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CH 4 Atomic Structure
CH 4.1 Studying Atoms
Ancient Greek Models of Atoms
Democritus
(approx. 2500 years ago)
 Believed that all matter consisted of
extremely small particles that could not be
divided called atoms from the Greek word
atomos, which means “uncut” or “indivisible.”
Aristotle
 Believed that there was no limit to the
number of times matter could be divided.
Dalton’s Atomic Theory
Dalton’s Theory
(1803)
 All elements are composed of atoms.
 All atoms of the same elements have the
same mass, and atoms of different
elements have different masses.
 Compounds contain atoms of more than
one element.
 In a particular compound, atoms of
different elements always combine in the
same way.
 Solid sphere model.
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Thomson’s Model of the Atom
Thomson’s Model
(1897)
 Thomson’s experiment was a tube with
metal disks at each end filled with gas that
created a beam when electricity was
applied.
 The beam was attracted to a positive plate.
 Provided the first evidence that atoms are
made of even smaller negative particles.
 Revised Dalton’s model to account for
these subatomic particles.
The “plum pudding” model:
Since atoms are neutral overall, the negative
charges were evenly scattered throughout
the atom filled with a positively charged
mass of matter.
(Looks like vanilla pudding with chocolate chips)
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Rutherford’s Atomic Theory
Rutherford’s Theory:
In 1899 Rutherford discovered that
uranium emits fast-moving particles that
have a positive charge. He named them
alpha particles.
In 1911, Ernest Marsden, a student of
Rutherford, conducted an experiment.
 Predicted that most particles would travel
in a straight path and a few would be
slightly deflected.
(see page 104)
The Gold Foil Experiment-Marsden aimed a
narrow beam of alpha particles at the gold.
 Many passed through without being
deflected.
 More were deflected than expected. Some
bounced straight back.
 The paths that were deflected came close to
another positively charged object.
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Rutherford concluded that the positive
charge of an atom is NOT
evenly spread throughout the atom.
It is concentrated in a very small,
central area that he called
the nucleus.
Rutherford’s Model
 All of an atom’s positive charge is
concentrated in its nucleus.
 The alpha particles that deflected more
than 90 degrees came very close to the
nucleus.
 The alpha particles whose paths were
not bent moved through without
coming close to any nucleus.
nucleus
(Remember the nucleus!!!)
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CH 4.2 Structure of the Atom
By 1920, Rutherford had evidence for the
existence of two subatomic particles and had
predicted the existence of a third subatomic
particle. (1.protons 2. electrons 3. neutrons)
Nucleus- The positively charged center of
an atom. Contains both protons and neutrons
(referred to as the nucleons). Contains most of
the mass of the atom. (See page 109)
Subatomic Particles
Protons (p+)-- Positively charged particles with
a relative mass equal to that of a neutron. (1 amu)
Each proton has an electrical charge of +1.
Electron (e-)-- Negatively charged particles
that surround the nucleus. 1/2000 of the
mass of a proton (relative mass is zero).
Each electron has an electrical charge of –1.
Neutron (n)-- Neutral (no charge) particles with
a relative mass equal to that of a proton. (1 amu)
Each neutron has an electrical charge of 0.
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Atomic Number and Mass Number
Atomic number- The number of protons in
the nucleus. {Example: Carbon has 6 protons;
therefore the atomic number is 6}
***Atoms are neutral by having the same
number of protons and electrons***
Mass number- The sum of the number of
protons and neutrons in an atom.
# of neutrons = mass # - atomic #
Isotopes
Isotopes- Atoms of the same element that
have different numbers of neutrons and
different mass numbers.
{Examples: Boron-11 and Boron-10}.
Atomic mass- An averaged value that
depends on the distribution of an element’s
isotopes in nature and the masses of those
isotopes.
Atomic mass unit (amu)- The unit of measure
of the particles in an atom. The mass of a proton
or a neutron equals 1u. (pg 109)
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4.3 Modern Atomic Theory
Bohr’s Model of the Atoms
Electron cloud- In the 1926 model of the
atom, the electrons move about in a cloud
that surrounds the nucleus rather than welldefined orbits as Niels Bohr, a Danish
scientist, depicted in 1913 that resembles a
solar system of planets. (See pages 114-115)
The electrons are so small that it is
impossible to determine where they are in
the electron cloud at any given moment.
{Example: See page 116. Propeller blades of an
airplane – When in motion, the blades spin so fast that
you only see a blur.}
Energy Levels- Levels (a.k.a. shells, shelves,
and orbits) within the electron cloud. The
electrons are at various distances from the
nucleus.
Electrons can move from one energy level
to another by gaining and losing energy.
{Example: closest-low energy is the ground state;
farthest-high energy is the excited state}.
Sci joke: “Why does hamburger yield a lower energy than steak?”
“Because it is in its ground state!”
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Each energy level has a maximum number
of electrons it can hold.
Energy Level
1
2
3
4
Number of
Orbitals
1
4
9
16
Maximum #
of Electrons
2
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8 or 18*
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*[A maximum of 8 electrons for the first 20 elements.
All other elements will have
9 to 18 electrons in the 3rd energy level].
Draw and label examples of energy levels.
(See notes posted in room 129.)
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