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					 Ch 3 – ATOM: THE BUILDING BLOCKS
             OF MATTER
I. HISTORY OF CHEMISTRY
   A. LEUCIPPUS, DEMOCRITUS – World is made of empty
      space and tiny particles called atoms (Greek - Atomos).
   B. ARISTOTLE – Matter was continuous substance called Hyle,
      and was not made up of smaller particles.
   C. Before 16th century
         Alchemy: Attempts (scientific or otherwise) to change
      cheap metal into gold
   D. 17th century
         ROBERT BOYLE
      a. First chemist to perform quantitative experiments
      b. Published “The Skeptical Chymist”
      c. Definition of element became generally accepted
   E. 18th century
         GEORGE STAHL (1660-1734) – “Phlogiston” flowed
      out of the burning material.
         JOSEPH PRIESTLEY (1733-1804) – Discovered oxygen
      gas and was found to support combustion.
   F. ANTOINE LAVOISIER
         Law of Conservation of Mass: Matter is neither created
         nor destroyed
         Combustion involved oxygen, not “phlogiston”
   G. JOSEPH PROUST
        Law of Definite Proportion – Specific substances always
        contain elements in the same ratio by mass
   H. JONS JACOB BERZELIUS
      1. Invention of a simple set of symbols for the elements,
         along with a system for writing formulas of compounds
      2. Discovered elements Cerium, Thorium, Selenium and
         Silicon
CH 3 – CHEMISTRY 1                                            1
II. DALTON’S ATOMIC THEORY (“BILLIARD BALL” MODEL)
    A. John Dalton (1808) – proposed an explanation for
       the law of conservation of mass, law of definite
       proportion and law of multiple proportions
   B. Postulates of Dalton’s Theory
      1. All matter is composed of extremely small particles
         called atoms
      2. Atoms cannot be subdivided, created or destroyed
      3. Atoms of a given elements are identical
      4. Atom of different elements combine in simple whole
         number ratios to form chemical compounds (Law of
         Multiple Proportion)
      5. In chemical reactions, atoms are combined,
         separated or rearranged
   C. Modern Atomic Theory
      1. Not all aspects of Dalton’s atomic theory have proven
         to be correct
      2. Dalton’s atomic theory had been modified
         a. All matter is composed of atoms, which are
            divisible into even smaller particles (subatomic
            particles)
         b. A given element can have atoms with different
            masses




CH 3 – CHEMISTRY 1                                             2
III. STRUCTURE OF THE ATOM
   A. Two Regions of the Atom
      1. Nucleus
         a. Located near the center of the atom
         b. Has a positive charge
         c. Occupies a very small part of the atom and
         d. Has high density
         b. Composed of protons and neutrons
      2. Electron Cloud
         a. Region occupied by electrons
         b. Surrounds the nucleus of the atom
         c. Occupy most of the volume of the atom
   B. Discovery of the Electron
      1. J.J. Thomson – Cathode Ray Research (1897)
         a. Cathode rays consist of small negatively charged
            particles called electrons.
         b. Determined the electron’s charge-to-mass ratio
         c. Plum Pudding Model – Electrons are like raisins
            dispersed in a pudding (the positive charge cloud)
      2. Robert Millikan – Oil Drop Experiment (1909)
         a. Determined the charge of the electron and
            confirmed that the electron carry a negative
            charge
         b. Calculated mass of electron – 1/1840 mass of
            Hydrogen atom (9.11 x 10-31 kg)
   C. Discovery of the Nucleus
      1. Gold Foil Experiment (1911) - by Ernest
         Rutherford, Hans Geiger and Ernest Marsden

