4th energy level # e goes to Ze, 3rd #protons goes to 9pt, 2nd symbol is F, 1st lowercase name of atom is Fluorine, which is an example.

Energy Level Diagrams

• Atomic # is the number of protons and total # of electrons.
• Period # is the number of energy levels.
• Group # is the last digit and # of electrons in the outer energy level.

Success of Bohr's Experiment

• Bohr's model explains Mendeleev's periodic law and the periodic table.
• Periods result from the filling of electron energy levels in an atom, and periods end when the maximum number of electrons in the outer level is full.
• Max # of electrons in each energy level is given by # of elements in each period of the periodic table.
• Last digit of group # = # of electrons in the valence energy level.
• Bohr did calculations as if electrons were in circular orbits, but the electrons' energy was the most important property (not motion).
• Bohr was able to explain the visible spectrum for hydrogen.

Problem with Bohr's Theory

• Works only for hydrogen atoms or ions with only 1 electron.
• Calculations of spectra lines using Bohr's theory for any atom/ion containing more than 1 electron did not work with experimental results.
• The more electrons, the more off the calculations were.

Experimental Evidence

• Mendeleev: there is periodicity of the physical and chemical properties of the elements.
• Mendeleev: there are 2 elements in the second period of the periodic table.
• Kirchhoff, Bunsen + Johann Balmer: emission and absorption line spectra, and not continuous spectra, exist for gaseous elements.

Theoretical Explanation

• A new period begins in periodic table when a new energy level of e is started in the atom.
• There are 2 electrons max in the 1st electron energy level and 8 in the next level.
• Therefore, the energy of light absorbed and emitted is quantized, and the energy of electrons in atoms is quantized.
• Electrons are responsible for producing light of line spectra.
• The hydrogen visible spectrum is the simplest.
• An electron "jumps" from one orbit (E level) to another (electron transition).