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Chapter 13 Notes. Electron Models. Evolution of Electron Models. The first model of the electron was given by J.J. Thompson—the electron’s discoverer. His was the “plum pudding” model. The Rutherford Model. With Rutherford’s discovery of the nucleus of an atom, the atomic model changed.
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Chapter 13 Notes Electron Models
Evolution of Electron Models • The first model of the electron was given by J.J. Thompson—the electron’s discoverer. His was the “plum pudding” model.
The Rutherford Model • With Rutherford’s discovery of the nucleus of an atom, the atomic model changed.
The Bohr Model • Niels Bohr introduced his model, which answered why electrons do not fall into the nucleus. • He introduced the concept of energy levels, where the electrons orbited similar to the way the planets orbit the sun.
Bohr Model and Energy Levels • In the Bohr model, electrons are in energy levels, or regions where they most probably are orbiting around the nucleus. • The analogy is that energy levels are like the rungs of a ladder—you cannot be between rungs, just like an electron cannot be between energy levels. • A quantum of energy is the amount of energy it takes to move from one energy level to the next.
Bohr Model and Energy Levels • The Bohr model worked well for explaining the behavior of electrons in hydrogen, but for all other elements, the equations he used to predict the electron location did not work.
Quantum Mechanical Model • In 1926, Erwin Schrodinger used the new quantum theory to write and solve mathematical equations to describe electron location.
The Quantum Mechanical Model, cont. • Today’s model comes from the solutions to Schrodinger’s equations. • Previous models were based on physical models of the motion of large objects. • This model does not predict the path of electrons, but estimates the probability of finding an electron in a certain position. • There is no physical analogy for this model!
Where are the electrons? • In an atom, principal energy levels (n) can hold electrons. These principal energy levels are assigned values in order of increasing energy (n=1,2,3,4...). • Within each principal energy level, electrons occupy energy sublevels. There are as many sublevels as the number of the energy level (i.e., level 1 has 1 sublevel, level 2 has 2 sublevels, etc.)
Where are the electrons? • There are four types of sublevels we will talk about—s,p,d and f. Inside the sublevel are atomic orbitals that hold the electrons. Every atomic orbital can hold two electrons. • S has one orbital, P has three, D has five and F has seven. How many electrons can each one hold?
Orbital Shapes s orbital = s sublevel = p sublevel + + pz orbital px orbital py orbital
Where are the electrons? • So how many electrons can each energy level hold? • Level 1 has an s sublevel=2 e- • Level 2 has an s and a p sublevel=8e- • Level 3 has an s, p and d sublevel=18e- • Level 4 has an s, p, d and f sublevel=32e-
Electron Configuration • In the atom, electrons and the nucleus interact to make the most stable arrangement possible. • The ways that electrons are arranged around the nucleus of an atom is called the electron configuration.
Pauli’s Exclusion Principal • An atomic orbital may describe at most two electrons. Hund’s Rule • When electrons occupy orbitals of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins. Aufbau Principal • Electrons enter orbitals of the lowest energy first.
1s He 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 5d 6p 7s 6d 7p 4f 5f
What does a wave look like? • With your partner, label all the parts of a wave you can remember.
A Quick Look at Waves • The number of waves to pass a point in a given time is called frequency (n) and is measured in 1/s or Hertz (Hz).
Electromagnetic Radiation (EMR) • According to the wave model, visible light consists of electromagnetic waves and is just a small fraction of waves classified as electromagnetic radiation. • Other EMR includes radio waves, microwaves, infrared, ultraviolet, X-rays, gamma rays and cosmic rays.
Electromagnetic Radiation (EMR) • All of these waves travel at the same speed, 3.0x108m/s! • The waves differ in their frequencies and wavelengths, and obey the equation: c = l n • This is an inverse relationship—as the frequency increases, the wavelength decreases.
Practice Problems • What is the wavelength of an electromagnetic wave with a frequency of 4.45x1015Hz? • What is the frequency of a light wave with a wavelength of 497nm? • What is the wavelength in nanometers of an electromagnetic wave with a frequency of 2.97x1014Hz?
Atomic Emission Spectra • When sunlight is broken down into the waves it is made of, it creates a continuous spectrum • Scientists used a hydrogen lamp to produce light, they expected a continuous spectrum but it wasn’t! They had an atomic emission spectrum.
When we previously found the electron configuration for elements, it was for electrons at ground state (the lowest energy possible). • As energy is added to atoms, they absorb the energy by electrons going from ground state to an excited state, where electrons are no longer in the lowest energy orbitals.
Electrons can then only go back to ground state by releasing the energy, usually in the form of light in discreet packets called photons. • These packets defied classical physics, that said electrons would go back to ground state continuously.
Max Planck • To understand why this points towards the concept of energy levels, we need to know about Max Planck’s discovery: • E = h n • Planck’s constant (h=6.6262x10-34Js)
Practice Problems • How much energy is associated with a wave with a frequency of 4.4x1014Hz? • An electromagnetic wave is found to have 1.18x10-19J of energy. What is its frequency? • How much energy is associated with a wave of red light with a wavelength of 697nm?
Putting It Together • So, if only specific frequencies of light are emitted when electrons fall back to ground state from being excited, then there are only certain energies that electrons can have. This explains atomic emission spectra!
Even Stranger… • Louis de Broglie predicts yet another property of electrons—that they have both a wave nature and a particle nature. • Any moving particle can be described to have a wave nature described by de Broglies equation: • l = h / mv
Even stranger still is the Heisenberg Uncertainty Principle. • It states that you cannot know both a particles exact position and exact velocity (the more you know about one the less you know about the other).
