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CH 2 Atoms_ Molecules_ and Ions

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					  CH 2: Atoms,
Molecules, and Ions
              Chapter Outline
 History   of chemistry (2.1 – 2.4)
     Chemical laws
     Path to the atom
 Modern  atomic structure (2.5)
 Molecules vs. Ions (2.6)
 Naming molecular and ionic compounds
  (2.8)
 Introduction to the periodic table (2.7)
          History of Chemistry
Greek Philosophers: 5th Century BCE
 (BCE = before the common era - replaces BC)
 The Greek philosophers were the first to
 reflect on the nature of matter.
     They proposed that all matter is made out of
      first 4 elements -- earth, air , water,
      fire. Aristotle added a fifth element, plasma
      (also called ether).
           Greek Philosophers
 Democritushad an alternate view of
 matter. He proposed that matter
 was made up of tiny particles called
 atoms.
     His "theory" was not well accepted at the time.
                       Alchemy
Alchemy: ~600-1600's CE
     (CE = common era, replaces AD)
     Alchemy developed at about the same time in
      China, India, and Greece. It spread into
      Europe in the 8th century.
                      Alchemy
Alchemists had two pursuits
  1.   Search for a means to convert “base” metals
       into gold
  2.   Search for the elixir of life
       •   Substance that would lead to immortality
                      Alchemy
Advances from Alchemy
     Many new substances where identified
      • Plaster of Paris, nitric acid….
     New lab techniques and equipment
      developed
     New medicines identified
    Modern Chemistry, ~1600 on
   First chemists/physicists to use scientific method
       Boyle - elements
       Lavoisier – law of conservation of matter
       Proust - law of definite proportions
       Dalton – law of multiple proportions, atomic theory
       Avogadro - hypothesis
       Thomson – charge to mass ratio for an electron
       Millikan – charge on the electron
       Bequerel and the Curies - radioactivity
       Rutherford – nuclear atom
           Modern Chemistry
Robert Boyle: ~1660
 Proposed a substance to be an element
  unless it can be broken down into simpler
  substances.
     Proposed one of the gas laws – CH 5
               Lavoisier: ~1760

Law of Conservation of Matter
     Matter is neither created nor destroyed in a
      chemical reaction.
             Proust: late 1700s

Law of Definite Composition/Proportions
     A given compound always contains the same
      proportion of elements by mass.
 Experimental      basis for the law of definite
 composition
     Proust found all samples of CuCO3 had the
      same relative composition of elements by
      mass:
       • 5.3 parts Copper
       • 4 parts Oxygen
       • 1 part Carbon
               John Dalton: ~1800

Law of Multiple Proportions
     When two elements form more than one
      compound, the ratios of the masses of the
      second element that combines with one gram
      of the first element can always be reduced to
      small whole numbers.

     Page 70: #27 is an example of experimental
      data consistent with this law
 Daltonalso proposed the first table of
 atomic masses
     Most masses later need revision


 Dalton   is best known for proposing Atomic
 Theory
           Dalton‟s Atomic Theory
1.       Elements are made up of tiny particles
         called atoms.
     •     Atoms are indivisible and indestructible


2.       Atoms of a given element are identical
          Atoms of different elements differ in some
           fundamental way(s)
           Dalton‟s Atomic Theory
3.       Compounds form when atoms of
         different elements combine with each
         other.
          A given compound always has the same
           relative number and types of elements.
           Dalton‟s Atomic Theory
4.       Chemical reactions occur when atoms
         change how they are bound to each
         other.
         Individual atoms are not changed, just
          rearranged
             Avogadro: 1811
Avogadro's Hypothesis
     At the same temperature and pressure equal
      volumes of gases contain the same number of
      particles.
      • Based on guy-Lussac‟s data
      • See pages 44/45
From Dalton to Atomic Structure
 Dalton‟s atomic theory lead to much
  research on the nature of the atom.
 This research showed the atom to made
  up of smaller particles.
                    J.J. Thompson
 Thomson measured the deflection of a
 cathode ray beam in electrical and
 magnetic fields of known strengths.
             Cathode ray                Applied electrical field

                                +
                                                               (+)
 (-)
 Metal                                                         Metal
 electrode                      -                              electrode


             Cathode ray tube experiment, pg 46/47
   Thompson found the cathode rays were
    attracted by the positive charge and repelled
    by negative

   These findings clearly indicated that the rays
    consisted of negatively charged particles.
     • Today we know these particles as electrons.
 Thompson  measured the deflection of the
 beam in a magnetic field and determined
 the charge:mass ratio for an electron

  e = - 1.76 x 108 C/g
  m

  e = charge on the electron in coulombs
  M = mass of an electron in grams
 Thomson   also found that the cathode ray
 particles were identical regardless of
 source.
     Concluded all elements contain these
      negative particles (electrons)
                 J.J. Thomson
Thomson:
     identified cathode ray beams as a stream of
      negatively charged particles
     calculated the charge to mass ratio for these
      negatively charged particles
     proposed the existence of positively charged
      particles
       • To balance the negative charge of the electrons
                Millikan ~1909
           oil drop experiment allowed him
 Millikan‟s
  to determine the charge on an electron
     This charge can be plugged into Thomson‟s
      formula and the mass of the electron
      calculated
       • Mass electron = 9.11 x 10-31 kg
                    Radioactivity
Becquerel, Marie and Pierre Curie: ~1896
 Henri Becquerel - observed the natural emission
  of energy/rays by uranium.

