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Atoms-Molecules-Ions

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					Chapter 2 Atoms Molecules and Ions

Atomic Theory of Matter

Democritus – Proposed the idea that matter is made up of small, indivisible
particles (atoms from the Greek atomos indivisible). This idea not accepted by
his contemporaries (e.g., Plato, Aristotle).

Evidence mounted for the existence of atoms Note => 1808; John Dalton
formulated his ideas (atomic theory); basically 3 postulates.

DALTON'S ATOMIC THEORY

Postulates

1.     Elements made up of ATOMS. All atoms of same element are alike;
       atoms of different elements are different.

2.     Chemical Reactions: Separation and joining of atoms. No atom is
       created or destroyed, no conversion of an atom from one element to
       another.

       –     explains LAW OF CONSERVATION OF MASS

3.     A chemical COMPOUND is the result of the combination of 2 or more
       elements in a simple numerical ratio.

     – explains LAW OF DEFINITE PROPORTIONS


Notes on Dalton’s Atomic Theory

1. Properties of 1 atom of Na different from 1 atom of Cl but the properties of 1
   atom of Na are no different than 1 gram of Na (we just have more of the
   same).
2. We can't have compounds like CO½; CH3/5; OH1/6; CO1/3.
3. Alternate statement for the law of conservation of mass
                                                                               2

Atomic Structure

 Based on Dalton, we can define an atom as the basis unit of an element that
  still retains the properties characteristic of the element (can enter into
  chemical combination).
 Atoms are not the smallest structures; they possess an internal structure.
 We now know that atoms can be split and created in nuclear reactions.
 Atoms consist of protons, neutrons, and electrons; these are the "sub–
  atomic“ particles.
 J.J. Thomson's cathode ray tube experiment determined the mass/charge
  ratio of the electron.
 Later experiments by R.A. Milliken estimated the electronic charge as 1.60 x
  10–19 C.

 Thomson's number i.e., – 1 g / 1.76 * 108C * 1.60 x *19 C
                                 = 9.09 * 10–28g


X–Rays and Radioactivity

 Radioactivity – Spontaneous emission of particles and/or radiation.
  Radioactive substances (radionuclides) break down or decay.

 Roentgen – Cathode rays striking glass and metals resulted in new and
  unusual rays (x–rays).

 Becquerel – Certain uranium compounds darkened photographic plates.
  The radiation resembled X–rays, but it didn't consist of particles [–rays
  (gamma)] (Marie Curie suggested the term "radioactivity” for this
  phenomenon).

     Note – units of radioactivity are the Becquerel (Bq) and the Curie (Ci)
     named in honour of these pioneers in radioactive chemistry.
                                                                                 3

Radioactive particles

      – particle – 24He (Mostly from Rutherford’s work)
       – particle – e–
       – ray – energy released when the metastable radionuclide returns to
       the ground state

Proton and Nucleus

 Thomson's Model – Uniform, positive sphere in which electrons are
  embedded.

 Rutherford's Model – the result of the Rutherford, Marsden, and Geiger
  experiment.

The  – Scattering Experiment

 Since most of the  – particles passed through the gold fail undeflected, he
  concluded that the nucleus was only a small fraction of the total volume of
  the atom, but it did contain most of the mass. The electrons occupied the
  large volume outside the nucleus.

 Mass of proton ~2000 times mass of electron.

 FACT: He: H problem Mass 4:1 (1840) charge 2:1

 Chadwick – 1932; solved the problem of the "missing nuclear mass” by
  proposing the existence of neutrons in the nucleus.

 FACT: Mass of neutrons = mass of protons, however, the neutrons possess
  no charge, while the proton possesses a charge of +1.

