Chemistry Chapter 4

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Chemistry Chapter 4

• Arrangement of Electrons in Atoms

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Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the
electron as a particle.
His son, George Thomson won the Nobel prize for
describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!
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• Properties of Light as a wave
• Light is called electromagnetic radiation
• Electromagnetic radiation form of energy which
  exhibits wave-like behavior as it goes through
  space
• Electromagnetic spectrum continuum of all of
  the forms of electromagnetic radiation
• Light has the characteristics of waves
• Wavelength distance between two crests or two
  troughs (symbol - ?, unit nm)
• Frequency the numbers of waves which passes a
  certain point in a certain amount of time (symbol
  f, unit Hz (Hertz))
• All electromagnetic radiation has the same
  velocity which is 3.0 X 108 m/s (speed of light
  c)
• Therefore, c ? x f indirect relationship

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Electromagnetic radiation moves through space as
a wave moving at the speed of light.
c ?f
C speed of light, a constant (3.00 x 108 m/s)
f frequency, in units of hertz (hz, waves/sec)
? wavelength, in nanometers (lambda)
As frequency increases, wavelength decreases and
vice versa.
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Wavelength Table
Long Wavelength Low Frequency Low ENERGY
Short Wavelength High Frequency High ENERGY
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Types of electromagnetic radiation
? Frequency increases
Wavelength increases ?
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• The Photoelectric Effect (Light as a Particle)
• Certain interactions between light and matter
  could not be explained by the wave theory of
  light
• 1900, Max Planck performed experiments with the
  light coming off hot objects which showed us that
  light does not emit energy continuously as would
  be expected from a wave and proposed
• E h x f (energy of a quantum of
  electromagnetic radiation)
• (h Plancks constant 6.626 X 10-34 J-s)
• Planck said that light emitted energy in small
  amounts called quanta
• Quantum min. quantity of energy that can be
  lost or gained by an atom
• 1905, Albert Einstein showed us that the
  Photoelectric Effect revealed that light was
  indeed not just a wave, but was a particle of
  energy with a particle frequency which determined
  the energy of that particle

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• Electromagnetic radiation was now not just a wave
  but also a particle (wave-particle duality)
• Photon a particle of light (electromagnetic
  radiation) that has zero mass and the energy of a
  quantum
• Ephoton h x f (energy of a photon) same as a
  quantum
• Photons can only be absorbed in whole number
  ratios basically all or none
• In order for an electron to be ejected from the
  surface of a metal, one photon alone must have
  the min. energy to knock the electron loose.

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The energy (E ) of electromagnetic radiation is
directly proportional to the frequency (?) of the
radiation.
E hf
E Energy, in units of Joules (kgm2/s2)
h Plancks constant (6.626 x 10-34 Js)
f frequency, in units of hertz (hz, waves/sec)
The high the frequency of light or EM radiation,
the greater the energy of the waves.
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• Hydrogen Atom Line-Emission Spectra
• Ground state lowest energy state of an atom
• Excited state energy level of higher potential
  energy
• When electricity is passed through a gas at low
  pressure, the electrons in the atoms move to
  excited states. When they return to the ground
  state, the excess energy is given off a light
  (ER) in the form of a photon
• When this light was sent through a prism, the
  scientists expected a continuous spectrum, but
  only got certain distinct wavelengths of light
  which made up the emitted light.
• This revealed that the energy states of a
  hydrogen electron must be distinct energy levels
  (specific values for the hydrogen atom)

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Spectroscopic analysis of the visible spectrum
white light produces all of the colors in a
continuous spectrum
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Spectroscopic analysis of the hydrogen spectrum
produces a bright line spectrum or emission
spectrum
Each line in bright line spectrum represents a
different transition of electrons from an excited
state to the ground state
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Electron transitionsinvolve jumps of definite
amounts ofenergy.
This produces bands of light with
definite Wavelengths each bands represents a
different transition from a higher energy level
to n2 (photons are in the visible region)
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The Bohr Model of the Atom
I pictured electrons orbiting the nucleus much
like planets orbiting the sun.
Niels Bohr

• Bohr's Hydrogen atom has an electron which runs
  like a train on isolated circular tracks.
• Photons are emitted or absorbed when the electron
  jumps from one track to another.

