Download Chapter 10: The Atom

Matter and Energy
Unit 4: Matter and Energy
Chapter 10: The Atom
 10.1 Atomic
Structure
 10.2
Quantum Theory of the Atom
 10.3
Nuclear Reactions
10.1 Investigation: The Atom
Key Question:
How is an atom organized?
Objectives:
Build atom models.
 Describe the relationship between the number of protons,
neutrons, and electrons in an atom to its atomic and mass
numbers.
 Infer that not all atoms of an element are identical.

Atomic structure

English scientist John Dalton (1766–1844) started experimenting with
gases in the atmosphere in 1787.

In 1808, Dalton published a detailed atomic theory that contained the
following four statements:
1.
Matter is composed of tiny, indivisible, and indestructible particles
called atoms.
2.
An element is composed entirely of one type of atom. The properties of
all atoms of one element are identical and are different from those of
any other element.
3.
A compound contains atoms of two or more different elements. The
relative number of atoms of each element in a particular compound is
always the same.
4.
Atoms do not change their identities in chemical reactions. They are
just rearranged into different substances.
Review: Elements and Compounds
 An
element is composed of one type of atom.
 A compound
contains atoms of more than one
element in specific ratios.
Thomson’s “plum pudding” model
 English
physicist J. J. Thomson
(1856–1940) observed that
streams of particles could be
made to come from different
gases placed in tubes carrying
electricity.
 Thomson
identified a
negatively-charged particle he
called the electron.
A gold foil experiment and the nucleus

In 1911, Ernest Rutherford (1871–
1937), Hans Geiger (1882–1945),
and Ernest Marsden (1889–1970)
launched fast, positively-charged
helium ions at extremely thin
pieces of gold foil.

They expected the helium ions
would deflect a small amount as a
We now know that every atom has
result of hitting gold atoms.

Instead, most of the helium ions
passed straight through the foil!
a tiny nucleus, that contains more
than 99 percent of the atom’s mass.
Three subatomic particles
 A proton
 An
is a particle with a positive charge.
electron is a particle with a negative charge.
 A neutron
is a neutral particle and has a zero
charge.
 The
charges on one proton and one electron are
exactly equal and opposite.
 Charge
is an electrical property of particles that
causes them to attract and repel each other.
Inside an atom
 The
mass of the nucleus
determines the mass of an
atom because protons and
neutrons are much larger
and more massive than
electrons.
 In
fact, a proton is 1,836
times heavier than an
electron.
Volume of an atom
 The
size of an atom is determined by how far the
electrons are from the nucleus.
 The
electrons define a region of space called the
electron cloud.
 If
an atom was the size of a football stadium, the
nucleus would be the size of a pea, and the electrons
would be like a few gnats flying around the stadium
at high speed.
Fundamental forces inside atoms
 Electrons
are bound to the
nucleus by electromagnetic
force, the attractive force
between electrons (-) and
protons (+).
 Because
of Newton’s first
law, the electrons do not fall
into the nucleus, because
they have inertia.
Fundamental forces inside atoms
 What
holds the nucleus together?
 There
is another force that is even stronger than
the electric force.
 We
call it the strong nuclear force.
Historical development of atomic force
 Henry
Cavendish (1731–
1810), a British scientist, was
the first to measure the
gravitational force between
two masses using a torsion
balance.
 Cavendish
detected a very
small torque between the
large and small spheres.
Historical development of atomic force
 The
unit of electric charge is the coulomb (C) in
honor of Charles-Augustin de Coulomb (1736–1806),
a French physicist who measured the
electromagnetic forces between charges in 1783.
Historical development of atomic force
 Hideki
Yukawa (1907–1981), was the first Japanese
to receive a Nobel Prize for his theory of the strong
nuclear force.
 This
theory predicted the meson, an elementary
particle that was discovered later.
Historical development of atomic force
 A theory
about the existence of the weak force was
first proposed by Enrico Fermi (1901–1954), an
Italian physicist who worked on the first nuclear
reactor and its applications.
 Fermi’s
decay.
theory was based on his observations of beta
How atoms of various elements differ
 The
atoms of different
elements contain different
numbers of protons in the
nucleus.
 Because
the number of
protons is so important, it is
called the atomic number.
How atoms of various elements differ
 Isotopes
are atoms of the
same element that have
different numbers of
neutrons.
 The
mass number of an
isotope tells you the number
of protons plus the number of
neutrons.
How are these carbon
isotopes different?
Average atomic mass
 Elements
in nature are
usually a mixture of
isotopes.
 The
element lithium has
an atomic mass of 6.94.
 On
average, 94% of
lithium atoms are lithium-7
and 6% are lithium-6.
Calculating average atomic mass
How many neutrons are present in an
aluminum atom that has an atomic number of
13 and a mass number of 27?
1.
Looking for: … number of neutrons in aluminum-27.
2.
Given: … atomic no. = 13; mass no. = 27
3.
Relationships: Periodic table says atomic no. = proton no.
and mass no. = protons + neutrons
4.
Solution: neutrons = mass no. – proton no.
no. neutrons = 27 – 13 = 14
Unit 4: Matter and Energy
Chapter 10: The Atom
 10.1 Atomic
Structure
 10.2
Quantum Theory of the Atom
 10.3
Nuclear Reactions
10.2 Investigation: Energy and the Quantum
Theory
Key Question:
How does an atom absorb
and emit light energy?
Objectives:

Distinguish between atoms in the ground and excited states.

Use the Photon and Lasers game to simulate the absorption
and emission of light from an atom.
The spectrum
 Each
different element has its
own characteristic pattern of
colors called a spectrum.
 The
colors of clothes, paint,
and everything else around
you come from this property of
elements to emit or absorb
light of only certain colors.
 Each
element emits a
characteristic color of light.
The spectrum
 Each
individual color in a spectrum is called a
spectral line because each color appears as a
line in a spectroscope.
 A spectroscope
is a device that spreads light into
its different colors.
Quantum theory and the Bohr atom
 Danish
physicist Neils Bohr
proposed the concept of
energy levels to explain the
spectrum of hydrogen.
 When
an electron moves from
a higher energy level to a
lower one, the atom gives up
the energy difference between
the two levels.
 The
energy comes out as
different colors of light.
Quantum theory

The Bohr atom led to a new way of thinking
about energy in atomic systems.

A quanta is a quantity of something that cannot be divided
any smaller.

One electron is a quanta of matter, because you can’t split an
electron.

Quantum theory says that when a particle, such as an
electron, is confined to a small space inside an atom, the
energy, momentum, and other variables of the particle
become quantized and can only have specific values.
Quantum theory
 In
1925, Erwin Schrödinger (1887–
1961) proposed the quantum model of
the atom we still use today.
 Quantum
theory says that when
things get very small, like the size of
an atom, matter and energy do not
obey Newton’s laws or other laws of
classical physics.
 The
electron is thought of as a “fuzzy”
cloud of negative charge called a
quantum state rather than as a
particle moving around the nucleus.
Electrons and energy levels
 In
the current model of the atom, we think of the
electrons as moving around the nucleus in an
area called an electron cloud.
 The energy levels occur because electrons in the
cloud are at different average distances from the
nucleus.
Pauli exclusion principle
 According
to the quantum
model, two electrons can
never be in the same
quantum state at the
same time.
 This
rule is known as the
Pauli exclusion
principle after Wolfgang
Pauli (1900–1958).
Planck’s constant
 The
“smearing out” of particles into fuzzy
quantum states becomes important when size,
momentum, energy or time become comparable
in size to Planck’s constant.
 If
you measure the momentum of an electron in a
hydrogen atom and multiply it by the size of the
atom, the result is about 1 × 10–34 joule·seconds.
The uncertainty principle
 The
work of German physicist Werner Heisenberg
(1901–1976) led to the uncertainty principle.
 According
to the uncertainty principle, a particle’s
position, momentum, energy, and time can never be
precisely known in a quantum system.
 The
uncertainty principle arises because the
quantum world is so small.
The uncertainty principle
Probability and quantum theory
 Because
electrons are so tiny, this type of calculation
is not possible. Instead, quantum theory uses
probability to predict the behavior of large numbers
of particles in a system.
 Probability
describes the chance for getting each
possible outcome of a system.
Wave function

In quantum theory, each quantum
of matter or energy is described by
its wave function.

The wave function mathematically
describes how the probability for
finding a particle is spread out in
space.

Quantum theory can only make
accurate predictions about the
behavior of large systems with
many particles.
Unit 4: Matter and Energy
Chapter 10: The Atom
 10.1 Atomic
Structure
 10.2
Quantum Theory of the Atom
 10.3
Nuclear Reactions
10.3 Investigation: Nuclear Reactions and
Radioactivity
Key Question:
How do nuclear changes
involve energy?
Objectives:

Determine the fraction of a radioactive sample that remains in
its original isotope after an integer number of half lives.

