Transcript Early Atomic History

The Structure &
Stability of Atoms
Early Atomic History
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There have been many different theories,
reflecting different times and cultures, to explain
the composition of matter.
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In addition, chemical reactions, refinements of
ores, purification of salt, etc. have been carried
out for thousands of years.
Early Atomic History
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The ancient Greek philosophers theorized that
matter is discrete, rather than continuous.
Some, notably Demokritos, suggested that there
is some small unit of matter that still retains the
properties of the larger sample. It was thought
that these smaller pieces of matter were
indivisible, and were given the name atomos from
which we get our modern word atoms.
Early Atomic Theory
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During the next 2000 years, a lot was learned
about matter. Several elements were discovered,
metals were refined, acids prepared, etc.
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In the mid-1600s, the scientific (rather than the
philosophical or applied) study of the nature
matter began to take shape.
Early Atomic Theory
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Since most laboratories contained rudimentary
equipment- burners and scales, many
experiments involved the measurement of
changes in volumes (for gases) and masses
during chemical reactions.
Based on measurements and observations,
several scientific laws were developed. These
laws form the basis for our understanding of the
composition of matter.
The Law of Conservation of Mass
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Antoine Lavoisier (1743-1794) measured the
masses of reactants and products for a variety of
chemical reactions. He determined that matter
is neither created nor destroyed during a
chemical reaction. This is known as the law of
conservation of mass.
The Law of Definite Proportion
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Joseph Proust (1754-1826) determined the
chemical composition of many compounds. He
found that a given compound always contains
the exact same proportion of elements by mass.
This is known as the law of definite proportion.
For example, all samples of water contain 88.8%
oxygen by mass, and 11.2% hydrogen by mass.
The Law of Multiple Proportions
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This chemical law applies when two (or more)
elements can combine to form different
compounds.
Common examples are carbon monoxide and
carbon dioxide, or water and hydrogen peroxide.
John Dalton (1766-1844) conducted
experiments on these types of compounds, and
determined that there is a simple relationship
between the masses of one element relative to
the others.
The Law of Multiple Proportions
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When two elements form a series of
compounds, the ratios of the masses of one
element that combine with a fixed mass of the
other element are always in a ratio of small
whole numbers.
The meaning of this law is difficult to
understand unless it is illustrated using a specific
series of compounds.
The Law of Multiple Proportions
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Consider the compounds of water and hydrogen
peroxide. At this point in history, chemists
knew the compounds were different, and that
they both contain (or can be broken down into)
the elements hydrogen and oxygen. They did
not yet know the formulas for either compound,
nor was the concept of atoms fully developed.
The Law of Multiple Proportions
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Analysis of 100 grams of the compounds produced the
following data:
Compound
Mass of
Mass of
oxygen/100g hydrogen/100g
of compound of compound
Grams of
oxygen/gram
of hydrogen
water
88.8 grams O
11.2 grams H
7.93 gO/gH
Hydrogen
peroxide
94.06 grams O
5.94 grams H
15.8 gO/gH
The Law of Multiple Proportions
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The Law of Multiple Proportions is illustrated when the
numbers in the last column are compared.
Compound
Grams of
oxygen/gram
of hydrogen
water
7.93 gO/gH
Hydrogen
peroxide
15.8 gO/gH
15.8/7.93 = 2/1
The small whole
number ratio suggests
that there is twice as
much oxygen in
hydrogen peroxide as
there is in water.
The Law of Multiple Proportions
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The key feature is that small whole numbers are
generated. The results support the hypothesis that
molecules consist of various combinations of atoms,
and that atoms are the smallest unit of matter. The
ratio doesn’t produce fractions, since there is no such
thing as a fraction of an atom.
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For the example cited, we would propose that hydrogen
peroxide contains twice as many oxygen
atoms/hydrogen atoms than does water. We cannot,
however, determine the actual formula of either
compound.
The Law of Multiple Proportions
Dalton’s Atomic Theory (1808)
1. Each element consists of tiny particles called atoms.
2. The atoms of a given element are identical, and differ
from the atoms of other elements.
3. Compounds are formed when atoms of different
elements combine chemically. A specific compound
always has the same relative number and types of
atoms.
4. Chemical reactions involve the reorganization of
atoms, or changes in the way they are bound together.
Sub-Atomic Particles
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The period from approximately 1900-1915
involved the study of the nature of the atom,
using two relatively new tools: electricity and
radioactivity.
Scientists knew that atoms of different elements
had different relative atomic masses and
different properties, and they wanted to find out
the reasons for the differences.
Sub-Atomic Particles
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J.J. Thomson (1856-1940) studied the properties of
cathode rays. The rays are produced in partially evacuated
tubes containing electrodes at either end.
The rays are invisible, unless a phosphorescent screen is
used.
