Chapter 6 Atomic Theory

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Chapter 6Atomic Theory
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Greek Idea

• Democritus
• Matter is made up of indivisible particles
• Dalton
• One type of atom for each element

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Thomsons Model

• Discovered electrons
• Atoms were made of positive stuff
• Negative electron embedded inside the atom
• Plum-Pudding model
• Blueberry muffin model

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Rutherfords Model

• Discovered dense positive piece at the center of
  the atom
• Nucleus
• Electrons moved around
• Mostly empty space

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Bohrs Model

• Why dont the electrons fall into the nucleus?
• Electrons move like planets around the sun
• In circular orbits at different levels
• Amounts of energy separate one level from another

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Bohrs Model
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Bohrs Model

• Further away from the nucleus means more energy
• There is no in between energy

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Bohrs Model
The Bohr model shows electrons circling around
the nucleus in definite orbits or paths. The
electrons orbits the nucleus much like planets
circle the sun. Electrons move from one orbit to
another. The further away the orbit from the
nucleus means more energy. The electrons can not
move in between orbits electrons can only
exist in an orbit at definite energy levels. The
energy absorbed or released when electrons change
energy levels is in the form of electromagnetic
radiation (light).
Energy Levels
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Energy

• When electrons jump orbits (move up or down)
  energy is needed or released respectively
• Energy is in the form of light
• It was through the study of light that several
  advances in science were made
• Physicists determined the properties of light
• The Quantum Mechanical model was produced

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Electromagnetic Radiation

• Light is just one component of the
  Electromagnetic Spectrum
• Composed of waves of different energies
• Waves travel through empty space as well as
  through air and other substances
• Electromagnetic radiation has a dual
  "personality
• Acts like waves and particles (photons)
• The photons with the highest energy correspond
  to the shortest wavelengths

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Electromagnetic Spectrum
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Electromagnetic Radiation
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Parts of the Wave
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Parts of a Wave
Crest

• Origin
• the base-line of energy
• Crest
• Highest point on a wave
• Trough
• Lowest point on a wave
• Amplitude
• Distance from origin to crest
• Wavelength
• Distance from crest to crest (abbreviated l -
  lambda)
• Frequency
• The number of waves passing a given point per
  second
• Units are cycles/sec or hertz (Hz)

Wavelength
Origin
Amplitude
Trough
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Frequency and the Spectrum

• Frequency and wavelength are inversely related
• n C
• l
• Frequency speed of light / wavelength
• Different frequencies of light have different
  energy levels
• E h? (Energy Plancks constant x frequency)
• High frequency light has high energy (violet
  light)
• Low frequency light has low energy (red light)
• White light is made up of all the colors of the
  visible spectrum, and all frequencies therein
• The whole range of colours is called a
  continuous spectrum one colour leads into the
  next with no break

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And now back to the atom

• Electrons occupy the lowest energy levels,
  making the atom stable (low energy content
  ground state)
• When electrons interact with energy (photon)
  they may absorb it and move away from the nucleus
• Referred to as an electron transition
• The electrons are no longer in the ground state
  they are in an excited state
• The electrons can return to the ground state by
  releasing quanta of energy
• Energy released is of a definite quantity
• The energy of electron transitions are quantized
  (fixed), quantum of energy is released

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Catch and Release

• Quantized energy is absorbed
• Electron is excited into a higher energy level

• Quantized energy is emitted (released)
• Electron jumps to a lower energy level

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Bohr Model
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Emission Spectra

• Light of different frequencies have different
  characteristic colours
• Low energy, low frequency light is seen as red
  light high frequency, high energy light is seen
  as violet light
• All possible jumps from one energy level to
  another can occur at the same time because of the
  number of atoms present, each giving off its
  characteristic colour
• The colour that we see is a mixture of colours
• Can be separated by using a diffraction grating
• Breaks up the predominant colour into bright
  lines of specific colours representing electron
  transitions

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Emission Spectra
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Hydrogen Emission Spectrum
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Iron Emission Spectrum
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The Quantum Mechanical Model

• Bohrs model of the atom introduced the concept
  of quantum energy levels, but couldnt explain
  how electrons are arranged in atoms
• Louis de Broglie (1892-1987)
• Proposed that if waves can have particle-like
  behaviour, particles of matter can behave like
  waves under appropriate conditions
• Suggested that as an electron moves about the
  nucleus, an appropriate wavelength is associated
  with it
• Just a few years later, experimental evidence
  supported de Broglie

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The Quantum Mechanical Model

• Werner Heisenberg (1901-1976)
• Limit to how precisely we can know both location
  and momentum of any object only important with
  subatomic particles
• Impossible to know both exact momentum and exact
  location of electron at any point in time
• Heisenberg Uncertainty Principle
• Erwin Schröndinger (1887-1961)
• Proposed an equation which leads to a series of
  wave functions
• Related the probability of finding the electrons
  to a particular volume of space (3D)
• Schröndingers wave equation

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The Plot for atomic orbital

• The orbital is classified by its shape.
• S orbital a sphere.
• P orbital is a dumbbell shape.

