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Transcript
The Periodic Law
Chapter 5
Sophia Nolas, Yen Dinh, Chris Fleming, Jane Smigiel
SECTION ONE (pg 123)
History of the Periodic Table
• Dmitri Mendeleev
He created a periodic table of the elements,
grouping them with other elements with
similar properties.
• Henry Moseley
His observations on his data led to both the
modern definition of atomic number and how
it is the basis for the organization of the
periodic table.
• Periodic Law: The physical and chemical
properties of the elements are periodic
functions of their atomic numbers.
• Periodic Table: An arrangement of the
elements where similar properties fall in the
same column, or group.
• Lanthanides: Are the 14 elements with
atomic numbers from 58 to 71. These
elements are so similar in chemical and
physical properties that it was a tedious task
for many chemists to identify them.
• Actinides: Are the 14 elements with atomic
numbers from 90 to 103.
Lanthanides and Actinides sit below the main
portion of the periodic table to save space.
SECTION TWO (pg 128)
Electron Configuration and the
Periodic Table
S sublevel = 2 electrons
P sublevel = 6 electrons
D sublevel = 10 electrons
F sublevel = 14 electrons
Electron Configuration:
• Carbon (C) = 1s2 2s 2p2
• Hydrogen (H) = 1s
• Copper (Cu) = 1s2 2s2 2p⁶ 3s2 3p⁶ 4s2 3d⁹
• Coefficient = the period where element is
located, (EXCEPTION = d orbitals are off by -1.)
• Letter = what sublevel
• Exponent = how many elements in the period
and sublevel to reach “goal” element (group
number).
Orbital sublevels (the letters in the
electron configuration)
Electron configuration explanation
http://www.khanacademy.org/video/more-onorbitals-and-electronconfiguration?playlist=Chemistry
A simpler way to show electron configuration:
• Take previous noble gas and put it in [ ].
• Then find the electron configuration from
there.
Example: Silver – Ag
Previous noble gas = Kr
[Kr] 5s2 4d⁹
Alkali Metals – Group 1
•
•
•
•
•
•
•
Group 1 of the periodic table
One electron in S sublevel
Do not occur in nature as free elements
Good conductors with low melting points
Ductile and malleable (can be cut with a knife)
Very reactive (with water to make hydrogen gas)
Usually held in kerosene due to the high reacting
level
• Used in lights, electricity technology
Alkaline-Earth Metals – Group 2
•
•
•
•
Group 2 of the periodic table
Contain electron pair in S sublevel
Harder and denser than alkali metals (AM)
Too reactive to be found in nature, but not as
reactive as AM
Halogens – Group 17
• Most reactive nonmetals
• React vigorously with metals
Questions!
• Name the group, period and block of [Xe]6s2
• Write the electron configuration of Group 1
element in the 3rd period
• Which element is more reactive?
Answers!
• Group = 2nd Period = 6th Block = s
• [Ne] 3s
• The first one
SAMPLE PROBLEMS
FOR MORE PRACTICE, GO TO PAGES 133, 136, 138, AND 139 IN THE
TEXTBOOK
• Give the group, period, and block for these elements (without periodic
table), and then determine the name of the element (with the periodic
table)
[Ne] 3s²3p1
[He] 2s1
[Ar] 4s2 3d9
• Write the outer electron configuration of these elements (without the
periodic table) and then identify the element (with the periodic table)
Group 2, Period 6
Group 7, Period 4
Group 14, Period 3
SECTION THREE (pg 140)
Electron Configuration and Periodic
Properties
• The way in which the elements arranged
on the periodic table corresponds to not
only their atomic number but also
electron configuration
– This creates periodic trends in the
properties of these elements
Atomic Radii
• One-half the distance between the nuclei of
identical atoms that are bonded together
• Trend: Increases from top to bottom and
from left to right (Towards francium)
– Due to the decrease of positive and negative
charges from less protons and electrons and
the attractive forces between them
– The electrons are pulled closer to the nucleus
from right to left & from bottom to bottom on
the periodic table (which is why the A.R.
decreases in these direction)
ATOMIC RADII CHART
Ionic Radius
• One half the diameter of an ion in an ionic
compound
– Cation: positive ions that have lost an electron,
thus are smaller than neutral atoms
– Anion: negative ions that have gained an electron,
thus are larger than neutral atoms
• Trend: Increases from top to bottom and
from right to left
Electronegativity
• How readily an atom will attract an electron(s)
in a chemical compound in order to achieve
noble gas configuration
– Combination of Ionization energy &n Electron
affinity
• Trend: Increases from bottom to top and from
left to right
ELECTRONEGATIVITY Chart
(Key: more red=more electronegative)
Ionization Energy
• The energy required in order to remove and
electron from an atom to create an ion
– A + energy
(A+) & e-
• Increases from bottom to top and from left to
right
– As you move down the groups, the more outer
shell electrons you have, creating a shielding
effect on the outermost electrons, making it easier
to remove Electrons from the atom
Electron Affinity
• The energy change that occurs when and
electron is added to a neutral atom
– Can be endothermic or exothermic
• Increases from bottom to top and from left to
right
– Second electron affinity is always endothermic: it
requires energy to force electron to bond to the
atom
SAMPLE PROBLEMS
FOR MORE PRACTICE GO TO PAGES 142, 146, AND 152
• List the following elements from biggest to smallest atomic
radius
– He, Fr, Zn, Ge, C
• List the following from highest to lowest electronegativity
– V, Os, Cd, Sn, Mg
• List the following from highest to lowest ionization energy
– H, Al, P, Sr, Re
• List the following from highest to lowest electron affinity
– O, As, Li, Au, Rb
FOR ADDITIONAL PRACTICE…
Chapter Summery + Review pg 155
Periodic Table
pg 130