Sub-Atomic Particles
Sub-Atomic Particles
Cathode Rays
(Cathode)
(Anode)
Sub-Atomic Particles
Thomson made the following observations:
1. The cathode rays had the same properties
regardless of the metal used for the cathode.
2. The rays traveled from the cathode (- charged) to
the anode (+ charged).
3. The rays were attracted to the positive plate of an
external electrical field, and repelled by the
negative plate.
Sub-Atomic Particles
Thomson concluded:
1. The cathode rays are a stream of negatively
charged particles called electrons.
2. All atoms contain electrons, and the
electrons from all elements are identical.
3. The atom must also contain matter with a
positive charge, as atoms are neutral in
charge.
Sub-Atomic Particles
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Thomson also carried out deflection
measurements, in which he applied a magnetic
field to deflect the beam along with an external
electrical field to straighten out the bent beam.
From his measurements, he was able to
calculate the charge/mass ratio of the electron:
e/m = -1.76x108 coulombs/gram
Sub-Atomic Particles
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Robert Millikan (1868-1963) published the
results of his Oil Drop Experiment in 1909. He
designed an apparatus that could be used to
determine the charge on an electron.
The device used a fine mist of oil drops that
had been exposed to ionizing radiation. The
radiation caused some of the oil drops to take
on one or more electrons.
Sub-Atomic Particles
The Charge of the Electron
Sub-Atomic Particles
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Millikin determined that the charge on the
electron is -1.60 x 10-19coulombs.
Using Thomson’s value for the charge to mass
ratio of the electron, the mass of the electron
could be calculated.
mass of e- = (-1.60 x 10-19 coulombs)
(-1.76 x 108 coulombs/gram)
= 9.11 x 10-28 grams
= 9.11 x 10-31 kilograms
Early Atomic Models
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J. J. Thomson had shown that all atoms contain
negatively charged particles called electrons.
Combined with the work of Millikan, they
discovered that the electron has very little mass.
Thomson proposed that the bulk of the atom is
a positively charged gel or cloud, with most of
the atomic mass and all of the positive charge
uniformly distributed throughout the gel.
Early Atomic Models
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The electrons were viewed as discrete, very small
particles that were stuck into the positively
charged gel or cloud “like raisins in a pudding.”
This model is often called the plum or raisin
pudding model of the atom.
The electrons could be knocked out of the gel if
enough energy is applied, and this is the source
of the cathode rays.
Early Atomic Models
One of the key
features of Thomson’s
atomic model is that
most of the atomic
mass and all of the
positive charge is
uniformly distributed
throughout the atom.
Early Atomic Models
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Thomson had a graduate student, Ernest
Rutherford, working for him. In 1911,
Rutherford, Geiger and Marsden performed an
experiment to confirm Thomson’s atomic
model.
They bombarded a thin gold foil with alpha (α)
particles. The α particles have twice the charge
of an electron and are positive in charge, with a
mass that is 7300 times greater than the mass of
an electron.
Early Atomic Models
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The α particles can best be thought of as a positively
charged, fast traveling atomic sized bullet. They created
a thin beam of α particles and directed the beam at a
very thin gold foil.
Early Atomic Models
If Thomson’s model is
correct, most of the α
particles should pass
right through the gold
atoms. Some slight
deflection might occur
if the positively charged
α particle travels near an
electron.
Gold Foil Experiment
Early Atomic Models
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The film that lined the apparatus showed that most α
particles went through the foil with little or no
deflection. However, some of the particles were
deflected at great angles.
Early Atomic Models
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The deflection of the α
particles was consistent
with a large
concentration of positive
charge and atomic mass.
This very small extremely
dense positively charged
area is called the nucleus.
Early Atomic Models
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The atom is mostly
empty space, with the
electrons found outside
of the nucleus. If the
nucleus was the size of a
pea, it would have a mass
of 250 million tons, and
the electrons would
occupy a volume
approximately the size of
a stadium.
Atomic Nucleus
Sub-Atomic Particles
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We now know that the positive charge of an
atom, contained in the nucleus, is due to
particles called protons.
Protons have a charge equal in magnitude to
that of an electron, but positive in charge.
The mass of a proton is roughly 1800 times
greater than the mass of an electron.
Sub-Atomic Particles
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The nuclei of atoms also can contain neutrons.
Neutrons are neutral in charge, with a mass similar
to that of a proton.
Neutrons are found in the nucleus of atoms, along
with protons.
Sub-Atomic Particles
Sub-Atomic Particles
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During chemical reactions, atoms may lose or
gain electrons to form charged particles called
ions.
Atoms of a given element may have differing
numbers of neutrons. These forms of the same
element are called isotopes.
It is the number of protons or the atomic number
that defines the identity of the atom.