   Marie and Pierre Curie studied “Becquerel's
    rays”.
       The Curies’ findings suggested that matter was
        composed of smaller particles than atoms.
       The Curies coined the term radioactivity to describe
        the rays emitted.
                  Radioactivity
 Three    types of radioactivity were identified:
     gamma rays - very high energy light
     beta particles - high energy electrons
     alpha particles - He+2 particles
       • 2 protons and 2 neutrons
 Inthe early 1900‟s the accepted model of
  the atom was called the plum pudding
  model of the atom
      Electrons (tiny and negatively charged) were
       pictured to be dispersed in a „cloud‟ of positive
       charge.
 Rutherford and the Nuclear Atom
 In1911 an experiment conducted in
  Ernest Rutherford‟s lab showed the “plum
  pudding” model to be incorrect.
      Experiment was conducted by Geiger and
       Marsden and the findings interpreted by
       Rutherford.


      See page 49
     The gold foil experiment
 What they did – see board
 What they found – see board
             Rutherford‟s Atom
 First   to propose a nuclear atom.
     An atom has a dense positive center
      containing all of positive charge and most of
      the mass of the atom – the nucleus
     Electrons are scattered in the empty space
      around the nucleus
       • Electrons occupy a volume that is huge as
         compared to the size of the nucleus.
     A New Model of the Atom

Expected based on
Plum pudding model

Rutherford‟s model
Based on ”his” results
       Modern Atomic Structure
 Rutherfordcontinued to study the atom
 and the positive matter of the atom.
     1919, + particle named the proton


 ~1932 James Chadwick proposed the
 existence of a third subatomic particle, the
 neutron.
              Subatomic Particles
Subatomic       Charge   Mass, amu   Location in atom
 Particle


 Electron         -1       0 amu     Outside of nucleus
   (e-)


Proton (p)        +1      ~1 amu          Nucleus



Neutron (n)       0       ~1 amu          Nucleus
  Mass of Subatomic Particles
 Protons    and neutrons have ~ the same
 mass (in the range of 10 -27 kg).
     Mass of each and of individual atoms is often
      expressed in amu rather than grams
       • Atomic mass unit (amu) – 1/12 the mass of a
         carbon-12 atom
  Mass of Subatomic Particles
 Themass of the electron (10-31 kg) is tiny as
 compared to that of the proton and
 neutron (10-27 kg) .
     Therefore, the electron‟s mass is considered
      to be ~0 amu when calculating the mass of an
      atom.
           Subatomic Particles and the
                   Elements
   Each   element has a unique number of
      protons.
          Number of protons defines the element.

Atomic # = # protons
                         6

                             C
          Subatomic Particles
 Since atoms are neutral, for every proton
 there is a/n _________.

 When    atoms interact to form compounds,
 it is their ___________ that interact.
                      Terms
 Mass number = sum of the # of protons
 and the # neutrons in the nucleus of an
 atom
     FOR MOST ELEMENTS THE MASS
      NUMBER IF NOT ON THE PERIODIC
      TABLE.
      • You will be given enough information to determine
        mass number or number of neutrons.
                     Terms
 Isotopes  = atoms of a given element that
 differ in mass number
     Isotopes have the same number of
      _____________.
     Isotopes differ in the number of _______.
                   Isotopes
 Writing atomic   symbols for isotopes
     pg 50


Mass #        11

              5
                B           Symbol for
                            element

  Atomic #
             FAQ - Isotopes
 When  is mass number found on the
 periodic table?

           atomic mass? Is it the same
 What‟s the
 as the mass number?
       84    (209)       6   12.0107


            Po
                             C
         Molecules and Ions (2.6)
 Atoms of different elements combine to
 form compounds
     Atoms in compounds are held together by
      chemical bonds.
       • Bonds involve interactions of the bonding atoms‟
         ________
                        Bonding
There are two types of bonds:
  1.   Covalent bonds – bonding atoms share
       electrons
       •   Atoms are always nonmetal atoms
       •   Covalently bonded atoms form molecules
       •   Ways to represent molecules
              Chemical formula; H2O
              Structural formula
                                           H
                                       O       O
                    Bonding
2.   Ionic bonds – attractive force among
     oppositely charged ions
     •   Bond formed between metal cations and
         nonmetal anions
     •   No molecules involved
                Ions - Terms
 Ion   – charged atom or group of atoms
     Formed when atoms gain or lose electrons


 Cation   – positively charged ion
     Formed when an atom _______ electrons
 Anion   – negatively charged ion
     Formed when an atom ______ electrons
                        Ions
 Describing   ion formation
     Cation example:




     Anion example:
      Naming Binary Compounds
 Binary compounds – compound composed
 of 2 elements
     NaCl

     CO

     CO2
  Types of Binary Compounds
 Type   I binary ionic compounds
     Metal forms only one ion
 Type   II binary ionic compounds
     Metal forms more than one ion
     Use roman numerals to indicate the charge
      on the ion
 Type   III binary covalent compounds
     Compound between 2 nonmetals
     Types I Binary Compounds
 Compound     between a metal and a
 nonmetal
    Metal forms only one ion
 Name   the cation and then the anion.
    Name of the cation is the name of the element
    Name of the anion is the name of the
     nonmetal with the ending changed to “ide”
  Monoatomic cations to know
Group # Charge on ion examples
  IA         +1      Na1+ sodium (ion)
                     K1+ potassium (ion)
  IIA        +2      Mg2+ magnesium (ion)

 IIIA        +3      Al3+   aluminum (ion)
metals
  Monoatomic anions to know
Group # Charge on ion examples
  VA         -3      N3-    nitride (ion)
                     P3-    phosphide (ion)
 VIA         -2      O2-    oxide (ion)
                     S2-    sulfide
 VIIA        -1      F1-    fluoride (ion)
                     Cl1-   chloride (ion)
                     Br1-    bromide (ion)
                     I1-      iodide (ion)
                 Practice
 Name    chemical formula

 Chemical   formula  name

				
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