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Atoms, Molecules and Ions
Chapter 2
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Democritus was a Greek philosopher who
lived between 470-380 B.C. He developed
the concept of the 'atom', Greek for
'indivisible'. Democritus believed that
everything in the universe was made up of
atoms, which were microscopic and
indestructible.
Democritus had many remarkable insights for his time.
He understood that the Milky Way was a large collection
of stars and also thought that space was limitless.
His insights were not supported by scientific evidence until the
18th century.
Aristotle rejected the idea of atoms.
He did not believe that empty
space could exist.
Because of his views, Democritus’
ideas were not accepted until many
centuries later.
Aristotle believed that empty space did not exist
and that matter is made of earth, fire, air and
water, not atoms.
John Dalton (1766-1844) formulated his
atomic theory to explain chemical reactions,
based on the concept that the atoms of
different elements are distinguished by
differences in their weights.
Dalton’s Atomic Theory (1808): The basic components of
matter are atoms, and only whole numbers of atoms can
combine to form compounds.
He first described color blindness,
and he kept a meteorological
journal with over 200,000
observations during his lifetime.
Dalton’s Atomic Theory (1808)
1.
Elements are composed of extremely small particles
called atoms. Atoms are indivisible.
2. All atoms of a given element are identical, having the
same size, mass and chemical properties. The atoms of
one element are different from the atoms of all other
elements.
3. Compounds are composed of atoms of more than one
element. In any compound, the ratio of the numbers of
atoms of any two of the elements present is either an
integer or a simple fraction. (Laws of D.P. and M.P.)
4. A chemical reaction involves only the separation,
combination, or rearrangement of atoms; it does not
result in their creation or destruction. (Law of C. of M.)
2.1
Theory:
An explanation supported by many
experiments; is still subject to new
experimental data, can be modified,
and is considered successful if it can
be used to make predictions that are
true.
The Law of Multiple Proportions
The atoms of different elements combine
in simple whole number ratios to form
compounds
Dalton’s Atomic Theory
2
2.1
Antoine Lavoiser
• Frequently referred to as the “Father
of Modern Chemistry”
• Stated the first version of the Law of
Conservation of Mass
1743-1794
• Credited with the discovery of the elements oxygen
and hydrogen
• Reformed chemical nomenclature
• Executed during the French Revolution
The Law of Conservation of Mass
Mass is neither created nor destroyed
during ordinary chemical reactions or
physical changes
16 X
+
8Y
8 X2Y
2.1
The Law of Definite Proportions
The size or source of a sample of a
compound has no effect on the proportion
of elements in that compound
e.g. H2O is always H2O
ATOM
The smallest particle of an
element that retains the chemical
properties of that element
What is the atom composed of?
Discovering the Electron
• Because of Dalton’s atomic theory, most
scientists in the 1800s believed that the atom
was like a tiny solid ball that could not be
broken up into parts.
• In 1897, a British physicist, J.J. Thomson,
discovered that this solid-ball model was not
accurate.
• Thomson’s experiments used a vacuum tube.
Discovering the Electron
• A vacuum tube has
had all gases pumped
out of it.
• At each end of the tube is a metal piece called
an electrode, which is connected through the
glass to a metal terminal outside the tube.
• These electrodes become electrically charged
when they are connected to a high-voltage
electrical source.
Cathode-Ray Tube
• When the
electrodes are
charged, rays
travel in the tube
from the negative
electrode, which
is the cathode, to
the positive
electrode, the
anode.
• Because these
rays originate at
the cathode, they
are called cathode
rays.
Cathode-Ray Tube
• Thomson found that
the rays bent toward
a positively charged
and away from a
negatively charged
plate.
• He knew that
objects with like
charges repel each
other, and objects
with unlike charges
attract each other.
Cathode-Ray Tube
• Thomson concluded that cathode rays are
made up of invisible, negatively charged
particles referred to as electrons.
• These electrons had to come from the matter
(atoms) of the negative electrode.
• For measuring the mass/charge of an electron,
J.J. Thomson won the 1906 Nobel Prize in
Physics.
