Transcript Atom

Atomic Structure
Chemistry 112
History of the Atom:
Time Line
• Find the following scientists in your textbook.
They are not given to you in the order they are
found in the text. That would be too easy! Some
may be found in other chapters besides 4 and 5!!
• Write down the years they were alive or the year
that their main discovery was made.
• Write down their contributions to the development
of today’s model of the atom.
• Create a creative, colorful, and informative
timeline!
Scientist to Include on Timeline:
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Rutherford
Millikan
Dalton
Marie and Pierre Curie
Democritus
Thompson
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Roentgen
De Broglie
Heisenberg
Chadwick
Bohr
Aristotle
Most of the following terms will
be a review for you
What is an Atom?
• Atom – The smallest particle of an element
that retains the properties of the element.
• How big is an atom? VERY SMALL!!!
• World Population
6,000,000,000
• Number of Atoms in a penny:
29, 000,000,000,000,000,000,000
• Scanning tunnel microscope allows us to
see atoms
Subatomic Particles found in the
nucleus of an Atom:
• Neutron: neutral (no charge) particle found
inside the nucleus of an atom.
• Proton: positive particle found in the
nucleus of an atom.
• Both the proton and neutron are
approximately the same mass (1 amu)
Subatomic particle outside of the
nucleus:
• Electrons are negatively charged particles
found outside the nucleus of an atom.
• The mass of an electron is negligible, only
1/1836th of an amu.
• The mass of the electrons do not enter into
calculating the atomic mass of an atom.
Atomic Mass Unit
• Abbreviated: amu
• Chemists developed a method of measuring the
mass of an atom without using very small numbers
in scientific notation.
• They chose an atomic standard: Carbon-12
• They agreed the mass of Carbon-12 was 12 amu.
• Therefore, 1 amu is 1/12 the mass of a Carbon-12
atom.
Atomic Number
• The whole number next to each element that
represents the number of protons in an
atom. (Proton number can never be altered).
• In a neutral atom the proton number will
equal the electron number. However atoms
can lose or gain electrons, even though they
can’t lose or gain protons under normal
circumstances.
Atomic Mass
• The atomic mass (when rounded to a whole
number) represents the # of protons + the
number of neutrons.
• If you subtract the atomic number (# of
protons) from the atomic mass (rounded to a
whole #) you will find the number of
neutrons in an atom.
Isotopes
• Atoms of the same element can be found in
nature with a different atomic mass than
others. This is because they can have
different #’s of neutrons
• An Isotope of an element is an atom of the
same element with a different # of neutrons.
Isotopes
• Isotopes can be expressed two ways:
– 1. C-14 (where 14 would be the whole number
mass)
– 2 146C (Where 14 is the mass and 6 is the
atomic number)
Why is the Atomic Mass not a
Whole Number?
• Atomic masses have decimals!
• Atomic mass – weighted average mass of
the isotopes of that element.
• For Example: Mass of Cl is 35.453
• Isotopes: 75% Chlorine-35
25% Chlorine-37
Finding Protons, Neutrons, and
Electrons
• Example: K-39
– This is a potassium atom with a mass of 39
– Look on the periodic and see the the atomic
number is 19.
– Therefore # of protons = 19 (equal to atomic #)
– Since it is a neutral atom (no charge) number of
electrons = 19
– Number of neutrons = 39 – 19 = 20
Ions
• Ions are charged atoms that have lost or
gained electrons.
• Two types of ions:
– Cations: positive ions that have been formed by
the loss of electrons. (Ca+2)
– Anions: negative ions that have been formed by
the gaining of electrons. (Cl-1)
4.4 Unstable Nuclei and
Radioactive Decay
• Late 19th century scientists noticed some
elements spontaneously emitted energy and
particles called radiation.
• Elements that give off radiation are said to
be radioactive.
• Thus, atoms are not unchangeable as Dalton
once thought.
• Nuclear reactions – a reaction that involves
a change in the nucleus of an atom.
• Radioactive decay – when nuclei are
unstable and gain stability by emitting
radiation.
• Fill out the Chart for Types of Radiation
Types of Radiation
Alpha
Beta
Gamma
What is it made of?
2 protons and 2
neutrons (like the
nucleus of a
Helium atom)
An electron
Electromagnetic
radiation
Mass
4 amu
1/1836th of an amu No mass; energy
Charge
positive
negative
none
Penetrating ability
low
medium
high
What stops it?
paper
Foil/metal
lead
Information from Chapter 25
• Transmutation – conversion of an atom of
one element to an atom of another element
by spontaneous emission of radiation.
• Induced Transmutation – nuclear reactions
produced artificially by striking a nucleus
with a high-velocity charged particle.
• Transuranium elements – all elements after
Uranium on Periodic Table.
– Produced in laboratory by induced
transmutation
– All are radioactive
Radioactive Decay Rates
• Half-life – the time required for one-half of
a radioisotope’s nuclei to decay into its
products. (see table 25-4)
– Ex. Strontium-90 half life is 29 years
• If today I have 10.0 g, in 29 years I would have 5.0g
• In another 29 years (total of 58), I would have 2.5 g
Nuclear Fission/Fusion
• HUGE amounts of energy produced by
Nuclear reactions.
• Nuclear Fission – splitting of a nucleus into
fragments to become more stable.
• Used in nuclear power plants (controlled)
• And nuclear bombs (uncontrolled)
• Nuclear Fusion – Combination of nuclei to
form a more stable nucleus.
• Large energy released!
• Sun powered by Fusion of hydrogen into
helium.
• Requires extremely high temperature to
occur.
• Scientist researching Cold Fusion
Atomic Model Development…
• Around the end of the 1700s Dalton
believed that an atom was a solid
indestructible mass and so did most
scientists.
• In the 1800s when JJ Thomson discovered
the electron (first subatomic particle) that
theory was shattered.
Model Continued
• The other subatomic particles were found and
further disproved the idea of the atom being a
solid mass.
• Niels Bohr, who was a student of Rutherford,
came up with the planetary model where electrons
followed specific paths around the nucleus. He
also stated that for an electron to move from one
energy level to another it gained or lost a quantum
of energy.
Model cont…
• Around 1925 Erwin Schrodinger came up
with the quantum theory of electron
placement which put electrons in clouds or
regions not in specific paths. These areas of
probability for electron location is based on
a mathematical calculation.
Identifying Where Electrons are Located in
the Atom
• The number and arrangement of electrons around
the nucleus determines the atoms chemical
properties. Specifically the outer electrons
(valence electrons) Therefore, scientists need a
shorthand way to write out electron arrangement.
• Electron Configuration
electrons in an atom.
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Arrangement
of
• Principle Quantum Number – Energy level the
electron occupies (n)
•
Examples: n= 1, 2, 3, 4, 5, 6, 7

