Chapter 4 Electron Configurations. Early thoughts Much understanding of electron behavior comes from...
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![Page 1: Chapter 4 Electron Configurations. Early thoughts Much understanding of electron behavior comes from studies of how light interacts with matter. Early.](https://reader036.fdocuments.us/reader036/viewer/2022062422/56649eff5503460f94c15552/html5/thumbnails/1.jpg)
Chapter 4 Electron Configurations
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Early thoughts
•Much understanding of electron behavior comes from studies of how light interacts with matter.
• Early belief was that light traveled as a wave, but some experiments showed behavior to be as a stream of tiny fast moving particles
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Waves
• Today scientists recognize light has properties of waves and particles
• Waves: light is electromagnetic radiation and travels in electromagnetic waves.
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4 Characteristics of a wave:• 1) amplitude - height of the wave. For
light it is the brightness• 2) Wavelength ()– distance from crest
to crest. • For light – defines the type of light • Visible light range – 400-
750nm
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Properties continued• 3) Frequency ()– measures how fast
the wave oscillates up and down. • It is measured in number per
second. • Hertz = 1 cycle per second• FM radio = 93.1 MHz or 93.1 x 106
cycles per second• Visible light = 4 x 1014 cycles per second
to 7 x 1014 cycles per second • 4) speed – 3.00 x 108 m/s
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• Short wavelength, high frequency
• Long wavelength, low frequency
• Visible Spectrum• ROY G BIV• Longer wavelength shorter wavelength
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Electromagnetic spectrum (meters)
• 10-11 gamma• 10-9 x-rays• 10-8 UV• 10-7 visible light• 10-6 infrared• 10-2 microwave• 1 TV
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Wavelength and frequency
• Wavelength and frequency are inversely related!!
= c/• Where is the
wavelength, c is the speed of light and is the frequency
• Speed of light = Constant = • 3.00 x 108m/sec
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Example
• Example: An infrared light has a wavelength of 2.42 x 10-6m. Calculate the frequency of this light.
• = c/ = 3.0 x 108m/sec =• 2.42 x 10-6m• = 1.2 x 1014 waves/sec
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Wavelength and frequency
•****Remember and are inverse. Therefore short wavelength = high frequency!!
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Quantum Theory
•Needed to explain why certain elements when heated give off a characteristic light (certain color)
•1900- Max Planck – idea of quantum
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Quantum Theory
• - The amount of energy (electromagnetic radiation) an object absorbs/emits occurs only in fixed amounts called quanta (quantum)
• - Quanta – finite amount of energy that can be gained or lost by an atom.
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1905 Einstein’s theory•Einstein explained photoelectric
effect by proposing that light consists of quanta of energy called PHOTONS
• Photon = discrete bit of energy
•Consider light traveling as photons
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Energy equation
• Amount of energy of a photon described as
• E = h• Where • E = energy• = frequency• h = Planck’s constant = 6.6262 x
10-34 J s• Joule = SI unit for energy
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Photoelectric effect –
• electrons are ejected from the surface of a metal when light shines on the metal. The frequency determines the amount of energy. The higher the frequency, the more energy per photon.
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Photoelectric effect
•Ex. X-rays have a high frequency; therefore can damage organisms while radio waves have a low frequency.
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•Dual nature of radiant energy
• Photons act both like particles and waves.
•Line Spectra: A spectrum that contains only certain colors, or wavelengths
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Question: How are electrons arranged in atoms?
• Note: All elements emit light when they are vaporized in an intense flame or when electricity is passed through their gaseous state.
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How are electrons arranged in atoms?• Explanation: Bohr atom: 1911• postulated that atoms have energy
levels in which the electrons orbit• energy levels and/or orbits are labeled
by a quantum number, n. • lowest energy level = ground state n=1
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How are electrons arranged in atoms?• when an electron absorbs energy, it
jumps to a higher level (known as the excited state) n = 2,3 or 4
• Bohr model of an atom • Hydrogen • Red- e falls from 3 to 2• Blue e falls from 4 to 2• Violet e falls from 6 to 2
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1924 – Louis DeBroglie
•– If waves of light can act as a particle, then particles of matter should act like a wave. Found to be true.
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DeBroglie
• Matter waves = wavelike behavior of particles.
• From equation – wave nature is inversely related to mass therefore we don’t notice wave nature of large objects. However, electrons have a small mass so they have a larger wave characteristic
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Schroedinger’s wave equation
•predicted probability of finding an electron in the electron cloud around nucleus. Gave us four quantum numbers to describe the position.
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Heisenberg’s Uncertainty Principle
The position and momentum of a moving object cannot simultaneously be measured and known exactly.
• Cannot know where it is and where its going at the same time.
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• Quantum mechanical model of an atom –
• Treats the electrons as a wave that has quantized its energy
• Describes the probability that electrons will be found in certain locations around the nucleus.
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Orbitals
• Orbitals – atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found (high probability)
• Have characteristic shapes, sizes and energies.
• Four different kinds of orbitals s, p, d, f.
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S and p orbitals
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Orbitals and energy• Principal Energy Level =
quantum number n. • Principal Quantum
number • Value of n =
1,2,3,4,5,6,7• Tells you the distance
from the nucleus
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sublevels
• Each level is divided into sublevels• # of sublevels = value of n• N=1 1 sublevel• n=2 2 sublevels • s,p,d or f shows the shape
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Orbital shape
• Each sublevel has a certain number of orbitals which are directed three dimensionally
• S = 1 sphere• P = 3 figure 8 (along the x,y or
z axis)• D = 5 figure 4-27 p. 145• F = 7
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• Each electron in an orbital will have a spin – 2 options clockwise, vs. counter clockwise.
• Electron configuration• Pauli Exclusion Principle – each orbital in
an atom can hold at most 2 electrons and their electrons must have opposite spin.
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•Aufbau Principle- electrons are added one at a time to the lowest energy orbitals available
•Hund’s Rule – electrons occupy equal energy orbitals so that the maximum number or unpaired electrons result.
•Occupy singly before pairing
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• Diagonal Rule: for order of sublevels:
• must remember 1s, 2s, 2p, 3s, 3p, 4s
• Exceptions to Aufbau Principle• Cr Z=24 • Cu Z = 29
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Energy levels
• Number of electrons per energy level
• 1st = 2• 2nd = 8• 3rd = 18• 4th = 32
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• Number of electrons per orbital• 2• Difference between paired and
unpaired electrons• Paired = 2 electrons in the
same orbital• Unpaired = 1 electron in the
orbital