CH 3 – CHEMISTRY 1                                           3
         a. A thin sheet of gold foil was bombarded with alpha
            particles
         b. Most of the alpha particles passed straight through
         c. Many particles were deflected at large angles
      2. Conclusions
         a. Most of the alpha particles passed directly through
            the foil because the atom is mostly empty space
         b. The deflected alpha particles are those that had a
            close encounter with the massive center of the
            atom
         c. Nuclear Atom Model – an atom with a dense
            center of positive charge (nucleus) with electrons
            moving around the nucleus
   D. Composition of the Nucleus
      1. Protons – positively charged subatomic particle.
         a. Eugene Goldstein (1886) – discovered a beam
            consisting of positive particles called “protons”
            using a modified cathode ray tube.
         b. J.J. Thomson – protons have the same amount of
            electrical charge as an electron but opposite in
            charge (+1)
            1. Calculated the mass of proton – 1840 times that
               of electron (1.673 x 10-27 kg)
      2. Neutrons – subatomic particles with no charge.
         a. James Chadwick (1932) – uncharged particles
            with a mass slightly greater than protons were
            formed when Beryllium was bombarded with high-
            energy alpha particles; called it neutrons
            1. Mass = 1.675 x 10-24 kg
CH 3 – CHEMISTRY 1                                            4
      3. Nuclear Force – a short-range proton-neutron,
         proton-proton, or neutron-neutron force that holds
         the nuclear particles together
      4. Nucleus has high density – 2 x 108 metric tons/cm3
      5. Properties of Subatomic particles – p.74
   E. Other Subatomic Particles
      1. Leptons (light)
         a. Electron, Positron, Neutrino, Muon, Tau
      2. Hadrons (heavy)
         a. Baryons – Protons, Neutrons
         b. Meson – Quarks (Up, Down, Charm, Strange,
             Bottom, Top)
             Quarks are held together by gluons
          c. Antiparticles – mirror image particles
             (Ex.: electron and positron)
   F. Sizes of Atoms
      1. Electron cloud – region occupied by the electrons;
         cloud of negative charge
      2. Atomic Radius – distance from the center of the
         nucleus to the outer portion of electron cloud
      3. Atomic radii range from 40 to 270 picometer (pm)
   G. Forces in Atoms
      1. gravitational           3. weak nuclear
      2. strong nuclear          4. electromagnetic
IV. COUNTING ATOMS
  A. Atomic Number (Z) – number of protons in the
     nucleus of each atom
     1. Determines the identity of the element
     2. Henry Moseley (1913)
CH 3 – CHEMISTRY 1                                        5
         a. Studied X-rays produced in X-ray tubes with
            anodes of different metals
         b. X-ray wavelength depend on number of protons
         c. Determined the atomic numbers of elements
   B. Mass Number (A) – the total number of protons and
      neutrons in the nucleus of the atom
   C. Isotopes – atoms of the same element that have the
      same number of protons but different number of
      neutrons
      1. Usually identified by specifying their mass number
         (hyphen notation)
         EX: Carbon-12,       Carbon-14
              Uranium-235, Uranium 238
      2. Nuclide – the general term for any isotope of any
         element
      3. Nuclear symbol – shows the composition of the
         nucleus          12    14            235    238
                        6   C   6   C      92   U   92   U
   D. Relative Atomic Masses
      1. Carbon-12 nuclide – the standard used for units of
         atomic mass
      2. Atomic Mass Unit – 1/12 the mass of a carbon-12
         atom
   F. Average Atomic Mass
      1. The weighted average of the atomic masses of the
         naturally occurring isotopes of an element
      2. Weighted average reflects both the mass and the
         relative abundance of the isotopes as they occur in
         nature.
CH 3 – CHEMISTRY 1                                             6
                              (%abundance A  AtMass A)  (%Abundance B  AtMass B)
           AveAtomicM ass 
                                                      100

                                                                               63
 The element copper is found to contain the naturally occurring isotopes       29 Cu   and
 65
 29  . The relative abundances are 69.1% and 30.9% respectively. Calculate the
      Cu

 average atomic mass of copper.




F. MASS, MOLE & MOLAR MASS
      1. Mole (mol) – amount of substance that contains as
         many particles as there are in 12 g of carbon-12.
      2. Avogadro’s Number (6.022 x 1023 units)
         a. Number of particles in exactly one mole of a pure
            substance
         b. Units can be atoms, ions, molecules, formula units
      3. Molar Mass (g/mol)
        a. Mass in grams of one mole of a pure substance
         b. It is equal to the atomic mass of the element
   G. GRAM/MOLE CONVERSIONS
         1 mol = Molar Mass in grams (Atomic Mass)
         1 mol = 6.022 x 1023 atoms
            1. Mol to Gram
           What is the mass in grams of 3.59 mol of the element Silicon, Si?




CH 3 – CHEMISTRY 1                                                                           7
      3. Gram to Mol
A chemist produced 15.0 g of aluminum, Al. How many moles of aluminum were
produced?




      3. Atoms to Mol
How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver, Ag?




      4. Mol to Atoms
How many atoms of Nickel, Ni, are in 5 mol of Ni?




      5. Atoms to Mass
What is the mass in grams of 7.5 x 1015 atoms of Zinc, Zn?




      6. Mass to Atoms
How many atoms of Sulfur, S, are in 12.5 g of sulfur?




CH 3 – CHEMISTRY 1                                                           8

				
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