CHAPTER #2 - Atoms, Molecules, and Ions
2.1 The Early History of Chemistry Ancient Greeks - thought matter was composed to 4 substances – earth,
air, fire, water (_________ vs. _________)
Alchemy - ______________________________
(discovered many elements; learned to prepare mineral acids)
Metallurgy - extraction of metals from ore
Robert Boyle - (1661 - Skeptical Chymist) - first quantitative experiments;
current concept of “element”
Georg Stahl – suggested “phlogiston” flowed out of burning material
Joseph Priestley - (1733-1804) “discovered” oxygen (not phlogiston)
2.2 - 2.3 Fundamental Chemical Laws/ Dalton’s Atomic Theory NATURAL LAWS:
1. Conservation of Mass - In any chemical reaction, mass is neither
created nor destroyed
2. Constant Composition – a given compound always contains exactly the
same proportion of elements by mass.
3. Multiple Proportions - When two elements form a series of compounds,
the ratios of the masses of the 2nd element that
combine with 1 gram of the 1st element can
always be reduced to small whole numbers.
Mass Oxygen
w/ 1 g C
Cpd I
1.33 g
Cpd II
2.66 g
Mass Nitrogen
w/ 1 g O
Cpd I
1.750 g
Cpd II
0.8750 g
Cpd III 0.4375 g
John Dalton (1808) “Father of Atomic Theory”
Essentials of his theory. . .
1. An element is composed of tiny particles called atoms. All atoms of a
given element show the same chemical properties.
2. Atoms of different elements have different properties. In an ordinary
chemical reaction, no atom of any element disappears or is changed
into an atom of another element.
3. Compounds are formed when atoms of two or more elements combine.
In a given compound, the relative numbers of atoms of each kind are
definite and constant. In general, these relative numbers can be
expressed as integers or simple fractions.
IN GENERAL Elements consist of tiny particles called _________, which retain their
identity in ____________________. In a compound, atoms of two or
more elements are combined in a fixed ratio of ___________________.
Gay-Lussac - (1809) Combining Volumes of Gases
• Performed experiments in which he measured (under same T & P) the
___________________ that reacted with each other.
Amadeo Avogadro - (1811) “Avogadro’s Hypothesis”
• Equal volumes of different gases contain the ______________________
at the same temperature and pressure.
2.4 Early Experiments to Characterize the Atom
JJ Thomson (1856 - 1940): English physicist
____________ Experiment (late 1890’s)
Important Observations
Inferences
This lead to the _____________ Model
Robert Milikan (1868 - 1953)
_____________ Experiment (1909)
Important Observations
Inferences
Sir James Chadwick (1891-1974): confirmed existence of ____________
Henri Bequerel (1896) French Chemist
essentially discovered _________________
(alpha, beta, and gamma particles discovered later)
Ernest Rutherford (1911) ________________ experiment
Important Observations
Inferences
2.5 - 2.6 Modern View of Atomic Structure/ Molecules and Ions COMPONENTS
Proton
Neutron
Electron
Nuclear Symbol
Relative
Mass
_____
_____
_____
Relative
Charge
_____
_____
_____
A
ZX
Location
_______
_______
_______
Actual
Mass
Actual
Charge
1.6727 x 10-24 g
1.6022 x 10-19 C
1.6750 x 10-24 g
__0____
-1.6022 x 10-19 C
9.1095 x 10-28 g
(X = element symbol)
• Z = Atomic Number = number of protons and is UNIQUE
• A = Mass Number = sum of neutrons and protons
is NOT unique
• Atoms and Molecules are electrically neutral ---> # protons = # electrons
Ions - Have an electrical charge due to gain/loss of electrons
cations anions isoelectric –
K+, Ca2+, P3-, S2-, Cl-, Ar
Isotopes Isotope
Mg-24
Mg-25
Mg-26
Mass (amu)
23.9850
24.9858
25.9826
Percent (Abundance)
78.99%
10.00%
11.01%
Determine the relative average atomic weight of Magnesium.
COMPLETE THE FOLLOWING TABLE:
31
# electrons
# protons
# neutrons
15P
_____
_____
_____
40
Ca2+
_____
_____
_____
74
As3_____
_____
_____
38
F2
_____
_____
_____
The Periodic Table
Naming – Ionic Cpds, Molecular Cpds, & Acids – see next
H3O+
_____
_____
_____