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• Bohr Model of the Atom
• 1913, Niels Bohr proposed that the electrons of
  an hydrogen atom have distinct energy states
  called orbits
• Electrons in these orbits have definite, fixed
  energies.
• Lowest energy level is closest to the nucleus
  lowest potential energy
• Energy levels are like the rungs of a ladder
  you cannot stand in between the rungs
• Each levels corresponds to a specific potential
  energy
• Depending on the amount of energy absorbed by the
  atom, this would determine to what excited state
  that the electron would jump.
• The difference between the higher state and the
  ground state determined the energy, and
  therefore, the frequency of the photons given
  off.
• Bohrs Model did not explain the spectra for
  anything but hydrogen and did not explain the
  chemical behavior of atoms.

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The Wave-like Electron
The electron moves through space as an energy
wave. To understand the atom, one must understand
the behavior of electromagnetic waves.
Louis deBroglie
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The electron acts as a standing wave

• Standing waves do not move through space
• Standing waves are fixed at both ends like the
  strings on a guitar

Only certain sized orbits can contain the
electrons standing waves whole number energy
levels
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Wave-like Electrons
Electrons are confined to certain spaces in the
atom called orbitals Electrons travel around
these spaces in a wave-like pattern.
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Heisenberg Uncertainty Principle
One cannot simultaneously determine both the
position and momentum/velocity of an electron.
You can find out where the electron is, but not
where it is going.
OR
You can find out where the electron is going, but
not where it is!
Werner Heisenberg
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Schrodinger Wave Equation

• Electrons exist in regions called orbitals 3D
  region around the nucleus that indicates the
  probable location of an electron
• His equation only provide the probability of
  finding an electron in a certain area.

Erwin Schrodinger
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• Electrons as Waves
• Two things contributed to electrons being
  compared to waves
• Photoelectric Effect
• Hydrogen Line-Emission Spectra
• de Broglie said that other types of waves
  confined to a certain are like sound in an organ
  pipe have only certain frequencies - harmonics
• We can consider the electron being confined to
  the area around the nucleus therefore,
  electrons can only exist in an atom at certain
  frequencies or energies by E h x f
• de Broglie confirmed the atom a wave by showing
  the following characteristics in electrons
• diffraction the bending of a wave around an
  object or field of energy
• Interference the constructive and destructive
  combination of more than one wave pattern
  resulting in decreases of energy in some areas
  and increases in others.
• Heisenberg (1927) involved in trying to detect
  the location of an electron at a certain point in
  time
• He discovered that it is impossible to determine
  both the position and the velocity of an electron
  or any other particle Heisenberg Uncertainty
  Principle
• Schrodinger (1926) developed an equation that
  explained the wave-particle behavior of the
  electron which led to what is now called Quantum
  Theory of the Atom
• His equation provided only the probability of
  finding an electron in a certain area.
• Electrons do not travel in neat orbits as Bohrs
  Model states instead they exist in regions
  called atomic orbitals
• Orbital a 3D region around the nucleus that
  indicates the probable location of an electron

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Quantum Numbers
Each electron in an atom has a unique set of 4
quantum numbers which describe its energy,
location, orientation, and spin in the atom.

• Principal quantum number energy level
• Angular momentum quantum number
• location/type of orbital
• Magnetic quantum number orientation
• Spin quantum number type of spin
• (/- ½)

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Pauli Exclusion Principle
No two electrons in an atom can have the same
four quantum numbers.
Each electron has a unique combination of quantum
numbers describing its energy level, type of
orbital, orientation, and spin.
Wolfgang Pauli
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Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
Ex. n 1 , 1st energy level, lowest energy
Number of electrons that can fit in a shell
2n2
Number of orbitals that can fit in a shell
n2
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Principal Quantum Number

• Principal Quantum Number (n) indicates the
  main energy level occupied by the electron (n
  1, represents the lowest energy level)
  (1,2,3,4,)
• More than one electron can have the same n value
  they are said to be in the same electron shell
• Total of orbitals in an energy level is equal
  to n2
• 1st level 1 orbital
• 2nd level 4 orbitals
• 3rd level 9 orbitals

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Angular Momentum Quantum Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell or
sublevel) in which the electron is located.
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Angular Momentum Quantum Number

