Chapter 2 Atoms_ Ions_ and Molec

							     Chapter 2: Atoms, Ions, and
             Molecules
                     Outline
2.1: Atoms/Atomic Theory   2.4: Molecules/Ions
2.2: Components of Atom    2.5: Ionic Compounds
2.3: Periodic Table        2.6: Names of
                                Compounds
    Building Blocks of Chemistry
Atoms: electrons, protons, and neutrons

Molecules: fundamental unit of compounds;
Identified by formulas and names

Ions: a charged species, either positive or negative,
as in ionic compounds
2.1: Atoms and the Atomic Theory
Dalton’s model of the atom (1808):
I. Element is composed of tiny particles called
atoms. All atoms of given element have same
properties. Atoms of different elements have
different properties
2.1: Atoms and the Atomic Theory
II. In an ordinary chemical reaction, atoms move
from one substance to another, but are not created
or destroyed or converted into an atom of another
element (law of conservation of matter).
III. Compounds are formed when atoms of two or
more elements combine. In a given compound, the
proportions of the atoms are definite and constant
(law of constant proportions).
       2.2: Components of the Atom
I. Electrons: First evidence observed in studies of
      conduction of electricity through gases at low
      pressures.
      a. J.J. Thomson (1897): cathode rays produced
      consisted of a stream of negatively charged particles.
      b. Electrons are common to all atoms, carry a unit
      negative charge of –1, and have a very small mass,
      roughly 1/2000 that of the lightest atom (Plum Pudding
      Model).
                Radioactivity
Roentgen discovered X-rays in 1895, while studying
 the glow produced in certain substances by cathode
 rays. Noted glow on piece of paper some distance
 from tube; remained glowing when taken into
 another room.
           Radioactivity (Cont’d)
 Fluoresence: chemicals which   continue to glow
  after being exposed to radiation (light).
 Becquerel studied fluoresence by wrapping
  photographic paper in black paper, placing a few
  crystals of material on the paper and exposing it to
  strong sunlight. Ordinary light would not pass
  through the black paper, however, X-rays would.
          Radioactivity (Cont’d)
Looking at Uranium compounds, He discovered that
 they fogged the paper even when not exposed to
 sunlight. His graduate student, Marie Sklodowska
 (Curie), called it radioactivity and they won the
 Nobel Prize in 1903. She won in again in 1911.
         3 Types of Radioactivity
Discovered by Ernest Rutherford.
 I. Alpha (): beam of positively charged particles.
 Mass 4 times that of a hydrogen atom and a charge
 twice the magnitude of, but opposite in sign to, an
 electron
 II. Beta (): negatively charged particles,
 electrons.
 III. Gamma (): not deflected by magnetic field
              Electrons (Cont’d)
II.   Millikan’s Oil Drop Experiment: Thomson able
      to measure mass to charge ratio for an electron
      but could not measure the mass or charge.
      a. Millikan discovered the charge in 1909.
      Mass measured using the rate at which the oil
      drops fall under the influence of gravity.
      b. Mass of electron determined to be 9.1 x 10-
      28g.
      2.2: Components of the Atom
I. Protons and Neutrons: Ernest Rutherford (1911)
     a. Bombarded a piece of thin gold foil with α-
     particles (helium minus its electrons)
     b. Most of the particles passed through the foil
     with no change in direction; some, however,
     were reflected back at acute angles, suggesting
     the mass of the atom was concentrated in the
     center and was positively charged
        Components of the Nucleus
I. Proton: mass nearly equal to that of a single hydrogen
      atom. Carries a charge of (+1), equal in magnitude to
      that of an electron (-1).
II. Neutron: uncharged particle with a mass slightly greater
      than that of a proton.
Because protons and neutrons are much heavier than
      electrons, >99.9% of the mass of the atom is in the
      nucleus, even though the volume of the nucleus is
      much smaller than that of the atom.
               Atomic Number
All the atoms of a particular element have the same
  number of protons. This number is a basic
  property of an element and is called its atomic
  number and is given the symbol Z.
                Z = number of protons
In a neutral atom, the number of protons in the
  nucleus is exactly equal to the number of electrons
  outside the nucleus
                   Examples
H atom: 1 proton, 1 electron; Z = 1
C atom: 6 protons, 6 electrons; Z = 6
Fe atom: 26 protons, 26 electrons; Z = 26
               Mass Numbers
The mass number of an element, A, is found by
 adding up the number of protons and neutrons in
 the nucleus:
          A = # of protons + # of neutrons
All atoms of a given element have the same number
 of protons, hence same atomic number. They may
 differ, however, from one another in mass and
 therefore, mass number
                    Isotopes
While the number of protons in an atom is fixed, the
 number of neutrons is not (elements are
 distinguished by the number of protons in the
 nucleus).
Atoms that contain the same number of protons, but
 a different number of neutrons are called isotopes.
           Examples of Isotopes
Z = 1: Hydrogen, A = 1 (no neutrons)
       Deuterium, A = 2 (1 neutron)
       Tritium, A = 3 (2 neutrons)
Z = 92: Uranium 235, A = 235 (143 neutrons)
        Uranium 238, A = 238 (146 neutrons)
               Nuclear Symbol
Composition of the nucleus is shown by its nuclear
 symbol, where A is the mass number, Z is the
 atomic number and X is the element symbol.

