Chapter 5

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# Chapter 5 - PowerPoint PPT Presentation

Chapter 5. Electrons in Atoms. The Bohr Model. An electron is found only in specific circular paths, or orbits, around the nucleus. Each orbit has a fixed energy. The orbits are called ‘energy levels.’. Energy Levels. Energy levels are like the rungs of a ladder:

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### Chapter 5

Electrons in Atoms

The Bohr Model
• An electron is found only in specific circular paths, or orbits, around the nucleus.
• Each orbit has a fixed energy. The orbits are called ‘energy levels.’
Energy Levels
• Energy levels are like
• the rungs of a ladder:
• You can move up or down by going from rung to rung.
• You can’t stand in-between rungs.
• For an electron to change energy levels it must gain or lose exactly the right amount of energy.
A Quantum
• A quantum of energy is the amount needed to move an electron from one energy level to another.
• The energy of an electron is said to be “quantized.”
• Energy levels in an atom are not all equally spaced.
An Airplane Propeller
• The blurry picture of an airplane propeller represents the area where the actual propeller blade can be found.
• Similarly, the electron cloud of an atom represents the locations where an electron is likely to be found.
The Model Quantum Mechanical
• Comes from the mathematical solution to the Schrodinger equation.
• Determines allowed energies an electron can have & how likely it is to find the electron in various locations around the nucleus.
• Uses probability
Atomic Orbitals
• A region in space in which there is a high probability of finding an electron.
• Energy levels of electrons are labeled by principal quantum numbers (n)

n = 1, 2, 3, 4 …

s Orbitals

are spherical

p Orbitals

are dumbbell- shaped

d Orbitals

4 out of the 5 d orbitals have clover leaf shapes

f Orbitals

are more complicated

Atomic Orbitals

The number and kinds of atomic orbitals depend on the energy sub level.

• N=1 has 1 sublevel called 1s
• N=2 has 2 sublevels called 2s and 2p
• N=3 has 3 sublevels called 3s, 3p, and 3d
• N=4 has 4 sublevels 4s, 4p, 4d, and 4f

The maximum number of electrons that can occupy a principle energy level is 2n2.

(n=principle quantum #)

Electron Configurations
• Electrons in an atom try to make the most stable arrangement possible (lowest energy)
• The Aufbau Principle, the Pauli Exclusion Principle, and Hund’s Rule are guidelines that govern electron configurations in atoms
Aufbau Principle
• Electrons occupy the orbitals of lowest energy first
Pauli Exclusion Principle
• An orbital can hold at most 2 electrons
• Does it make sense that two negatively charged particles will ‘want’ to share the same space?
• This phenomenon is made possible because electrons possess a quantum mechanical property called spin
Electron Spin
• Spin may be thought of as clockwise or counter-clockwise
• An arrow indicates an electron and its direction of spin
• An orbital containing paired electrons is written
Hund’s Rule
• When filling orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with similar spin
Hund’s Rule
• How would you put 2 electrons into a p sublevel?
• How would you put 7 electrons into a d sublevel?
Light
• Now that we understand how electrons are arranged in atoms, we can begin to look at how the frequencies of emitted light are related to changes in electron energies
Light
• Light waves properties:
• Amplitude – the wave’s height from zero to crest
• Wavelength – the distance between crests
• Frequency – the number of wave cycles to pass a given point per unit of time (Usually Hz = 1/s)
Light
• Wavelength has the symbol (λ) lambda.
• Frequency has the symbol (ν) nu.
• The speed of light is a constant (c) = 3x108 m/s
• c = λν
Light
• How are wavelength and frequency related?
• They are inversely related. As one increases, the other decreases
• How long are the wavelengths that correspond to visible light?
• 700-380 nanometers
Electromagnetic Spectrum
• Visible light is only a tiny portion of the electromagnetic spectrum which also includes radio waves, microwaves, infrared, visible light, ultra violet, X-rays, and gamma rays.
• If the entire electromagnetic spectrum was a strip of professional 16 mm movie film stretching from Los Angeles to Seattle, the portion of visible light would be only ONE frame of film.
Atomic Spectra
• When atoms absorb energy, electrons move to higher energy levels
• Electrons then lose energy by emitting light as they return to lower energy levels
• Atoms emit only specific frequencies of light that correspond to the energy levels in the atom
• The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element
Atomic Spectra
• An electron with its lowest possible energy is in its ground state
• The light emitted by an electron is directly proportional to the energy change of the electron.
• E = hν
• Atomic spectra are like fingerprints: no two are alike!