 Protons and electrons have the same magnitude of charge, but opposite sign
                                                                               4

Mass Relationships of Atoms

 Atomic Symbols
      A
      Z   Symbol
 where
     A = mass number = number of protons + neutrons
     Z = atomic number = number of protons
     number of neutrons = A – Z

Examples
                                       79
                                       35   Br
      A bromine atom with 35 protons and (79–35) = 44 neutrons
                                       99
                                       43   Tc
      Technetium–99 (a synthetic element) – used heavily in nuclear
       medicine; 43 protons and 56 neutrons

                                       97
                                       43   Tc
      Technetium – 97 – technetium atom with 43 protons 54 neutrons

 Isotopes – atoms of a given element that differ in the number of neutrons,
  and hence, the mass number A.
 An atom of a specific isotope is called a nuclide

e.g. three isotopes of hydrogen

      1
      1   H      Protium – name is rarely used for atomic hydrogen.
      2
      1   H      Deuterium – the D in D2O or heavy water.
      3
      1   H      Tritium – radioactive isotope of hydrogen.
                                                                              5

Molecules and Chemical Formulas

 Molecules – aggregates of atoms joined together in a definite arrangement
  by chemical forces
 Chemical Formula – combination of symbols representing the atoms in a
  molecule of a compound. Describes not only the type of atoms but also the
  number of atoms, e.g., H2O2.
 Molecular Formula – the exact # of atoms of each element in a molecule;
  CO2; N2O; H2O; O2; N2; H2; HBr.
 diatomic molecules – two – atoms; e.g., O2; N2; H2
 polyatomic – molecules with more than two atoms; e.g., MnO2; MnCl2; O3;
  NH3
 Allotropes – different forms of the same element
            O2; O3 (diamond, graphite, "buckyballs")

  Molecule vs. compound

 Molecule – unit of substance composed of two or more atoms of the same
  element or of different elements.
 Compound – a substance composed of atoms of two or more elements. 64
      Cl2; H2 – they are molecules but not compounds.
      NH3 – molecule and compound.
 Empirical Formula – Expresses the simplest whole number ratio of the
  atoms in a substance.
            H2O2 – ratio of H:O = 2:2 or 1:1; empirical formula = HO
            N2H4 – ratio of H:N = 4:2 or 2:1; empirical formula = NH2
 Many molecules have the same molecular (or true) formula and the empirical
  formula, e.g., H2O.
 Often we write the molecular formula showing how its atoms are joined
  together – structural formula

Ions and Ionic Compounds
 Ions – atoms or groups of atoms that possess a charge. Ions are formed
   when electrons are add or removed form a neutral atom (or molecule).
 Example        Li atom, Z = 3, 3 protons 3 electrons
 Li ion
     +
                 3 protons 2 electrons; cation – positively charged ion.
                                                                                  6

   Ca2+       (20 protons 18 electrons)
   Br(atom) 35 protons;         35 electrons
        –
   Br         35 protons; 36 electrons;     anion – a negatively charged ion.
   monatomic ions – ions that contain only one type of atom
   polyatomic ions – ions containing two or more types of atoms, e.g., SO42–
    (sulfate ion), CO32– (carbonate), NH4+ (ammonium ion) SO32– (sulphite), OH–
    (hydroxide ion). (note: Table 2.4 in text, very important).
   Compounds containing cations and anions are ionic compounds e.g. solid
    sodium chloride (NaCl) contains an equal # of anions (Cl–) and cations (Na+),
    i.e., the solid compound is electrically neutral.
   In most cases, ionic compounds contain an metallic ion as the cation and a
    non–metallic ion as the anion. exception NH4+ (IMPORTANT EXCEPTION)
   Na  Na+ +e– Na+ + Cl– NaCl (ionic compound) note that discrete units of
    NaCl don't really exist.
   We generally write NaCl for the ionic compound formed between Na and Cl;
    the formula unit = NaCl

       NOTE: when we write the empirical formula for an ionic compound, we
        must know the changes of the ions comprising the compound. All we
        have to do is keep the total # of negative charges equal to the total # of
        positive charges.

             NaCl      1 Na+ 1 Cl–
             BaCl2     1 Ba2+: need 2 Cl– ions to to preserve electrical neutrality
             CaCl2     Ca  Ca2+ + 2 e–; Cl + e–  Cl–

				
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