Atoms, Sub-Atomic Particles

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Unit 3

• Atoms, Sub-Atomic Particles Nuclear Chemistry

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The Particle Theory of Matter

• In 400 B.C. Democritus, a Greek philosopher,
  first proposed the idea of a basic particle of
  matter that could not be divided any further.
• He called this particle the atom, based on the
  Greek word atomosmeaning indivisible.
• This early theory was not backed up by
  experimental evidence and was ignored by the
  scientific community for nearly 2000 years.

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Foundations of Atomic Theory

• By the late 1700s, experiments with chemical
  reactions led to the discovery of 3 basic laws
• The Law of Conservation of Mass.
• This law states that mass is neither created nor
  destroyed during ordinarychemical reactions or
  physical changes.
• Formulated by Antoine Lavoisier in 1789.

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Foundations of Atomic Theory
The Law of Conservation of Mass
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Foundations of Atomic Theory

• The Law of Definite Proportions
• States that 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.
• Discovered by Joseph Proust in 1797.

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Foundations of Atomic Theory

• The Law of Multiple Proportions
• States that if different compounds are composed
  of the same 2 elements, then the ratio of the
  masses of the elements is always a ratio of
  small whole numbers.
• Published by John Dalton in 1804.

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Daltons Atomic Theory

• In 1808, John Dalton proposed an explanation for
  the three laws. His atomic theory states
• All matter is composed of atoms.
• Atoms of the same element are identical atoms of
  different elements are different.
• Atoms cannot be subdivided, created, or
  destroyed.
• Atoms of different elements combine in simple
  whole-number ratios to form chemical compounds.
• In chemical reactions, atoms are combined,
  separated, or rearranged.

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Corrections to Daltons Theory

• Dalton turned Democrituss idea into a scientific
  theory that could be tested by experiment.
• But not all aspects of Daltons theory have
  proven to be correct. We now know that
• Atoms are divisible into even smaller particles.
• A given element can have atoms with different
  masses.

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Daltons Atomic Model

• An atom is the smallest particle of an element
  that has all the properties of that element.
• Atoms are too small to seeeven through the
  most powerfulmicroscope!!
• Dalton thought atoms were solid balls of matter
  and were indivisible.

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Discovery of the Electron

• In 1897, Joseph John (JJ) Thomson showed that
  cathode rays are composed of identical negatively
  charged particles, which were named
  electrons.
• The electron was the first subatomic particle to
  be discovered.

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Thomsons Cathode Ray Tube Experiment
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Charge and Mass of the Electron

• In 1909, Robert Millikan measuredthe charge on
  the electron during his oil drop experiment.
• Using the charge-to-mass ratio, scientists were
  able to figure out the mass of the electron
  about 1/2000 the mass of a hydrogen atom.

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Millikans Oil Drop Experiment
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Thomsons Plum Pudding Model

• After the work of Thomson andMillikan, the
  accepted model of the atom was called the plum
  pudding model.
• The atom was viewed as a ball of
  positively- charged material with tiny
  negatively-charged electrons spread
  evenly throughout.

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Discovery of the Atomic Nucleus

• More detail of the atoms structure was provided
  in 1911 by Ernest Rutherford and his associates
  Hans Geiger and Ernest Marsden.
• The results of their gold foil experiment led to
  the discovery of a very densely packed bundle
  of matter with a positive electric charge.
• Rutherford called this positive bundle of matter
  the nucleus.

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The Gold Foil Experiment
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Rutherfords Atomic Model

• After Rutherfords gold foil experiment, the
  accepted model of the atom looked like this
• A small, positively-charged nucleus with negative
  electrons surrounding it at some distance away.
  Most of the atom is empty space.

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Structure of the Atom

• Rutherford proposed that the nucleus had
  particles with the same amount of charge as an
  electron but the opposite sign, called protons.
• Relative charge 1
• Relative mass 1 amu
• For an atom to be neutral there must be equal
  numbers of protons and electrons.
• Throughout the 1920s scientists accepted an
  (incorrect) model of the atom composed of protons
  and electrons.

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Some Problems

• How could beryllium have 4 protons stuck together
  in the nucleus?
• shouldnt they repel each other?
• If a beryllium atom has 4 protons, then it should
  weigh 4 amu but it actually weighs 9.01 amu!
  Where is the extra mass coming from?
• Each proton weighs approximately 1 amu.
• The electrons mass is only about 0.00055 amu and
  Be has only 4 electrons, so they dont account
  for the extra 5 amu of mass.

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There Must Be Something Else There!

• These questions were answered in 1932 by James
  Chadwick (a student of Rutherfords), who
  discovered another particle in the nucleus, which
  he called a neutron.
• Charge 0 (no charge).
• Relative mass 1 amu.

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Subatomic Particles

• The nucleus is made up of at least one
  positively charged particle called a proton and
  usually one or more neutral particles called
  neutrons.
• Protons, neutrons, and electrons are often
  referred to as subatomic particles.

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Atomic Number

• The atomic number (Z) of an element is the number
  of protons of each atom of that element.
• Atoms of the same element all have the same
  number of protons.

