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Transcript
The Periodic
Table
Chapter
6
Introduction
• The periodic table represents an
organizing principle which allows the
prediction of the properties of each
element based on their position in the
periodic table.
• The elements of the periodic table can be
classified into different categories.
• There are trends that appear in the
periodic table that allow us to make
predictions about an atoms size, ionization
energy, and electronegativity.
Organizing Elements
(Section 6.1)
• Searching for
an Organizing
Principle
• Mendeleev’s
Periodic Table
• The Periodic
Law
• Metals,
Nonmetals, and
Metalloids
Defining the Periodic Table
Periodic Table: An arrangement of elements
in which the elements are separated into
groups based on a set of repeating
properties.
· A periodic table allows you to easily
compare the properties of one element
(or group of elements) to another
element (or group of elements).
I.) Searching for an Organizing
Principle
• Copper, gold, and silver have been known for
thousands of years.
• Only 13 elements were known of by the year
1700.
• Between 1765-1775 five more elements were
discovered.
• Problem was how to ID new elements and how
many new ones there actually were.
The Dobereiner System
• Organized known
elements into triades.
• Recognized a
relationship between
atomic weights and
chemical properties
• Triade: A set of three
elements with similar
properties (ex. Cl, Br, I)
• Not all the known
elements could be
grouped into triades
J.W. Dobereiner
German Chemist
1780-1849
The Dobereiner Triades
1 element in each
trade
tended to have
properties
with values that fell
midway between the
other two.
Here we see atomic
and mass numbers
II.) Mendeleev’s Periodic Table
• From 1829-1869 many
other systems were
proposed but none
gained wide
acceptance.
• Mendeleev created his
table while working on
a text book for his
students
• Beat a competitor
because he was better
able to explain the
table’s usefulness.
Dmitri Mendeleev
Russian Chemist/Teacher
1834-1907
• Arranged the elements
in his table in order of
increasing atomic mass.
• There was a close
match between the
predicted properties of
unknown elements and
the actual properties of
the elements.
• This organizational
method had its
problems and does not
account for all elements
(ex. atomic masses for
Te and I).
III.) The Periodic Law
• There was a problem
with organizing
elements by atomic
masses.
• Organizing elements
by increasing atomic
number was more
useful.
• Moseley determined
the atomic number for
the known elements
at the time
Henry Moseley
British Physicist
1887-1915
Expression of the Periodic Law
When elements are arranged
in order of increasing atomic
number, there is a periodic
repetition of their physical
and chemical properties.
The Modern Periodic Table
Horizontal Rows = Periods
Vertical columns = Groups
Periods in the Periodic Table
• These are 7 rows extending horizontally across
the periodic table.
– Period 1 = 2 elements
– Period 2/3 = 8 elements
– Period 4/5 = 18 elements
– Period 6/7 = 32 elements
• Each period corresponds to a principle energy
level.
– More elements in the higher periods because there are
more orbitals in the higher energy levels
• Properties of elements within elements within a
period change as you move from left to right.
Groups in a Periodic Table
• These are the 18 columns that run up
and down the periodic table.
• There are 3 different ways that the
groups are numbered.
• These groups also possess names.
• Elements within a group in the
periodic table have similar physical
and chemical properties.
IV.) Metals, Nonmetals, &
Metalloids
•
•
We saw how the periodic table can be divided
into 7 periods and 18 groups.
We can also divide the table into three broad
classes based on the general properties of the
elements.
1. Metals
2. Nonmetals
3. Metalloids
•
Across the periods, the properties of elements
become less metallic and more nonmetallic.
The Three Classes of Elements
Metals
• This is the most numerous class.
• Characteristics of metals:
– Good conductors of heat and electric current
– Possess luster and sheen
– Solids at room temperature (except Hg)
– Ductile (i.e. can be drawn into wires)
– Most are malleable (i.e. can be hammered
into thin sheets)
Examples of Metals
Copper is ductile.
Aluminum is malleable and has
luster and sheen.
Nonmetals
• Less numerous than the metals
• There is greater variation in the characteristic of
nonmetals.
– One general characteristic: They are not metals.
– Poor conductors of heat and electric current
(carbon is an exception)
– Solids tend to be brittle.
• Most nonmetals are gases at room
temperature, a few are solids, and 1 is a liquid
(Br).
Examples of Nonmetals
Chlorine is a gas
Bromine is a liquid
Carbon is a solid and is
a good conductor of
electricity.
Metalloids
• Least numerous elements
• Have properties that are similar to those of
metals and nonmetals.
• The behavior is often controlled by changing
the conditions.
Examples of Metalloids
Silicon is not a good
conductor of electricity
until mixed with boron.
Arsenic has luster and
sheen like metals.
Classifying Elements
(Section 6.2)
• Squares in the
Periodic Table
• Electron
Configuration in
Groups
• Transition
Elements
I.) Squares in the Periodic Table
• All periodic tables
display at least the
symbol, the atomic
number, and the mass
number of the elements.
• Some periodic tables
provide more information
such as physical state,
electron configuration,
and classification of
each element.
