Chapter 3: Atoms
I.
Foundations of Atomic Theory
A. Atom - smallest particle of an element that retains the chemical
identity of that element.
1. Originally proposed in 400 B.C. by the Greek philosopher,
Democritus. He called these particles atomos, from the Greek for
“indivisible”.
2. This theory was not believed by Aristotle, a Greek philosopher
whose ideas formed the basis of western scientific thought for
generations.
B. In the late 1700’s scientists began to focus on quantitative analysis.
Improved equipment allowed better analysis and the discovery of
several laws.
1. Law of Conservation of Mass - mass is neither created nor
destroyed during ordinary chemical reactions of physical changes.
2. Law of Definite Proportions - a chemical compound contains
the same elements in exactly the same proportions by mass
regardless of the size of the sample or source of the compound.
3. Law of Multiple Proportions - if two or more different
compounds are composed of the same two elements, then the
ratio of the masses of the second element combined with a certain
mass of the first element is always a ratio of small whole numbers.
C. Dalton’s Atomic Theory
1. Formulated by John Dalton, English schoolteacher, in 1808. He
proposed his theory as an explanation for the above laws. His
theory consists of the following 5 postulates:
a. All matter is composed of extremely small particles called
atoms.
b. Atoms of a given element are identical in size, mass, and other
properties; atoms of different elements differ in size, mass, and
other properties.
c. Atoms cannot be subdivided, created, or destroyed.
d. Atoms of different elements combine in simple whole number
ratios to form chemical compounds.
e. In chemical reactions, atoms are combined, separated, or
rearranged.
2.
Today, we know that not all of these postulates are 100 % true,
but is a good starting point.
II.
Structure of the Atom
A. Atoms consist of two regions: the nucleus, which contains protons and
neutrons; and the region surrounding the nucleus, called the electron
cloud, where the electrons are. Protons, neutrons, and electrons are
also called subatomic particles.
1. Protons - Positively charged particles.
2. Neutrons - neutrally charged particles.
3. Electrons - negatively charged particles.
B. Discovery of the Electron
1. The electron was the first subatomic particle to be discovered.
a. When current is passed through a cathode-ray tube the surface
of the tube opposite the cathode glowed. This glow was
caused by a stream of particles called a cathode ray.
b. Observations of this beam led to the hypothesis that the
particles of the beam are negatively charged.
2. English physicist J.J. Thomson was able to measure the
charge/mass ratio of cathode ray particles.
a. He found that this ratio is always the same, regardless of the
metal used to produce them.
b. He concluded that all atoms contain electrons and all electrons
are the same.
c. He proposed the “plum pudding model” (chocolate chip cookie)
model of the atom. Negative electrons were spread
throughout the positive charge of the rest of the atom, like
chocolate chips in the cookie dough.
3. American physicist Robert A. Millikan measured the negative charge
of the electron.
a. He showed that the mass of the electron is about 1/2000th the
mass of the simplest type of hydrogen atom. (More accurate
measurements show the mass to be 9.109 * 10-31 kg or
1/1837 the mass of the same atom.)
b. Because atoms are electrically neutral, there must be a positive
charge to balance.
c. Because electrons have so much less mass than atoms, atoms
must contain other particles that account for most of the mass.
C. Discovery of the Atomic Nucleus
1. Ernest Rutherford developed the “gold foil experiment”.
2. At the time it was thought that all of the mass and charge was
evenly distributed throughout the atom.
3. He shot high energy, positively charged alpha particles (we will
discuss them more later) at a thin piece of gold foil.
4.
5.
6.
7.
Based on the chocolate chip cookie model, they expected the
particles to pass through with only a slight deflection. For most of
the particles this was the case.
About 1 in 8000 particles were redirected back toward the source.
Rutherford said it was “as if you had fired a 15 inch shell at a piece
of tissue paper and it came back and hit you!”
This proved that atoms are mostly empty space. He suggested
that the electrons surrounded the positively charged nucleus.
If an atom is the size of a large football stadium, the nucleus would
be a little smaller than a dime, and an electron would be a little
smaller than FDR’s eye on that dime.
D. Composition of the Atomic Nucleus
1. All atomic nuclei are made of two kinds of particles, protons and
neutrons.
2. A proton has a mass of 1.673 * 10-27 kg, which is 1836 times
greater than the mass of an electron. A neutron has a mass of
1.675 * 10-27 kg, slightly larger than a proton.
3. Nuclei of atoms of different elements differ in the number of
protons they contain.
4. The nucleus is held together by nuclear forces, which are shortrange proton-neutron, proton-proton, and neutron-neutron forces.
5. Most of the volume of the atom is due to the electron cloud.
III. Counting Atoms
A. Atomic Number
1. Atomic Number (Z) - the number of protons in the nucleus of
each atom of that element.
