1. Arrangement of Electrons in Atoms

2. Development of a New Atomic Model
What prevents negative electrons from being drawn into the positive nucleus?
A relationship between light and electrons led to further understanding.

3. Properties of Light The Wave Description of Light
Electromagnetic radiation (a form of energy that exhibits wavelike behavior as it travels through space).
Electromagnetic spectrum (Together, all the forms of electromagnetic radiation).
Wavelength (the distance between corresponding points on adjacent waves).
Frequency (the number of waves that pass a given point in a specified time, usually one second, often expressed as hertz).
Frequency and wavelength are mathematically related. As wavelength increases its frequency decreases and vice versa.

4. Photoelectric Effect
Refers to the emission of electrons from a metal when light shines on the metal.
Particle description of light
Quantum of energy: minimum quantity of energy that can be lost or gained by an atom
Einstein furthered this by explaining that light can be thought of as a stream of particles. (photons)
A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.
Electromagnetic radiation is absorbed by matter only in whole numbers of photons.
Electrons in different metals are bound more or less tightly, so different metals require different minimum frequencies to exhibit the photoelectric effect.

5. Hydrogen-Atom Emission-Line Spectrum
The lowest energy state of an atom is the ground state.
A state in which an atom has higher potential energy than it has in its ground state is an excited state.
When an excited atom returns to its ground state or lower energy excited state, it gives off the energy it gained in the form of electromagnetic radiation (neon signs).

6. When investigators passed electric current through H gas at low pressure (cathode tube) a pinkish glow was emitted.
When this was then passed through a prism 4 bands of light were identified and this is part of hydrogen’s emission-line spectrum.
Classical theory predicted that the hydrogen atoms would be excited by whatever energy amount was added to them and thus expected a continuous range of frequency emissions of electromagnetic radiation (continuous spectrum).

7. Quantum Theory
Whenever an excited H atom falls to its ground state or to a lower-energy state, it emits a photon of radiation. This energy is equal to the difference in energy between the atom’s initial state and final state.
This means that atoms exist in only very specific energy states.

8. Bohr Model 1913 Danish physicist
Proposition: hydrogen-atom model that linked the atom’s electron to photon emission. The electron can circle the nucleus only in allowed paths or orbits. When the electron is in one of these orbits, the atom has a definite, fixed energy.
The energy of the electron is higher when the electron is in orbits that are successively farther than from the nucleus.

9. Bohr’s model continued
An electron can only be in one orbit or another but not in between or in two at one time.
While in a given orbit, the electron is neither gaining or losing energy.
Electron’s can move to a higher-energy orbit by gaining an amount of energy equal to the difference in energy between the higher-energy orbit and the initial energy orbit.
Emission – when the electron falls to a lower energy a photon is emitted.
Absorption – energy is added to move an electron from a lower energy orbit to a higher energy orbit.

10. Problem
It was soon recognized that Bohr’s model did not explain the spectra of atom with more than one electron….
Nor did it explain the chemical behavior of atoms……