Chemistry Unit I: Chapter 4

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Chemistry Unit IChapter 4 Atoms
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Chp 4.1 Atomic Structure, p. 113-118

• Our understanding of atoms took many centuries.
• The 4th century Greek philosopher, Democritus,
  suggested that the universe was made of invisible
  units called atoms. The word atom means that
  which cannot be divided.
• The Greek idea of atoms had to wait 2000 years
  before it became accepted.

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John Dalton

• In 1808, an English schoolteacher named John
  Dalton proposed his own atomic theory. He used
  evidence such as the law of definite
  proportions.
• His theory proposed
• Every ELEMENT is made of tiny, unique particles
  called atoms that cannot be subdivided.
• Atoms of the same element are exactly alike.
• Atoms of different elements can join to form
  molecules.

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J.J. Thomson

• In 1897, British scientist, J.J. Thomson, found
  that atoms contain negatively charged particles.
  These particles were named electrons.
• Since atoms are neutral, Thomson reasoned that
  atoms must contain some sort of positive charge.
  (Cathode Ray Tube experiment)
• His model, like a plum pudding, had negative
  electrons (plums) embedded in a positive sphere
  (pudding).

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Ernest Rutherford

• In 1911, Ernest Rutherford, discovered that atoms
  contain a dense, positively charged nucleus
  through his Gold Foil Experiment.
• Later Rutherford name the positive particles
  protons.
• In his model, atoms were mostly empty space with
  electrons moving around a small, very dense,
  positively-charged nucleus in the center of the
  atom.

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One additionthe neutron

• In 1932, British scientist, James Chadwick,
  discovered another particle in the nucleus of an
  atom. The particle was nearly the same mass as a
  proton.
• This particle is called a neutron, because it
  is electrically neutral.

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Chapter 4.2 The Structure of Atoms, p.119-127

• Three main subatomic particles are distinguished
  by mass, charge and location in the atom.
• Proton, charge 1, located in the nucleus, mass
  1.67x10-27
• Neutron, charge 0, located in the nucleus, mass
  1.67x10-27
• Electron, charge -1, located in orbitals, mass
  9.11x10-31

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Atoms

• Each element has a unique number of protons.
• Each atom of the element will have the same
  number of protons.
• Atoms are electrically neutral overall. Atoms
  have the same number of protons as electrons.
  These charges exactly cancel leaving no charge.
• Positive and negative charges attract each other
  with a force known as electric force.

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Ions

• Atoms may undergo a process called ionization,
  in which they lose or gain electrons in order to
  fill their valence energy level.
• An ion is an atom (or group of atoms) that has
  lost or gained one or more electrons, therefore
  has a net charge.
• Removing electrons, causes a positive ion to
  form, aka cation.
• Adding electrons, causes a negative ion to form,
  aka anion.

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Atomic Nucleus

• The atomic number tells you how many protons are
  in an atom.
• Every element has a different atomic number.
• Atoms may or may not have the same number of
  neutrons as protons.
• The mass number tells you how many protons plus
  neutrons are in the nucleus of an atom all
  together.

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Isotopes

• Some atoms contain a different number of neutrons
  thereby making them heavier or lighter than
  the average atom. These atoms are called
  isotopes.
• Isotopes have the same number of protons just a
  different number of neutrons.
• To calculate the number of neutrons, subtract the
  atomic number (Z) from the mass number (A) like
  this
• A Z number of neutrons

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Atomic Masses

• The mass of a single atom is extremely small!
• Atomic masses are usually expressed in unified
  atomic mass units.
• A unified atomic mass unit (u) is equal to
  one-twelfth of the mass of a carbon-12 atom. Aka
  amu
• The average atomic mass for an element is a
  weighted average and appears as a decimal number
  on the periodic table.

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Moles

• Scientists often deal with very large numbers of
  small particles. They use a counting unit called
  a mole.
• A mole is 602213670000000000000000 !!!
• This number is usually written as
• 6.02 x 10 23 /mole and is called Avogadros
  constant. This is how many particles are in a
  mole of any substance.
• Like the word dozen means 12, a mole of a
  substance contains 6.02 x 10 23 particles.

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Molar Mass

• The mass in grams of 1 mol of a substance is
  called its molar mass.
• Example, one mole of lithium has a mass of 6.941
  g.
• Because mass and amount of a substance are
  related, it is possible to convert moles to grams
  and vice versa.

