Atomic Structure

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Atomic Structure

• Ms. Titolo
• Chemistry 112

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History of the AtomTime Line

• Find the following scientists in your textbook.
  They are not given to you in the order they are
  found in the text. That would be too easy! Some
  may be found in other chapters besides 4 and 5!!
  
• Write down the years they were alive or the year
  that their main discovery was made.
• Write down their contributions to the development
  of todays model of the atom.
• Create a creative, colorful, and informative
  timeline!

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Scientist to Include on Timeline

• Roentgen
• De Broglie
• Heisenberg
• Chadwick
• Bohr
• Aristotle

• Rutherford
• Millikan
• Dalton
• Marie and Pierre Curie
• Democritus
• Thompson

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Most of the following terms will be a review for
you
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What is an Atom?

• Atom The smallest particle of an element that
  retains the properties of the element.
• How big is an atom? VERY SMALL!!!
• World Population 6,000,000,000
• Number of Atoms in a penny
• 29, 000,000,000,000,000,000,000
• Scanning tunnel microscope allows us to see atoms

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Subatomic Particles found in the nucleus of an
Atom

• Neutron neutral (no charge) particle found
  inside the nucleus of an atom.
• Proton positive particle found in the nucleus of
  an atom.
• Both the proton and neutron are approximately the
  same mass (1 amu)

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Subatomic particle outside of the nucleus

• Electrons are negatively charged particles found
  outside the nucleus of an atom.
• The mass of an electron is negligible, only
  1/1836th of an amu.
• The mass of the electrons do not enter into
  calculating the atomic mass of an atom.

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Atomic Mass Unit

• Abbreviated amu
• Chemists developed a method of measuring the mass
  of an atom without using very small numbers in
  scientific notation.
• They chose an atomic standard Carbon-12
• They agreed the mass of Carbon-12 was 12 amu.
• Therefore, 1 amu is 1/12 the mass of a Carbon-12
  atom.

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Atomic Number

• The whole number next to each element that
  represents the number of protons in an atom.
  (Proton number can never be altered).
• In a neutral atom the proton number will equal
  the electron number. However atoms can lose or
  gain electrons, even though they cant lose or
  gain protons under normal circumstances.

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Atomic Mass

• The atomic mass (when rounded to a whole number)
  represents the of protons the number of
  neutrons.
• If you subtract the atomic number ( of protons)
  from the atomic mass (rounded to a whole ) you
  will find the number of neutrons in an atom.

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Isotopes

• Atoms of the same element can be found in nature
  with a different atomic mass than others. This
  is because they can have different s of
  neutrons
• An Isotope of an element is an atom of the same
  element with a different of neutrons.

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Isotopes

• Isotopes can be expressed two ways
• 1. C-14 (where 14 would be the whole number mass)
• 2 146C (Where 14 is the mass and 6 is the atomic
  number)

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Why is the Atomic Mass not a Whole Number?

• Atomic masses have decimals!
• Atomic mass weighted average mass of the
  isotopes of that element.
• For Example Mass of Cl is 35.453
• Isotopes 75 Chlorine-35
• 25 Chlorine-37

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Finding Protons, Neutrons, and Electrons

• Example K-39
• This is a potassium atom with a mass of 39
• Look on the periodic and see the the atomic
  number is 19.
• Therefore of protons 19 (equal to atomic )
• Since it is a neutral atom (no charge) number of
  electrons 19
• Number of neutrons 39 19 20

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Ions

• Ions are charged atoms that have lost or gained
  electrons.
• Two types of ions
• Cations positive ions that have been formed by
  the loss of electrons. (Ca2)
• Anions negative ions that have been formed by
  the gaining of electrons. (Cl-1)

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4.4 Unstable Nuclei and Radioactive Decay

• Late 19th century scientists noticed some
  elements spontaneously emitted energy and
  particles called radiation.
• Elements that give off radiation are said to be
  radioactive.
• Thus, atoms are not unchangeable as Dalton once
  thought.

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• Nuclear reactions a reaction that involves a
  change in the nucleus of an atom.
• Radioactive decay when nuclei are unstable and
  gain stability by emitting radiation.
• Fill out the Chart for Types of Radiation

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Types of Radiation
Alpha Beta Gamma
What is it made of? 2 protons and 2 neutrons (like the nucleus of a Helium atom) An electron Electromagnetic radiation
Mass 4 amu 1/1836th of an amu No mass energy
Charge positive negative none
Penetrating ability low medium high
What stops it? paper Foil/metal lead
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Information from Chapter 25

• Transmutation conversion of an atom of one
  element to an atom of another element by
  spontaneous emission of radiation.
• Induced Transmutation nuclear reactions
  produced artificially by striking a nucleus with
  a high-velocity charged particle.