Explain how probability and half life are related concepts.

Describe the three different types of radioactive decay (alpha,
beta, and gamma decay).
Chemical vs. Nuclear Reactions
 The
involvement of energy in
chemical reactions has to do
with the breaking and forming of
chemical bonds.
 A nuclear
reaction involves
altering the number of protons
and/or neutrons in an atom.
 The
total amount of mass and
energy is conserved in nuclear
reactions.
Nuclear reactions and energy
 Protons
and neutrons
are attracted by the
strong nuclear force and
release energy as they
come together.
 Nuclear
reactions often
involve huge amounts of
energy, as protons and
neutrons are rearranged
to form different nuclei.
Nuclear reactions
and energy
 A nuclear
reaction that
changed 1 kg of uranium
into 1 kg of iron would
release 130 trillion J of
energy.
Isotope notation
 In
a nuclear reaction, each atom is represented using
isotope notation.
 In
this notation, the element symbol is given along
with its mass number and atomic number.
How many protons,
neutrons and electrons
are found in this
isotope?
Fusion reactions

Fusion reactions (the combining of
atomic nuclei) only release energy if the
final nucleus has lower energy than the
initial nuclei.

The fusion reaction to make magnesium
from carbon actually goes through a
nuclear changes.

The end result is that 56 TJ are released
as the protons and neutrons in 1 kg of
carbon-12 are rearranged to make 1 kg
of magnesium-24 nucleus.
Fusion
 Nuclear
fusion occurs in the Sun and the resulting
energy released provides Earth with heat and light.
Fission reactions

For elements heavier than iron,
breaking the nucleus up into
smaller pieces (fission) releases
nuclear energy

A fission reaction can be started
when a neutron bombards a
nucleus and makes it unstable

The fission of 1 kg of uranium into
the isotopes molybdenum-99 and
tin-135 releases 98 TJ.
Chain reactions
 A chain
reaction occurs
when the fission of one
nucleus triggers fission of
many other nuclei.
 The
increasing number of
neutrons causes even more
nuclei to have fission
reactions and releases
enormous amounts of
energy.
Fission
 A nuclear
reactor is a power plant that uses fission
to produce heat.
Radioactive Decay

The products of fission usually
have too many neutrons to be
stable and are radioactive.

Radioactive means the nucleus
continues to change by
ejecting protons, neutrons, or
other particles.

A process of radioactive decay
results in an unstable,
radioactive isotope like carbon14 becoming the more stable
isotope nitrogen-14.
Radioactive Decay

There are three types of
radioactive decay:
1.
alpha decay,
2.
beta decay, and
3.
gamma decay.
Radioactive Decay

In alpha decay, the nucleus ejects two protons and two
neutrons.
 Beta decay occurs when a neutron in the nucleus splits into a
proton and an electron.
 Gamma decay is not truly a decay reaction in the sense that
the nucleus becomes something different.
Atomic Number
 Remember,
the atomic
number is the number of
protons all atoms of that
element have in their
nuclei.
 If
the atom is neutral, it will
have the same number of
electrons as protons.
Half-life
 The
half-life of carbon-14
is about 5,700 years.
 If
you start out with 200
grams of C-14, 5,700
years later only 100
grams will still be C-14.
 The
rest will have
decayed to nitrogen-14.
Half-life
 Most
radioactive materials
decay in a series of
reactions.
 Radon
gas comes from
the decay of uranium in
the soil.
 Uranium
(U-238) decays
to radon-222 (Ra-222).
Carbon dating

Living things contain a large amount of carbon.

When a living organism dies it stops exchanging carbon
with the environment.

As the fixed amount of carbon-14 decays, the ratio of C-14
to C-12 slowly gets smaller with age.
Applications of radioactivity
 Many
satellites use radioactive decay
from isotopes with long half-lives for
power because energy can be produced
for a long time without refueling.
 Isotopes
with a short half-life give off
lots of energy in a short time and are
useful in medical imaging, but can be
extremely dangerous.
 Smoke
detectors contain a tiny amount
of americium-241, a radioactive isotope.
Indirect Evidence and Archaeology

Using indirect evidence is
transforming the field of
archaeology.

Using remote sensing techniques,
archaeologists can locate and
describe features of ancient
civilizations before a shovel ever
touches the soil.

Centuries of foot travel and camel
caravans packed down the desert
floor so that 1,700 years later,
they still reflected infrared
radiation differently than the
surrounding terrain.