Chapter 4

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Chapter 4Atomic Structure
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Section 4.1 Defining the Atom

• Greek philosopher Democritus
• (460 B.C. 370 B.C.) suggested
• existence of atoms (Greek word
• atomos)
• Believed atoms were indivisible and
  indestructible
• His ideas agreed with later scientific theory,
  but didnt explain chemical behavior - was not
  based on scientific methods only philosophy

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Daltons Atomic Theory (experiment based!)

• All elements composed of tiny indivisible
  particles called atoms
• Atoms of same element identical. Atoms of any
  one element are different from all other
  elements.

John Dalton (1766 1844)

1. Atoms of different elements combine in simple
   whole-number ratios to form chemical compounds
2. In chemical reactions, atoms are combined,
   separated, or rearranged but never changed into
   atoms of another element.

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Sizing up the Atom

• Elements subdivided into smaller particles
  called atoms, and they still have properties of
  that element
• 1.0 x 108 copper atoms in a single file, they
  would be approximately 1 cm long
• individual atoms are observable with instruments
  such as scanning tunneling (electron) microscopes

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Section 4.2Structure of the Nuclear Atom

• One change to Daltons atomic theory - atoms are
  divisible into subatomic particles
• Electrons, protons, and neutrons

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Discovery of Electron
1897 - J.J. Thomson used cathode ray tube to
deduce presence of negatively charged
particle.the electron
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Modern Cathode Ray Tubes
Television
Computer Monitor

• CRTs pass electricity through gas contained -
  very low pressure.

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Mass of the Electron
Mass of the electron is 9.11 x 10-28 g
The oil drop apparatus
1916 Robert Millikan determines mass of
electron 1/1840 the mass of hydrogen atom has
one unit of negative charge
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Conclusions from the Study of the Electron

1. Cathode rays have identical properties regardless
   of element used to produce them. All elements
   must contain identically charged electrons.
2. Atoms are neutral, so there must be positive
   particles in atom to balance negative charge of
   electrons
3. Electrons have so little mass that atoms must
   contain other particles that account for most of
   mass

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Conclusions from the Study of the Electron

• Eugen Goldstein in 1886 observed what is now
  called the proton - particles with a positive
  charge, and a relative mass of 1 (or 1840 times
  that of an electron)
• 1932 James Chadwick confirmed the existence of
  neutron particle with no charge, but mass
  nearly equal to proton

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Subatomic Particles
Particle Charge Mass (g) Location
Electron (e-) -1 9.11 x 10-28 Electron cloud
Proton (p) 1 1.67 x 10-24 Nucleus
Neutron (no) 0 1.67 x 10-24 Nucleus
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Thomsons Atomic Model
J. J. Thomson
Believed electrons were like plums embedded in
charged pudding, called plum pudding model.
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Ernest RutherfordsGold Foil Experiment - 1911

• Alpha particles - helium nuclei w/ charge -
  The alpha particles were fired at thin sheet of
  gold foil
• Particles that hit on the detecting screen
  (film) were recorded

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Rutherfords problem
In the following pictures, there is a target
hidden by a cloud. To figure out the shape of the
target, we shot some beams into the cloud and
recorded where the beams came out. Can you figure
out the shape of the target?
Target 2
Target 1
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The Answers
Target 2
Target 1
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Rutherfords Findings

• Most particles passed through
• Few deflected
• VERY FEW greatly deflected

Like howitzer shells bouncing off of tissue
paper!
Conclusions

1. Small nucleus
2. Dense nucleus
3. charge nucleus

The Atom Song
Atoms song - Mark Rosengarten
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The Rutherford Atomic Model

• His experimental evidence
• atom mostly empty space
• All positive charge, almost all mass in small
  center. Nucleus
• protons and neutrons make nucleus!
• electrons distributed around nucleusoccupy most
  volume
• His model called nuclear model

Rutherfords Atom 308
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Section 4.3Atomic Number

• All atoms composed of identical protons,
  neutrons, and electrons
• How then are atoms of one element different from
  another element?
• Elements different b/c they contain different
  of PROTONS
• atomic number of element is number of protons
  in nucleus
• protons in atom electrons

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Atomic Number
Atomic number (Z) of element is of protons in
nucleus of each atom of that element.
Element of protons Atomic (Z)
Carbon 6 6
Phosphorus 15 15
Gold 79 79
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Mass Number
Mass number is of protons and neutrons in
nucleus of an isotope
Mass p n0
Nuclide p n0 e- Mass
Oxygen - 10
- 33 42
- 31 15
18
8
8
18
Arsenic
75
33
75
Phosphorus
16
15
31
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Complete Symbols

• Contain symbol of element, mass number atomic
  number.

Mass number
X
Superscript ?
Atomic number
Subscript ?
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Symbols

• Find each of these
• number of protons
• number of neutrons
• number of electrons
• Atomic number
• Mass Number

80
Br
35
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Symbols

• If an element has an atomic number of 34 and a
  mass number of 78, what is the
• number of protons
• number of neutrons
• number of electrons
• complete symbol

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Symbols

• If an element has 91 protons and 140 neutrons
  what is the
• Atomic number
• Mass number
• number of electrons
• complete symbol

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Symbols

• If an element has 78 electrons and 117 neutrons
  what is the
• Atomic number
• Mass number
• number of protons
• complete symbol

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Isotopes

• Dalton was wrong about elements of same type
  being identical
• Atoms of same element can have different numbers
  of neutrons.
• different mass numbers
• isotopes

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Isotopes

• Frederick Soddy (1877-1956) proposed idea of
  isotopes in 1912
• Isotopes - atoms of same element with different
  masses, b/c varying s of neutrons
• Won 1921 Nobel Prize in Chemistry
• has a small crater named for him on the far side
  of the Moon.

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Naming Isotopes

• We can also put mass number after name of the
  element
• carbon-12
• carbon-14
• uranium-235

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Isotope Protons Electrons Neutrons Nucleus
Hydrogen1 (protium) 1 1 0
Hydrogen-2 (deuterium) 1 1 1
Hydrogen-3 (tritium) 1 1 2
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Isotopes
Elements occur in nature as mixtures of isotopes.
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Atomic Mass

• How heavy is an oxygen atom?
• Depends, b/c different kinds of oxygen atoms
  exist.
• Were more concerned with average atomic mass.
• Based on abundance () of each variety of that
  element in nature.
• Dont use grams - numbers tooooo small.

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Measuring Atomic Mass

• Atomic Mass Unit (amu)
• one-twelfth mass of a carbon-12 atom.
• Carbon-12 chosen b/c of its isotope purity.
• Each isotope has own atomic mass
• we determine average from abundance.

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To calculate the average

• Multiply atomic mass of each isotope by abundance
  (decimal), then add results.
• If not told otherwise, mass of isotope expressed
  in atomic mass units (amu)

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Atomic Masses
Atomic mass is average of all naturally occurring
isotopes of that element.
Atomic mass (amu)
12
13.00
14.00
Isotope Symbol Composition of the nucleus in nature
Carbon-12 12C 6 protons 6 neutrons 98.89
Carbon-13 13C 6 protons 7 neutrons 1.11
Carbon-14 14C 6 protons 8 neutrons lt0.01
What is the average atomic mass of Carbon?
12.01
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- Page 117
Question
Knowns and Unknown
Solution
Answer
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End of Chapter 4