Download atoms

The Early History of Chemistry

Greeks 400 B.C.
Four fundamental substances – fire, earth, water, a
air.
-

Democritus – uses term “atomos” (atoms) to
describe small, indivisible matter. No experiments to
support the idea, so it is dropped.
Before 16th Century
–
Alchemy: Attempts (scientific or otherwise) to
change cheap metals into gold
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
1
The Early History of Chemistry

17th Century
–
Robert Boyle: First “chemist” to perform
quantitative experiments (pressure/volume)
–- Incorrectly believed that the alchemist’s view that
metals were not true elements and that a way would
eventually be found to change one metal into another.

18th Century
–
George Stahl: Phlogiston flows out of a burning
material.
–
Joseph Priestley: Discovers oxygen gas,
“dephlogisticated air.”
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
2
Law of Conservation of Mass
 Discovered
 Mass
by Antoine Lavoisier
is neither created nor destroyed
 Combustion
involves oxygen, not
phlogiston
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
3
Other Fundamental Chemical Laws
Law of Definite Proportion
A
given compound always contains exactly
the same proportion of elements by mass.
– always 39.34% Cl and 60.66% Na
(mass)
 NaCl
 Carbon
tetrachloride is always 1 atom
carbon per 4 atoms chlorine.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
4
Other Fundamental Chemical Laws
Law of Multiple Proportions
 When
two elements form a series of
compounds, the ratios of the masses of the
second element that combine with 1 gram of
the first element can always be reduced to
small whole numbers.
 The ratio of the masses of oxygen in H2O
and H2O2 will be a small whole number
(“2”).
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
5
Law of Multiple Proportions
Mass of oxygen that
combines with 1 g of
Carbon
Compound 1 (CO)
1.33 g
Compound 2 (CO2)
2.66 g
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
6
Dalton’s Atomic Theory (1808)
 Each element is made up of tiny particles
called atoms.
 The atoms of a given element are
identical; the atoms of different elements
are different in some fundamental way or
ways.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
7
Dalton’s Atomic Theory
(continued)
 Chemical compounds are formed when
atoms combine with each other. A given
compound always has the same relative
numbers and types of atoms.
 Chemical reactions involve reorganization
of the atoms - changes in the way they are
bound together. The atoms themselves are
not changed in a chemical reaction.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
8
Avogadro’s Hypothesis (1811)
At the same temperature and pressure, equal
volumes of different gases contain the same
number of particles.
• 5 liters of oxygen
• 5 liters of nitrogen
• Same number of particles!
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
9
Figure 2.5: A representation of combining
gases at the molecular level. The spheres
represent atoms in the molecules.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
10
Early Experiments to
Characterize the Atom

J. J. Thomson - postulated the existence of
electrons using cathode ray tubes.

Ernest Rutherford - explained the nuclear
atom, containing a dense nucleus with electrons
traveling around the nucleus at a large distance.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
11
Figure 2.7: A cathode-ray tube. The fastmoving electrons excite the gas in the tube,
causing a glow between the electrodes.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
12
Figure 2.8: Deflection of cathode
rays by an applied electric field.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
13
Figure 2.10: A schematic representation of
the apparatus Millikan used to determine
the charge on the electron.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
14
The Electron
Tiny, negatively charged particle
 Very light compared to the mass of
an atom – 1/1837th the mass of a H
atom
 Move extremely rapidly within the
atom.

Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
15
The Modern View of Atomic
Structure

:
Electrons

Protons: found in the nucleus, they have a
positive charge equal in magnitude to the
electron’s negative charge.

Neutrons: found in the nucleus, virtually same
mass as a proton but no charge.

The nucleus contains: protons and neutrons.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
16
Figure 2.9: The plum pudding
model of the atom.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
17
Figure 2.12: Rutherford's experiment on
-particle bombardment of metal foil.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
18
Figure 2.13: (a) The expected results of
the metal foil experiment if Thomson's
model were correct. (b)Actual results.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
19
Figure
2.14: A
nuclear
atom
viewed in
cross
section.
Note that
this drawing
is
not to scale.
The Mass and Change of the
Electron, Proton, and Neutron
Particle
Mass (kg)
31
Electron
9.11  10
Proton
1.67  1027
Neutron
1.67  10
27
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
Charge
1
1+
0
21
Atomic Number / Atomic Mass

The atomic number is equal to the
number of protons.

The atomic mass is equal to the
number of protons + neutrons
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
22
Atomic Number / Atomic Mass

Essentially, all of the mass of the
atom is considered to reside in the
nucleus.

In a neutral atom, the number of
protons equals the number of
electrons.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
23
Isotopes
All atoms of an element have the
same number of protons.
 The number of protons = the atomic
number
 Atoms of an element with different
numbers of neutrons are called
isotopes.

Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
24
Isotopes continued…
All isotopes of an element are
chemically identical
 Isotopes have different masses
 Isotopes are identified by their mass
number.

Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
25
The Chemists’ Shorthand: Atomic
Symbols
Mass number 
Atomic number 
39
K
19
 Element Symbol
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
26
Figure 2.15: Two isotopes of sodium. Both have eleven
protons and eleven electrons, but they differ in the
number of neutrons in their nuclei.
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
27
Isotopes
Isotope
Atomic
Number
Mass
Number
# of Protons
# of
Electrons
# of
Neutrons
20
18
22
U-234
Na-
24
S2-
32
Copyright©2000 by Houghton
Mifflin Company. All rights reserved.
28