Atomic Symbols
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The periodic table lists the elements in order of
increasing atomic number (the number of
protons).
The atomic number, represented by the letter Z,
is linked with the atomic symbol. For example,
oxygen is atomic number 8, and any atom
containing 8 protons, regardless of the number
of neutrons or electrons, is represented by the
symbol O.
Atomic Symbols
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To indicate a specific isotope, the atomic symbol
must also contain the mass number.
The mass number is the number of neutrons plus
protons for a particular isotope. The mass
number is never found on the periodic table.
Since the mass number is the number of
particles (neutrons + protons) in the nucleus, it
is always an integer.
Isotopes of Sodium
Mass number
Atomic number
Atomic Symbols
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For example, there are three isotopes of carbon:
12C, 13C and 14C
The mass number, if specified, appears in the
upper left corner of an atomic symbol. Since all
carbon atoms have 6 protons (carbon is atomic
number 6 on the periodic table), atoms of
carbon may have 6, 7 or 8 neutrons in the
nucleus.
The isotopes are called carbon-12, carbon-13 and
carbon-14.
Atomic Symbols
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If the atom has lost or gained electrons, the
charge is written in the upper right corner of the
atomic symbol.
The atomic number, though optional, may be
written in the lower left corner of the symbol.
37Cl1This ion of chlorine contains 17 protons, 20
neutrons, and 18 electrons.
Relative Atomic Masses
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Once the formulas of simple gases and
compounds could be determined, scientists
could also determine the relative masses of the
elements.
For example, since equal volumes of gases
contain equal numbers of particles (at the same
T and P), the masses of gases could be
compared to hydrogen, the lightest gas.
Stoichiometry
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Stoichiometry is a Greek word that means using
chemical reactions to calculate the amount of
reactants needed and the amount of products
formed.
Amounts are typically calculated in grams (or
kg), but there are other ways to specify the
quantities of matter involved in a reaction.
Relative Atomic Masses
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As the early chemists explored the nature of
matter, they discovered that atoms of the
elements had different masses.
Avogadro’s Hypothesis which states that under
the constant temperature and pressure equal
volumes of gases contain an equal number of particles
could be used to determine relative atomic
masses for gaseous elements.
Relative Atomic Masses
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Equal volumes of gases contain an equal number of
particles.
Although the number of particles (atoms or
molecules) in a liter of gas (at a specific T and P)
wasn’t known, Avogadro’s Hypothesis said that
a liter of any other gas under the same
conditions would contain the same number of
particles.
Relative Atomic Masses
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Equal volumes of gases contain an equal number of
particles.
Since the masses of the gaseous samples could
be determined, a comparative or relative scale of
atomic and molecular masses could be derived.
Relative Atomic Masses
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Equal volumes of gases contain an equal number of particles.
For example, if the masses of a liter of oxygen (O2),
chlorine (Cl2) and hydrogen (H2) were compared under
identical conditions, the hydrogen sample has the
smallest mass, and the chlorine sample has the largest
mass.
The ratio of the masses of the 1 liter samples is:
35.5/16.0/1.00
Cl2/ O2/ H2
Relative Atomic Masses
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Equal volumes of gases contain an equal number of particles.
The ratio of the masses of the 1 liter samples is:
35.5/16.0/1.00
Cl2/ O2/ H2
Since all three gases are diatomic, we can say that an
oxygen atom is 16.0 times heavier than a hydrogen
atom, and that a chorine atom is 35.5 times heavier
than a hydrogen atom.
Relative Atomic Masses
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A scale of relative atomic mass was devised.
Individual atoms are much too small to weigh,
but the masses of large collections of atoms
could easily be compared.
The relative masses of the atoms are listed on
the periodic table. An arbitrary unit, the atomic
mass unit (amu) is used for relative masses.
Relative Atomic Masses
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Eventually, the carbon-12 isotope (12C) was
assigned an atomic mass of exactly 12 atomic
mass units, and all other atomic masses are
expressed relative to this assignment.
The atomic mass for carbon, found on the
periodic table, is 12.01 amu, and not 12.000
amu. This is because the periodic table lists
the average relative atomic mass for all
isotopes of the element.
Relative Atomic Masses
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Carbon exists as three isotopes:
12C has a relative mass of exactly 12 amu
13C has a relative atomic mass of 13.003 amu
14C has a relative atomic mass of 14.0 amu
The value found on the periodic table, 12.01
amu, reflects the relative abundance of the
isotopes. The majority of carbon (98.89%) is
12C,
with 1.11% 13C, and a trace of 14C.
Relative Atomic Masses
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Chorine exists as two isotopes:
35Cl, with a relative atomic mass of 35.0 amu
and 37Cl, with a relative atomic mass of 37.0
amu.
What does the atomic mass of chlorine on
the periodic table tell you about the relative
abundance of the two isotopes?