                     A
                       X
                     Z
   Introduction to the Periodic Table
An element is a substance all of whose atoms have
 the same number of protons and thus the same
 atomic number. The periodic table is a listing of
 all known elements and their atomic numbers.
 a. Horizontal rows are called periods.
 b. Vertical columns are called families (groups).
     i. Designated by numbers 1-18 (IUPAC, 1985)
                   Periodic Table
Elements falling in groups 1, 2, 13, 14, 15, 16, 17, and 18
  (formerly groups 3A-8A) are referred to as main-group
  elements.
Elements in groups 3-12 are called transition metals.
Elements in group 1 are called the alkali metals.
Elements in group 2 are called the alkaline earth metals.
Elements in group 7 are called the halogens.
Elements in group 8 are called the noble gases.
        Trends in the Periodic Table
Elements in the same group have similar chemical
  properties, i.e. Li, Na, and K all react vigorously with
  water to produce hydrogen gas.
  He, Ne, and Ar do not react with any other substances.
The periodic table is an arrangement of elements, in order of
  increasing atomic number, and in horizontal rows of such
  a length, that elements with similar chemical properties
  fall directly beneath one another in vertical groups
  (Mendeleev).
           Metals and Nonmetals
The diagonal line that starts from boron separates
 metals and nonmetals; elements along this line are
 called metalloids (B, Si, Ge, As, Sb, Te).

Metals: conductive of heat and electricity,
 malleable, ductile, shiny, solids.
Non-metals: not always solids, poor conductors of
 heat/electricity
              Molecules and Ions
Molecules and Ions result from the combination of atoms
 from two or more elements.
 a. A molecule is uncharged and is usually composed of
 nonmetallic elements. The atoms are held together by
 covalent bonds (shared electrons).
 b. Represented by molecular formula, the number of
 atoms of each element indicated by a subscript written
 after the elemental symbol.
             e.g. H2O CO2 CH4 NH3 CH3OH
                        Ions
If an atom gains or loses electrons, charged particles
   called ions are formed.
   a. Metals lose electrons to form positively charged
   ions called cations.
       e.g. Na0  Na+ + 1 electron
   b. Nonmetals typically gain electrons to form
   negatively charged particles called anions.
       e.g. Cl0 + 1 electron  Cl-1
                  Ions (Cont’d)
If an ion is derived from a single atom, they are said
   to be monatomic.
Many of the most important ions in chemistry are
   polyatomic, that is, containing more than 1 atom.
             e.g. OH-1        NH4+1

A polyatomic ion is essentially a charged molecule.
              Ionic Compounds
Ionic compounds are compounds in which the ions
  are held together by an ionic bond (transfer of
  electrons).
The molecular formula for an ionic compound is
  illustrated the same way as for molecules.
      e.g. NaCl        CaCl2      KBr
Note that the metal is always shown first.
 Cations and Anions With Noble Gas
              Structure
The charges of ions formed by main-group elements can be
  predicted in a straight-forward manner.

