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**Ch. 13 Electrons in Atoms Ch. 13.1 Models of the Atom**

Ch Electron Arrangement in Atoms Ch Physics and the Quantum Mechanical Model

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**Ch. 13.1 Models of the Atom Evolution of Atomic Models**

John Dalton – atomic theory JJ Thomson – plum pudding model Ernst Rutherford - nuclear atom Niels Bohr – planetary model Fixed energy levels Quantums of energy Schrodinger – quantum mechanical model

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**Ch. 13.1 Models of the Atom The quantum mechanical model**

Primarily mathematical model Restricts energy of electrons Estimates probability if finding an electron in a certain position 90% probability

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**Ch. 13.1 Models of the Atom Atomic orbitals**

Designates energy levels of electrons with principal quantum numbers (n) n=1, n=2, n=3, n=4 and so on The principal energy levels have a specific number of sublevels The sublevels are designated s, p, d, and f Sublevels have specific numbers of orbitals The orbitals have specific shapes that correspond to the path of the electrons Nodes are areas where the probability of finding electrons is low

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**Ch. 13.2 Electron Arrangement in Atoms**

Electron Configurations The ways in which electrons are arranged around the nucleus of an atom Three rules explain how to find the electron configuration of atoms Aufbau principle – e- enter the lowest energy level first Pauli exclusion principle – an atomic orbital may hold at most 2 e- (with opposite spin) Hund’s rule – when e- occupy orbitals of equal energy, one e- enters each orbital until all the orbitals contain one e- with parallel spins Exceptional electron configurations – Cr, Cu

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**Ch. 13.3 Physics and the Quantum Model**

Light and atomic spectra Electromagnetic radiation Amplitude – height of wave from origin Frequency – measured in hertz (Hz) Wavelength – distance between crests (nm) Speed – speed of light (a constant, c) v = c/l Frequency = speed/wavelength Visible spectrum (ROY G BIV) Atomic emission spectra (fingerprint)

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**Ch. 13.3 Physics and the Quantum Model**

The quantum concept and the photoelectric effect Max Planck Amount of energy absorbed or emitted is proportional to the frequency of the radiation E = hv (energy = Planck’s constant x frequency) h = x Js Albert Einstein Light can be described as quanta of energy (photons) Dual nature of light (waves and particles) First to explain the photoelectric effect Energy must meet a threshold value to eject e- from the surface of metal

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**Ch. 13.3 Physics and the Quantum Model**

An explanation of atomic spectra When energized, e- move from the ground state to an excited state (n = 2,3,4, etc) The same amount of energy that was absorbed is then emitted as light and the e- falls back down to a lower energy level Results in Lyman, Balmer and Paschen series An upper limit for the frequency of light emitted exists because a very excited e- will escape the atom Bohr’s model of the atom was not able to explain bonding and was eventually replaced

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**Ch. 13.3 Physics and the Quantum Model**

Quantum mechanics De Broglie’s equation predicts that all matter exhibits wave-like motions True for all matter, but motion depends on the size of the object Wavelengths for objects visible to the naked eye are not measurable Wavelengths for extremely small objects are measurable Heisenberg’s uncertainty principle It is impossible to know exactly both the position and velocity of a particle at the same time The more precise the measured velocity, the less precise the position, and vice versa

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