2. The atomic number identifies the element.
B. Isotopes
1. Isotopes - atoms of the same element that have different masses.
2. They have the same number of protons and a different number of
neutrons.
3. The three isotopes of hydrogen are protium (0 n), deuterium (1 n),
and tritium (2 n).
C. Mass Number
1. Mass Number - the total number of protons and neutrons in the
nucleus of an isotope.
2. Protium has a mass number of 1, deuterium = 2, tritium = 3.
D. Designating Isotopes
1.
2.
Isotopes of elements can be referred to in two ways.
a. Name of element - mass number. For example (uranium-235).
b. Symbol with subscript (at #) and superscript (mass #).
Nuclide - a general term for any isotope of any element.
E.
Relative Atomic Mass
1. Masses of atoms expressed in grams are very small. (For example,
oxygen-16 has a mass of 2.657 * 10-23 g.) To make these
numbers easier to use we use relative atomic mass.
2. Carbon-12 has been assigned to be the arbitrary base number. A
single Carbon-12 is said to weigh 12 atomic mass units (amu).
This is an approximation!!
3. Although isotopes have different masses, they do not differ
significantly in their chemical behavior.
4. The masses of the subatomic particles can be expressed in amu as
well. 1 p = 1.007276 amu, 1 e = 0.0005486 amu, 1 n = 1.008665
amu. Actual mass and relative mass are different because of these
deviations.
F.
Average Atomic Masses of Elements
1. Average atomic mass is the weighted average of the atomic
masses of the naturally occurring isotopes of an element.
2. Calculating the average atomic mass is done by multiplying each
atomic mass by its relative abundance than summing up all of the
isotopes.
G. Relating Mass to Numbers of Atoms - done using several terms.
1. The mole is the amount of a substance that contains as many
particles as there are atoms in exactly 12 g of carbon-12.
2. Avogadro’s Number has been determined to be the number of
atoms in 12 g of carbon-12. This number has been experimentally
determined to be 6.0221367 * 1023 atoms in 12 g of carbon-12.
Therefore, 6.0221367 * 1023 atoms = 1 mole of anything.
3. Molar mass is the number of grams in 1 mole of a pure
substance. It usually has units of grams/mole (g/mol). This
number is numerically equal to the atomic mass of the element in
amu (from the periodic table). A molar mass of an element
contains one mole of atoms.
4. We can use this information to convert from grams to moles of a
substance (g/mol).
5. We can use Avogadro’s Number to find the number of atoms or
molecules in a sample of a substance (mol/atoms or molecules).
Chapter 21: Nuclear Chemistry
I.
The Nucleus
A. Mass Defect and Nuclear Stability
1. Protons and neutrons are collectively known as nucleons.
2. Mass Defect: the difference between the mass of the atom
and the sum of the masses of its subatomic particles.
3. Nuclear Binding Energy: energy released when a nucleus is
formed from nucleons. Sometimes the binding energy is given
as binding energy per nucleon.
4. Binding energy can be calculated using E=mc2.
B. Nucleons and Nuclear Stability
1. The neutron/proton ratio is a good indicator of nuclear
stability.
2. For smaller atoms 1:1 works best. As atoms get larger best
values increase up to about 1.5:1. As the number of protons
increases, the repulsions increase. More neutrons are required
to stabilize the nucleus.
3. Beyond atomic number 83 (Bi) no stable nuclei exist.
4. A graph of # of neutrons vs. # of protons (Figure 2 on page
683) shows the band of stability.
5. Stable nuclei tend to have an even number of nucleons.
Things are even more stable if there is an even number of both
protons and neutrons.
6. The most stable nuclei are those having 2, 8, 20, 28, 50, 82, or
126 total nucleons. These are called magic numbers.
II. Radioactive Decay
A. Radioactive Decay: the spontaneous disintegration of a nucleus
into a slightly lighter nucleus accompanied by the emission of
particles, electromagnetic radiation, or both.
B. Nuclear Radiation: particles or EM radiation emitted from the
nucleus during radioactive decay.
C. There are several types of radiation involved in radioactive decay:
1. Alpha Particle(): 2 protons and 2 neutrons, essentially a
helium nucleus.
2. Beta Particle(): an electron emitted from the nucleus
during radioactive decay.
3. Positron: a particle with the same mass as an electron, but
has a positive charge, that is emitted from the nucleus.
4. Gamma Rays(): High energy EM waves emitted from a
nucleus as it changes from an excited state to a ground state.
