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Ch. 4: Atoms and Elements by HC120727123534

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									Ch. 4: Atoms and Elements

     Dr. Namphol Sinkaset

    Chem 152: Introduction to
       General Chemistry
       I. Chapter Outline
I. Introduction
II. Atomic Theory
III. The Nuclear Atom
IV. Elements
V. The Periodic Table
VI. Ions
VII. Isotopes
VIII. Atomic Mass
I. Introduction
        • Atoms are the
          building blocks of
          everything we
          experience.
        • What we smell, what
          we feel, what we see.
        • In this chapter, we
          trace the history of
          the atom and learn
          about its makeup.
            II. The Greeks
• From out perspective, matter can be
  infinitely divided.
• However, Leucippus and Democritus
  (5th century B.C.) believed there was a
  limit.
   Eventually, you will reach something that
    was “atomos” or “indivisible.”
• Unfortunately, their idea was not
  accepted.
     II. Revival of the Atom
• The idea of the atom lay dormant for
  over 2000 years.
• John Dalton revived the idea in order to
  explain 3 natural laws that puzzled
  everyone at the time.
• Dalton’s Atomic Theory worked so well
  that it was quickly accepted.
II. Postulates of Dalton’s Theory
1. Each element is composed of tiny,
   indestructible particles called atoms.
2. All atoms of a given element have the
   same mass and other properties that
   distinguish them from the atoms of
   other elements.
3. Atoms combine in simple, whole-
   number ratios to form compounds.
             II. Atoms
• Today, overwhelming evidence points
  towards the existence of atoms.
• Atoms can be imaged and arranged!
         III. Not “Atomos”
• Dalton’s theory treated atoms as
  permanent, indestructible building
  blocks that composed everything.
• However, J.J. Thomson discovered
  electrons, which were much smaller
  than an atom and negatively charged!
• Since atoms are neutral, where’s the
  positive charge?
           III. Plum Pudding

• J.J. Thomson
  proposed the
  plum pudding
  model of the
  atom.
   Electron “raisins”
   “Pudding” of
    positive charge
    III. Rutherford likes Plum
             Pudding
• Ernest Rutherford was a student of J. J.
  Thomson.
• He tried to prove the plum pudding
  model by shooting a-particles at gold
  foil.
• Note that a-particles are 7000x more
  massive than an electron and have a
  positive charge.
III. Rutherford’s Expectation
III. Rutherford’s a-Particle
        Experiment
      III. Conclusions from
     Rutherford’s Experiment
• Most of an atom’s mass and all of its
  positive charge exist in a nucleus.
• Most of an atom is empty space,
  throughout which tiny electrons are
  dispersed.
• By having equal numbers of positively-
  charged particles (protons) and electrons,
  an atom remains electrically neutral.
III. Rutherford’s Interpretation
       III. The Nuclear Atom
• Surprisingly, an atom is mostly empty space!
• The nucleus holds 99.9% of the atom’s mass.
  III. Components of an Atom
• Protons. Positively-charged particles in
  the nucleus. Mass of 1.67262 x 10-27 kg
  or 1.0073 amu.
• Neutrons. Neutral particles in the
  nucleus. Mass of 1.67493 x 10-27 kg or
  1.0087 amu.
• Electrons. Negatively-charged particles.
  Mass of 9.1 x 10-31 kg or 0.00055 amu.
               III. Charge

• Charge is a
  fundamental property.
• To designate charge,
  the sign GOES
  AFTER the
  magnitude, e.g. 2+.
• Matter is charge
  neutral.
       IV. An Atom’s Identity
• The number of protons in an atom
  determines its elemental identity.
   IV. Referring to Elements
• Since each element has a unique # of
  protons, we could refer to elements
  using Z, the atomic number, which
  equals the # of protons in an atom.
   e.g. The Z = 2 element.
• More commonly, we use an element’s
  name or chemical symbol.
   e.g. The element helium, or He.
      IV. Chemical Symbols
• Chemical symbols are a one or two
  letter abbreviation of an element’s
  name.
• First letter always capitalized; second
  letter is LOWERCASE.
• Some symbols are based on historical
  names: e.g. Au from aurum.
IV. The Periodic Table
        IV. Sample Problem
• Find the name and atomic number of
  the following elements.
  a)   V
  b)   N
  c)   Hg
  d)   Rh
  e)   Mo
 V. Organizing Chemical Info
• Dmitri Mendeleev was the first to organize
  information of elements according to periodic
  law, i.e. when arranged properly, elements
  show repeating properties.
V. Mendeleev’s Breakthrough

