Download Chapter 2 ATOMS AND ELEMENTS

ATOMS AND ELEMENTS
Democritus
400 B.C.
Believed that matter was
composed of invisible
particles of matter he
called atoms
According to Democritus,
atoms could not be broken
into smaller particles.
1. Law of Conservation of Mass
Antoine Lavoisier (1743-1794)
• The total mass of substances
present at the end of a chemical
process is the same as the mass of
substances present before the
process took place.
2. Law of Constant Composition
Joseph Proust (1754–1826)
• Also known as the law of definite
proportions.
• The elemental composition of a
pure substance never varies.
Atomic Theory of Matter
• The theory that atoms are
the fundamental building
blocks of matter
reemerged in the early
19th century, championed
by John Dalton.
• Using two scientific laws
discovered in the late
1700’s, Dalton built his
atomic theory
Dalton’s Postulates (1803)
Dalton’s Atomic Theory not only explained
the law of conservation of mass and law of
constant composition as they applied to the
atom and their compounds, but also predicted
the law of multiple proportions.
law of multiple proportions:
If two elements A and B
combine to form more than
one compound, the masses
of B that can combine with a
given mass of A are in the
ratio of small whole
numbers.
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J. J. Thomson
(1897)
1. Two different compounds are
formed by the elements carbon and
oxygen. The first compound
contains 42.9% by mass carbon and
57.1% by mass oxygen. The second
compound contains 27.3% by mass
carbon and 72.7% by mass oxygen.
Show that the data are consistent
with the Law of Multiple
Proportions.
Using a cathode ray tube, he
determined the charge-to-mass ratio for
the electron as:
1.76 x 108 C/g
Milikan Oil Drop
Experiment (1909)
The Atom, circa 1900:
• “Plum pudding”
model, put forward by
Thomson.
• Positive sphere of
matter with negative
electrons imbedded in
it.
• Using voltage and change
in the rate of fall of
charged oil drops, he was
able to determine the
charge on each drop.
• From Thompson’s charge
to mass ratio, Milikan
determined the charge and
mass of an electron.
Radioactivity
Corrected to modern
instrumentation:
• One of the pieces of evidence for the fact
that atoms are made of smaller particles
came from the work of Henri Becquerel
and Marie Curie.
• In 1892, Henri Becquerel discovered the
spontaneous loss of nuclear energy from
uranium salts.
• This lead Marie Curie to the discovery of
radioactivity, the spontaneous
disintegration of some elements into
smaller pieces.
Mass of an electron
9.10939 x 10-28 g
Charge of an electron
-1.602 x 10-19 C
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Rutherford’s Gold Foil Experiment
(1910)
Ernest
Rutherford
Results:
1. No holes
2. α- particles deflected at specific angles –
some backwards
• Discovered alpha,
beta and gamma
radiation.
James Chadwick
(1932)
Rutherford (~1911)
Nuclear Model
• Combination of Millikan’s
Findings and the Au Foil
Experiment Lead to
Rutherfords Model
. .
heavy
central
(+) nucleus
. .
. .
.
. .
.
• Further developed the atomic
model by theorizing that alpha
and beta radiation results from the
decomposition of a neutral particle
found in the nucleus, the neutron
e- “about”
nucleus
“sea of e-”
• H atoms - 1 p; He atoms - 2 p
• mass He/mass H should = 2
• measured mass He/mass H = 4
The Bohr Model
(1913)
So, what is a neutron?
Possibly, a proton and electron
held together by smaller particles
called neutrinos (no charge)
central (+)
nucleus
Neutrinos could account for the
mass difference between a proton +
electron and a neutron along with
Electromagnetic Radiation.
. . .. .
. .. . ..
.
Planetary Model
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e- in
allowed
orbits
From Bohr’s model, atoms could
then be described as gaining charges
Therefore, an atom that has
lost an electron would have one
more proton than electron and
would have a net +1.602 x 10-19 C
As we have seen, electrons have a
charge of -1.602 x 10-19 C.
charge, or +1 charge.
Likewise, protons have a charge of
+1.602 x 10-19 C.
Atoms with uneven numbers of
electrons and protons are called ions
for convenience, charges are
reported as whole multiples of their
charges or +1, -1; known as an atoms
electronic charge
anion (-)
cation (+)
Atomic Mass
Modern View of the Atom
(mid 1920s)
•
Because the masses of atoms are so small, the
units of grams is much too large to be used
conveniently. Therefore, the Atomic Mass
Unit (amu) is used.
