Download The Periodic Table and Periodic Law

Chapter 6
The Periodic Table and
Periodic Law
Vocabulary
alkali metal
metalloid
alkaline metal
noble gas
electronegativity
nonmetal
group
octet rule
halogen
period
inner transition metal periodic law
ion
representative element
ionization energy
transition element
metal
transition metal
Development of the Periodic
Table
• In the 1790’s
Antoine Lavoisier
complied of all
known elements
into four classes
• This listing
contained only 23
elements
• 1800’s during the
industrial revolution,
new developments
such as electricity
and the
spectrometer lead to
the identification
more elements
• By 1870- about 70
elements were known
• Along with their discovery also
came large amounts of
information about the new
elements
• Scientist needed a tool for
organizing this information
• Developing a common method to
calculate atomic mass helped the
beginnings of the organization
process
• 1864, English
chemist, John
Newlands
proposed an
organization
scheme for the
elements by
arranging them
by increasing
atomic mass and
using the Law of
Octaves

• Law of Octaves- states that
chemically similar elements
occur every eight element much
like the octaves of a musical
scale
Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
K
Ca
Cr
Ti
Mn
Fe
Co, Ni
Cu
Zn
Y
In
As
Se
Br
Rb
Sr
La,Ce
Zr
Nb, Mo
Ru, Rh
Pd
Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
K
Ca
Cr
Ti
Mn
Fe
Co, Ni
Cu
Zn
Y
In
As
Se
Br
Rb
Sr
La,Ce
Zr
Nb, Mo
Ru, Rh
Pd
• This model had some
problems, it left out some
elements and did not leave
room for undiscovered
elements
• German chemist,
Lothar Meyer and
Russian chemist,
Dmitri Mendeleev
each
demonstrated
a connection
between
atomic mass
and elemental
properties
• Mendeleev’s published table
arranged the elements in order of
increasing atomic mass and into
columns with similar properties
• Mendeleev
left blanks in
his table for
possible
undiscovered
elements
• Using his table, he was able to
predict not only the existence but
also the properties of these
undiscovered element
• There were some problems
with Mendeleev’s table
• Several elements in the table
were not in the correct order,
their properties did not
match the properties of the
other elements in the same
group
• English chemist,
Henry Moseley
discover the problem
when he determined
that each element
contained a unique
number of protons
(atomic number) in
their nuclei
• By rearranging the elements
in order of increasing atomic
number the problems were
solved
• Periodic law states that there is
a periodic repetition of chemical
and physical properties of the
elements when they are arranged
by increasing atomic number
Modern Periodic Table
• Consists of Oxygen
boxes each
8
containing
information
O
about each
15.999
element
Element
Atomic
number
Symbol
Atomic
mass
• The element boxes are arranged
in order of increasing atomic
number into a series of rows and
columns
• The rows are called periods
• There are seven periods on the
periodic table
1
2
3
4
5
6
7
• The columns are called groups or
families
• The groups of the period
table are numbered using two
different methods
– They can simply be numbered 118
– They can be divided in groups A
and B with the numbers 1-8
1
One way is to just number 18
1-18.
2
13 1415 1617
3 4 5 6 7 8 9 10 11 12
1A
Another way is to use A &
B
groupings
and
numerals.
2A
3A 4A 5A 6A 7A
3B 4B 5B 6B 7B
8B
1B
2B
8A
• Groups 1,2,13,-18 are called the
main group or representative
elements
• They are called this because the
elements in have a wide range of
chemical and physical properties
A
A
• Groups 3 – 12 are the
transition elements
B
• Three main classifications
of elements
Metals
Metalloids
nonmetals
Hydrogen
• most common
element in universe
• reacts rapidly
• rarely in free state
• common in
compounds
• Can gain or lose 1
electon
• Metals – elements that are
shiny, solids, good
conductors of heat and
electricity, ductile and
malleable
• Most of the elements on
the periodic table are
metals
• Groups 1 and 2 are the alkali
metals and the alkaline earth
metals, respectively, are very
chemically reactive
2
Alkaline earth metals
Alkali metals
1
Group 1: Alkali Metals
• High reactivity—they are the MOST
reactive metals
• Explosive in water
• Very soft metals
• React with oxygen in the air
• Have 1 valence electron
• Lose 1 electron to become stable and then
have a
• +1 charge
• Excellent conductors of electricity
• Not found free
Group 2: Alkaline-Earth Metals
• Reactive, but not as reactive as the
Gr. 1A
• Harder and stronger than Gr. 1A
• Have 2 valence electrons
• Lose 2 electrons and then have a
• +2 charge
• Not found free
• All of the B groups are metals
• They consists of two sets the
transition metals and the inner
transition metals
Transition
metals
Inner transition metals
Transition Metals
• Found between Groups 2
and 13
• Valence electrons vary
• They all LOSE electrons
• Some common transition
metals are
copper, silver, gold, zinc,
• Nonmetals are elements that are
generally gases or brittle, dulllooking solids which are poor
conductors of heat and electricity
• Bromine is the only nonmetal that is
a liquid a room temperature
Br
• Group 7A, Halogens, is a group of
highly reactive nonmetals
• Group 8A, the Noble Gases, is an
extremely unreactive group of
elements
8A
7A
Groups 13: Boron Family
The METALS of this group (all of them
except for boron)
• Have 3 valence electrons
• Usually lose 3 electrons
• +3 charge
Group 14: Carbon Family
• Have 4 valence electrons
• Carbon, silicon, and germanium share
electrons – do not gain or lose electrons
Group 15: Nitrogen Family
• Have 5 valence electrons
• GAIN 3 electrons and have a
• -3 charge
Group 16: Oxygen Family
(The Chalcogens)
•
•
•
•
Have 6 valence electrons
GAIN 2 electrons to have a
-2 charge
Also called the chalcogens
Group 17: The Halogens
• Highly reactive -- they are the
most reactive nonmetals.