• Mendeleev placed
  elements with
  similar properties in
  vertical columns.
• He left blank spaces
  where he thought
  elements should
  exist.
V. Three Types of Elements
         V. Sample Problem
• Categorize the elements below as either
  a metal, nonmetal, or metalloid.
  a)   Ru (ruthenium)
  b)   Se (selenium)
  c)   I (Iodine)
  d)   Ba (barium)
  e)   Es (einsteinium)
  f)   Kr (krypton)
V. Main Group vs. Transition
V. Families of Elements
         V. Sample Problem
• To which group (new numbering
  system) does each of the following
  elements belong? If the group has a
  name, indicate that as well.
  a)   Br (bromine)
  b)   N (nitrogen)
  c)   Cs (cesium)
  d)   Mn (manganese)
VI. Atoms Can Lose/Gain e-’s
• In chemical reactions, it’s common for
  atoms to lose or gain electrons and
  become ions.
• ion: a particle that has a charge
• Examples:
   Na  Na+ + e-
   I + e -  I-
    VI. Origin of the Charge
• The charge arises from the different
  number of protons and electrons in the
  atom.
   Ion Charge = # protons - # electrons
• A neutral Na atom has 11 protons and
  11 electrons. If it loses and electron…
   Ion Charge = 11 – 10 = 1+
     VI. Cations and Anions
• An ion is fundamentally different than a
  neutral atom, so it needs a different
  name.
• cation: a positively-charged ion
• anion: a negatively-charged ion
• Note that cations and ions have
  different properties than their parent
  atoms, e.g. Na vs. Na+.
       VI. Sample Problem
• Determine the charges of the ions
  described below.
  a) A chromium atom that has lost 3
     electrons.
  b) A sulfur atom that has gained 2 electrons.
  c) An iron atom (Fe) that has 24 electrons.
  d) A phosphorus atom (P) that has 18
     electrons.
VI. Ions and the Periodic Table
• The charge of an ion can be predicted
  by the position of its parent element on
  the periodic table IF it’s a main group
  element.
• Simply count the number of spaces to
  the nearest noble gas (forward or
  backward).
• If you go forward, it’s an anion; if you go
  backward, it’s a cation.
VI. Predicting Ion Charge
         VI. Sample Problem
• What are the ions that form from atoms
  of the following elements?
     aluminum (Al)
     tellurium (Te)
     rubidium (Rb)
     oxygen (O)
            VII. Isotopes
• Protons are the only thing that
  determines the identity of an atom.
• Therefore, it’s possible for atoms of the
  same element to have different masses
  due to differing number of neutrons.
• isotopes: atoms with the same number
  of protons, but different numbers of
  neutrons
VII. Percent Natural Abundance
• The different types and amounts of each
  isotope is determined by nature.
• Note that in an isotope, the # of neutrons
  varies which makes the mass number (A)
  vary as well.
   VII. Referring to Isotopes
• Isotopes can be represented using the
  A, Z, X symbol.
    VII. Referring to Isotopes
• Alternatively, the X, A notation can be
  used.
      VII. Sample Problem
• How many protons and neutrons are in
  a potassium isotope with a mass
  number of 39? What are the three ways
  to represent this isotope?
  VIII. What’s the Mass of an
             Atom?
• It depends!
• Are we talking about the mass of a
  specific atom, i.e. a given isotope?
   If so, it’s just approximately the mass
    number.
• Are we talking about in general?
   Then it’s more complicated…
        VIII. Atomic Mass
• Not all atoms of the same element have
  the same mass, but we can calculate an
  average.
• The atomic mass is the weighted
  average mass of an element which
  accounts for all isotopes and their
  percent natural abundances.
   VIII. Calculating Atomic Mass

  • The equation below enables calculation
    of atomic mass.

Atomic mass  (fractionisotope1  mass isotope1) 
             (fractionisotope 2  mass isotope 2) 
             (fractionisotope 3  mass isotope 3)  
       VIII. Sample Problem
• Calculate the atomic mass of
  magnesium using the information in the
  table below.

    Isotope      Mass (amu)   Natural Abundance
    Mg-24          23.99           78.99%
    Mg-25          24.99           10.00%
    Mg-26          25.98           11.01%

								
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