Chapter 3 Main Ideas, Modern Chemistry Holt, Rinehart, & Winston, pp 67-87
I. Atomic Theory
A. Contributors
1. Democritus (400 B.C.) Said “atom” as nature’s basic particle, but did no experiments
a. Countered by (and lost to) Aristotle
2. Dalton (1790s)
a. Developed atomic theory based on quantitative analysis of experiments
All matter is composed of atoms. Atoms of same element are identical to each
other & unique from others. Atoms cannot be subdivided, created, nor
destroyed *** Atoms of different elements combine in whole-number ratios to
form cpds; Atoms are combined, separated, or rearranged in chemical rxns
B. Laws upheld
1. Law of conservation of mass
a. Chemical rxns = the combination, separation, or rearrangement of atoms
b. C + O2  CO2; 2 C + O2  2 CO
2. Law of definite proportions
a. CO always equals carbon monoxide; CO2 always equals carbon dioxide
3. Law of multiple proportions
a. CO2 always has a 2:1 ratio of O atoms to C atoms; CO has a 1:1 O:C ratio
b. In CO2, there are 2x as many O atoms as there are in CO
II. Structure of Atom
A. Discoverers
1. E. Goldstein (1886)
a. Cathode-ray tube

Gas at low pressure fills tube; currents pass from (-) cathode to the (+) anode

“Canal rays” traveled in opposite direction from anode to cathode
b. Found positively charge particles, now called “protons”
c. Now know mass of proton = 1 amu = 1.67 x 10-27 kg
2. Joseph John Thomson (1897)
a. Cathode-ray tube

Current was negatively charged particles; rays deflected from (-) charges
b. Found “electrons” – negatively charged particles w/ large charge: mass ratio
c. Plum pudding model
3. Robert A. Millikan (1909)
a. Measured charge & mass of electron (~1/2000 mass of proton)
b. Now know that mass of electron = 1/1837 mass of proton, or 9.109 x 10-31 kg
4. Ernest Rutherford (1911)
a. Gold-foil experiment with alpha particles; Hans Geiger & Ernest Marsden
b. “Nucleus” – positively charged center comprising bulk of atom’s mass
c. Nucleus: atom = marble on 50-yd line compared to the whole stadium
5. Niels Bohr (1913)
a. Electrons were in orbitals surrounding the nucleus like planets & our sun
b. Explained atomic emission-line spectrum of hydrogen
6. James Chadwick (1932)
a. Found “neutrons” as particles with no charge but with mass ~ proton = 1 amu
B. Atomic nucleus
1. Composition
a. Protons – 1+ charge, 1 amu, 1.673 x 10-27 kg
1. Number of protons = atomic number = Z; is unique for each element
b. Neutrons – No charge, 1 amu, 1.675 x 10-27 kg
1. Number of neutrons of the same element can vary, e.g., isotopes

Mass # - Atomic # = # of neutrons

Nuclides of elements = specific isotope of an element

H’s nuclides: protium (1amu);deuterium (2amu);tritium (3amu)

Written as element symbol-hyphen-atomic mass; H-1, H-2, H-3
2. Forces
a. Electric force = 1 of the 4 basic forces of the universe

Opposites attract and like repels, as a rule; occurs at long distances.

Is 100x weaker than the weak nuclear force (another basic force)
b. Weak nuclear force – attraction between nuclear particles at close range.

Between protons & neutrons

Between neutron & neutron

Between proton & proton (up to 83 in one nucleus)
c. Strong nuclear force – attraction between the quarks

Are 6 unique quarks with “flavors”, “colors”, & “direction”

Quarks comprise neutrons (3 each) & protons (different 3 each)
3. Size of atom
a. Radius = Distance between nucleus and outer realm of electron cloud
b. Expressed in picometers, i.e., 10-12 m
c. Range = 40 – 270 pm
d. Size of nucleus = 1 fm = 10-15 m or 0.001 pm
e. Density of nucleus = 2 x 108 metric tons/ cm3 [WOW!!!!]
C. Counting atoms
1. Atomic number = # protons
2. Mass number = # protons + # neutrons
3. Average Atomic Mass
a. Weighted average of masses of isotopes
b. If 69.15% is Cu-63 @ 62.930 amu & 30.85% is Cu-65 at 64.928 amu,
then 0.6915(62.930 amu) + 0.3085 (64.928 amu) = 63.55 amu for Cu
4. Relative atomic mass
a. C-12 is standard for all masses of atoms for elements of periodic table
b. 1 atomic mass unit (amu) = 1/12 of C-12 mass
c. Because there are an equal number of protons & neutrons, each is ~1 amu.

Proton’s mass = 1.007276 amu

Neutron’s mass = 1.008665 amu

Electron’s mass = 0.000 5486 amu
5. Relating mass to numbers of atoms
a. Avogadro’s number = # of particles in a mole = 6.022 x 1023 atoms/ mole
b. Mole = # particles as in exactly 12 g of C-12
c. Molar mass = mass of 1 mole of a substance *** Listed on periodic table***

1 mol of C = 12.01 g; 1 mol of H = 1.008 g

1 mol of Li = 6.941 g
d. Molecular masses on periodic table, therefore, are used as “grams/ mole”
e. Can use Avogadro’s number to calculate # atoms per grams (or moles) of X

Given: 3.01 x 1023 atoms of Ag; ? # of moles of Ag

3.01 x 1023 atoms/ 6.022 x 1023 atoms/ mole = # moles

= 0.500 mole of Ag

0.500 mole of Ag (197 grams/ mole) = 98.5 grams