Electrons in Atoms

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Electrons in Atoms
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• Why do ions have the charges they have? Like
  Al3 or Fe2 or Fe3 or O2-
• Why does an atom become an ion in the first
  place?
• Why are the BrINClHOF elements the only ones that
  make molecules with themselves? Why dont any
  other elements do that?
• How do fireworks make colors when they explode?
• How do fluorescent lights work?
• How do we know how hot the sun is?
• How is there life on this planet?
• The electrons in atoms can answer all of these
  questions and so many others.
• It all happens because of electrons..

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Daltons Model of the Atom1803

• Atoms are tiny, indestructible spheres
• No internal structure

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Thomsons Model1897

• Referred to as the plum-pudding model.
• The whole atom is a sphere of positive charge,
  with little negative electrons embedded in it.

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Rutherfords Model1911

• Small, dense core of positive charge.
• Electrons circle the nucleus in fixed orbits.

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Rutherfords Model

• Electrons revolve around the nucleus like planets
  around the sun (fixed orbits).
• This model failed to explain some properties of
  atoms.

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Niels Bohrs Model1913

• Electrons orbit the nucleus in specific orbits a
  fixed distance away.

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Neils Bohrs Model (1913)

• They orbit at a particular energy level. They
  can move to a higher level, but they need energy
  to do so.
• A quantum of energy is the required amount to
  move an e- to a higher level. Exactly this
  amount, no in-between.

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Neils Bohrs Model (1913)

• This model of the atom had shortcomings. It
  failed to explain some phenomenon in nature.
• So, a better version was still out there waiting
  to be discovered..

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Waves

• To understand the electronic structure of atoms,
  one must understand the nature of electromagnetic
  radiation.
• The distance between corresponding points on
  adjacent waves is the wavelength (?).

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Waves

• The number of waves passing a given point per
  unit of time is the frequency (?).
• For waves traveling at the same velocity, the
  longer the wavelength, the smaller the frequency.

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Light
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Electromagnetic Radiation

• All electromagnetic radiation travels at the same
  velocity the speed of light (c),
• 3.00 ? 108 m/s.
• Therefore,
• c ??
• This all suggests light is a wave.

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The Nature of Energy

• The wave nature of light does not explain how an
  object can glow when its temperature increases.
• Max Planck explained it by assuming that energy
  comes in packets called quanta.
• Equantum h?
• Einstein used this assumption to explain the
  photoelectric effect.

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Photoelectric Effect

• The emission of electrons from a metal when light
  is shined upon the metal.
• Depending on the metal used, only light of a
  certain wavelength (color) would cause an
  electron to be emitted.

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Photoelectric Effect

• Einstein concluded that energy is proportional to
  frequency
• Ephoton h?
• where h is Plancks constant, 6.626 ? 10-34 J-s.
• This suggests light as a particle.

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Nature of Energy

• Therefore, if one knows the wavelength of light,
  one can calculate the energy in one photon, or
  particle, of that light
• c ??
• E h?

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The Nature of Energy

• Another mystery in the early 20th century
  involved the emission spectra observed from
  energy emitted by atoms and molecules.

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The Nature of Energy

• For atoms and molecules one does not observe a
  continuous spectrum, as one gets from a white
  light source.
• Only a line spectrum of discrete wavelengths is
  observed.

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The Nature of Energy

• Niels Bohr adopted Plancks idea of quanta and
  explained these phenomena in this way
• Electrons in an atom can only occupy certain
  orbits (corresponding to certain energies).
• Electrons in permitted orbits have specific,
  allowed energies these energies will not be
  radiated from the atom.
• Energy is only absorbed or emitted in such a way
  as to move an electron from one allowed energy
  state to another the energy is defined by
• E h?

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The Nature of Energy

• The energy absorbed or emitted from the process
  of electron promotion or demotion can be
  calculated by the equation
• ?E -Rh(1/nf2 - 1/ni2)
• where RH is the Rydberg constant, 2.18 ? 10-18 J,
  and ni and nf are the initial and final energy
  levels of the electron.

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The Wave Nature of Matter

• Louis de Broglie posited that if light can behave
  with material properties (photons), matter should
  exhibit wave properties.
• He demonstrated that the relationship between
  mass and wavelength was
• ? h/mv
• In other words, if light waves can act like
  particles, then things can move like waves.

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Heisenbergs Uncertainty Principle

• Heisenberg showed that the more precisely the
  momentum of a particle is known, the less
  precisely is its position known
• (?x) (?mv) ? h/4p
• For regular-sized objects, the uncertainty is
  practically zero, but in many cases, our
  uncertainty of the whereabouts of an electron is
  greater than the size of the atom itself!

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Quantum Mechanics

• Erwin Schrödinger developed a mathematical
  treatment into which both the wave and particle
  nature of matter could be incorporated.
• It is known as quantum mechanics.
• Amazingly accurate in describing electrons and
  microscopic behaviors, but also exceedingly
  strange.