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Calculating with Moles

• Using a conversion factor allows you to change
  from mass to amount and back again.
• The numerator of the conversion factor contains
  the unit you want to change to. The denominator
  of the conversion factor contains the unit you
  start with.
• The numerator and denominator equal 1 since they
  are equal.
• Example you could use either
• 10 gumballs/ 21.4 g
• 21.4 g/ 10 gumballs

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Calculating with moles

• Converting between amount of an element in moles
  and its mass in grams you can use this chart
• Amount (mol) x molar mass (g)/ 1 mol of element
  Mass (g)
• Mass (g) x 1 mol of element/ molar mass of
  element (g) Amount (mol)

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Chp 4-3 Modern Atomic Theory p.128-132

• Modern Model of the Atom
• Electrons can be found only in certain energy
  levels, not between levels.
• Furthermore, the location of electrons cannot be
  predicted precisely.

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Niels Bohr

• In 1913, Niels Bohr, a Danish scientist, revised
  the Rutherfords model.
• Bohr showed that electrons could only have
  specific amounts of energy, leading them to move
  in certain orbits.
• This model resembled planets orbiting the sun.

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Cloud of Electrons

• In the 1920s, scientists again revised the
  model. They found that electrons can be anywhere
  in a cloudlike region around the nucleus. This
  region is called an orbital.
• Electrons behave more like waves on a vibrating
  string than like particles.
• The different energy levels contain different
  numbers of electrons depending on the energy
  level size and electron energy.

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Electron Energy Levels

• The number of energy levels that are filled in an
  atom depends on the number of electrons.
• The electron(s) in the outermost energy level of
  an atom are called valence electrons.
• The most valence electrons that an atom can have
  is eight (8).
• Atoms with a filled valence energy level are the
  stable and dont react chemically.
• Electron dot diagrams show the elements symbol
  and valence electrons as dots.

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Electrons and Orbitals

• Electrons may occupy four different kinds of
  orbitals within atoms
• S is the simplest orbital, a sphere shape,
  lowest energy.
• p is dumbbell-shaped, and can be oriented 3
  ways, and are the next higher energy levels.
• d and f orbitals are much more complex. There
  are 5 possible d orbitals and 7 possible f
  orbitals. These are the highest energy levels.
• These orbitals are all different shapes and can
  hold a maximum of two electrons each.

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Electron Transition

• Electrons jump between energy level when an atom
  gains or loses energy.
• Ground state is the lowest state of energy of an
  electron.
• If an electron gains energy, it moves to an
  excited state.
• It gains energy by absorbing a particle of light
  called a photon.
• It will release the photon when it moves back to
  its ground state.

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Stop for Now!!!! Chapter 4 content complete

• (the following slides are from Chp 5please skip
  for now)

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Chapter 5.3 Families of Elements, p 156-165

• The two main groups on the periodic table are
  metals and non-metals.
• Most metals are shiny solids that can be
  stretched (ductile) and shaped (malleable). The
  are good conductors of heat and electricity.
• Non-metals can be solid, liquid or gas. They are
  typically dull, brittle and are poor conductors.
• Between the two groups are semiconductors, or
  metalloids, which share characteristics of both
  groups.

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Alkali Metals

• Group 1 of the periodic table contains elements
  with one valence electron, which can be easily
  removed to form a positive ion.
• They are all highly reactive and are not found in
  nature as elements.
• Examples include sodium and potassium. These
  easily combine with non-metals to form salts.

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Alkaline Earth Metals

• Group 2 of the Periodic Table are the alkaline
  earth metals.
• They contain 2 valence electrons and react (not
  as violently as group 1) losing their 2 electrons
  to become a positive ion then forming compounds.
• Examples in calcium and magnesium.

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Transition Metals

• Groups 3 through 12 contain the transition
  metals.
• They are less reactive than groups 1 or 2 but
  some still combine with non-metals to form
  compounds. Some have many versions of cations.
• Examples include gold, silver, copper mercury etc.

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Non-metals

• Except for hydrogen, non-metals are found on the
  right side of the periodic table including some
  elements in groups 13 thru 16 and all elements in
  groups 17 and 18.
• Non-metals tend to share or gain electrons
  (anion) and are plentiful on Earth.
• Examples include hydrogen, carbon, oxygen,
  sulfur, nitrogen, etc.

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Halogens

• Group 17 contain non-metals called halogens.
  These elements contain 7 valence electrons and
  are highly reactive.
• Examples include fluorine, chlorine, iodine, etc.

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Noble Gases

• Group 18 of the periodic table contain the noble
  gases. Because these atoms have filled valence
  energy levels, they do not react with other
  elements to form compounds.hence the name
  noble.
• They all exist as single gas atoms.
• Examples include helium, neon, argon, etc.

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Semiconductors

• Located along the staircase line,
  semiconductors or metalloids have properties of
  both metals and non-metals.
• Semiconductors are able to conduct heat and
  electricity under certain conditions, some are
  hard, some are shiny and so on.
• The 6 elements are
• Boron, silicon, germanium, arsenic, antimony, and
  tellurium. Sometimes polonium and astatine are
  also included.