Section 2.2 and Chapter 7 Electron Configurations and Waves

Waves

 Frequency: number of waves in a given unit of time Wavelength = speed/frequency

Waves  Electromagnetic waves travel at the speed of light regardless of wavelength  c = 3.0*10 8 m/s Wavelength = c/frequency

Electromagnetic Spectrum  Visible Light is between 400 (violet) to 750 (red) nanometers  1 nanometer = 1*10 -9 meters

Quantum Theory  Proposed by Max Planck in 1900  The energy emitted or absorbed by an object is restricted to “pieces” of particular size Each piece is called a quantum.

Quantum Theory  The energy in a quantum is proportional to the frequency of the radiation (wave). E = hf E = Energy (J) h = Planck’s constant (6.626 x J-s) f = frequency

Quantum Theory  Planck’s theory proposes that energy absorbed or emitted from an atom is quantized  Why do we not notice this quantization of energy?

Quantum Theory  Planck’s constant is extremely small. In the everyday world, the quantization of energy is unnoticeable.  On the atomic level, it is extremely significant.

The photoelectric effect  When light is shined on a metal, electrons are ejected from the surface Each metal has a minimum frequency of light required to release electrons.  Example: Sodium metal Releases no electrons from red light Easily releases electrons from violet light  How can this be explained?

The photoelectric effect  Albert Einstein proposed that light is composed of quanta of energy that behave like tiny particles Photons Each photon carries an amount of energy determined by Planck’s equation

The photoelectric effect  When photon strikes an atom, it transfers its energy to an electron “All or nothing”  If the photon has a high enough frequency, it will have the energy required to eject the electron

Dual Nature of Light  Light behaves as both waves and particles (photons) Acts as a wave because it has a frequency and wavelength Acts as particles because they carry definite amounts of energy

The Bohr Model atom/hydrogen-atom_en.jnlp

Line Spectra  Each element will absorb energy and release energy only in specific wavelengths and frequencies  Atomic fingerprint  Sodium will produce yellow light  Lithium produces red, potassium purple

So waves act like matter…can matter act like waves?  It turns out that matter does act like waves in that it moves with wavelengths and frequencies.  Not noticed on macroscopic level, but it is on the atomic level.  Electrons move in wave-like forms

Heisenberg’s Uncertainty Principle  Proposed in 1927 by Werner Heisenberg  The position and momentum of an electron cannot simultaneously be measured and known exactly You can never know exactly where an electron is or how it is moving

Electron Orbitals  Electron orbitals do not show the exact path of an electron. Shows where an electron with a given energy is likely to be located.  Electrons are divided into energy levels Each energy level is divided into sublevels of different orbitals  4 basic shapes (types) of orbitals s, p, d, and f

s orbitals  Lowest energy  Holds up to 2 electrons  1 orbital per energy level

p orbitals  3 per energy level  Each can hold 2 electrons  6 total electrons per energy level  First p orbitals are in the second energy level

d and f orbitals  d orbitals: 5 orbitals, 2 electrons each 10 total electrons in each energy level Do not show up until the 3 rd energy level  f orbitals 7 orbitals, 2 electrons each 14 total electrons in each energy level Do not show up until the 4 th energy level

Electron Configurations  The distribution of electrons throughout the orbitals of an atom  Electrons will occupy the lowest energy orbitals available

Electron Spin  Two electrons in an orbital will have opposite spins. One may spin counterclockwise, the other will spin clockwise The direction is not important. What’s important is that they spin opposite each other.

Orbital Notation  Same as electron configuration except it shows the actual orbital and spin of each electron.