1. Atomic Spectra and Atomic Energy States Reminder: A.S. 13.1.5-13.1.7 due Monday 3/23/15 A.S. 13.1.8-13.1.13 due Tuesday 3/24/15 WebAssign Quantum #3 due Tuesday 3/24/15

2.  When a gas is heated to a high temperature, or exposed to a large electric field  Electrons in the atoms absorb the energy  When the electrons fall back down to the original electron energy the energy is emitted as Electromagnetic radiation  To the right: hydrogen gas, exposed to high electrical potential Emission Spectra

3.  Direct light through a diffraction grating, or a prism  Light at different wavelengths will diffract or refract different amounts  The wavelengths that combine to give us the color we see emitted can be separated in this way  Hydrogen, for example, is composed of 4 main wavelengths of light: How do we see spectral lines?

4.  Sometimes, we see spectra showing us which wavelengths were absorbed by a sample of gas:  Interestingly, the wavelengths of light ABSORBED by the gas are the SAME as those EMITTED by the gas… Absorption Spectra

6. “Balmer Series”

7.  Rutherford came up with the planetary model of the atom:  There is a central, dense, positively charged nucleus  Electrons occupy a large space outside the nucleus  Electrons occupy “orbits”, much like planets orbit the sun (our center of the solar system)  WHY doesn’t this work? Review: Planetary Model

8.  Combining the ideas of Balmer and Rutherford, Niels Bohr made an attempt to “correct” the fundamental flaw of the planetary model using the following assumption:  Electrons exist with discrete energy in each orbit (energy level)  In order to move between energy levels, a discrete amount of energy must be absorbed by or released from the electron Electron Energy Levels

9.  Electrons exist at specific radii from the nucleus—energy levels  Quantitatively, the energy of the electron in that energy level can be determined using the following relationship: Bohr Model of the Atom

10.  As n increases, the energy levels become closer together (unlike the diagram on the previous slide)  As n approaches infinity, the total energy of the electron approaches 0  As E approaches zero, the force keeping the electrons bound to the nucleus decreases  Ionization Energy: The energy that must be added to an electron in order to release it from the atom Characteristics of Electron Energy Levels

11.  Significantly increasing the temperature  Bombarding it with additional electrons (high velocity collisions)  Subjecting it to a very high electric potential  Causing photons to fall on the atoms Ways of ionizing an atom:

12.  Describes the behavior of the electron in a Hydrogen atom really well…however:  Does NOT treat any atom with more than one electron  Assumes circular orbits  Cannot predict INTENSITIES of emitted light—only wavelength  Does not predict the division of energy levels (i.e. the p, d, f orbitals all have subdivisions) Limitations to Bohr’s Model

13.  Schrodinger Theory:  Assumptions:  Electrons in the atom can be described by wave functions  Wave functions fit boundary conditions in 3 dimensions, allowing for multiple “modes” that have a discrete energy state  Electron has an undefined position, but there is a probability that the electron exists in a position So…now what?

14.  Wavefunction (ψ): a function of position and time  Mathematically the probability that an electron will be in a particular position at a particular time can be determined by the square of the absolute value of the wavefunction at that time.  In other words, there are places where electrons are most likely to be found…not just circular orbits! Electron Wavefunctions

15.  For each energy level for Hydrogen, there is a probability curve describing how likely it is that an electron can exist in that position. Hydrogen electron probability

16.  Fundamental idea: wave-particle duality  Since particles sometimes act like waves, and waves sometimes act like particles, there isn’t a perfect, clean way to divide physical objects into one category or the other.  Misconception alert! This has nothing to do with experimental uncertainties!  It’s all about measuring things with an indefinite precision (remember those distribution graphs we just saw? ) Uncertainty Principle

17.  It is not possible to simultaneously measure both the position and the momentum of a particle.  The more sure we are about the position of a particle, the less certain we are about its momentum, and vice- versa. Heisenberg’s Uncertainty Principle

18.  We can also describe the uncertainty principle in terms of Energy and Time: Another variation…