Download Chapters 4 and 5

Chapters 4 and 5
The Structure of the Atom
And
Electrons in Atoms
Early Theories of Matter

Democritus (460-370 B.C.)

Named atom (atomos)
Early Theories of Matter

Aristotle (384-322 B.C.)
Early Theories of Matter

John Dalton (1766-1844)

First Atomic Theory
Defining an Atom



The smallest particle
of an element that
retains the
properties of the
element.
About 1 X 10-10 m in
diameter.
Can be seen with a
scanning tunneling
microscope.
Discovering the Electron

William Crookes (1800’s)
Discovering the Electron

J.J. Thomson (late 1890’s)



Determined the charge-to-mass ratio
Mass must be less than a hydrogen
atom
Plum Pudding Model of atom
Discovering the Electron

Robert Millikan
(1909)
Determined charge of
electron
 1/1840 mass of a
hydrogen atom

The Nuclear Atom

Ernest Rutherford (1911)
The Nuclear Atom

Atom contains:



Mostly empty space
Tiny, dense nucleus
which is positively
charged
Creates nuclear
model of atom
Other Subatomic Particles

Rutherford (1920)



Concluded nucleus contains proton
Proton as equal but opposite charge of
electron
James Chadwick (1932)


Discovered neutron
Neutron has no charge
Subatomic Particles
How Atoms Differ

Moseley (shortly after Gold Foil)


Atoms of each element contain a unique
number of protons
Atomic Number= #protons

Identifies the atom
Isotopes



Isotopes – atoms that contain the same
number of protons but different number of
neutrons.
Most elements contain a mixture of
isotopes.
The relative abundance of each isotope is
constant.
Isotopes

Mass Number = #protons + #neutrons
Simple Practice
Atomic Mass
# of
# of
# of
Number Number Protons neutrons electrons
Mg
12
25
Zn
30
Be
4
9
Hg
80
200
12
65
30
4
80
13
12
35
30
5
4
120
80
Mass of Atoms


Atomic mass unit – 1/12 of a carbon-12
atom.
Atomic Mass – weighted average mass of
the isotopes of that element.
Calculating Atomic Masses
X has mass of 6.015 amu and
abundance of 7.50%. 7X has mass of
7.016 amu and abundance of 92.5%.
 6

(6.015)(.0750) + (7.016)(.925) = 6.94
amu
More Challenging Problems!


Cu-63 has a mass of 62.940 amu and
an abundance of 69.17%. Find the
mass and abundance of the other
isotope.
Boron has two isotopes with the masses
of 10.013 amu and 11.009 amu. Find
the abundance of each isotope.
Radioactivity

Nuclear Reactions – changes an atom’s
nucleus.




Atom changes into a new element
Due to unstable nuclei
Radiation contains rays and particles emitted
from a radioactive material.
Radioactive decay is the spontaneous
emission of radiation.
Types of Radiation
Types of Radiation
Radiation
Type
Symbol
Mass
(amu)
Charge
Alpha
 or
4
2+
1/1840
1-
0
0
4
He
2
Beta
e- or
0

1
Gamma
0
0
Nuclear Reactions


Mass numbers and Atomic numbers on both
sides of the reaction must be equal
Practice Problem:
14
0
C
  _______
6
1
Chapter 5
Electrons in Atoms
Electromagnetic Radiation


Electromagnetic Radiation is a form of energy
that has wave-like behavior.
4 properties of waves: wavelength,
amplitude, speed and frequency.
Properties of Waves


Frequency()- number of waves that pass a
given point per second. (hertz or 1/s or s-1)
Speed (c)- is constant for all waves. 3 x 108
m/s
Calculating Properties of
Waves



c=
What is the frequency of light with a
wavelength of 5.80 x 10-7 m?
A radio station broadcasts with a frequency
of 104.3 MHz. What is the wavelength of
the broadcast?
Particle Nature of Light

Max Planck (1900) discovered that matter
can gain or lose energy in small, specific
amounts called quanta.

Equantum= h

Planck’s Constant (h)=6.626 x 10-34J·s
Practice Problems



What is the energy of a wave with a
frequency of 6.25 x 1019Hz?
What is the frequency of a wave that
contains 8.64 x 10-18J of energy?
A wave contains 4.62 x 10-15J of energy.
Determine its wavelength.
Photoelectric Effect




Photoelectric effect – electrons are emitted from a
metal’s surface when light of a certain frequency
shines on it.
Frequency (color) of light, not brightness of light
determines if electrons are emitted.
Einstein (1905)- light has wave-like properties but
is also a stream of tiny particles or bundles of
energy called photons.
Photon – a piece of EM with no mass and carries a
quantum of energy.
Atomic Emission Spectrum


When atoms absorb energy they become
excited.
Atomic Emission Spectrum- unique set of
frequencies emitted by excited atoms.
Bohr Model of the Atom

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Bohr (1913) proposed why the emission
spectrum of hydrogen is not continuous.
Electrons can have only certain “energy states”
Ground State - the lowest allowable energy
state.
Excited State – energy state of an electron
when it gains energy
Bohr Model of the Atom
Electrons as Waves


Louis de Broglie (1924) thought Bohr’s
model had electrons having similar
properties to waves.
h

de Broglie equation:
m
 Predicts that all moving particles have
wave properties.
Heisenberg Uncertainty
Principle


When viewing an
electron, a photon of
light hits it and changes
the velocity and position
of the electron.
It is impossible to know
precisely both the
velocity and position of
a particle at the same
time.
Quantum Mechanical Model of
the Atom



Schrödinger (1926) derived an
equation that treated hydrogen’s
electron as a wave.
Allows electron to have only
certain energy but does not give
path of electron.
Atomic orbital – a 3-D region
around the nucleus in which the
electron can be found 90% of
the time.