Electronic structure

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Electronic structure
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Atomic structure
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Atoms neutral

• Protons have a positive charge and electrons have
  a negative charge.
• Atoms are neutral so they have the same number
  of electrons as protons. 
• The protons are in the nucleus and do not change
  or vary except in some nuclear reactions.
• The electrons are in discrete pathways or shells
  around the nucleus.

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Shells - Orbitals
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ELECTRONIC STRUCTURE AND ATOMIC ORBITALS

• In any introductory chemistry course you will
  have come across the electronic structures of
  hydrogen and carbon drawn as

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Hydrogen's electron - the 1s orbital
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Names of orbitals

• Each orbital has a name.
• The orbital occupied by the hydrogen electron is
  called a 1s orbital. The "1" represents the fact
  that the orbital is in the energy level closest
  to the nucleus.
• The "s" tells you about the shape of the
  orbital. s orbitals are spherically symmetric
  around the nucleus - in each case, like a hollow
  ball with the nucleus at its centre.

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2s orbitals
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A p orbital is rather like 2 identical balloons
tied together at the nucleus. The diagram on the
right is a cross-section through that
3-dimensional region of space. Once again, the
orbital shows where there is a 95 chance of
finding a particular electron.
P orbitals
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P orbitals

• At any one energy level it is possible to have
  three absolutely equivalent p orbitals pointing
  mutually at right angles to each other. These are
  arbitrarily given the symbols px, py and pz. This
  is simply for convenience - what you might think
  of as the x, y or z direction changes constantly
  as the atom tumbles in space.

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The electronic structure of hydrogen

• Hydrogen has only one electron and that will go
  into the orbital with the lowest energy - the 1s
  orbital.
• Hydrogen has an electronic structure of 1s1. We
  have already described this orbital earlier.
• The electron of a hydrogen atom travels around
  the proton nucleus in a shell of a spherical
  shape

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Helium

• Helium, element number two, has two electrons in
  the same spherical shape around the nucleus.
• The first shell only has one subshell, and that
  subshell has only one orbital, for electrons.
• Each orbital has a place for two electrons. The
  spherical shape of the lone orbital in the first
  energy level has given it the name s orbital.

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The electronic structure of carbon

• Carbon has six electrons. Two of them will be
  found in the 1s orbital close to the nucleus. The
  next two will go into the 2s orbital.
• The remaining ones will be in two separate 2p
  orbitals. This is because the p orbitals all have
  the same energy and the electrons prefer to be on
  their own if that's the case.
• The electronic structure of carbon is normally
  written 1s22s22p2

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Carbon
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Carbon

•  Name Carbon
• Atomic Number 6
• Chemical Symbol C
• Electronic Configuration1s22s22p2
• Atomic Mass 12.0
• Atomic Ionization Energies
• 2p 1086 kJ mol-1
• 2s 1601 kJ mol-1
• 1s 27788 kJ mol-1  

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Orbitals, Shells

• Physicists and chemists use a standard notation
  to describe the electron configurations of atoms
  and molecules. For atoms, we use atomic orbital
  labels (eg, 1s, 3d, 4f) with the number of
  electrons assigned to each orbital placed as a
  superscript. For example
• hydrogen has one electron in the s-orbital of the
  first shell, so its configuration is written 1s1.
  
• Lithium has two electrons in the 1s-subshell and
  one in the (higher-energy) 2s-subshell, so its
  configuration is written 1s2 2s1 (pronounced
  "one-ess-two, two-ess-one").
• Phosphorus (atomic number 15), is as follows
  1s2 2s2 2p6 3s2 3p3.

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Electronic configuration
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Electronic configuration
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Electronic configuration
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Shells

• The shells or energy levels are numbered or
  lettered, beginning with K.
• So K is one, L is two, M is three, N is four, O
  is five, P is six, and Q is seven.
• As the s shells can only have two electrons and
  the p shells can only have six electrons, the d
  shells can have only ten electrons and the f
  shells can have only fourteen electrons.

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Shells - Orbitals
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Shells

• The electron shells are labelled K, L, M, N, O,
  P, and Q or 1, 2, 3, 4, 5, 6, and 7 going from
  innermost shell outwards.
• Electrons in outer shells have higher average
  energy and travel further from the nucleus than
  those in inner shells, making them more important
  in determining how the atom reacts chemically and
  behaves as a conductor, etc, because the pull of
  the atom's nucleus upon them is weaker and more
  easily broken.

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Subshells

• Each shell is composed of one or more subshells,
  which are themselves composed of atomic orbitals.
  For example, the first (K) shell has one
  subshell, called "1s" the second (L) shell has
  two subshells, called "2s" and "2p" the third
  shell has "3s", "3p", and "3d" and so on.3 The
  various possible subshells are shown in the
  following table

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Orbitals, Shells,subshells
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Orbitals, Shells, subshells
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Shells and subshells
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Shells and subshells
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Number of electrons in each shell

• Each s subshell holds no more than two electrons
• Each p subshell holds no more than six electrons
• Each d subshell holds no more than ten electrons
• Each f subshell holds no more than fourteen
  electrons
• Therefore, the K shell, which contains only an s
  subshell, can hold up to 2 electrons the L
  shell, which contains an s and a p, can hold up
  to 268 electrons and so forth. The general
  formula is that the nth shell can in principle
  hold up to 2n2 electrons.

