Chapters 4 and 5 The Structure of the Atom And Electrons in Atoms.
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Transcript of Chapters 4 and 5 The Structure of the Atom And Electrons in Atoms.
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Chapters 4 and 5
The Structure of the AtomAnd
Electrons in Atoms
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Early Theories of Matter
Democritus (460-370 B.C.) Named atom (atomos)
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Early Theories of Matter
Aristotle (384-322 B.C.)
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Early Theories of Matter
John Dalton (1766-1844) First Atomic Theory
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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.
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Discovering the Electron
William Crookes (1800’s)
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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
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Discovering the Electron
Robert Millikan (1909) Determined charge
of electron 1/1840 mass of a hydrogen atom
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The Nuclear Atom
Ernest Rutherford (1911)
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The Nuclear Atom Atom contains:
Mostly empty space
Tiny, dense nucleus which is positively charged
Creates nuclear model of atom
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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
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Subatomic Particles
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How Atoms Differ
Moseley (shortly after Gold Foil) Atoms of each element contain a
unique number of protons Atomic Number= #protons
Identifies the atom
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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.
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Isotopes
Mass Number = #protons + #neutrons
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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
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Mass of Atoms Atomic mass unit – 1/12 of a carbon-
12 atom. Atomic Mass – weighted average
mass of the isotopes of that element.
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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
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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.
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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.
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Types of Radiation
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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
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Nuclear Reactions
Mass numbers and Atomic numbers on both sides of the reaction must be equal
Practice Problem:
_______1
0
6
14
C
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Chapter 5
Electrons in Atoms
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Electromagnetic Radiation
Electromagnetic Radiation is a form of energy that has wave-like behavior.
4 properties of waves: wavelength, amplitude, speed and frequency.
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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
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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?
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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
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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.
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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.
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Atomic Emission Spectrum When atoms absorb energy they
become excited. Atomic Emission Spectrum- unique set of
frequencies emitted by excited atoms.
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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
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Bohr Model of the Atom
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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.
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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.
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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.