D. There are 5 types of radioactive decay that we will discuss.
1. Alpha Emission: an alpha particle is released from the
nucleus. (ex. 210Po --> 206Pb + 4He)
2. Beta Emission: a beta particle is released from the
nucleus. (ex. 14C --> 14N + 0)
3. Positron Emission: a positron is released from the nucleus.
(ex. 38Kr --> 38Ar + 0)
4. Electron Capture: an electron is absorbed by the nucleus.
(ex. 106Ag + 0e --> 106Pd)
5. Gamma Emission: gamma rays are released. Usually occurs
after another type of decay.
E.
Half-life (t1/2): Time required for one half of the atoms of a
radioactive nuclide to decay.
1. After 1 half life, 50% of the original radioactive atoms remain.
The other 50% have decayed.
2. After 2 half lives, 25% of the original radioactive atoms remain.
3. After 3 half lives, 12.5% of the original radioactive atoms
remain.
4. After 4 half lives, 6.25% of the original radioactive atoms
remain.
5. If we know the length of a half life & the % of the sample that
is radioactive, we can calculate the age of the sample. This is
the process used in radioactive dating.
F.
Decay Series
1. A series of radioactive nuclides produced by successive
radioactive decays until a stable nuclide is reached.
2. The heaviest nuclide of each decay series is called the parent
nuclide.
3. The nuclides produced by the decay of the parent nuclides are
called daughter nuclides.
G. Artificial Transmutations
1. Artificial Radioactive Nuclides: radioactive nuclides not
found naturally on Earth.
2. Artificial Transmutations: bombardment of nuclei with
charged and uncharged particles.
3. Charged particles require very large amounts of energy of to
overcome repulsive forces. This is accomplished by using
particle accelerators.
III. Nuclear Radiation
A. Radiation Exposure
1. The units used to measure radiation exposure are the
Roentgen (R) and the Roentgen Equivalent, Man (rem).
B. Biological effects of radioactivity
1. No detectable effects at public maximum levels to about 50
times public maximum levels.
2. Effects to high levels of exposure: Short term: 50x-200x:
reduced red and whit blood cell count; 200x-600x: nausea,
vomiting; 600x-1200x: loss of hair, trouble fighting infections;
more than 1200x: severe central nervous system damage,
death. Long term: Cancers, leukemia, cataracts, genetic
defects, potential sterility.
3. Alpha is stopped by layer of dead skin or piece of paper.
Beta/positron is stopped by 100-150 pieces of paper, or 0.51 cm of water. 50% of gamma rays are stopped by 1 cm of
lead, or 5-15 cm of water.
4. Driving a car is “safe” (1:6K), living at home is “safe” (Falls:
1:20K, Fires: 1:50K, poisoning: 1:40K, total: 1:10K). Radiation
is “safe” (1:20K).
C. Radiation Detection
1. Many different devices are used.
2. Film Badges: use exposure of film to measure the approximate
radiation exposure of people working with radiation.
3. Geiger Counters: instruments that detect radiation by counting
electric pulses carried by gas ionized by radiation.
4. Scintillation Counters: instruments that convert scintillating light
to an electric signal for detecting radiation. (Scintillate: absorbs
ionizing radiation and emit visible light.)
D. Uses of radiation
1. Radioactive Dating
2. Medical: X-rays, cancer treatments, radioactive tracers.
3. Power generation: Nuclear power plants. 16% of worldwide
energy production is nuclear. 19.8% in U.S.
4. Food irradiation: Kills bacteria and viruses in foods. Food does not
become radioactive itself.
E. Nuclear Waste
1. Containment of Nuclear Waste: difficult because of the half life.
2. Storage of Nuclear Waste: Commonly spent fuel rods from nuclear
reactors. Each reactor has storage pools on site. When the pools
are full, the waste is removed to dry storage areas until disposal.
3. Disposal of Nuclear Waste: There are 77 sites around the U.S.
IV. Nuclear Fission and Fusion
A. Nuclear Fission
1. Process by which a very heavy nucleus splits into more stable
nuclei of intermediate mass.
2. Releases large amounts of energy. Mass is converted directly into
energy.
3. A Chain Reaction is a reaction in which the material that starts
the reaction is also one of the products and can start another
reaction.
4. The minimum amount of nuclide that provides the number of
neutrons needed to sustain a chain reaction is called the critical
mass.
5. Nuclear power plants use energy from nuclear reactors to produce
electrical energy.
6. Each reactor has control rods that are neutron absorbing rods that
help control the reaction by limiting the number of neutrons
produced by the fission.
B. Nuclear Fusion
1. Low mass nuclei combine to form a heavier, more stable nucleus.
2. Releases much more energy than fission. This is the reaction that
powers our sun and stars.
3. Requires temperatures around 108 K to induce fusion.
4. Currently, fusion requires much more energy input than is
generated by the reaction, making it an inefficient process.