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Chapter 2: Atoms
(1) Dalton’s Atomic Theory
* John Dalton’s proposal by 1808 to explain major Laws
* Consists of several postulates:
I: Atoms are the smallest, indivisible particles of matter
still retaining a chemical identity.
II: All atoms of an element are identical.
(not correct!)
III: Atoms of different elements are chemically different.
IV: Atoms can combine in integer ratios to form compounds.
V: Chemical reactions involve a rearrangement of atoms,
not a change in the atoms themselves.
(2) Atomic Theory Applications
* Helped explain the following major Laws:
I: Law of Conservation of Mass:
* Discovered by Antoine Lavoisier in 1774.
* The total mass of reactants always equals the total mass of
products during a chemical reaction.
II: Law of Constant Composition:
* Discovered by Joseph Proust in 1799.
* The mass ratio of elements in a compound are always the same,
regardless of how the compound was obtained.
III: Law of Multiple Proportions:
* Derived from Dalton’s Atomic Theory
* Certain elements can combine to form more than one compound,
but always with integer atomic ratios.
* Example of N with O: N2O, NO, NO2 or N2O5
* The Atomic Theory led to major experimental efforts to find the atom.
(3) Atomic Discoveries
Ernest Rutherford
1909
Gold foil experiment
Ernest Rutherford
1919
+, 1.007 amu
nucleus
James Chadwick
1932
neutral, 1.008 amu
subatomic particles
p, n and e–
J.J. Thomson
1897
– , 0.0006 amu
Cathode ray tubes
(4) Isotopes:
* Same # protons and electrons; different # neutrons
* Use isotopic notation (2 types)
Mass Number, A: # nucleons (protons + neutrons)
235
U
92
or
U–235
Atomic Number, Z: # protons in nucleus
# neutrons = A – Z
Example #1: How many of each subatomic particle does an isotope
of Am–241 have (used in smoke alarms).
# protons = 95
# electrons = 95
(from the Periodic Table)
(same as #p for a neutral atom)
# neutrons = 241 – 95 = 146
(calculated)
Example #2: How many of each subatomic particle does the following
isotope have (used in bone scans): 85Sr2+
38
# protons = 38
# electrons = 36
(from the notation or Periodic Table)
(lost 2 e– due to positive cation)
# neutrons = 85 – 38 = 47
(calculated)
(5) The Periodic Table
* There are several ways to organize the Periodic Table:
I: Representative Elements (Main-Group or Type A)
Transition Metals (Type B)
Inner Transition Metals (Lanthanides & Actinides)
II: Groups (columns) and Periods (rows)
III: Metals, Nonmetals and Metalloids
IV: Families:
Group IA: Alkali Metals
Group IIA: Alkaline Earth Metals
Group VIA: Chalcogens
Group VIIA: Halogens
Group VIIIA: Noble Gases
(6) The Mole
* Used for counting large # of particles
* Avogadro’s Number = 6.022 x 1023 particles mol–1
* A conversion factor between # particles and moles
(7) Atomic Mass
* Weighted average over all isotopes
* Conversion factor between mass and # atoms or mol
* Units are g/mol or amu/atom
Example: What mass would 3.20 x 1024 atoms of iron have:
3.20 x 1024 Fe atom
1 mol Fe
6.022 x 1023 Fe atoms
55.845 g Fe = 296 .75 g Fe
1 mol Fe
3 SF
Avogadro’s #
Atomic
Mass
297 g Fe
Determining Atomic Masses
* Averaged over all isotopes and based on 2 factors:
1) Isotopic Masses (mi)
2) Fractional Abundance (fi) = % / 100.
Avg atomic mass = Σ fimi
Example: What is the average atomic mass of Silicon, provided:
Si–28 = 27.97693 u and is 92.23 %; Si–29 = 28.97649 u and is 4.67 %;
and Si–30 = 29.97376 u and is 3.10 %. What is the average atomic mass of Si?
at. mass = 92.23 27.97693 u + 4.67 28.97649 u + 3.10 29.97376 u
100
100
100
at. mass =
25.80¦31 u + 1.35¦32 u + 0.929¦18 u
= 28.08¦54 u -----> 28.09 u