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Mass Number

• The mass number is the total number of protons
  and neutrons in the nucleus of an atom.
• Atoms of the same element can
  havedifferentmass numbers.

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Isotopes

• Isotopes are atoms of the same element that have
  different masses.
• Isotopes have the same number of protons and
  electrons but different numbers of neutrons.
• Most of theelements consist of mixtures
  ofisotopes.

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Designating Isotopes

• Hyphen notation The mass number is written with
  a hyphen after the name of the element.
• uranium-235
• Nuclear symbol The superscript indicates the
  mass number and the subscript indicates the
  atomic number.

Mass number
Atomic number
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Calculating Neutrons

• The number of neutrons is found by subtracting
  the atomic number from the mass number.
• mass number - atomic number number of neutrons
• Nuclide is a general term for a specific isotope
  of an element.

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Calculating Subatomic ParticlesSample Problem

• How many protons, electrons, and neutrons are
  there in an atom of chlorine-37?
• Solution
• Number of protons
• Number of electrons
• Number of neutrons

atomic number (on periodic table)
17
number of protons
17
mass number - protons
20
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The Atomic Mass Unit

• The standard used by scientists to compare units
  of atomic mass is the carbon-12 atom.
• One atomic mass unit, or 1 amu, is exactly 1/12
  the mass of a carbon-12 atom.
• The atomic mass of any atom is determined by
  comparing it with the mass of the carbon-12
  atom.

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Average Atomic Mass

• Average atomic mass is the weighted average of
  the atomic masses of the naturally occurring
  isotopes of an element.
• The average atomic mass of an element
  depends on both the mass and the
  relative abundance of each of the
  elements isotopes.

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Calculating a Weighted Average

• If you have the following grades, what would your
  marking period average be?
• First, change percents to decimals.
• Next, multiply each grade by its decimal percent.
• Finally, add up all the products.

Category Percent of Grade Grade
Tests 45 72
Quizzes 10 78
Labs 25 84
HW/CW 10 98
Project 10 94
100 0.45
x 32.4
x 7.8
100 0.10
100 0.25
x 21.0
x 9.8
100 0.10
x 9.4
100 0.10

80.4
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Calculating Average Atomic MassSample Problem 1

• Copper consists of 69.15 copper-63, which has an
  atomic mass of 62.929 601 amu, and 30.85
  copper-65, which has an atomic mass of 64.927 794
  amu. What is the Average Atomic Mass of Copper?
• Solution
• Change percents to decimals.
• Multiply the atomic mass of each isotope by its
  relative abundance.
• Add up all of the products.

Relative Abundance Mass
Cu-63 0.6915 62.93
Cu-65 0.3085 64.93
x 43.52
x 20.03

63.55
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Calculating Average Atomic MassSample Problem 2

• A student believed that she had discovered a new
  element and named it mythium. Analysis found it
  contained two isotopes. The composition of the
  isotopes was 19.9 of atomic mass 10.013 and
  80.1 of atomic mass 11.009. What is the average
  atomic mass, and do you think mythium was a new
  element?
• Solution
• Average Atomic Mass(.199 x 10.013) (.801 x
  11.009) 10.811
• Because the atomic mass is the same as the atomic
  mass of boron, mythium was not a new element.

Round off to
10.8
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Forces in the Atom

• Electrons and protons attract because of opposite
  electrical charges, but protons and protons repel
  since they have the same charge.
• The nucleus is heldtogether by a mysterious
  force called the strong nuclear force which
  only exists between nucleons (protons and
  neutrons) whichare very close together.

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Valley of Stability
for Z 1 ? 20, stable N/Z 1
for Z 20 ? 40, stable N/Z approaches 1.25
for Z 40 ? 80, stable N/Z approaches 1.5
for Z gt 83, there are no stable nuclei
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Naturally Radioactive Elements

• All of the elements beyond atomic number 83 are
  unstable and thus radioactive.

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Nuclear Reactions

• Large, unstable nuclei spontaneously break apart
  to form smaller, more stable nuclei.
• A nuclear reaction is a reaction that affects the
  nucleus of an atom.Example
• A transmutation is a change in the identity of a
  nucleus as a result of a change in the number of
  its protons.

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Nuclear ReactionsSample Problem

• Identify the products that balance the following
  nuclear reactions
• a.
• b.
• Solution
• Atomic Mass 212 4 _____
• Atomic Number 84 2 _____
• Atomic Mass 22 ____ 22
• Atomic Number 11 ____ 10

208
82
0
-1
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Radioactive Decay

• Radioactive decay is the spontaneous
  disintegration of a nucleus into a lighter
  nucleus, accompanied by nuclear radiation.
• Nuclear radiation is particles and/or
  electromagnetic radiation emitted from the
  nucleus during radioactive decay.

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Types of Radioactive Decay

• Alpha Emission
• An alpha particle (a) is two protons and two
  neutrons bound together and is emitted from the
  nucleus during some kinds of radioactive decay.
• The atomic number decreases by two and the mass
  number decreases by 4.