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• Transuranium elements all elements after
  Uranium on Periodic Table.
• Produced in laboratory by induced transmutation
• All are radioactive

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Radioactive Decay Rates

• Half-life the time required for one-half of a
  radioisotopes nuclei to decay into its products.
  (see table 25-4)
• Ex. Strontium-90 half life is 29 years
• If today I have 10.0 g, in 29 years I would have
  5.0g
• In another 29 years (total of 58), I would have
  2.5 g

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Nuclear Fission/Fusion

• HUGE amounts of energy produced by Nuclear
  reactions.
• Nuclear Fission splitting of a nucleus into
  fragments to become more stable.
• Used in nuclear power plants (controlled)
• And nuclear bombs (uncontrolled)

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• Nuclear Fusion Combination of nuclei to form a
  more stable nucleus.
• Large energy released!
• Sun powered by Fusion of hydrogen into helium.
• Requires extremely high temperature to occur.
• Scientist researching Cold Fusion

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Atomic Model Development

• Around the end of the 1700s Dalton believed that
  an atom was a solid indestructible mass and so
  did most scientists.
• In the 1800s when JJ Thomson discovered the
  electron (first subatomic particle) that theory
  was shattered.

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Model Continued

• The other subatomic particles were found and
  further disproved the idea of the atom being a
  solid mass.
• Niels Bohr, who was a student of Rutherford, came
  up with the planetary model where electrons
  followed specific paths around the nucleus. He
  also stated that for an electron to move from one
  energy level to another it gained or lost a
  quantum of energy.

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Model cont

• Around 1925 Erwin Schrodinger came up with the
  quantum theory of electron placement which put
  electrons in clouds or regions not in specific
  paths. These areas of probability for electron
  location is based on a mathematical calculation.

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Identifying Where Electrons are Located in the
Atom

• The number and arrangement of electrons around
  the nucleus determines the atoms chemical
  properties. Specifically the outer electrons
  (valence electrons) Therefore, scientists need a
  shorthand way to write out electron arrangement.
• Electron Configuration Arrangement of electrons
  in an atom.

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• Principle Quantum Number Energy level the
  electron occupies (n)
• Examples n 1, 2, 3, 4, 5, 6, 7
• Ø      The larger the value of n, the farther
  away the electrons are from the nucleus and the
  higher the energy of the electron. Electrons are
  lazy, so they want to be as close to the nucleus
  as they can so they can expend the LEAST amount
  of energy.

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• Ø      Only a certain amount of electrons can fit
  in each energy level. The way to find the
  maximum in each level is 2n2
• o       Example n 1 Can have 2(1)2
  electrons 2 electrons in energy level 1
• o       You solve for n 2
• o       n 3
• o       n 4

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• Ø   How many energy levels does hydrogen have?
• Helium?
• Lithium?
• Sodium?
• Calcium?
• Xenon?

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Sublevels

• Sublevels further explain where the electrons
  are located in each energy level.
• Ø The names of the sublevels we use are s, p,
  d, and f

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Sublevels
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• Ø      The number of sublevels in each energy
  level is equal to n
• o       n 1 has one sublevel s
• o       n 2 has two sublevels s, p
• o       n 3 has three sublevels s, p, d
• o       n 4 has four sublevels s, p, d, f
• o       n 5 has five sublevels but the fifth
  one is not used in ground state elements.
• Notice, energy levels must have an s sublevel
  before a p!

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Orbitals

• Orbitals Each sublevel has specific orbitals
  the electrons can be in
• Each orbital can hold 2 electrons. Therefore,
  the maximum number of electrons in the s orbital
  is 2.
• Ø      S has 1 orbital maximum electrons 2
• Ø      P has 3 orbitals maximum electrons 6
• Ø      D has 5 orbitals maximum electrons 10
• Ø      F has 7 orbitals maximum electrons 14

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Order of Filling

• http//lectureonline.cl.msu.edu/mmp/period/electr
  on.htm

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• Three Main Rules Electrons Abide By
• Ø      Aufbau principle (Electrons are Lazy)
  each successive electron occupies the lowest
  energy orbital available.
• Ø      Exclusion Principle Pauli (Two to
  Tango) maximum of two electrons may occupy a
  single atomic orbital, but only if the electrons
  have opposite spins.
• Hunds Rule (Why share if you dont have to?)
  Single electrons with the same spin must
  occupy each equal-energy orbital before
  additional electrons with opposite spins can
  occupy the same orbital.

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• Orbital Diagrams
• Ø Used to show how electrons are distributed
  within sublevels and to show the direction of
  spin.
• Ø  Boxes are used to represent orbitals.
• Ø  Arrows are used to represent electrons.

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• Ø The first electron in the orbital is
  represented by an arrow pointing up, ?, meaning
  clockwise spin.
• Ø The second electron in the orbital is
  represented by an arrow pointing down, ?, meaning
  counterclockwise spin.
• ØAn electron configuration notation is an
  abbreviated form of the orbital diagram.

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• Electron Dot Diagrams
• ØValence Electrons electrons in the atoms
  outermost energy level.
• ØThese electrons are involved in forming chemical
  bonds.
• Ø They are represented visually by an electron
  dot structure.
• ØAlso known as Lewis Dot Formulas.
• Examples