Download Chapter 4 Atomic Structure

Chapter 4
“Atomic Structure”
Introduction to the Atom
and Atomic Models
DEMOCRITUS
400
BC
Democritus believed all things
consisted of tiny indivisible units.
He called these tiny units he
called atomos. The Greek word
for “can not be cut” or
“indivisible”
Ancient philosopher: Father of the Atom
John Dalton
(1799)
 Developed what is considered to
be the 1st Atomic Theory
 Was born into a modest Quaker
family in England
 Began lecturing in public at the age
of 12
Dalton’s Model (1799)

Dalton's model was that the atoms were tiny,
indivisible, indestructible particles and that each one
had a certain mass, size, and chemical behavior that
was determined by what kind of element they were.
Dalton’s Model
Dalton’s model of the
atom was similar to a
tiny billiard ball.
Dalton’s model of the
atom was solid and had
no internal structure.
Dalton’s Atomic Theory

elements consisted of tiny particles called
atoms.
all atoms of an element are identical
 atoms of each element are different from one
another; they have different masses.
 compounds consisted of atoms of different
elements combined together.
 chemical reactions involved the rearrangement of
combinations of those atoms.

Flaws in Dalton’s Model

Dalton’s falsely believed that the atom was
the most fundamental particle.


We now know the atom is made up of even
smaller particles we call the proton, neutron
and electron.
Dalton’s theory could also not account for
the formation of ions (charged particles)
Daltons Atomic Model Summary
Called: Billiard Ball Model
 Could account for
 Atoms of different atomic masses
 Elements were tiny particles
 Could NOT account for
 Though atom was smallest particle
 Did not have an internal structure
 The formation of charged particles

John J. Thomson (1897)

Discovered the electron using the
Cathod Ray Tube (CRT)

Thomson found that the beam of
charge in the CRT was attracted
to the positive end of a magnet
and repelled by the negative end.
Thomson’s Hypothesis

Concluded that the cathode beam was a stream of negative
particles (electrons).

He tested several cathode materials and found that all of
them produced the same result.

He also found that the charge to mass ratio was the same for
all electrons regardless of the material used in the cathode or
the gas in the tube.

Thomson concluded that electrons must be part of all
atoms.
Thomson’s atomic model

Called the “plum-pudding” it was the most
popular and most wildly accepted model of the
time.
Thompsons atomic model could
account for…..



the atom having an internal
structure
Light given off by atoms
Atom with different atomic
masses
Thompsons atomic model could NOT
account for…..


Empty space (had atom filled
with positive pudding)
Formation of ions
Gold Foil Experiment

Conducted by students of Rutherford.

Proved that all atoms had a tiny, positively
charged center.

Confirmed that atom’s were mostly empty
space.
Rutherford ~ early 1900s

α-particle interaction with matter
studied in gold foil experiment
Rutherford's Nuclear Model
1. The atom contains a tiny dense center
 the volume is about 1/10 trillionth the volume
of the atom
2. The nucleus is essentially the entire mass
of the atom
3. The nucleus is positively charged
 the amount of positive charge of the nucleus
balances the negative charge of the
electrons
4. The electrons move around in the empty
space of the atom surrounding the nucleus




Rutherford’s atomic model (1911)

Could account for:
Empty space
 Ions
 Internal structure
 Light given off when heated
to high temperature.


Could not account for:

Stability
Philipp Lenard (1903)

Aluminum foil experiment

Lenard found that a beam of electrons was able
to pass through a sheet of Al foil with almost no
deflection.

Lenard correctly concluded that the majority of
an atom’s volume is empty space.
Lenard’s atomic model


Lenard’s model was composed of
dynamids.
Lenard calculated the size of a dynamid
based on his experimental results and
found it to be 1 billionth
(1/1,000,000,000) the size of the atom.
Lenard’s model could account for:

Different atomic masses (based on the number of
“dynamids”).