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Atomic Orbitals

• Represents the likelihood of finding an electron
  at a particular point
• Densest near the nucleus less dense with
  increasing distance
• Indicates most probable location around the
  nucleus
• Scientists arbitrarily draw the orbital surface
  to contain 90 of the total probability
  distribution
• To assign relative sizes and energies to
  orbitals, the quantum mechanical model assigns
  Principal Quantum Numbers (n)
• n specifies the atoms major energy levels
  Principal Energy Levels

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Atomic Orbitals

• Within each energy level the complex math of
  Schrödingers equation describes several shapes
• Each type of shape is identified as a sublevel
  containing a specific atomic orbital
• Energy level 1 (n1) has one sublevel that
  contains one s orbital
• Every succeeding energy level contains an s
  sublevel with an s orbital
• The s orbital is spherically shaped
• Each s sublevel is identified by its principal
  quantum number n
• Called 1s, 2s, 3s,
• Each s sublevel has one s type orbital, and can
  hold 2 electrons

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Atomic Orbitals

• Energy level 2 (n2) has 2 sublevels, the s and
  p sublevel (The sublevel name comes from the type
  of orbital it contains.)
• Every succeeding energy level contains a p
  sublevel with three p orbitals
• The p orbital is shaped like a dumbbell and
  labelled depending on its orientation
• One p orbital is oriented along the x axis (px)
• One p orbital is oriented along the y axis (py)
• One p orbital is oriented along the z axis (pz)
• Each p orbital can hold 2 electrons, so the p
  sublevel can hold a maximum of six electrons

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Atomic Orbitals

• The third energy level (n3) has three sublevels
• s sublevel with one s orbital
• p sublevel with three p orbitals
• d sublevel with five d orbitals
• The d orbitals are more complex in shape and are
  oriented along planes not axes
• Each of the five orbitals can hold two electrons
  so the d sublevel can hold a maximum of ten
  electrons

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Atomic Orbitals

• The fourth energy level (n4) has four sublevels
• s sublevel with one s orbital
• p sublevel with three p orbitals
• d sublevel with five d orbitals
• f sublevel with seven f orbitals
• The f orbital are even more complex in shape
• Each of the seven orbitals can hold two
  electrons so the f sublevel can hold a maximum of
  fourteen electrons

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Orbitals Summary

• Atomic orbital pictures
• Atomic orbital density representations

Principal quantum number (n) Number of sublevels Number of orbitals Total number of electrons
n n n2 2n2
1 1 (s) 1 2
2 2 (s, p) 4 8
3 3 (s, p, d) 9 18
4 4 (s, p, d, f) 16 32
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Electron Configuration

• The Electron Configuration is the arrangement of
  the electrons within an atom
• The atomic orbitals do not fill up in a neat
  order the energy levels overlap
• Large complex orbital shapes cause the electron
  to be, on average, further from the nucleus than
  the simple orbital shapes from greater energy
  levels
• Arrangements follow specific principles
• aufbau principle
• Electrons enter the lowest available energy
  level or sublevel first
• This causes difficulties because of the overlap
  of orbitals of different energies

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Electron Configuration

• Pauli Exclusion Principle
• At most 2 electrons per orbital
• Electrons have the same charge and will repel
  each other
• The repulsion can be reduced if the electrons
  have different spins (paired spins) leading to
  opposite magnetic force fields
• Hunds Rule
• When electrons occupy orbitals of equal energy
  (orbitals within the same sublevel) they dont
  pair up until the sublevel is half full

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n 4 With more sublevels the energy
level n 3 gets
wider (leads to an overlap). n 2 n 1
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Electron Configuration
Electron Configuration Electron Configuration Electron Configuration Electron Configuration Electron Configuration
7p 6d 5f ?
7s 6p 5d 4f ?
6s 5p 4d ?
5s 4p 3d ?
4s 3p ?
3s 2p ?
2s ?
1s ? Start Start Start
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Determining Electron Configuration

