Transcript Periodic Table - Ms. Lisa Cole-

Periodic Table
Antoine Lavoisier (1700’s)
• Listed all known
elements (33) at the
time
• 4 groups: gases,
metals, nonmetals,
and earths
Dobereiner (early 1800’s)
• Dobereiner arranged the elements into
triads (groups of three elements) based on
similarities in properties.
John Newlands (1864)
• Arranged elements by
increasing atomic
mass (70)
• Noticed a repeating
pattern of properties
• Created the law of
octaves (repeating
patterns at every
eighth element)
Newlands
Lothar Meyer (1869)
• Identified and proved
that there was a
connection between
atomic mass and the
property of the
element
• Arranged the
elements by
increasing atomic
mass (added the new
ones)
Dmitri Mendeleev (1869)
• Proved a connection
between atomic mass
and element
properties
• Arranged elements by
increasing atomic
mass
• Predicted the
existence and
properties of
elements yet to be
discovered
Henry Moseley (1913)
• Discovered atomic
number
• Arranged elements by
increasing atomic
number
• By doing this a
pattern of properties
was discovered and
fixed previous
problems
Periodic Law
• When elements are
arranged by
increasing atomic
number, there is a
periodic repetition of
physical and chemical
properties.
Modern Periodic Table
• Periods (rows)contain a variety of
elements ranging Periods

from metals to
nonmetals to Noble
gases. There are 7.
• Groups or Family
(columns)- contain
elements that share
similar properties.
There are 18.
Representative (Main) Elements
• Marked by “A” on
most groups.
• Elements in the
‘s’ and ‘p’ block
• Wide range of
characteristics
• This is Newlands
octaves
Transition Elements (B)
• Consists of only
metals.
• Found in the
center of the
period table.
• Elements of the
“d” block
Metals, Nonmetals, and Metalloids
Metals
• Make up most of the
periodic table
• Solid at room
temperature (except
Mercury)
• Good conductors of
heat and electricity
• Ductile and malleable
• Have Luster (shiny)
Nonmetals
• Gases or solids at
room temperature
(except Br, it is a
liquid)
• Poor conductors of
heat or electricity
• Brittle
• Dull
Metalloids
• Combination of
characteristics of both
metals and nonmetals
• Silicon and
Germanium both
used in computer
chips
Answer the following questions about iron:
1. In what period is iron found?
2. In what group is iron found?
3. How many protons does an atom of iron have?
4. Write the electron configuration for iron.
5. Write the dot notation for iron.
6. Is iron a representative or transition element?
7. Is iron a metal or nonmetal?
8. List three properties of iron.
9. Describe one way in which classification was
used in biology.
Color Code the Periodic Table
Elements
Color
Hydrogen
Aqua
Group 1: Alkali Metals
Red
Group 2: Alkali Earth Metals
Yellow
Groups 3 -12: Transition Elements
Blue
Group 13: Boron Group
Grey
Group 14: Carbon Group
Peach
Group 15: Nitrogen Group
Brown
Group 16: Oxygen Group
Purple
Group 17: Halogens
Orange
Group 18: Noble Gases
Green
Rare Earth Elements: Lanthanide
Series
Rare Earth Elements: Actinide Series
Pink
Magenta
S- Block
• Alkali Metals
– 1 valence e-. This
makes them highly
reactive
– Exist only as
compounds
– Silvery white in color
– Often bond with
halogens
– Used in salts and
batteries
– Forms ions with a 1+
charge.
• Alkaline Earth Metals
– 2 valence e-. Makes
them highly reactive
– Ca and Mg are
important components
of living cells
– Silvery in color
– Used to make laptop
casings
– Forms ions with a 2+
charge.
S- Block
P- Block (Families 13-18)
• Boron Family (13)
– 3 valence e– Tends to give its
valence e- away
– Most are metals
– Not as reactive as
group one and two
– Forms ions with a 3+
charge
• Carbon Family (14)
– 4 valence e– Can either give away
its valence electrons
or take additional
electrons
– Sn and Pb will form
ions with 4+charges
P- Block
• Nitrogen Family (15)
– 5 valence e-, but will
form 3- (it prefers to
gain 3 e- rather than
give away 5)
– N and P are reactive
and found in many
molecular compounds
• Oxygen Family (16)
– 6 valence e– Forms 2- ions. It
prefers to gain 2 erather than give away
6)
– O and S are reactive
and found in many
compounds
P- Block
• Halogens (17)
– 7 valence e– Form 1- ions (gains 1
e-)
– Highly reactive
nonmetals
– Will often bond with
metals to make salts
• Noble Gases (18)
– 8 valence e-, full p
sublevel
– Does not form ions
– Inert gases
(unreactive)
– They do not bond with
other elements
because they do not
need anymore e-
P- Block
D- Block (3-12)
•
•
•
•
All transition metals
Most are hard metals
All can exist as free elements in nature
Will form a variety of charged ions due to
the fact that the s and d sublevels are
close in energy amounts
D- Block (3-12)
F- Block (Period 6 and 7)
• Lanthanide Series
– Shiny metals
– Highly reactive
– Fits in period 6
• Actinide Series
– Radioactive
– The first 4 are
naturally occurring the
rest are lab created
– Fits in period 7
F- Block
Identify each of the following
elements described below:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
Nonmetal of the second period and group 4A.
The noble gas in period 3.
This element has two more protons than phosphorus.
The only nonmetal in group 1A.
Metal in period 7 with two valence electrons.
The element whose electron configuration ends with 3p1.
The nonreactive element consisting of 4 energy levels.
The metalloid with three valence electrons.
The only noble gas that does not have 8 valence electrons.
The element in group 2A that has fewer energy levels than
magnesium.
Periodic Trends
• Patterns in the periodic table that can be
determined by comparing a period or a
group
• Atomic Radius, Ionic Radius, Ionization
energy, Electronegativity
Atomic Radius
increases
• Size of the atom
• Half the distance
between two identical
nuclei
increases
Atomic Radius Increases:
• Top to Bottom within a group because as
you move down a column, the number of
energy levels is increasing.
• Right to Left because as you move left
across a period the atomic number
decreases which results in a larger atomic
radius
Ionization Energy
increases
• Energy needed to
remove the
outermost electron.
• Follows this trend
because small
atoms have a high
ionization energy
increases
Formation of Ions
• A positive ion (called a cation) results when an
atom loses electrons.
• Metals have low I.E. and therefore, form positive
ions.
• Positive ions are smaller than the parent atom.
• A negative ion (called an anion) results when an
atom gains electrons.
• Nonmetals have high I.E. and therefore, gain
electrons.
• Negative ions are larger than the parent atom.
Octet Rule
• Every Element wants 8 valence e-
Ionic Radius
Increases
• The
formation of
an ion
• When atoms
lose e- they
become
smaller
• When atoms
gain e- they
become
larger
increases
•
Ionization Energy
• Energy
required to
remove an efrom a
gaseous
atom. (Pulling
an e- off)
Electronegativity
• The ability of
an atom to
attract an
electron while
in a chemical
bond