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Types of Radioactive Decay (continued)

• Beta Emission
• A beta particle (ß) is an electron emitted from
  the nucleus during some kinds of radioactive
  decay (a neutron can be converted into a proton
  and an electron.)
• The atomic number increases by one and the mass
  number stays the same.

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Types of Radioactive Decay (continued)

• Positron Emission
• A positron (ß) is a particle that has the same
  mass as an electron, but has a positive charge
  (to decrease the number of protons, a proton can
  be converted into a neutron by emitting a
  positron.)
• The atomic number decreases by one and the mass
  number stays the same.

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Types of Radioactive Decay (continued)

• Electron Capture
• In electron capture, an inner orbital electron is
  captured by the nucleus of its own atom. (An
  inner orbital electron combines with a proton to
  form a neutron.)
• The atomic number decreases by one and the mass
  number stays the same.

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Types of Radioactive Decay (continued)

• Gamma Emission
• Gamma rays (?) are high-energy electromagnetic
  waves emitted from an unstable nucleus.
• Atomic number and mass number both stay the
  same because gamma rays have no charge and no
  mass.
• They are pure energy, and very dangerous to
  living things.

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Comparing Alpha, Beta and Gamma

• Alpha particles are big andslow. They cant
  penetrate skin or paper.
• Beta particles have about 100 x the penetrating
  power of alpha. They can be stopped by clothing,
  wood, or aluminum foil.
• Gamma rays have the greatest penetrating ability.
  They can only be stopped by a thick layer of
  lead or concrete. They cause a lot of damage to
  living cells.

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Half-Life

• Half-life is the time required for half the
  atoms of a radioactive nuclide to decay.
• Each radioactive nuclide has its own half-life.
  
• More-stable nuclides decay slowly and have
  longer half-lives.

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Half-LifeSample Problem 1

• The half-life of radon-222 is 4 days. After what
  time will ¼ of a given amount of radon remain?
• Solution
• Determine the number of half-lives that it takes
  to cut a sample to ¼ of the original amount.
• Then multiply that number by the half-life.

8 days
2
x
4 days

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Half-LifeSample Problem 2

• Phosphorus-32 has a half-life of 14.3 days. How
  many
• milligrams of phosphorus-32 remain after 57.2
  days if
• you start with 4.0 mg of the isotope?
• Solution
• Determine the number of half-lives in 57.2 days.
• For each half-life, multiply the original amount
  by ½

57.2 days
4 half-lives

14.3 days

0.25 mg
x
½
x
½
4.0 mg
x
½

x
½
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Uses of Radiation

• Radioactive dating scientists can determine the
  approximate age of an object based on the amount
  of certain radioactive nuclides present.

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Uses of Radiation (continued)

• Radioactive tracers are radioactive atoms that
  are incorporated into substances so that movement
  of the substances can be followed by radiation
  detectors.
• Radioactive tracers can be used by doctors to
  diagnose diseases.
• Radioactive tracers are also used in agriculture
  to determine the effectiveness of fertilizers.

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Uses of Radiation (continued)

• Irradiated Food nuclear radiation is used to
  prolong the shelf life of food.

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Uses of Radiation (continued)

• Nuclear Power Plants use energy as heat from
  nuclear reactors to produce electrical energy.
  They have five main components
• 1. Shielding radiation-absorbing material used
  to decrease exposure to radiation from nuclear
  reactors.
• Fuel usually Uranium-235.
• 3. Coolant usually water, it absorbs excess
  heat energy.
• 4. Control rods neutron-absorbing rods that
  limit the number of free neutrons
• 5. Moderator used to slow down the fast
  neutrons produced by fission.

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Nuclear Power Plant
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Whole-Body Radiation Exposure

• Rem The unit used to measure the biological
  effects of absorbed radiation in humans.

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Radiation Detection

• Film badges - use exposure of film tomeasure the
  approximate exposure of people working with
  radiation.
• Geiger-Müller counters - instruments that detect
  radiation by counting electric pulses carried by
  gas ionized by radiation.
• Scintillation counters - instruments that convert
  scintillating light to an electric signal for
  detecting radiation.

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Fission vs. Fusion

• Nuclear Fission very heavy nuclei split into
  smaller, more stable nuclei.
• Can occur spontaneously or when nuclei are
  bombarded by particles.
• Controlled fission chain reactions are used in
  nuclear power plants. Uncontrolled fission
  chain reactions are used in nuclear bombs.

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Fission vs. Fusion (continued)

• Nuclear Fusion low-mass nuclei combine to form
  a heavier, more stable nucleus.
• Fusion releases even more energy per gram of
  fuel than fission.
• Scientists are not yet able to control fusion
  reactions, so we cant use them in power plants.
• Fusion is the primary process that fuels our sun
  and the stars.

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Nuclear Waste

• Radioactive waste produced in nuclear reactors
  can take hundreds of thousands of years to decay.
• Disposal of nuclear waste is done with the
  intentionof never retrieving it.
• There are 77 disposal sites around the country.
  A new one (Yucca Mountain) is being developed for
  the permanent disposal of much of our nuclear
  waste beginning in 2017.