The larger the value of n, the farther away the
electrons are from the nucleus and the higher the
energy of the electron. Electrons are lazy, so they
want to be as close to the nucleus as they can so
they can expend the LEAST amount of energy.

Only a certain amount of electrons can fit in
each energy level. The way to find the maximum
in each level is: 2n2
•o
•o
•o
•o
Example: n =1 Can have 2(1)2 electrons =
2 electrons in energy level 1
You solve for n = 2
n=3
n=4
 How many energy levels does hydrogen
have?
• Helium?
• Lithium?
• Sodium?
• Calcium?
• Xenon?
Sublevels
• Sublevels – further explain where the
electrons are located in each energy level.
 The names of the sublevels we use are:
s, p, d, and f
Sublevels

The number of sublevels in each energy
level is equal to n
•o
n = 1 has one sublevel
s
•o
n = 2 has two sublevels
s, p
•o
n = 3 has three sublevels
s, p, d
•o
n = 4 has four sublevels
s, p, d, f
•o
n = 5 has five sublevels but the fifth one is
not used in ground state elements.
• Notice, energy levels must have an s sublevel
before a p!
Orbitals
• Orbitals – Each sublevel has specific orbitals the
electrons can be in:
• Each orbital can hold 2 electrons. Therefore, the
maximum number of electrons in the s orbital is 2.

S has 1 orbital maximum electrons = 2

P has 3 orbitals maximum electrons = 6

D has 5 orbitals maximum electrons = 10

F has 7 orbitals maximum electrons = 14
Order of Filling
• http://lectureonline.cl.msu.edu/~mmp/perio
d/electron.htm
• Three Main Rules Electrons Abide By:

Aufbau principle: (“Electrons are Lazy”)
each successive electron occupies the lowest
energy orbital available.

Exclusion Principle Pauli (“Two to Tango”):
maximum of two electrons may occupy a single
atomic orbital, but only if the electrons have
opposite spins.
• Hund’s Rule (“Why share if you don’t have to?”)
: Single electrons with the same spin must occupy
each equal-energy orbital before additional
electrons with opposite spins can occupy the same
orbital.
• Orbital Diagrams
 Used to show how electrons are
distributed within sublevels and to show the
direction of spin.
 Boxes are used to represent orbitals.
 Arrows are used to represent electrons.
 The first electron in the orbital is
represented by an arrow pointing up, ,
meaning clockwise spin.
 The second electron in the orbital is
represented by an arrow pointing down, ,
meaning counterclockwise spin.
An electron configuration notation is an
abbreviated form of the orbital diagram.
• Electron Dot Diagrams
Valence Electrons – electrons in the atom’s
outermost energy level.
These electrons are involved in forming
chemical bonds.
 They are represented visually by an electron dot
structure.
Also known as Lewis Dot Formulas.
• Examples: