Chapter 2 Atoms, Molecules Ions

							Chapter 2 Atoms, Molecules & Ions
2.1 The Atomic Theory of Matter
Dalton’s Theory 1807
1) Elements are composed of small particles called atoms.
2) Atoms of an element are identical; atoms of different elements are different.
3) Atoms cannot be created or destroyed during chemical reactions.
4) Compounds are formed when atoms of different elements combine.
Law of Conservation of mass: during a chemical reaction, mass is conserved.
Total mass of the reactants = Total mass of the products
Law of Multiple proportions: If 2 elements A and B combine to form more than 1
compound, the mass of B which combines with a mass of A is a ratio of small whole
numbers. E.g. H2 O and H2 O2
2.2 The Discovery of the Atomic Structure
Electrons
1. Results of Thomson's experiments (1890's) on the behavior of cathode rays in
    electric & magnetic fields:
   • cathode rays consist of negatively charged particles called electrons
   • charge to mass ratio of e- = 1.76 x 108 C/g
2. Results of Millikan’s oil drop experiment:
    • charge of e- = 1.60 x 10-19 C; mass of e- = 9.11 x 10-28 g
Radioactivity: spontaneous emission of radiation
1. Becquerel discovered radioactivity in a Uranium pitchblende sample (1896)
2. Madame Curie discovered polonium and radium (early 1900’s)
3. Results of Rutherford's experiments on radiation (1910-1920)
   • An atom has a nucleus - a small, dense, positively charged region.
   • Electrons are located outside of the nucleus. Most of the atom is empty
     space.
   • Three types of radiation exist:
     alpha: high mass particles with +2 charge (helium nuclei)
     beta: low mass particles with -1 charge (electrons)
     gamma: neutral, high energy radiation similar to x-rays
2.3 The Modern View of Atomic Structure
Nuclear Model:
1) Atoms consist of protons, electrons & neutrons (3 subatomic particles).
2) Protons & neutrons are located in the nucleus; it contains most of the mass.
   Size analogy: If an atom were the size of a football stadium, nucleus would be
   the size of a marble.
3) Electrons move rapidly in region outside of the nucleus.
•   Masses of atoms are so small that we use the atomic mass unit (amu) to scale
    up the numbers.         1 amu = 1.66054 x 10-24 g
Electron: -1 charge, mass = 9.11 x 10-28 g = 5.486 x 10-4 amu, Thomson (1897)
proton: +1 charge, mass = 1.672 x 10-24 g = 1.0073 amu, Rutherford (1919)
neutron: no charge, mass = 1.675 x 10-24 g = 1.0087 amu, Chadwick (1932)
Isotopes
• The number of protons in an atom defines what an element is. This is the
   atomic number, Z. (top # on periodic table)
Isotopes are atoms of an element that have a different number of neutrons.
Isotopes of an element have the same atomic number, but a different mass
number.
Mass number = A = number of nucleons (protons & neutrons) in nucleus
nuclide: atom of a specific isotope
•   Isotope symbol for element X        Z X
                                        A
                                                        A = p + n; Z = p
•   For neutral atoms: # protons = # electrons
E.g. 3 H isotopes: normal H: 1 H
                               1           deuterium: 2 H
                                                        1           tritium: 3 H
                                                                             1