•
The amu is defined by assigning a mass of
exactly 12 to and atom of the carbon 12
isotope:
Heisenberg, deBroglie, Schroediner
e- in regions
defined by math
functions
. . ..
. ..
.. . .
1 amu = 1.660 538 73 x 10-24 g
Quantum Mechanical Model
We will revisit this shortly
ATOM
COMPOSITION
ATOMIC COMPOSITION
• Protons
 + electrical charge
 mass = 1.67262158 x 10-24 g
 relative mass = 1.0073 (amu)
• Electrons
 negative electrical charge
 mass = 9.10938188 x 10-28 g
 relative mass = 0.0005486 amu
• Neutrons
 no electrical charge
 mass = 1.67492716 x 10-24 g
 mass = 1.0087 amu
The atom is mostly
empty space
• protons and neutrons in the nucleus.
• the number of electrons is equal to the number of
protons.
• electrons in space around the nucleus.
• extremely small. One teaspoon of water has 3 times
as many atoms as the Atlantic Ocean has teaspoons
of water.
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Example:
Atomic symbols
How many protons, neutrons
and electrons are in the
following atom?
• Nuclear symbol - describes the
number of particles in the nucleus
of an atom.
A
Z
X
• Atomic # (Z)
number of protons in the nucleus
• Mass # (A)
total number of protons and
neutrons in the nucleus
23
11
Na
23
Na
Na  23
Hyphen notation
Masses of Isotopes
Isotopes
• determined with a mass spectrometer
• Atoms of the same element
(same Z) but different mass
number (A).
• Boron-10 (10B) has 5 p and 5 n
• Boron-11 (11B) has 5 p and 6 n
11B
10B
Bone scans with
radioactive
technetium-99.
The average atomic mass is found
by calculating the mass in grams of
each isotope in relationship to its
naturally occurring abundance.
Average Mass
• Because in the real world we use large amounts of
atoms and molecules, we use average masses in
calculations.
• Average mass is calculated from the isotopes of an
element weighted by their relative abundances.
• Boron is 20% 10B and 80% 11B. That is, 11B is 80
percent abundant on earth.
• For boron, its atomic weight
= 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu
Example:
A sample of chlorine gas is
75.53% 35Cl
24.47% 37Cl
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Atomic Mass
Remember, the amu is defined by assigning a
mass of exactly 12 to and atom of the carbon 12
isotope. Therefore, 1amu = 1/12 the mass of a
Carbon-12 isotope, or 1/12 the mass of 6
neutrons and 6 protons (electrons are negligible)
Try it and see if you get the same number
as below:
1 amu = 1.66054 x 10-24 g
Average Atomic Mass (35.45)
What is the mass in amu of 1 Gram of matter?
Average Atomic Mass vs.
Average Atomic Weight
Periodic Table
• Dmitri Mendeleev developed
the modern periodic table. He
argued that element properties
are periodic functions of their
atomic masses.
• We now know that element
properties are periodic
functions of their ATOMIC
NUMBERS.
• Groups vs. Periods
•The atomic weight describes the average
mass of the naturally occurring isotopes
multiplied by their relative percentages in
atomic mass units
•The average atomic mass describes the
average mass of the naturally occurring
isotopes multiplied by their relative
percentages in grams.
Regions of the
Periodic Table
and Element
Abundance
Hydrogen
Shuttle main engines
use H2 and O2
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The Hindenburg crash,
May 1939.
Group 2A: Alkaline Earth Metals
Group 1A: Alkali Metals
Magnesium
Reaction of
potassium + H2O
Magnesium
oxide
Cutting sodium metal
Group 3A: B, Al, Ga, In, Tl
Group 4A: C, Si, Ge, Sn, Pb
Quartz, SiO2
Aluminum
Diamond
Boron halides
BF3 & BI3
Group 6A: O, S, Se, Te, Po
Group 5A: N, P, As, Sb, Bi
Ammonia, NH3
Sulfuric acid
dripping from
snot-tite in cave
in Mexico
White and red
phosphorus
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Sulfur from
a volcano
Group 8A:
He, Ne, Ar, Kr, Xe, Rn
Group 7A:
F, Cl, Br, I, At
• Lighter than air balloons
• “Neon” signs
XeOF4
Colors of Transition Metal
Compounds
Transition Elements
Lanthanides and
actinides
Iron in air gives
iron(III) oxide
Iron
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Cobalt
Nickel
Copper
Zinc