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Richard Feynman on Quantum Mechanics (1965)

• There was a time when the newspapers said that
  only twelve men understood the theory of
  relativity. But after people read the paper a
  lot of people understood the theory of
  relativity. On the other hand I think I can
  safely say that nobody understands quantum
  mechanics.

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Schrodingers Wave Equations

• The wave equation is designated with a lower case
  Greek psi (?).
• The square of the wave equation, ?2, gives a
  probability density map of where an electron has
  a certain statistical likelihood of being at any
  given instant in time.

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Some Folks are Really Into It.
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Quantum Numbers

• Solving the wave equation gives a set of wave
  functions, or orbitals, and their corresponding
  energies.
• Each orbital describes a spatial distribution of
  electron density.
• An orbital is described by a set of three quantum
  numbers.

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Principal Quantum Number, n

• The principal quantum number, n, describes the
  energy level on which the orbital resides.
• The values of n are integers 1.
• n also describes the relative size of the
  orbital, 2 larger than 1, and so on.

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Angular Momentum Quantum Number (l)

• This quantum number defines the shape of the
  orbital.
• Allowed values of l are integers ranging from 0
  to n - 1.
• We use letter designations to communicate the
  different values of l and, therefore, the shapes
  and types of orbitals. This is where s, p, d f
  come into play

Value of l 0 1 2 3
Type of orbital s p d f
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Magnetic Quantum Number (ml)

• The magnetic quantum number describes the
  three-dimensional orientation of the orbital.
• Allowed values of ml are integers ranging from -l
  to l
• -l ml l.
• Therefore, on any given energy level, there can
  be up to 1 s orbital, 3 p orbitals, 5 d orbitals,
  7 f orbitals, etc.

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Magnetic Quantum Number (ml)

• Orbitals with the same value of n form a shell.
• Different orbital types within a shell are
  subshells.

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s Orbitals

• The value of l for s orbitals is 0.
• They are spherical in shape.
• The radius of the sphere increases with the value
  of n.

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s Orbitals

• Observing a graph of probabilities of finding an
  electron versus distance from the nucleus, we see
  that s orbitals possess n-1 nodes, or regions
  where there is 0 probability of finding an
  electron.

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p Orbitals

• The value of l for p orbitals is 1.
• They have two lobes with a node between them.

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d Orbitals

• The value of l for a d orbital is 2.
• Four of the five d orbitals have 4 lobes the
  other resembles a p orbital with a doughnut
  around the center.

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Energies of Orbitals

• For a one-electron hydrogen atom, orbitals on the
  same energy level have the same energy.
• That is, they are degenerate.

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Energies of Orbitals

• As the number of electrons increases, though, so
  does the repulsion between them.
• Therefore, in many-electron atoms, orbitals on
  the same energy level are no longer degenerate.

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Spin Quantum Number, ms

• In the 1920s, it was discovered that two
  electrons in the same orbital do not have exactly
  the same energy.
• The spin of an electron describes its magnetic
  field, which affects its energy.

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Spin Quantum Number, ms

• This led to a fourth quantum number, the spin
  quantum number, ms.
• The spin quantum number has only 2 allowed
  values 1/2 and -1/2.

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Pauli Exclusion Principal

• No two electrons in the same atom can have
  exactly the same energy.
• Therefore, no two electrons in the same atom can
  have identical sets of quantum numbers.

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Electron Configurations

• This shows the distribution of all electrons in
  an atom.
• Each component consists of
• A number denoting the energy level,
• A letter denoting the type of orbital,
• A superscript denoting the number of electrons in
  those orbitals.

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Orbital Diagrams

• Each box in the diagram represents one orbital.
• Half-arrows represent the electrons.
• The direction of the arrow represents the
  relative spin of the electron.

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Hunds Rule

• For degenerate orbitals, the lowest energy is
  attained when the number of electrons with the
  same spin is maximized.

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Periodic Table

• We fill orbitals in increasing order of energy.
• Different blocks on the periodic table (shaded in
  different colors in this chart) correspond to
  different types of orbitals.

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Atomic Orbitals

• Region of space where there is a high probability
  of finding an electron.
• Principal Quantum (n) --denotes the energy level
  of electrons (1,2,3,4,etc.)
• Also denotes the of sublevels at that energy
  level (s,p,d,f)
• Sublevels describe the shapes and sizes of
  orbitals where e- may be found.

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Shapes of Orbitals

• s-spherical, with nucleus at the center
• p-dumbbell, or figure-8, with nucleus at the
  center
• das shown
• fas shown
• As you increase energy levels, the shape of each
  remains the same, but size gets larger.

