Download Periodic Table - Buford High School Chemistry

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
Newland’s
octaves
Transition Elements (B)
• Consists of only
metals.
• Found in the
center of the
period table.
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
• B, Si, Ge, As, Sb, Te,
Po, At
• Silicon and
Germanium are 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.
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 any more 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
–
–
–
–
elements 57-70
Fits in period 6
Shiny metals
Highly reactive
• Actinide Series
–
–
–
–
elements 89-102
Fits in period 7
Radioactive
The first 4 are
naturally occurring the rest are lab
created
Lanthanides and Actinides
• Electrons fill the f orbitals in an unpredictable
manner and there are many exceptions to
electron configuration rules.
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.
Octet Rule
• Every Element wants 8 valence e-
Formation of Ions
• A positive ion (called a cation) results when an
atom loses electrons.
• A negative ion (called an anion) results when an
atom gains electrons.
Periodic Trends
• Patterns in the periodic table that can be
determined by comparing a period or a
group
• Atomic Radius, Reactivity, Density
Electronegativity
Atomic Radius
• The “size of atom”
• The outer edge of an atom is not clearly
defined; there is no definite edge
• Atomic radius is therefore half the distance
between two identical nuclei
Atomic Radius Trend
• Decreases as you go across the period - as
you go across a period, more protons in
nucleus, greater positive charge. Added
electrons in same principal energy level, so
they get pulled closer to nucleus.
Atomic Radius Trend
• Increases as you go down the group – as you
go down a group, more protons are added
increasing nuclear charge. However added
electrons are in successively higher energy
levels which are further from the nucleus.
• Electrons are NOT pulled as tightly toward the
nucleus.
Atomic Radius Trend
Atomic Radius Trend
Practice
• Which has the largest atomic radius, Mg, Si, S?
• Which has the smallest?
• The atoms below are helium, krypton, and
radon.
• Which one is krypton? How do you know?
Practice
• Determine which element in each pair has the
largest atomic radius.
 period 2, group 1; or period 3 group 18
 period 5, group 2; or period 3 group 16
 period 3, group 14; or period 6 group 15
 period 4, group 18; or period 2 group 1
Electronegativity
 the ability of an atom to attract an electron
while in a chemical bond
 the larger the electronegativity value, the
better an atom is at attracting electrons while
in the chemical bond
Electronegativity
Trend:
 Increases across the period – atoms are
smaller going across a period, outer shell is
closer to being full, and it is closer to the
positive nuclear charge
Electronegativity
• Decreases down the group – atoms are larger
going down a group, outer orbitals further
from positive nuclear charge, not good at
attracting electrons
Electronegativity
Practice
• Which element, Cl or Mg, would have a
higher electronegativity?
• How many more electrons does each
need to be stable?
• Why is the last group not included in the
electronegativity chart?
Ionization Energy
• Energy required to remove an electron
from a neutral atom.
• The greater the ionization energy the
stronger hold an atom has for its electrons
Ionization Energy Trend
• Across a Period (moving from left to
right)
Ionization energy increase from left to right
Ionization Energy
• From Top to bottom (in a group)ionization energy decreases