Download Chapter 2 Atoms, Ions, and Molecules

Chapter 2
Atoms, Ions, and Molecules
What is matter made of? Continuous or particulate?
Ancient view: Four elements (fire, air, earth, water)
Democritus (460–370 BC): “father of atomism”.
Aristotle (384–322 BC) held that Democritus’ views were impossible. Atomic concept suppressed for
~2000 years. But by the 1700s, things had started to change.
Observations Leading to the Atomic View of Matter
Law of Conservation of Mass: total mass of substances does not change during a chemical reaction.
[Matter can not be created or destroyed.]
–Lavoisier, 1789
Law of Definite Composition: no matter its source, a substance is composed of the same elements in
the same fractions (or ratio) by mass.
–Proust, 1799
mass of element
 fraction by mass =
total mass of substance
 percent by mass =
mass of element
total mass of substance
100%
Example: A 15.00 gram sample of water contains 13.32 grams of oxygen. Calculate the % by
mass of oxygen and hydrogen in water. (Water contains only hydrogen and oxygen.)
Law of Multiple Proportions: If elements A, B form a series of substances, the different masses of B
that combine with a fixed mass of A will be in a ratio of small whole numbers.
–Dalton, 1803
Example: Two compounds of copper and bromine with different properties
Substance
mass of bromine that
combines with 1g copper
I
1.26 g
II
2.52 g
properties
yellowish green solid
MP = 504 °C, BP = 1345 °C
light sensitive, water insoluble
black solid
MP = 498 °C, BP = 900 °C
not light sensitive, is water soluble
bromine = element B, copper = element A, ratio of B w/fixed mass of A = 2.52:1.26 or 2:1
2-1
Dalton’s Atomic Theory (1808)
1. matter is made up of microscopic, indivisible
particles (atoms) that cannot be created or destroyed
Our Understanding Today
2. all atoms of an element are identical in mass, and
have identical physical and chemical properties
3. atoms of different elements have different masses,
physical properties, and chemical properties
4. atoms of different elements combine in simple
whole numbers to form compounds
5. atoms of an element cannot be converted into atoms of other
elements; chemical reactions involve reorganization of the atoms
But Dalton’s Atomic Theory couldn’t explain it all—
 Why do atoms combine the ways they do?
 What about the electrically charged particles that were being observed?
Observations Leading to the Nuclear Model of the Atom
 Cathode Ray Tube Experiments (Figure 2.2): J.J. Thomson, 1897
The particles discovered were called electrons. Thomson determined the mass/charge ratio of the
electron in a related experiment. Another experiment was needed to determine the mass or charge.
 Millikan’s Oil Drop Experiment ( Figure 2.3), 1909
Control of oil drop by electric field allowed for calculation of charge of 1 electron:
e– charge = –1.602 10–19 C
Since the mass/charge ratio was already known, the mass could also be calculated:
e– mass = 9.109 10–28 g
So, where’s the positive charge? Speculation was that it was spread throughout the atom.
 Rutherford’s Alpha Particle Scattering Experiment ( Figure 2.6), 1910
To test this speculation, Rutherford shot α particles (positively charged and relatively massive) at
gold foil. Most of the α particles passed through the foil relatively unaffected, but a few bounced
back nearly right back at the source. Conclusion: Atoms have a dense, positively charged core
called the nucleus.
By 1932, Chadwick discovered the neutron (neutral, not electrically charged).
Figure 2.6
2-2