• Lets determine the electron configuration for
  phosphorus
• Need to account for 15 electrons
• The first two electrons go into the 1s orbital
  (1s2)
• Remember electrons have opposite spins
• The next two electrons go into the 2s orbital
  (2s2)
• The next six electrons go to the 2p sublevel
  (2p6)
• 3 paired spins.
• Next is the 3s orbital with 2 electrons (3s2)
• Total is now 12, only 3 more to go
• The next three electrons go 3px1, 3py1, 3pz1
• The Electron Configuration for phosphorus is
• 1s2 2s2 2p6 3s2 3px1, 3py1, 3pz1
• 1s2 2s2 2p6 3s2 3p3 (not showing the individual p
  orbitals)

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Determining Electron Configuration

• Try for
• Na
• C
• Cl
• Start off with How many electrons to place?
• Add electrons to the orbitals starting at 1s and
  following through the table

Electron Configuration Electron Configuration Electron Configuration Electron Configuration Electron Configuration
7p 6d 5f ?
7s 6p 5d 4f ?
6s 5p 4d ?
5s 4p 3d ?
4s 3p ?
3s 2p ?
2s ?
1s ? Start Start Start
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Valence Shell and Valence Electrons

• The highest numbered energy level is given a
  special name the valence shell
• The electrons in this valence shell are referred
  to as valence electrons
• Only the outer s, and p electrons

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Valence Shell and Valence Electrons
Electron Configuration Electron Configuration Electron Configuration Electron Configuration Electron Configuration
7p 6d 5f ?
7s 6p 5d 4f ?
6s 5p 4d ?
5s 4p 3d ?
4s 3p ?
3s 2p ?
2s ?
1s ? Start Start Start

• As electrons fill the sublevels, we see
• 1s2 2s2 2p6 3s2 all fill up first
• 3p6 4s2 ? now were in the 4th level, so we
  start in a new valence shell
• 3d10 4p6 5s2 ? again, a new shell
• The d and f electrons never see valence
  because a new level is started before theyre
  filled

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Valence Shell and Valence Electrons

• The highest numbered energy level is given a
  special name the valence shell
• The electrons in this valence shell are referred
  to as valence electrons
• Only the outer s, and p electrons
• Identify each element, state the valence shell
  and the number of valence electrons for
• 1s2 2s2 2p6
• 1s2 2s2 2p6 3s2 3p6 4s1
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

Ne K Ge Sr
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Exceptions to the Electron Configuration

• Orbitals fill in order of lowest to highest
  energy
• The most stable sublevel or energy level is
  full
• The next most stable is a half full
• Adding electrons can change the energy of the
  orbital or sublevel
• More than half filled sublevels or less than
  half filled sublevels are higher in energy than
  those that are exactly half full
• The most stable is a full sublevel (least
  energy)
• Sometimes electrons will be promoted/demoted to
  help the atom stay in the lowest energy state
  (most stable) possible

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Electron Exceptions

• Copper has 29 electrons so we expect
  1s22s22p63s23p64s23d9
• The electron configuration for the copper is
  also 1s22s22p63s23p64s13d10
• One 4s electron has been promoted to a 3d
  orbital, leaving a valence shell with only 1
  electron and a full 3d
• Copper exists in both states
• 1s22s22p63s23p64s23d9
• 1s22s22p63s23p64s13d10

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Copper Ions

• Because copper can exist in both states, it can
  form 2 different cations
• Cu2
• 1s22s22p63s23p64s03d9 4s2 electrons are lost
• Cu
• 1s22s22p63s23p64s03d10 4s1 electron is lost
• Many transition metals can form multiple cations

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Electron Exceptions

• Iron with 26 electrons has an electron
  configuration of 1s22s22p63s23p64s23d6
• When reacting, it can lose the 2 valence 4s2
  electrons creating a 2 cation
• 1s22s22p63s23p64s03d6
• Iron can also lose one 3d electron in addition,
  to yield a half full 3d sublevel 3d5, creating a
  3 cation
• 1s22s22p63s23p64s03d5
• most stable

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Silver (Ag)

• Electron configuration should be
  1s22s22p63s23p64s23d104p65s24d9
• Its actually 1s22s22p63s23p64s23d104p65s14d10
• Results in a lower overall energy, and higher
  stability
• Full d sublevel partially filled s sublevel