Robert Millikan (1868-1953) won the
1923 Nobel Prize in physics for his
work on the elementary electric
charge (of an electron) and on the
photoelectric effect.
The photoelectric effect is the
emission of electrons from a metal
when light shines on the metal,
creating an electric current.
The Electron (cont.)
• In the early 1910s, Robert Millikan used the
oil-drop apparatus shown below to
determine the charge of an electron.
Measured mass of e(1923 Nobel Prize in Physics)
e- charge = -1.60 x 10-19 C
Thomson’s charge/mass of e- = -1.76 x 108 C/g
e- mass = 9.10 x 10-28 g
2.2
J.J. Thomson (1856-1940) demonstrated
that electric currents are composed of
negatively charged particles, which were
later named electrons.
He won the 1906 Nobel prize for this
discovery, and subsequently discovered isotopes.
Atoms were known to be neutral (no charge), but
contained positive charges to counterbalance the
negative charges.
The “Plum Pudding” model of the atom allowed
for the charges to spread out as far as possible.
2.2
Ernest Rutherford (1871-1937), won the
1908 Nobel prize for his for his
investigations into the disintegration of the
elements, and the chemistry of radioactive
substances .
• Aim was to support the “Plum
Pudding” model of the atom
• Bombarded a thin gold foil with fast
moving alpha particles (positively
charged Helium ions)
• Discovered the atomic nucleus
(1908 Nobel Prize in Chemistry)
 particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
1. atoms’ positive charge is concentrated in the nucleus
2. proton (p) has opposite (+) charge of electron (-)
3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)
2.2
The atom contains a small dense
“core”
 Most
alpha (+) particles went through
the gold foil
 A small number of alpha particles
were deflected by the gold foil
 Atoms must contain a “nucleus” that
is positively charged
 Additional experiments proved that
the nucleus contains both protons (+)
and neutrons
Rutherford’s Model of
the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
“If the atom is the Houston Astrodome, then
the nucleus is a marble on the 50-yard line.”
2.2
Discovery of the proton
Henry Moseley used cathode rays to knock
electrons out of their orbits. When an
electron jumped back into orbit, it emitted
electromagnetic radiation.
According to atomic theory, an electron jumping to the
innermost orbit of a heavy atom emits X-rays.
From the wavelength of the X-rays, Moseley could
calculate the charge of the nucleus.
He proved that the charge on the nucleus was
equal to its atomic number.
Discovery of the neutron
James Chadwick bombarded a sheet of Be
with α particles, and found that a type of
radiation was emitted that “knocked”
protons from paraffin. These particles
were identified as the elusive neutron.
J. Chadwick
1935 Nobel Prize
in Physics
Chadwick’s Experiment (1932)
H atoms have 1 proton, He atoms have 2 protons
Mass He/mass H should = 2
But, the measured mass He/mass H = 4
 + 9Be
1n
+ 12C + energy
neutron (n) is neutral (charge = 0)
n mass ~ p mass = 1.67 x 10-24 g
2.2
mass p = mass n = 1840 x mass e2.2
Atomic number, Mass number
and Isotopes
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different
numbers of neutrons in their nuclei
Mass Number
Atomic Number
1
1H
235
92
A
ZX
2
1H
U
Element Symbol
(D)
238
92
3
1H
U
(T)
2.3
Isotopes
The Isotopes of an element have
different mass numbers because they
have different numbers of neutrons,
but they all have the same atomic
number.
The Isotopes of Hydrogen
2.3
Calculating Atomic Mass
from Isotope Data
Calculating Atomic Mass
from Isotope Data
•Copper exists as a mixture of 2 isotopes
•The lighter isotope (Cu-63), with 29 protons
and 34 neutrons makes up 69.17% of copper
atoms
•The heavier isotope (Cu-65) with 29 protons
and 36 neutrons, constitutes the remaining
30.83% of copper atoms
Calculating Atomic Mass
from Isotope Data
•The atomic mass of Cu-63 is 62.930 amu
and the atomic mass of Cu-65 is 64.928 amu
•Use the data above to calculate the average
atomic mass of copper
Calculating Atomic Mass
from Isotope Data
First calculate the contribution of each
isotope to the average atomic mass; be sure
to convert each percent to a fractional
abundance.