• Usually combine with most
metals to form salts (“Halogen”
is Greek for ‘salt former’.)
• Have 7 valence electrons
• GAIN 1 electron to have a
• -1 charge
Group 18: Noble Gases
• Do not react with other elements (don’t gain
or lose electrons)
• 8 valence electrons
• All are gases
• Metalloids, or semimetals, are the
elements that border the stair-step
line
• Metalloids are elements with
physical and chemical properties of
both metals and non-metals
Rare Earth Metals
• Lanthanide Series: (also called the Lanthanoid
series) -- atomic #58 – 71
• They are shiny, reactive metals that are often used
to make alloys.
• Actinide Series: (also called the Actinoid series)
-- atomic #90 – 103
• Have unstable arrangements or protons and
neutrons
• All are radioactive and most are man-made
Organizing the elements by
electron configuration
• Atoms in the same group have
similar chemical properties because
they have the same number of
valence electrons
• Group IA electron configurations
– H 1s1
All have s1
– Li 1s2 2s1
valence
– Na 1s2 2s2 2p6 3s1
electron
– K 1s2 2s2 2p6 3s2 3p6 4s1
• The energy level of an
element’s valence electrons
indicates the element’s period
– H 1s1
found in period 1
– Li 1s2 2s1
found in period 2
– Na 1s2 2s2 2p6 3s1 … period 3
– K 1s2 2s2 2p6 3s2 3p6 4s1…
period 4
• In the A groups the number of
valence electrons and the groups
number are related
• In group 1A, elements have 1 valence
electron
• Group 2A- 2 valence electrons
• Group 3A- 3 valence electrons
• Group 4A- 4 valence electrons
• … Group 8A- 8 valence electron (except
He has only 2 electrons)
• The periodic table is divided into
sections or blocks which reflect the
sublevel being filled by valence
electrons
s-block
f-block
d-block
p-block
s-block elements
• can hold a maximum of 2 electrons
• the s-block portion of the periodic
table spans two groups 1A and 2A
and
includes
hydrogen
and
helium
H
1A 2A
He
p-block
• can hold a maximum of 6 electrons
• The p-block portion of the periodic
table spans 6 groups8A
Groups 3A-8A.
3A 4A 5A 6A 7A
p-block
• The noble gases found in group
8A are unique in that they are
very stable.
• This stability is due to having
completely filled s and p orbitals 8A
d-block
• can hold a maximum of 10 electrons
• d-block spans over 10 groupsGroups 1B-8B
– these groups are called the transition
elements
8B
3B 4B 5B 6B 7B
Transition
elements
1B 2B
d-block
• The d-block orbitals are one energy
level less than the previously filled
s-block orbitals
Sc- [Ar] 4s23d1
Tc- [Kr] 5s24d5
8B
3B 4B 5B 6B 7B
1B 2B
Sc
Tc
Transition elements
f-block
• can hold a maximum of 14 electrons
• The f-block includes the
inner transition elements
Inner transition elements
Periodic trends
• Atomic radius
– Is half the
distance between
the nuclei of
adjacent atoms in
either a crystal
or molecule of
the element
• Within a period the atomic radii
generally decreases
– This is due to the increasing positive
charge of the nucleus while electrons
are added to the same energy level
decreases
• Down a group the atomic radii
generally increases
– This is due to the addition of
principal energy levels
– The inner electrons shield the outer
electrons from the attraction of the
nucleus
• Ionic radius
– Ions form when an
atom loses (cation)
or gains electron(s)
(anion)
– Positive ions lose electrons becoming
smaller
– Negative ions gain electrons becoming
larger
• Within a period positive ions
gradually decrease until Group 5A
or 6A negative ion greatly
increase then gradually decrease
decreases
5A
Increases then
decreases
• Down a group ionic radius generally
increase in both positive and
negative ions
Ionization energy
• the energy required to
remove an electron from a
gaseous atom
• an indication of how strongly
an atom’s nucleus holds onto
its electrons
– A high ionization energy
indicates an atoms has a strong
hold on its electrons
– These atom are less likely to
form positive ions
– Atoms with low ionization
energy values lose electrons
readily and form positive ions
• Within a period ionization
energy increases from left to
right
increases
• Down a group ionization
decreases
• Octet rule– States that atoms will gain, lose, or
share electrons in order to acquire a
full set of eight valence electrons
– This fills the s and p orbitals so that
resemble the electron configuration
of a noble gas
– Metals tend to lose electrons while
nonmetals tend to gain or share
electrons to follow the octet rule
• Electronegativity
– Indicates the relative ability of
atoms to attract electrons in a
chemical bond
– The unit of electronegativity is
call the Pauling.
– The greater electronegativity
value the more strongly the
atom attracts bonding
electrons
• Within a period electronegativity
generally increases from left to
right
increases
• Down a group electronegativity
generally decreases
Summary of Periodic Table Trends
Moving Left --> Right
•Atomic Radius Decreases
•Ionic Radius Decreases large increase
in Group 5A or 6A then decreases
•Ionization Energy Increases
•Electronegativity Increases
Moving Top --> Bottom
•Atomic Radius Increases
•Ionic Radius Increases
•Ionization Energy Decreases
•Electronegativity Decreases