Atoms that are close to a noble gas (Group 18) in the
  periodic table form ions that contain the same number of
  electrons as the closest noble gas atom.

The unreactivity of noble gases suggest a stable electronic
  configuration, which other atoms would like to achieve.
Trends in the Formation of Ions of the
        Main Group Elements

Group   No. of Electrons in Atom Charge Examples
  1      1 more than noble gas     +1    Na, K
  2      2 more than noble gas     +2    Mg, Ca
  16     2 less than noble gas     -2    O, S
  17     1 less than noble gas     -1    F, Cl
               Polyatomic Ions
There are only 2 common polyatomic cations:
                 NH4+, Hg2+
All other cations considered here will be derived
 from individual metal atoms.

Most of the polyatomic anions contain 1 or more
 oxygen atoms. These species are referred to as
 oxoanions
          Names of Compounds
A compound can be identified by one of 2 ways:
  chemical formula or chemical name
Nomenclature of Ions: Monoatomic cations take the
  name of the metal from which they are derived.
  e.g. Na+ sodium                K+ potassium
Complications arise from metals which can adopt
  more than one ion (oxidation state)
Metals With Multiple Oxidation States
In order to distinguish between these different ions, a
  roman numeral is used to indicate the charge
  e.g. Fe2+ iron(II)        Fe3+ iron(III)
An older system uses the endings –ic for the ion of
  higher charge and –ous for the ion of lower charge.
  e.g. Fe2+ ferrous         Fe3+ ferric
            Monoatomic Anions
Monoatomic anions are named by adding the suffix
 -ide to the stem of the name of the nonmetal from
 which they are derived.
 e.g. N3- nitride      O2- oxide       H- hydride
                       S2- sulfide     F- fluoride
                       Se2- selenide Cl- chloride
                       Te2- telluride Br- bromide
                                       I- iodide
         Multiple Polyatomic Ions
When a nonmetal forms two oxoanions, the suffix
 -ate is used for the anion with the larger number of
 oxygen atoms. The suffix –ite is used for the anion
 containing the fewest number of oxygens.
When a nonmetal forms more than 2 oxoanions, the
 prefixes per- (largest number of oxygen atoms)
 and hypo- (fewest oxygen atoms) are used.
             Ionic Compounds
The name of an ionic compound has two parts: the
  first names the cation and the second names the
  anion (same order in which the ions appear).
In naming compounds containing transition metals,
  the charge is indicated by a Roman numeral:
  Cr(NO3)3 chromium(II) nitrate
  SnCl2      tin(II) chloride
     Binary Molecular Compounds
When two nonmetals combine, the product is often a
 binary molecular compound (covalent compound).

There is no simple way to deduce the formulas of
 covalent compounds; however, there is a
 systematic way of naming them
         Nomenclature of Covalent
              Compounds
I.    The first word gives the name of the element
      that appears first in the formula; a greek prefix
      (2 = di, 3 = tri, 4 = tetra, etc) is used to show
      how many atoms of that element are present
II.   The second word consists of the appropriate
      greek prefix indicating how many atoms of the
      second element are present; the stem of the
      name of the second element and the ending
      –ide.
        Examples
N2O5

NO2

PCl3

Cl2O7
  Common Covalent Compounds
H2O               water
H2O2              hydrogen peroxide
NH3               ammonia
C2H2              acetylene
PH3               phosphine
NO                nitric oxide
CH4               methane
                      Acids
Some covalent compounds containing H atoms
  ionize in water to form H+ ions.
Such compounds are called acids. Bases ionize to
  form OH-.
An example of an acid is hydrogen chloride (HCl).
  In water, HCl ionizes to H+ + Cl-. In solution, it
  is referred to as hydrochloric acid.
                Acids (Cont’d)
Most acids contain oxygen in addition to hydrogen.
 Such acids are called oxoacids. Two of the most
 common are sulfuric (H2SO4) and nitric (HNO3).

The names of the oxoacids are derived from the
 names of the corresponding oxoanions. The –ate
 ending of the anion is replaced by –ic in the acid
 and the –ite is replaced by –ous.