• Angular Momentum Quantum Number - (l)
  indicates the shape of each orbital called
  sublevels (0, 1, 2, 3, n-1)
  Sublevel
• l 0 s
• 1 p
• 2 d
• 3 f

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Magnetic Quantum Number (ml )
The magnetic quantum number, generally symbolized
by ml, denotes the orientation of the electrons
orbital with respect to the three axes in space.
(x, y, and z axes 3D coordinates)
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Magnetic Quantum Number (ml)

• Magnetic Quantum Number (ml) indicates the
  orientation of the orbital around the nucleus
• s m 0 (centered around nucleus)
• p m -1, 0, 1 (3 orbitals)
• d m -2, -1, 0, 1, 2 (5 orbitals)
• f m -3, -2, -1, 0, 1, 2, 3 (7
  orbitals)
• Each orientation of each orbital type can hold
  two electrons so therefore
• s 2 electrons
• p 6 electrons (3 orientations)
• d 10 electrons (5 orientations)
• f 14 electrons (7 orientations)

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Spin Quantum Number (ms)
Spin quantum number (ms) denotes the behavior
(direction of spin) of an electron within a
magnetic field.
Possibilities for electron spin
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Assigning the Numbers

• The first three quantum numbers (n, l, and m) are
  integers.
• The principal quantum number (n) cannot be zero
  n must be 1, 2, 3, etc.
• The angular momentum quantum number (l) can be
  any integer between 0 and n - 1.
• For n 3, l can be either 0, 1, or 2.
• The magnetic quantum number (m) can be any
  integer between -l and l.
• For l 2, m can be either -2, -1, 0, 1, or 2.

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Principle, angular momentum, and magnetic quantum
numbers n, l, and ml
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s orbital shape
S orbital shape
The s orbital has a spherical shape centered
around the origin of the three axes in space.
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Sizes of s orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases
Nodes are regions where electrons are not likely
to exist.
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P orbital shapes
There are three dumbbell-shaped p orbitals in
each energy level except n1, each assigned to
its own axis or has its own orientation (x, y
and z) in space.
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d orbital shapes
D orbital shapes
Things get a bit more complicated with the five d
orbitals that are found in the d sublevels
beginning with n 3. To remember the shapes,
think of double dumbbells and a dumbbell
with a donut!

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F Orbital Shapes
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Types of Electron Configuration Notation

• Electron Configuration Notation
• Ex. F 1s2 2s2 2p5
• Coefficient represents the energy level
• Letter type of orbital
• Superscript - of electrons in that orbital
• Orbital Notation
• Ex. F ?? ?? ?? ?? ? .
• 1s 2s 2px2py2pz
• Noble Gas Notation
• Ex. F He 2s2 2p5
• In brackets, show the Noble Gas from the row
  above your element
• Then, show only the notation for the period in
  which the element is located only the valence
  electrons
• Valence electrons electrons in the outside
  energy level

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Rules in Writing Electron Configurations

• Aufbau Principle an electron occupies the
  lowest energy orbital that can receive it.
• Pauli Exclusion Principle no two electrons have
  the same set of quantum s
• Each orbital can hold 2 electrons (one of each
  spin)
• Hunds Rule Orbitals of equal energy are each
  occupied by one electron before any orbital is
  occupied by a second electron, and all electrons
  in orbitals with one electron must have the same
  spin.
• Ex. ? ? ? ? ? ??
  ? ? .
• 2p 2p
  2p

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Rules in Writing Electron Configurations

• What is the Atomic of the element?
• Ex. N 7 electrons
• Fill the orbitals in order following the rules
  noted previously.
• Ex. N - 1s2 2s2 3p3
• Count to make sure you have used all the
  electrons

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Orbital filling table
In the 4th period, you can see that the 4s
orbital is filled before the 3d. D orbitals are
higher in energy than the s orbitals. D orbitals
are always filled one period below (3d in 4th
period)
Noble gases have a full outside energy level (no
empty of half full orbitals)
The of valence electrons is the same as the
Group of the element (Transition Elements all
have 2 valence electrons)
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Electron configuration of the elements of the
first three series
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Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to half fill its 3d
sublevel
Copper steals a 4s electron to FILL its 3d
sublevel