Moles
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Many chemical reactions are carried out using a
few grams of each reactant. Such quantities
contain huge numbers (on the order of 1023) of
atoms or molecules.
A unit of quantity of matter, the mole, was
established. A mole is defined as the number of
carbon atoms in exactly 12 grams of 12C.
Avogadro determined the number of particles
(atoms or molecules) in a mole.
Moles
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Avogadro’s number = 6.022 x 1023 particles/mole
Atoms are so small, that a mole of most substances can
be easily held in ones hand.
Cu
I2
Al
Hg
S
Fe
Moles
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If we consider objects we can see, a mole of
sand would cover the entire planet and be
several miles deep! However, the collection of
atoms, called a mole, is very convenient in the
laboratory (just like dozens are useful in buying
eggs or pencils).
Moles
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A mole of any atom has a mass equal to the
element’s atomic mass expressed in grams.
A mole of iron atoms has a mass of 55.85 grams;
a mole of iodine molecules (I2) has a mass of
(126.9) (2) = 253.8 grams.
Moles
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The masses of each molar sample are provided
below.
I2=253.8 g
S=32.07g
Cu = 63.55g
Al=26.98g
Hg = 200.6 g
Fe=55.85 g
Molar Mass
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For compounds, once the formula is known, the
mass of a mole of the substance can be
calculated by summing up the masses of all the
atoms in the compound.
For example, hydrogen peroxide has the
formula H2O2:
2H +2O = 2(1.008g) + 2(16.00g) = 34.02 g/mol
A mole of hydrogen peroxide has a mass of
34.02 grams.
Nuclear Chemistry
Many nuclei are unstable, and may emit
particles and/or energy until they reach a more
stable combination of protons and neutrons.
A nuclear reaction differs from a chemical
reaction, and results in the formation of new
elements as an unstable nucleus decays.
Nuclear Stability
For the lighter elements in the periodic table,
most of the stable nuclei contain approximately
an equal number of protons and neutrons. For
heavier elements, the neutron to proton ratio
generally increases in the stable isotopes.
Nuclear Stability
Nuclear Stability
The mass of a nucleus is less than the sum of
the masses of its neutrons and protons. A very
small amount of mass is converted into the
nuclear binding energy.
During nuclear reactions, some of this
energy is released. A very small loss of mass
results in a huge release of energy.
Nuclear Binding Energy
Nuclear Stability
Smaller atoms will undergo fusion, and
combine to form heavier atoms. Larger atoms
will undergo fission, and break into smaller
nuclei and particles.
Nuclear Equations
Nuclear equations represent only the nuclei
of the particles, rather than the atom complete
with electrons.
Nuclear equations must be balanced. This
means that the sum of the mass numbers on the
left side equals the sum of the mass numbers on
the right side.
Nuclear Equations
Likewise, the sum of the atomic numbers on
the left must equal the sum of the atomic
numbers on the right.
Note that the identity of the elements will
change during a nuclear reaction.
Nuclear Reactions
Unstable nuclei can decay in a variety of
ways. If the nucleus contains too many
protons/neutron, βemission may occur. A β
particle is an electron, but it results from a
neutron disintegrating into a proton and an
electron. The proton remains in the nucleus of
the newly formed atom, and the electron is
ejected.
βEmission
The net result of β emission is an increase in the
number of protons in the nucleus.
βEmission
Nuclear reactions must be balanced in terms of the total
of the atomic numbers and the total of the mass
numbers.
α Emission
Some unstable nuclei emit an α particle. α
particles are the same as a helium nucleus. The
alpha particle contain two protons and two
neutrons. The net result is a decrease in both
atomic number (by 2) and mass number (by 4).
α Emission
Alpha emission is common in heavier
elements. All elements above 209Bi are
radioactive. They often emit alpha particles
(along with other forms of radiation) until they
attain a stable nucleus.
α Emission
γ Emission
Gamma (γ) emission is emission of a very
high energy form of radiation. It accompanies
both α and β emission, and represents a release
of energy. Since it is not a particle, it is often
omitted from nuclear equations.
Positron Emission
Nuclei with a neutron to proton ratio which
is too low, may undergo positron emission. A
positron, β+, has the same mass as an electron,
but it is positive in charge.
The positron results from a proton decaying
into a neutron and ejecting the positron. As a
result, the number of neutrons increases, and the
number of protons decreases.
Positron Emission
A positron has the same mass as an electron but an
opposite charge. It can be thought of as a “positive
electron.”
Positron,  +
Electron Capture
Nuclei with a neutron to proton ratio which
is too low, may undergo electron capture. In
this process, a proton in the nucleus combines
with an electron outside the nucleus to form a
neutron.
The net result is an decrease in atomic
number and an increase in the mass number.
Electron Capture