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Tin Sn
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spdf notation

• Electronic Configuration 1s22s22p63s23p64s23d104s
  1

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Spdf notation
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Electrons in boxes
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The order of filling orbitals
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Orbitals shellsEnergy levels

• The circles show energy levels - representing
  increasing distances from the nucleus. You could
  straighten the circles out and draw the
  electronic structure as a simple energy diagram

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Energy levels
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Energy levels
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Ground state

• The electrons are housed in shells around the
  nucleus. There is a ranking or hierarchy of the
  shells, usually with the shells further from the
  nucleus having a higher energy.
• As we consider the electron configuration of
  atoms, we will be describing the ground state
  position of the electrons.
• Place electrons in lower energy shells.
• Start from shell nearest to the nucleus, i.e 1s2

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Order of filling atomic orbitals
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Order of filling orbitals
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Ground state

• Electrons fill low energy orbitals (closer to the
  nucleus) before they fill higher energy ones.
  Where there is a choice between orbitals of equal
  energy, they fill the orbitals singly as far as
  possible.
• The diagram (not to scale) summarises the
  energies of the various orbitals in the first and
  second levels.
•  
• Notice that the 2s orbital has a slightly lower
  energy than the 2p orbitals. That means that the
  2s orbital will fill with electrons before the 2p
  orbitals. All the 2p orbitals have exactly the
  same energy

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Quantum

• Electrons are able to move from one energy level
  to another by emission or absorption of a quantum
  of energy, in the form of a photon. Because of
  the Pauli exclusion principle, no more than two
  electrons may exist in a given atomic orbital
  therefore an electron may only leap to another
  orbital if there is a vacancy there. The word
  comes from the Latin "quantus," for "how much.

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Principal quantum number

• An electron shell is the set of atomic orbitals
  which share the same principal quantum number, n
  (the number before the letter in the orbital
  label) hence the 3s-orbital, the 3p-orbitals and
  the 3d-orbitals all form part of the third shell.
  
• An electron shell can accommodate 2n2 electrons,
  ie the first shell can accommodate 2 electrons,
  the second shell 8 electrons, the third shell
  18 electrons, etc.

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The periodic table

• The form of the periodic table is closely related
  to the electron configuration of the atoms of the
  elements. For example, all the elements of group
  2 have an electron configuration of E ns2
  (where E is an inert gas configuration), and
  have notable similarities in their chemical
  properties.
• The outermost electron shell is often referred to
  as the "valence shell" and determines the
  chemical properties. It should be remembered that
  the similarities in the chemical properties were
  remarked more than a century before the idea of
  electron configuration.

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History

• The existence of electron shells was first
  observed experimentally in X-ray absorption
  studies. Shells were first labeled with the
  letters K, L, M, N, O, P, and Q.
• The origin of this terminology was alphabetic.
• Later experiments indicated that the K absorption
  lines are produced by the innermost electrons.
• These letters were later found to correspond to
  the n-values 1, 2, 3, etc.
• K1 ,L 2 , M 3, N 4, etc.

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Quantum theory

• The "quantum" theory was proposed more than 90
  years ago, and has been confirmed by thousands of
  experiments. Science and education has failed to
  clearly describe the energy level concept to
  almost four generations of citizens. This
  experiment is an exercise aimed at throwing a
  little more light on the subject.
• Atoms have two kinds of states a ground state
  and an excited state. The ground state is the
  state in which the electrons in the atom are in
  their lowest energy levels possible (atoms
  naturally are in the ground state). This means
  the electrons have the lowest possible values for
  "n" the principal quantum number.

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Spectrum, spectra
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Wavelength and frequency
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The visible spectrum with respect to infrared and
ultraviolet radiation. 
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UV and visible light

• Ultraviolet (200-400 nm) and visible (400-800 nm)
  radiation are found towards the short wavelength,
  high frequency end of the electromagnetic
  spectrum. Figure 1 shows the portion of the
  electromagnetic spectrum where UV-Vis radiation
  exists.

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White light

• You need to know that white light is the
  combination of all colors of the spectrum. 

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Separation of white light spectrum
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Continuous,emission ,absorption
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Different kinds of spectra

• Thus, emission spectra are produced by thin
  gases in which the atoms do not experience many
  collisions (because of the low density). The
  emission lines correspond to photons of discrete
  energies that are emitted when excited atomic
  states in the gas make transitions back to
  lower-lying levels.
• A continuum spectrum results when the gas
  pressures are higher. Generally, solids, liquids,
  or dense gases emit light at all wavelengths when
  heated.
• An absorption spectrum occurs when light passes
  through a cold, dilute gas and atoms in the gas
  absorb at characteristic frequencies since the
  re-emitted light is unlikely to be emitted in the
  same direction as the absorbed photon, this gives
  rise to dark lines (absence of light) in the
  spectrum.
•  

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Emission spectrum

• An element's 'emission spectrum' is the relative
  intensity of electromagnetic radiation of each
  frequency it emits when it is heated (or more
  generally when it is excited).
• When the electrons in the element are excited,
  they jump to higher energy orbits. As the
  electrons fall back down, and leave the excited
  state, energy is re-emitted, the wavelength of
  which refers to the discrete lines of the
  emission spectrum. Note, however, that the
  emission extends over an area of frequencies, an
  effect called spectral line broadening. This
  spectrum is caused by hydrogen.

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Emission spectrum...

• The emission spectrum can be used to determine
  the composition of a material, since it is
  different for each element of the periodic table.
  One example is identifying the composition of
  stars by analysing the received light

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Electronic configuration of Na
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Emission spectrum for Na
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Absorption spectra
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Flame test

• A flame test is a procedure used in chemistry to
  detect the presence of certain metal ions, based
  on each element's characteristic emission
  spectrum. The color of flames in general also
  depends on temperature see flame color.
• The test involves introducing a sample of the
  element or compound to a hot, non-luminous flame,
  and observing the color that results. Samples are
  usually held on a platinum wire cleaned
  repeatedly with hydrochloric acid to remove
  traces of previous analytes

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Flame test
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Flame test
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Shells - Orbitals