Studying atoms is difficult
because they are too small to see
or directly observe even with the
best scientific tools. Write a
similar example of something that
can not be studied directly.
CHAPTER 4: ATOMIC
STRUCTURE
Learning Targets1. I can describe the structure of atoms.
2. I can describe how structure of an
atom affects it’s properties.
3. I can create a timeline that shows the
developments that lead to the current
model of the atom.
ANCIENT GREEK MODELS OF ATOMS
460-370 BC
Democratis- all matter consisted of extremely small particles that
could not be divided he called them atoms from the Greek word
Atomos which means “uncut” or “indivisiable”
384-322 BC
Aristotle- no limit to the number of times that matter could be divided
Accepted until the 1800s
THE EXISTENCE OF ATOMS WASN’T SCIENTIFICALLY
PROVEN UNTIL THE EARLY 1800S.
John Dalton, 1766-1844,
English chemist and teacher
• Studied the behavior of gases
in air
• Concluded that gas contains
individual particles
• No matter how large or small
the sample the ratio of the
elements in compounds is
always the same
DALTON’S ATOMIC THEORY
•
•
All elements are
composed of
atoms
All atoms of the
same element
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
DALTONS ATOMIC MODEL
Tiny solid
spheres
with
different
masses
Dalton’s theory explained data from
many experiments and thus
became widely accepted
ALTHOUGH ATOMS ARE INCREDIBLY SMALL THEY
CAN NOW BE OBSERVED WITH A SCANNING
TUNNELING MICROSCOPE
CHARGED MATERIALS
• Some materials when rubbed gain
the ability to attract or repel other
materials
• Gain either a positive or negative
charge
• Object with like charges repel or
push apart, objects with opposite
charges attract or pull together
• Charged particles can flow
creating an electric current
THOMSON’S MODEL OF AN ATOM
J.J. Thomson, an English physicist 1856-1940
Wires connect the metal disks at opposite ends of a empty glass tube, one
disk becomes negatively charged and one becomes positively charged
A glowing beam appears in the space between the plates
The beam is repelled by a negatively charged metal disk or attracted by a
positively charged plates brought near it
THOMSON’S MODEL
Thomson hypothesized the ray was a stream of negatively
charged particles contained inside atoms, now called
electrons, which are part of all atoms and carry a charge of
-1.
No matter what metal he used he got the same particles
The mass was always 2000 times smaller than the mass of
hydrogen atoms (a proton)
THOMSON’S MODEL
First evidence that atoms are made of even smaller particles
Atom is neutralnegative charges
scattered
throughout an
atom filled with
a positively
charged mass
of matter
ERNEST RUTHERFORD’S GOLD FOIL EXPERIMENT
Ernest Rutherford- 18711937
Atom was believed to have its
positive charge spread
throughout.
Rutherford shot alpha
particles (large 2 + atoms)
at a very thin sheet of gold
foil.
 If the current model of the
atom was correct the alpha
particles should pass though
gold the mass and charge
being too small to deflect the
alpha particles.
RUTHERFORD’S ATOMIC THEORY
Most alpha particles actually
passed straight through a
small fraction bounced off
the gold foil at large angles
 Atoms are mostly empty space
 Positive charge is concentrated
in the nucleus, not evenly
distributed, which contains
the protons and neutrons and
has a positive charge
 Positive charge varies among
elements
 Each nucleus must contain at
least one proton, each proton
is assigned a charge of +1
RUTHERFORD’S ATOMIC MODEL
All of an atoms
positive charge is
concentrated in
the dense
nucleus
Electrons are
outside the
nucleus
JAMES CHADWICK
1932
James Chadwick, English physicist did experiments that proved the
neutrons existed
Concluded that they were neutral because they were not effected by a
charged particle
Neutrons are contained in the nucleus and have a mass equal to that
of a proton
Compare the mass,
location and charge
of protons,
neutrons, and
electrons.
Rutherford's atomic model couldn’t explain chemical properties of
elements
 required knowledge of electron behavior
Niels Bohr (1885-1962)
Focused on Electrons
Agreed with Rutherford that the nucleus of and atom was surrounded
by a large volume of space
THE BOHR MODEL
Electrons are only found in
specific circular pathsorbits around the
nucleus
 Each electron orbit has a
fixed energy or energy
level
 An electron cannot exist
between energy levels
 The energy level closest to
the nucleus is the lowest
 An electron in an atom
can move from one energy
level to the next when it
gains or loses energy
 Energy lost can be in the
form of light
THE QUANTUM MECHANICAL MODEL
Rutherford-Bohr described the path of an electron as a large object would behave
which was inconsistent with theoretical calculations and experimental results
Electrons move in a much less predictable way then planets in a solar system.
Erwin Schrodinger 1887-1961
Devised and solved a mathematical equation describing the behavior of the electron
in hydrogen atom
QUANTUM MECHANICAL (ELECTRON CLOUD)
MODEL
 Electron Cloud- visual model of the
most likely locations for electrons in
an atom
 Based on the probability of finding an
electron with in a certain volume of
space surrounding the nucleus is
described as a fuzzy cloud where the
electron is 90% of the time
 More dense- probability high
 Less dense- probability low
ATOMIC ORBITALS
•
The electron cloud represents all the orbitals in an atom
•
An orbital is a region of space around the nucleus where an electron is likely to be
found
•
Each orbital can hold 2 electrons
•
The lowest energy level has one orbital, the second has four, the third 9 and the
fourth 16
•
Electron configuration- the arrangement of electrons in the orbitals of an atom
•
The most stable electron configuration is the one in which the electrons are in
orbitals with the lowest possible energies this is called the ground state
DISTINGUISHING AMONG ATOMS
How are atoms of Hydrogen
different than atoms of
oxygen?
Element of different atoms are
contain different numbers of
protons
Atomic Number- the
number of protons in
the nucleus of an atom
of a given element
-Identifies an Element
-Atoms are electrically neutral so
the number of electrons are also
equal to the atomic number
Mass Number- the number of
protons and neutrons in the nucleus
•Example- carbon has 6 protons
and 6 neutrons so the mass
number is 12
•Most the mass of an atom is
concentrated in the nucleus
•The number of neutrons in an
atom is the difference between the
mass number and the atomic
number.
•# neutrons = mass # - atomic #
Elements can be
represented in the
following
shorthand
notation:
symbol
Mass #
Au
197
79
Atomic #
Or Gold-197
Isotopes- atoms that
have the same number
of protons but a
different number of
neutrons
•Different mass
numbers
•Chemically alike- same
protons and electrons
which are responsible
for chemical behavior
Atomic Mass
•The mass of atoms are givin in comparison to carbon-12
•An atomic mass unit (amu) = 1/12 the mass of a carbon-12
atom
•Carbon 12 amu
•Flourine- actual mass 3.155 x 10 –23 g, atomic mass- 18.998
amu
•1 proton or neutron- 1 amu
•Atomic mass- weighted average mass and relative abundance of
isotopes as they occur in nature
Atomic
•Example- Hydrogen –1, 99 %, 1.0078
massHydrogen- 2, 1% heaver
1.0079
Hydrogen- 3