Mass contribution = mass of isotope X abundance of isotope
For Cu-63:
Mass contribution = 62.930 X 0.6917 = 43.529
For Cu-65:
Mass contribution = 64.928 X 0.3083 = 20.017
Calculating Atomic Mass
from Isotope Data
The average atomic mass of the element is
the sum of the mass contributions of each
isotope.
Atomic mass of Cu = mass contribution Cu-63
+ mass contribution Cu-65
Atomic mass Cu = 43.529 + 20.017 = 63.546
Calculate the average atomic
mass of Germanium.
Radioactivity
W. Roentgen
1901 Nobel
Prize in Physics
In 1895 the German physicist Wilhelm
Roentgen discovered that invisible
rays were emitted when electrons
bombarded the surface of certain
materials. The emitted rays were
discovered because they caused
photographic plates to darken. He
called these invisible high energy
emissions “X”-rays because he didn’t
know what they were.
First medical X-ray;
Mrs. Roentgen’s hand
Not long after Roentgen’s
discovery, French physicist
Antoine Henri Becquerel
was studying minerals that
emit light after being exposed
to sunlight, a phenomenon
called phosphorescence.
Building on Roentgen’s
work, Becquerel wanted to
determine whether
phosphorescent minerals
also emitted X rays.
A.H. Becquerel
1903 Nobel Prize
in Physics
Image of
Becquerel’s
photographic
plate which
has been
fogged by exposure to radiation
from a uranium salt. The shadow
of a metal Maltese Cross placed
between the plate and the
uranium salt is clearly visible.
Becquerel accidentally discovered that
phosphorescent uranium
salts, even when not
exposed to light,
produced spontaneous
emissions that darkened
photographic plates.
Like X rays, the rays from the uranium compound
were highly energetic and could not be deflected by a
magnet, but they differed from X rays because they
arose spontaneously.
As a student of Becquerel, Marie
Curie suggested the name
radioactivity to describe the
spontaneous emission of particles
and/or radiation.
M. Slodowska-Curie
1903 Nobel Prize in
Physics; 1911 Nobel
Prize in Chemistry
The rays and particles emitted by a radioactive
source are called radiation.
(Uranium compound)
2.2
Types of Radiation
The three most common types of radiation are alpha
(α), beta (β), and gamma (γ).
Type of
Radiation
Alpha (α)
Beta (β)
Gamma (γ)
Composition
Alpha Particles
Beta Particles
High energy EM
radiation
Description
Helium nuclei
Electrons
Photons
+2
−1
0
Charge
Symbol
Relative
Penetrating
Power
4
2
He..or..24 
0
1
e..or..10 
Blocked by Paper Blocked by Metal
Foil
γ
Not completely Blocked
by Lead or Concrete
Types of Nuclear Reactions
Alpha emission (alpha decay)
Types of Nuclear Reactions
Beta Emission (beta decay)
Types of Nuclear Reactions
Gamma emission (gamma decay)
Gamma rays do not have mass, so
there is no change in the nuclear
composition of an element undergoing
gamma decay.
238
92
U*
→
238
92
U +
0
0
γ
Do You Understand
Isotopes?
How many protons, neutrons, and electrons are
14
in 6 C ?
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are
11
in 6 C ?
6 protons, 5 (11 - 6) neutrons, 6 electrons
2.3
Do You Understand
Ions?
How many protons and electrons are in
27 3+
13 Al ?
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in
78
2Se
?
34
34 protons, 36 (34 + 2) electrons
2.5
What is an allotrope?
An allotrope is a variant of a substance
consisting of only one type of atom.
Common substances with allotropes are carbon,
phosphorous, sulfur and oxygen.
Example: allotropes of carbon
amorphous (coal, soot)
graphite (pencil lead)
diamond
buckminsterfullerene (Bucky balls)