Transcript CP Chemistry

CP Chemistry
Chapter 11
Patterns in the Periodic Table
Development of Periodic
Table
• 1800 – scientists began to see
patterns in properties of
elements
• 1829
• Dobereiner – German chemist –
first periodic table
• Put elements with similar
properties in groups of three –
called “triads”
• Example of one of his triads:
chlorine, bromine, iodine
• Did not work so well with the
other elements known at that
time
• 1864 Newlands
• Suggested that elements could
be put in groups of eight –
“octaves”
• Did not explain the properties of
all 62 elements
First version of modern
periodic table
• 1869
• Mendeleev
• Arranged according to
increasing atomic mass
• Placed elements with similar
properties in same column
• He used the word “periodic”
to describe his table
• Problems with that setup
• Certain elements didn’t “fit”
• He predicted that elements
would be discovered – left
room
Moseley
• X-ray experiments
• Concept of atomic number
• Rearranged the table according
to increasing atomic number
• Fixed the problems!
Modern Periodic Table
•Elements arranged
according to
increasing ATOMIC
NUMBER
Periodic Law
• States that when the elements
are arranged by increasing
atomic number, there is a
periodic repetition of their
chemical and physical
properties
Arrangement of table
• Groups (columns) or
“families”
• Vertical
• 1-18
• Periods
• Horizontal
• 1-7
• Left of “step” are metals
• Right of “step” are
nonmetals
• On the “step” are
metalloids
“Representative
elements”
•Groups 1,2, and 13-18
•(1A, 2A, 3A-8A)
•All “A” groups
Group “B” Elements
• Transition metals
–Groups 3 - 12
• Inner Transition elements
–Lanthanide and Actinide
series
Metals
• Malleable
• Ductile
• Solids
• *Hg* exception – is liquid
• Good conductors-heat and
electricity
• Lustrous (shiny)
•Tend to lose
electrons when
they bond with
nonmetals
• If start out neutral, and lose
electrons, what charge do they
then have?
• Positive!!!!!!!!!!!!!!!!
• Positive ions called cations
Group 1
Alkali Metals
•VERY REACTIVE
•OCCUR IN NATURE
ONLY IN
COMPOUNDS
Group 2
“Alkaline Earth Metals”
•Also very reactive
•But Group 1 beats it!!
Nonmetals
• Solids and gases
• Solids – brittle, lackluster
• Liquid - bromine
• Poor conductors of heat
electricity
NONMETALS
• Tend to GAIN electrons when
they bond with metals to form a
compound
• If neutral and then GAIN
electrons, what charge does it
have?
• Negative!!!!
• Negative ions called anions
• Share electrons when they
bond with other nonmetals
to form a compound
HALOGENS – Group 17 (7A)
• Very reactive nonmetals
• Occur in nature only in
compounds
• Fluorine – most reactive
nonmetal
Halogens
• Only group on the table that has
elements in all three phases
• F, Cl – gases
• Br – liquid
• I – solid
• Are all diatomic when by themselves
• F2, Cl2, Br2, I2
• Group 18 Noble gases
• Unreactive
• Anyone know the reason
why??
Valence electrons
• Valence electrons and
group number
• Valence electrons and
period number
“Blocks” of the periodic
table
•
•
•
•
s
p
d
f
f block
• Called the “Inner
Transition metals”
• Lanthanides (a.n. 58-71)
• Actinides (a.n. 90-103)
Periodic trends
• Atomic radius
–Increases down a group
–Decrease across a
period – “shielding
effect”
Ionic radius
• Metals – ions are smaller than
the corresponding atom
• Lost electrons
• Ionic radius is smaller that its
atomic radius
Ionic radius - nonmetals
• Ions have gained
electrons
• So, ionic radius is larger
than its atomic radius
Ionization energy
• Amount of energy needed to
remove an electron from an
atom
• Metals – low
• Nonmetals – high
• So, IE increases as you go from
left to right on the periodic table
• IE
• Decreases as you go
down a group
Electronegativity
• The ability of an atom to
attract electrons in a
chemical bond
• F has the highest
electronegativity
Electronegativity
• Increases as you go from
left to right on the
periodic table
• Decreases as you go
down a group (as atomic
number increases)