Chapter 4 Atomic Structure
Section 4.1 Early Atomic Theory
A. ________________________ (460-370 BC)
• First person to propose idea that matter was not infinitely divisible
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He believed matter was made of atomos
Aristotle (continuous - divide in half infinitely)
Neither had experimental evidence, Aristotle more respected, thus believed true
B. ____________________________ (1766-1844)
• His Atomic Theory
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1) All matter is made of extremely small particles called atoms
2) All atoms of a given element are identical.
Atoms of a specific element are different from those of any other element
3) Atoms cannot be created, divided into smaller particles, or destroyed
4) Different atoms combine in simple whole number ratios to form compounds
5) In a chemical reaction, atoms are separated, combined or rearranged
Model was called the ____________ __________.
Picture
All atoms of a given element are identical was proven wrong because of the existence of
__________
Atoms cannot be created, divided into smaller particles or destroyed was proven wrong because
of the existence of _____________, ________________ and ______________.
C. Defining the Atom
• Atom -
Section 4.2 Subatomic Particles and the Nuclear Atom
A. ________________________ (1856-1940)
• Experiment called Cathode Ray Tube experiment
• sealed tube, electrode at each end (cathode= -) (anode=+)
• beam emitted could be bent with magnet (NOT light)1
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Led to discovery of 1st subatomic particle called the ___________, small negative particle
Determined the mass of the charged particle was much less than that of a hydrogen atom, the
lightest known atom
Disproved Dalton’s theory that an atom was indivisible.
Model was called the ____________ __________.
Picture
B. __________________________ (1868-1953)
• Experiment called Oil Drop Experiment
• small negatively charged oil droplets, fell through
chamber (adjusted for gravity)
• measured charge of electron- (-1), which led him to calculate the mass 1/1837th of H atom
C. __________________________ (1871-1937)
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Experiment called Gold Foil Experiment
• Diagram of Experiment
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Conclusions from experiment
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2)
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In 1920, Rutherford refined his concept of the nucleus. He concluded that the nucleus contained
a subatomic particle called the __________. The charge of this particle was equal but opposite
of the electron; that is it has a positive one (+1) charge.
D. _______________________ (1891-1974)
• discovered the third subatomic particle called the ____________: high energy, no charge, same
mass as proton
E. Atom and Subatomic Particles
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Section 4.2 How Atoms Differ
A. _______________________ (1887-1915)
• discovered that the atoms of each element contain a unique positive charge in their nucleus. The
number of protons in an atom is referred to as the element’s atomic number.
1) Atomic Number:
The number of __________ in an atom
Identifies the element
The number of protons equals the number of ___________ in a neutral atom
Found on the periodic table
2) Mass Number:
Mass of the _________
Equal to the number of _________ plus _________
Units are a.m.u. (atomic mass unit)
Must be given or calculated CANNOT BE FOUND ON THE PERIODIC TABLE
B. Isotopes
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same number of protons (atoms of the same element) but different number of neutrons and
different mass number
Isotopic Notation
1) The element name is given with a dash and then the mass number.
Ex. Carbon-13
2) Write the symbol and then to the left of the symbol the mass number is given as a superscript
and the atomic number is given as a subscript.
Ex. 136C
For Example
Determine the number of protons, neutrons, electrons, mass number, atomic number and give the
isotopic symbol of the following isotopes:
Potassium -39
Potassium – 40
Potassium-41
Protons
Neutrons
Electrons
Atomic Number
Mass Number
Isotopic Symbol
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B. Ions
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an atom with a charge, which means ___________ are not equal to the ______________.
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positive ions are called _____________ In a positive ion, the __________ are greater
than the _________
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negative ions are called _____________ In a negative ion, the __________ are greater
than the ____________
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To determine the charge, subtract the ____________ from the _________________.
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Equation: Charge = protons – electrons
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The charge is shown as a superscript to the right of the symbol.
Ex. As-3
24
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11𝑁a
33 -2
16𝑆
240 +6
92𝑈
Protons
Neutrons
Electrons
Atomic
Number
Mass Number
Name
C. Average Atomic Mass
• carbon-12 standard for average atomic mass (perfect 12 amu)
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amu = 1/12 mass of carbon-12 atom
average mass of all isotopes of atom
use a mass spectrometer to figure out the masses and abundances of each type of isotope
average mass is weighted by abundance and the reason why most have decimals
found on the periodic table.
Equation:
Sample Problem - Average Atomic Mass
Naturally occurring lead is found to have natural percentage abundances of 1.48% 204Pb, 23.6%
206Pb, 22.6% 207Pb, and 52.3% 208Pb. Calculate the average atomic mass of lead.
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Chapter 6 and 7 PERIODIC TABLE PART I
DMITRI MENDELLEV (Russian, 1834-1907)
 Arranged the known elements in order of increasing atomic mass
Noticed a repeating pattern in the properties of the elements
Designed a table with rows and columns to show the repeating pattern
 Left blank spaces in the table to represent elements not known in his time
Predicted the chemical properties of those elements
MODERN PERIODIC TABLE
 Henry Moseley(British, 1887-1915) determined the nuclear charge, also called the atomic
number, of the atoms of the elements
Arranged the elements in a table by order of atomic number
 Elements are still arranged by increasing atomic numbers today
 “periodic” name given because of the similar properties that repeat every so many elements
The periodic law:
When the elements are arranged in order of increasing atomic number, there is a periodic pattern in
their physical and chemical properties.
 Table is arranged in columns (groups or families) and rows (periods)
COLUMNS
 Each column called a group
 18 groups, numbered 1 to 18
 Groups 1,2 and 13-18 are called main-group, or representative elements. They are also
labeled IA, IIA and IIIA-VIIIA respectively.
 Groups 3-12 are called transition metals
Note: the IUPAC (International Union of Applied Chemistry) suggested the numbering 1-18 in
order to avoid confusion. The Roman numerals and As and Bs explained above are used in North
America, but most tables show both conventions.
 Groups reflect the number of electrons in the outer shell(valence electrons). Group IA has 1
valence electron, Group IIA and the Bs have 2 valence electrons, Group IIIA has 3…
Lewis Dot Notation – way to represent the number of valence electrons using the symbol
Ex. Sulfur is in Group 16 or 6A… It has 6 valence electrons
Dot notation would be:
 Elements in each group resemble each other; they react in similar ways to other substances
 Elements in the same group are similar because they have the same number of electrons in the
outer shell.
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Examples:
 Group IA Alkali metals
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Group IIA Alkaline earth metals
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Group VIIA Halogens
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Group VIIIA Noble Gases
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Transition metals
ROWS
 7 rows (or periods)
 rows 6 and 7 are too long to fit in the table, so part of them are placed at the bottom of the table.
rows correspond to energy level, n
 Atomic numbers increase along each row
 Ordered according to the filling of successive shells with electrons.
 The two rows at the bottom of the table are called the inner-transition metals. The first row is
called the lanthanides and the second row is called the actinides.
METALS, NONMETALS AND METALLOIDS
 A stairstep line on the periodic table separates metals from nonmetals
 Metals are to the left of the stairstep line
 Nonmetals are to the right of the stairstep line
Properties of Metals
Properties of Nonmetals
Metalloids
 All elements on either side of the stairstep line EXCEPT Aluminum
 They have properties of both metals and non metals
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