The internal structure of an atom

The fact that most of the atom was empty space.
Flaws in Lenard’s model

Formation of Ions (gaining or losing charge)

The light given off by materials when heated to a
high temperature.
Hantaro Nagaoka
• First to present an atomic model close
to the presently accepted model.
• He came up with his model in 1903.
Nagoka’s Model

Nagoka’s model of the atom was unstable.

According to the laws of planetary motion, the
atom would collapse over time.
Planetary model


Planetary model used to explain
electrons moving around the tiny, but
dense nucleus
Nucleus contains
Protons- existence proposed in 1900s
 Neutrons- existence proposed in 1930s

Successes of Nagoka’s model




Atoms were able to give
off electrons to form ions.
Accounted for the
experimental fact that
atom’s were mostly empty
space
Explained different atomic
weights.
Explained the light given
off when heated to high
temperatures.
Bohr

Questioned ‘planetary model’ of atom


Electrons located in specific levels from
nucleus (discontinuous model)
Proposed electron cloud model based on
evidence collected with H emission
spectra
Bohr’s Atomic Model (1913)

Bohr was a student of Rutherford.

Improved Rutherford’s model by proposing electrons are
found only in specific fixed orbits.

These orbits have fixed levels of energy

This explained how electrons could give off light (gain or
lose energy)
BOHR MODEL

Electrons are placed in energy levels
surrounding the nucleus
8e8eNucleus
(p+ & n0)
2e-
Bohr’s Atomic Model

Could account for
Internal structure
 Atoms of different masses
 Atom being mostly empty space
 Light given off
 Formation of positive ions


Flaws

Only really worked for Hydrogen
Chadwick (1932)

Discovered the neutron by bombarding Be with
beta radiation.

Nuclear fission released a neutron.
chart
Review

Describe each of the 6 different atomic models.
Give the
Scientist Name
 Name of model
 What they could account for
 What they could not account for (flaws)

Subatomic particle summary
Particle
Discovery by
Year
experiment
Proton
Rutherford
1911
Gold Foil Experiment
Electron
Thompson
1887
The response of cathode
ray tube to a magnetic
and electric fields
Neutrons
Chadwich
1932
Bombarded Be with beta
radiation and a neutron
was released
Subatomic Particles
Name
Symbol
Relative
Charge mass
Actual
mass (g)
Electron
e-
-1
1/1840
9.11 x 10-28
Proton
p+
+1
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
Subatomic Particles (cont.)

All atoms of an element have the same # of
protons
protons identify an atom  atomic #

Atoms are electrically neutral
#p = #e-

Only neutrons and protons contribute to an
atoms mass
#n + #p = atomic mass
ISOTOPES
= atoms
with the same
number of protons but
DIFFERENT
numbers of neutrons
Mass Number
Atomic Number
Element Symbol
Ex.
Na-23 or Sodium-23 or 23 Na or 23Na11
C-14 or Carbon-14 or 14C or 14C
6
B-10 or Boron-10 or 10B or 10B5
Isotope Practice
Element has
a. 6 p+, 8 n and 6 eb. 6 p+, 6 n and 6 ec. 19 p+, 21 n and 10 ed. __p+, __ n and __ ee. __p+, __ n and __ e-
Symbol
a. .
b. .
c. .
d. Cu-65
e. H-2
Ions





Ion is an atom that has gained or lost one or more
electrons and now has a charge
Metals always lose electrons to form POSITIVE
ions
Non-metals always gain electrons to form
NEGATIVE ions
The charge of the element is show on the right
side of the symbol as a super script
Example: Na+1, Zn+1, S-2, N-3
Ions practice
Element
Ca+2
Br-1
Al+3
I-1
Number of electrons
?
?
?
?
D. Average Atomic Mass



weighted average of all isotopes
on the Periodic Table
round to 2 decimal places
D. Average Atomic Mass

EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O, and
0.20% 18O.
D. Average Atomic Mass

EX: Find chlorine’s average atomic mass if
approximately 8 of every 10 atoms are chlorine-35 and
2 are chlorine-37.