                  stable, 1 p, 0 n         stable, 1 p, 1 n         radioactive, 1 p , 2 n
2.4 The Periodic Table -1st table 1869
Features of modern periodic table:
1) Elements arranged in order of increasing atomic number.
2) Horizontal Rows in periodic table are called periods. 7 periods exist
3) Vertical Columns are groups or families; elements have similar properties.
Group names:      Group 1A: alkali metals       Group 2A: alkaline earth metals
                  Group 7A: halogens            Group 8A: noble gases
4) representative elements: A Group; transition elements: B Group
5) Metals are located to the left of the stair-step line. Nonmetals are located to
   the right of the stair-step. Elements located at the stair-step are intermediate
   in character - semiconductor or metalloid elements: B, Si, Ge, As, Sb, Te, At
Physical state of elements at 25 °C & 1 atm:
      Gases: O2, N2, H2, F2, Cl 2, and Noble gases
      Liquids: Br2, mercury
      Solids: everything else
2.5 Molecules and Molecular Compounds
molecule: 2 or more atoms bonded together; discrete entities.
•   Many elements exist as diatomic molecules: H2, N2, O2, F2, I2, Cl 2, Br2
Molecular compounds consist of nonmetal elements.
Molecular formulas give the actual numbers and types of atoms in a molecule.
E.g. CH4, H2O2, C2H4, C6 H12 O6
Empirical formulas give the smallest whole number ratio of atoms in a molecule.
E.g. CH4, HO, CH2, CH2 O
2.6 Ions and Ionic Compounds
Many chemical reactions involve transfer of electrons between atoms:
      Metal atoms tend to lose electrons & form + charged cations.
      Nonmetal atoms tend to gain electrons & form - charged anions.
Generally, atoms gain or lose enough electrons to have same number of electrons
as nearest noble gas:
                 Group            1A    2A   3A    5A    6A  7A
                 Charge of ion    1+    2+   3+    3-    2-  1-
Ionic compound: consists of metals and nonmetals (or polyatomic ions); ionic
compound is a long 3-D array of cations & anions; not individual molecules.
Ionic formulas: the number of electrons lost & gained must be equal, so + and -
charge cancel out.
Rules for writing ionic formula:
1) Write down formulas of ions
2) Combine the smallest # of ions to give the charge sum equal to 0; if the
   charges are not equal, find the lowest common multiple
      E.g. Predict the formula for the compound formed from the following elements:
      Ca & O:       Ca 2+O2- → CaO
      Mg & N        Mg2+N3- → Mg3 N2
      Al & Cl       Al 3+Cl - → AlCl 3
Ions have a different # of electrons & protons

Ex.   23
      11   Na+ : 11p; 12 n; 10 e-            35
                                             17   Cl - : 17 p; 18 n, 18 e-

NOMENCLATURE- IONIC COMPOUNDS
A. Naming Cations:
1. Fixed charge metals: Cations have same name as the metal element. Groups
   1A, 2A, Al, Ag, and Zn are fixed charge metals – cations that have 1 specific
   charge. E.g. Ag+ silver ion     Zn2+ zinc ion          Al 3+ Aluminum ion
           Li lithium ion
             +
                                   Ca 2+
                                         calcium ion
2. Variable charge metals: If the metal can form more than 1 cation, the charge is
   indicated by a Roman numeral in parenthesis after the metal name. Most of the
   transition metals are variable charge metals.
      E.g. Common metals which exist in more than one positive state:
      Fe 2+ iron(II)   Au+ gold(I)            Cu+ copper(I)   Hg 2 + mercury(I)
                                                                 2

      Fe 3+ iron(III)  Au3+ gold(III)    Cu2+ copper(II) Hg2+ mercury(II)
3. Polyatomic Cations: consist of nonmetals
     H3O+ hydronium        NH + ammonium
                               4

B. Naming Anions
1. monoatomic anions: change ending to -ide
     E.g. oxygen → oxide sulfur → sulfide                   hydrogen → hydride
2. Polyatomic anions: most end in -ate or -ite; usually contain O (oxy)
     Know polyatomic anions on handout.
     a. Rule for naming oxy series anions:
     per-....-ate     1 more O than -ate
            -ate
            -ite      1 less O than -ate
     hypo-...-ite     2 less O than -ate
     b. If H+ is added to a polyatomic ion, write hydrogen (or bi-) in front of name.
          −
     HCO 3            hydrogen carbonate or bicarbonate
           H2PO −
                4     dihydrogen phosphate
NOMENCLATURE RULES

I. IONIC COMPOUNDS contain cations & anions
1) Name metal cation.
2) Include Roman numeral in parenthesis ONLY IF metal has variable charge. Fixed
    charge metals: Group 1A, 2A, Ag, Zn, and Al; others are variable.
3) Name anion.
    E.g.   MgBr2 magnesium bromide           PbS lead(II) sulfide
           barium nitride Ba N → Ba3N2
                             2+ 3-
                                             iron(III) sulfite Fe 3+SO 3 − → Fe 2 (SO3)3
                                                                       2