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Electron Configurations

• Orbitals of an atom will fill so that the atom is
  in its most stable state. There are 3 rules that
  govern this
• Aufbau Principle- e- occupy lowest-energy
  orbitals first
• Pauli Exclusion Principle- 2 e- in same orbital
  must have opposite spin
• Hunds Rule- e- occupy orbitals of the same
  energy so that theres a max of same spin e-

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Exceptions to Aufbau

• If you did the configuration for Cu according to
  the three rules, it would look like this
• 1s22s22p63s23p64s23d9
• In actuality, it is this
• 1s22s22p63s23p64s13d10

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Another

• Chromium, Cr, also is an exception to the Aufbau
  Principle
• According to Aufbau, Cr should have this
  configuration
• 1s22s22p63s23p64s23d4
• But it actually has this
• 1s22s22p63s23p64s13d5

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Why Would an Atom Do This?

• Because a filled shell is the most stable
  arrangement, and a half-filled shell is the next
  best arrangement.

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Valence Electrons

• The electrons that exist in the outermost energy
  level of an atom are valence electrons.
• A full shell or a half-filled shell is the most
  stable arrangement.
• Noble gases always have a full valence, or a full
  outer shell, which is what every other element is
  trying to achieve. (Max. of 8 valence electrons)

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• What does the term orbital describe?
• A region around the nucleus where an electron is
  most likely to be found.
• What does an elements electron configuration
  describe?
• All of the orbitals that the elements electrons
  occupy, and how those electrons are distributed.

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• We do not need to focus on all the electrons that
  an atom has, we really only need to focus on the
  valence electrons. Why?
• Because they are the outermost electrons, and
  they are the only electrons that can possibly
  interact with other atoms.
• How many valence electrons does Oxygen have?
• 6

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• Why are the alkali metals so reactive?
• They all have an s1 electron (1 valence electron)
  that they are trying to lose.
• Noble gases are also called inert gases. Why are
  the noble gases so unreactive?
• Because they have a full outer shell (eight
  valence electrons) and do not need any more or
  less electrons.

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First Periodic Table

• In 1869, the first table having elements
  organized by their properties was published by a
  Russian chemist and professor named Dmitri
  Mendeleev.
• He listed them in order of atomic mass.

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Gallium and Germanium Discovered in 1875 1886
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• Mendeleev arranged the elements in order of
  increasing mass.
• In the 1860s, the proton was not yet discovered.
• In 1913, British physicist Henry Moseley arranged
  the elements in order of increasing atomic number
  ( of protons).

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Some Vocabulary

• Vertical columns are called
• groups or families.
• Horizontal rows are called periods.
• How many elements are in period 2?
• 8
• How many elements are in period 6?
• 32
• How many elements are in group 2?
• 6

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The Periodic Law (Cont.)

• Elements within a column of a group have similar
  properties.
• Properties in a period change as you move across
  a period from left to right.
• The pattern of properties within a period repeats
  as you move form one period to the next.
• Periodic Law When elements are arranged in order
  of increasing atomic number, there is a periodic
  repetition of their physical and chemical
  properties.

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Electron Configurations in Groups
Helium (He) 1s2
Neon (Ne) 1s22s22p6
Argon (Ar) 1s22s22p63s23p6
Krypton (Kr) 1s22s22p63s23p63d104s24p6
Noble Gases
Lithium (Li) 1s22s1
Sodium (Na) 1s22s22p63s1
Potassium (K) 1s22s22p63s23p64s1
Alkali Metals
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Blocks of Elements
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Metals

• 80 of elements

Nonmetals
Metals
Metalloids
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Metals (Cont.)

• Conductors of heat
• Conductors of electric current
• High luster
• Ductile
• Malleable
• Solids _at_ room temp. (except Hg)

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Nonmetals

• Most are gases _at_ room temp
• Poor conductors of heat
• Poor conductors of electric current
• Solid nonmetals are brittle

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Metalloids

• Properties similar to those of
• metals and nonmetals
• Behaviors can be controlled by changing the
  conditions.
• Example Silicon

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Classifying the Elements

• Group 1A Elements Alkali Metals
• Group 2A Elements Alkaline Earth Metals
• Group 7A Elements Halogens
• Group 8A Elements Noble Gases
• Group B Elements Transition Metals
• Below the Main Body- Inner Transition Metals

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Periodic Trends

• Atomic Size
• Atomic radius one half the distance between the
  nuclei of two atoms of the same element
• Increases from top to bottom within a group
• Decreases from left to right across a period

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Atomic Radius vs. Atomic Number
Increase within Group 1 Shielding Effect
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Periodic Trends in Atomic Size
Size decreases
Size Increases
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Trends in Ionization Energy

• Ionization Energy
• The energy required to remove an electron from an
  atom.
• First Ionization Energy
• The energy required to remove the first electron
  from at atom.
• First ionization energy tends to decrease from
  top to bottom within a group and increase from
  left to right across a period.

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First Ionization Energy vs. Atomic Number
Why is the 1st ionization energy for the noble
gases higher?
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Periodic Trends in Ionization Energy
Energy Increases
Energy Decreases
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• Electronegativity- The ability of an atom of an
  element to attract electrons when the atom is
  BONDED to another atom in a compound.

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Periodic Trends

• Metallic propertiesas shown. As you approach
  the nonmetals, metallic properties decrease.