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Chapter 6: The Periodic Table and Periodic Law
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Development of the Periodic Table
• 1790s – Antoine Lavoisier composed a list of the 23
known elements– Included gold, silver, carbon, and oxygen
• Electricity which is used to break down compounds into elements led to an “explosion” in chemistry as did the spectrophotometer and the industrial revolution.
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• 1864– John Newlands proposed an organization
scheme for the elements– Arranged by increasing atomic mass and
noticed that the properties of the elements repeated after every 8th element (PERIODIC)
– See fig 6.2 page 153 for picture
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Meyer, Mendeleev, and Moseley, OH MY!!!
• Lothar Meyer and Dmitri Mendeleev showed a connection between atomic mass and elemental properties
• Mendeleev published first!!!• Left spaces on the Periodic Table for the
unknown elements • By noting trends in the periodic table, he was
able to predict the properties of yet to be known elements.
• Mendeleev organized the periodic table by atomic mass
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Mendeleev’s Periodic Table• http://http://z.about.com/d/chemistry/1/0/0/W/mendeleevperiodic.jpg
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• Mendeleev was not completely correct – More accurate measurements of atomic mass
• Mosely (1913)- arranged elements in order of increasing atomic number – Resulted in clearer patterns of properties
• PERIODIC LAW:There is a periodic repetition of chemical and physical
properties of the elements when they are arranged by increasing atomic number
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The Modern Periodic Table
• Groups - the columns of the periodic table(Sometimes called families)
• Periods – the rows of the periodic table
• SEE PAGE 154 Fig. 6.4
• Representative Elements (labeled 1A-8A)• Transition Elements (labeled 3B-12B)
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Classifying the Elements (3 Types)
1. Metals (solid, shiny, good conductors)• Group 1A: Alkali Metals
• Most reactive of all metals
• Group 2A: Alkaline Earth Metals• Also very chemically reactive but not as much as
the alkali metals
• Group 3A: Transition Metals (main part of table) and Inner Transition Metals (bottom two rows, lanthanide and actinide series)
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2. Nonmetals – generally gases, dull, brittle, poor conductors– Group 7A is called the halogens and are very reactive – Group 8A is called the Noble Gases and are
unreactive
3. Metalloids – Phys. and chem. properties of both metals and
nonmetals– Border on the stair-step line– Silicon and Germanium are two most important
(comp. chips)
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6.2: Classification of the Elements
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• Valence electrons • Found in highest principle energy level• All elements of group 1A have the same number of
valence electrons; therefore, have same chemical properties
• Valence electrons by period• The energy level by the valence electrons are
found reveals the period
• Valence electrons by group• The group number corresponds the number of
valence electrons
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The s, p, d, and f elements
• Review pages 160-161 in case you had trouble or may be a little confused
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6.3 PERIODIC TRENDS
This is a very important section!!!!!
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• Many properties on the periodic table change in a very predictable manner
• Includes:– Atomic Radius– Ionic Radius– Ionization Energy– Electronegativity
• YOU MUST MEMORIZE THESE!!!!!!
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1. Atomic Radius
• Atomic size is based on how closely an atom is to it’s neighboring atom
• Because the neighboring atom can vary from one substance to another, the size itself tends to vary
• For sodium, The atomic radius is defined as half the distance between adjacent nuclei in a crystal of an element
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Atomic Radius CONT’D
• Trends within the periods– As you move left-right, atomic size decreases– Caused by the increasing positive charge in a
nucleus– Each successive element increases in
number of electrons and protons– Remain in same principal energy level– The increased nuclear charge pulls the
outermost electrons in closer to the nucleus
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Atomic Radius CONT’D
• Trends within groups– Increase as you move down a group– The nuclear charge increases and electrons are
added to higher principal energy levels– Outer electrons are farther from the nucleus
1
2
3
4 5
6
7
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2. Ionic Radius
• Atoms can gain or lose electrons to form ions• Because electrons are negatively charged the
change in quantity causes there to be a change in the net charge
• ION- an atom or bonded group of atoms that has a positive or negative charge– When atoms lose electrons, they become positive
and, therefore, are smaller– When atoms gain electrons, they become negative
and, therefore, are larger
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• Ionic Radius– Cations (+)
• lose e-
• smaller
© 2002 Prentice-Hall, Inc.
– Anions (–)
• gain e-
• larger
Ionic Radius CONT’D
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Ionic Radius CONT’D
• Trends within periods:– Decrease as you
move left to right
• Trends within groups:– Increase as you move – down a group
1
2
3
4 5
6
7
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3. Ionization Energy
• Defined as the energy required to remove an electron from a gaseous atom
• A high ionization value indicates that the atom has a strong hold on its electrons therefore, tend to not form positive ions
• Trends within periods:– Increase left to right
• Trends within groups:– Decrease down a group
• OCTET RULE:– States that atoms tend to lose or gain electrons in
order to achieve a set of 8 valence electrons
1
2
3
4 5
6
7
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4. Electronegativity
• Indicates the relative ability of atoms to attract electrons in a chemical bond
• Calculated based on many factors and are expresses in terms of a value of 4.0 or less
• Units are called Paulings• Fluorine is the most electronegative with a
value of 3.98 and Francium is the least electronegative with a value of 0.70.
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Electronegativity CONT’D
• The greater the electronegativity, the more strongly it attracts the bond’s electrons
• Trends within periods:– Increases left to right
• Trends within groups:– Decreases down the
group
1
2
3
4 5
6
7