           Ca(ClO2)2 calcium chlorite        Cr2(CO 3)3 chromium(III) carbonate

II. Binary Molecular compounds: contain 2 nonmetals
1) Name 1st element & use a prefix (table 2.6) to indicate the number of atoms.
     Note that mono- is never used for the first element.
3) Name 2nd element & include prefix for number of atoms (see table 2.6).
4) Change ending of 2nd element to –ide.
     E.g.    N2O5 dinitrogen pentoxide               ICl 3 iodine trichloride
             tetraphosphorus hexasulfide P4S6        dibromine heptaoxide Br2O7

III. Acid: substance that yields H + ions in aqueous solution
A. Binary Acids (Nonoxy acids): contain H & 1 nonmetal
1) Name hydrogen as hydro-.
2) Name nonmetal & change ending to -ic acid.
   E.g.     HCl (aq) hydrochloric acid                 H2S(aq) hydrosulfuric acid
            hydrophosphoric acid H P → H3P(aq)
                                       + 3-


            hydroselenic acid H+Se2- → H2Se(aq)

B. Ternary Acids (Oxy Acids): contain H & polyatomic ion
1) no hydro prefix
2) Change ending of polyatomic ion:
      -ate → -ic acid
      -ite → -ous acid
      E.g. HNO3(aq) nitric acid                  H2SO4(aq) sulfuric acid
            H2SO3(aq) sulfurous acid             H3PO4(aq) phosphoric acid
            chlorous acid H ClO 2 → HClO2
                             +    −
                                                 perbromic acid H+BrO − → HBrO 4
                                                                         4




                                        ELEMENTS
                   st
Know names of 1 36 elements (H – Kr) &
Ag silver    Au gold            Ba barium           Pb lead        Pt platinum
I iodine     Sn tin             Hg mercury          U uranium      Cd cadmium
Know these Polyatomic ions

Cations:

Ammonium         NH +
                    4


Hydronium        H3O+


Anions:

Acetate          C2H3O −
                       2


Carbonate        CO 3 −
                    2
                                      Bicarbonate          −
                                                       HCO 3

Oxalate          C2O 2 −
                     4


Cyanide          CN-

Thiocyanate      SCN-                 Perchlorate      ClO −
                                                           4


Hydroxide        OH-                  Chlorite         ClO −
                                                           2


Chlorate             −
                 ClO 3                Hypochlorite     ClO-
                    −
Nitrate          NO 3                 Nitrite          NO −
                                                          2


Sulfate          SO 2 −
                    4                 Sulfite          SO 3 −
                                                          2



Phosphate        PO 3 −
                    4                 Phosphite        PO 3 −
                                                          3


Permanganate     MnO −
                     4


Chromate         CrO 2 −
                     4


Dichromate       Cr2O 7 −
                      2




*Note: Names of ions on the right can be derived from the lefthand group. The
oxy-anions for bromine and iodine can be named in a manner analogous to that
listed for chlorine.
                                          NOMENCLATURE Worksheet

A. Write the correct formula:

1. iron (III) sulfide            Fe3+S2- →         Fe2S3

2. silver dichromate Ag+Cr2O 7 − → Ag2Cr2O7
                             2




3. sodium phosphide              Na+P3- →          Na3P

4. cobalt (III) nitrite          Co3+NO − → Co(NO2)3
                                        2



5. tin(IV) perchlorate           Sn+4ClO −
                                         4        → Sn(ClO4)4

6. diphosphorus pentasulfide              P2 S5

7. calcium phosphite             Ca2+PO 3 − → Ca3(PO3)2
                                        3



8. magnesium permanganate Mg2+MnO − → Mg(MnO4)2
                                  4



9. chlorous acid                 H+ClO − → HClO2(aq)
                                       2



10. hydrosulfuric acid           H+S2- →           H2S(aq)

B. Write the correct name:

1. S 4O8                  tetrasulfur octoxide

2. AlH3                   aluminum hydride

3. Cr(SCN)3               chromium(III) thiocyanate

4. PbO2                   lead(IV) oxide

5. HBr(aq)                hydrobromic acid

6. Zn(HSO4)2 zinc bisulfate (or zinc hydrogen sulfate)

7. MnC 2O4                manganese(II) oxalate

8. NH4C2H3O2              ammonium acetate

9. H2CO3(aq)              carbonic acid

10. Fe(BrO)2              iron(II) hypobromite