Electrons in Atoms Ch. 5 Notes

Light & Quantized Energy Nuclear Model was incomplete How are electrons arranged? Why aren’t e- pulled into the nucleus? Why do elements behave differently? Early 1900s Certain elements emit visible light when heated in a flame

Electromagnetic Radiation A form of energy that exhibits wavelike behavior Waves can be described… Wavelength (l) Frequency (n) Amplitude

Wavelength – The shortest distance between equivalent points on a wave Frequency – The # of waves that pass a given point in one second Amplitude – Height from the origin to the crest/trough

Wave Diagram

All electromagnetic radiation travels at the speed of light (c). c = 3.00 x 108 m/s The velocity of a wave is equal to the wavelength times frequency c = ln

Though the speed of electromagnetic radiation is always the same both the frequency and the wavelength can vary. The electromagnetic spectrum encompasses all forms of electromagnetic radiation. Only difference in portions are wavelength and frequency

Considering light as a wave does not explain everything Different colors of light correspond to different frequencies and wavelengths.

Quantum – The minimum amount of energy that can be gained or lost. 1900 Max Planck trying to figure out why certain types of light are emitted from heated objects. Discovered matter can only gain or lose energy in small specific amounts (quanta). Quantum – The minimum amount of energy that can be gained or lost.

Equantum = hn h – Planck’s constant h = 6.626 x 10-34 Photoelectric effect – photo e- are emitted from the surface of a metal when light with a certain frequency shines on its surface. Wave model could not explain

1905 Albert Einstein Electromagnetic radiation has both wavelike and particlelike natures. Photon – Particle of electromagnetic radiation with no mass that carries one quantum of energy Ephoton = hn

Atomic emission Spectrum – the set of frequencies of electromagnetic radiation emitted by atoms of an element Each element has a unique spectrum

Electron Configurations The arrangement of electron’s in an atom Tend to assume lowest possible energy (ground state) Rules that tell how e- are arranged Aufbau Principle Pauli Exclusion Principle Hund’s Rule

All orbitals of an energy sublevel are of equal energy Aufbau Principle Each electron must occupy lowest energy level orbital available Fig. 5-17 p. 133 All orbitals of an energy sublevel are of equal energy Energy sublevels within a principal energy level have different energies. s < p < d < f Orbitals from one sublevel and principal level can over lap with those from another principal level

. Pauli Exclusion Principle (Wolfgang Pauli) Hund’s Rule Max of 2 electrons can occupy one orbital Iff they have opposite spins Hund’s Rule Single electrons w/ the same spin must occupy each equal-energy orbital before opposite spin electrons can occupy the same orbitals

Electron-dot structure (G.N. Lewis) Valence electrons Electrons in the outermost orbitals Highest principal energy level Chiefest in determining chemical properties of an element Electron-dot structure (G.N. Lewis) Element’s symbol surrounded by dots Symbol = nucleus and inner electrons Dots = valence electrons