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Chapters 4 and 5
The Structure of the AtomAnd
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
AtomicNumber
MassNumber
# of Protons
# of neutrons
# of electrons
Mg 25 12
Zn 30 35
Be 4 9
Hg 120 80
12 13 12
65 30 30
4 5 4
80 200 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
6X has mass of 6.015 amu and abundance of 7.50%. 7X has mass of 7.016 amu and abundance of 92.5%.
(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+
Beta e- or 1/1840 1-
Gamma 00 0 0
He2
4
1
0
Nuclear Reactions
Mass numbers and Atomic numbers on both sides of the reaction must be equal
Practice Problem:
_______1
0
6
14
C
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 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.
de Broglie equation:
m
h
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.