Early Atomic History

download report

Transcript Early Atomic History

Atoms & Elements
The building blocks of matter
Early Atomic History
There have been many different theories,
reflecting different times and cultures, to explain
the composition of matter.
In addition, chemical reactions, refinements of
ores, purification of salt, etc. have been carried
out for thousands of years.
Early Atomic History
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
During the next 2000 years, a lot was learned
about matter. Several elements were discovered,
metals were refined, acids prepared, etc.
In the mid-1600s, the scientific (rather than the
philosophical or applied) study of the nature
matter began to take shape.
Early Atomic Theory
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
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
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
This chemical law applies when two (or more)
elements can combine to form different
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
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
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
Analysis of 100 grams of the compounds produced the
following data:
Mass of
Mass of
oxygen/100g hydrogen/100g
of compound of compound
Grams of
of hydrogen
88.8 grams O
11.2 grams H
7.93 gO/gH
94.06 grams O
5.94 grams H
15.8 gO/gH
The Law of Multiple Proportions
The Law of Multiple Proportions is illustrated when the
numbers in the last column are compared.
Grams of
of hydrogen
7.93 gO/gH
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
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.
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
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
4. Chemical reactions involve the reorganization of
atoms, or changes in the way they are bound together.
Sub-Atomic Particles
The period from approximately 1880-1915
involved the study of the nature of the atom,
using two relatively new tools: electricity and
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
In the late 1880s, J.J. Thomson (1856-1940) studied the
properties of cathode rays. The rays are produced in
partially evacuated tubes containing electrodes at either
The rays are invisible, unless a phosphorescent screen is
Sub-Atomic Particles
Cathode Rays
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
Sub-Atomic Particles
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.
Sub-Atomic Particles
From his measurements, he was able to
calculate the charge/mass ratio of the electron:
e/m = -1.76x108 coulombs/gram
Sub-Atomic Particles
Around the same time as Thomson (1886),
Eugen Goldstein observed that if a cathode
ray tube contained very small amounts of gas,
a glowing substance travelled toward the
Goldstein called the glowing substance “canal
rays”, and observed that they were positive in
charge, as they are attracted toward the
negative cathode.
Sub-Atomic Particles
The apparatus contained a perforated cathode
that contained many small holes. When an
electric current is applied, the reddish glow
forms in the stream of electrons, and travels
toward the negative cathode and through the
small holes.
Sub-Atomic Particles
It was several years before Goldstein could
explain his observations. The rays were quite
different from cathode rays:
Unlike cathode rays, canal rays were barely
deflected by a magnetic field or external
electric field.
The properties of the rays varied with the
gas contained.
Sub-Atomic Particles
Later studies determined that hydrogen gas
produced the “ray” with the largest charge to
mass ratio (ie., the smallest mass). This particle,
produced when a hydrogen atom loses its
electron, was identified as the proton.
Sub-Atomic Particles
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
The Charge of the Electron
Sub-Atomic Particles
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
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
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
Thomson had a graduate student, Ernest
Rutherford. 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
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
Gold Foil Experiment
Gold Foil Experiment
Early Atomic Models
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
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
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
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
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
During chemical reactions, atoms may lose or
gain electrons to form charged particles called
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
The periodic table lists the elements in order of
increasing atomic number (the number of
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
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
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
The isotopes are called carbon-12, carbon-13 and
Atomic Symbols
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.
The Periodic Table
metal/non-metal line
The Periodic Table
The modern periodic table was developed in
1872 by Dmitri Mendeleev (1834-1907). A
similar table was also developed independently
by Julius Meyer (1830-1895).
The table groups elements with similar
properties (both physical and chemical) in
vertical columns. As a result, certain properties
recur periodically.
The Periodic Table
It is important to note that these periodic
tables were developed prior to any knowledge of
atomic theory. The proton, electron and
neutron had not yet been detected.
The tables were based solely on observed
properties and trends in chemical and physical
behavior and properties.
The Periodic Table
Mendeleev left empty spaces in his table for
elements that hadn’t yet been discovered. Based
on the principle of recurring properties, he was
able to predict the density, atomic mass, melting
or boiling points and formulas of compounds
for several “missing” elements.
The Periodic Table
The Periodic Table
The Periodic Table
Keep in mind:
 Elements along the metal/non-metal dividing
line are called semi-metals or metalloids. These
elements sometimes behave like metals, and
sometimes exhibit non-metallic properties and
 Hydrogen, though in group IA, is not a metal.
It is sometimes also placed in group 7A.
Ionic Charges
The position of an element on the periodic
table can sometimes be used to predict the
charge of its most stable ion. Metals lose
electrons to form cations. Non-metals gain
electrons to form anions.
When a metal and a non-metal combine,
electrons are transferred to form an ionic
Typical Ionic Charges
Relative Atomic Masses
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
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
Equal volumes of gases contain an equal number of
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
Relative Atomic Masses
Equal volumes of gases contain an equal number of
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
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
The ratio of the masses of the 1 liter samples is:
Cl2/ O2/ H2
Relative Atomic Masses
Equal volumes of gases contain an equal number of particles.
The ratio of the masses of the 1 liter samples is:
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
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
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
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
with 1.11% 13C, and a trace of 14C.
Relative Atomic Masses
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
What does the atomic mass of chlorine on
the periodic table tell you about the relative
abundance of the two isotopes?
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.
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.
If we consider objects we can see, a mole of
pennies would cover the entire planet and be
300 meters 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).
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.
The masses of each molar sample are provided
I2=253.8 g
Cu = 63.55g
Hg = 200.6 g
Fe=55.85 g
Molar Mass
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.
Determine the number of silver atoms in a 10.0
gram sample of silver.