Alkali
Metals
Noble Gases
Alkaline
Earth
Metals
Halogens
II.) Electron Configuration
Groups
• Elements can be sorted into
separate groups based on their
electron configuration
–Noble Gases
–Representative Elements
–Transition Metals
–Inner Transition Metals
The Noble Gases
• These are the elements located in Group
18 (the farthest column to the right)
• These are the “inert” gases because they
rarely participate in reactions.
• The “s” and “p” orbitals of the highest
occupied energy levels are filled for all
noble gases.
Let’s show this for helium, neon, argon, and
krypton by writing out the electron configuration
for each.
1. Helium
2. Neon
3. Argon
4. Krypton
The Representative Elements
• These elements display a wide range of
physical and chemical properties.
• The atoms of the representative elements
have “s” and “p” orbitals of the highest
occupied energy levels that are not full.
• For any element of this group, the group
number (the American and European
numbering system) equals the number of
electrons in the highest occupied energy level.
Let’s Look at the Electron
Configuration of Some Representative
Elements
Lithium
Sodium
Carbon
Silicon
Transition Elements
• Two kinds of transition metals: transition
and inner transition metals – classification
is based on the electron configuration of
an element.
• Transition Metals: Atoms have highest
occupied sublevels that have electrons in
the “s” and “d” orbitals.
• Inner Transition Metals: Atoms of these
metals have the highest occupied “s”
orbital and nearby “f” orbitals that contain
electrons.
Let’s Look at the Electron
Configuration of Some Transition
Elements
Iron
Silver
Nickel
Chromium
Let’s Look at the Electron
Configuration of Some Inner
Transition Elements
Cerium
Uranium
The Divisions Based on Electron
Configuration
The electron configuration and
the position of an element in
the periodic table gives a
particular pattern. This pattern
are the blocks we see here.
Using the periodic table to write
electron configurations.
• Based upon the blocks that were described in
the previous slide, we can write the electron
configuration for any element based on its
location in the periodic table.
• The steps:
1.
2.
3.
4.
5.
6.
7.
Find the element on the periodic table.
Start counting from hydrogen.
Move towards the right
At the end of each row drop down one row
Begin counting towards the right again
Each row represents an energy level.
Each square represents an electron.
Noble Gas Configuration
This is a shorter way to write the shorthand
notation for electron configurations.
There are four easy steps:
1.) Locate the element in the periodic table.
2.) Find the noble gas that precedes it.
3.) Place the symbol for this gas in brackets ([ ]).
4.) Write the remaining electron configuration.
Write the noble gas configuration for rubidium (Rb).
Let’s try this. Write the extended
and Noble gas configuration for the
following elements.
1.) chlorine
2.) lead
Section 6.3
Periodic Trends
• Trends in Atomic
Size
• Ions
• Trends in
Ionization Energy
• Trends in Ionic
Size
• Trends in
Electronegativity
Atomic Size
Atomic radius: One half the distance
between the nuclei of two atoms of the same
element when the atoms are joined.
Distance between Nuclei
Atomic Radius
Ionization Energy
The energy required to remove an electron
from an atom.
Ions: An atom or a group of atoms that has a
positive or negative charge resulting from a
loss or gain of electrons, respectively.
Cation: An ion with a positive charge
Anion: An ion with a negative charge
Electronegativity
The ability of an atom of an element to
attract electrons when the atom is in a
compound.
This property is related to chemical bonding
and is best understood by examining chemical
bonds.
I.) Periodic Trend in Atomic
Radius
In general,
atomic
size
increases
from top to
bottom
within a
group and
increases
from right
to left
across
a period.
These trends depend upon the number of protons
and electrons being added as we move through
the periodic table.
Question: Based on the data for alkali metals and
halogens, how does the atomic size change within
a group? Why?
Within a group:
• Increasing # of
occupied orbitals.
• Shielding of outer
electrons
increases
the atomic radius.
Within a period:
• Electrons are being added
to the same energy level.
• Shielding is constant
• Increasing nuclear charge
pulls all electrons closer
II.) Ions
An atom or a group of atoms that has a
positive or negative charge resulting from a
loss or gain of electrons, respectively.
Cations
• These are positively charged ions,
resulting from a loss of an electron.
• Metals tends to lose electrons from
their highest occupied energy levels
to become cations.
• The charge for a cation is written as a
number followed by a plus sign.
 Na1+
 Ca2+
Representing the Formation of
Cations
• Atoms tend to lose their electrons
from the outer most energy levels
to become cations.
• Na → Na+ + e• Ca → Ca2+ + 2e• Why would the atoms lose their
electrons to become cations?
Anions
• These are negatively charged ions
resulting from a gain of electrons.
• Nonmetals tend to add electrons into
their highest occupied energy levels
to become anions.
• The charge for an anion is written with
a number followed by a negative sign.
–Cl1–F1-
Representing the Formation of
Anions
• Atoms tend to accept electrons
into their highest occupied energy
levels to become anions.
• e- + Cl → Cl
• e- + F → F
• Why would these atoms accept
electrons to become anions?
Ionic Compounds
Fe2(SO4)3
FeS2
Fe2O3
III.) Trends in Ionization Energy
The energy required to remove an
electron from an atom.
• This energy is measured when an element
is in the gaseous state.
• There are ionization energies for every
electron in an atom
• The 1st ionization energy is the energy
needed to remove the first electron from an
atom.
• Each successive ionization energy
increases dramatically.
Periodic Trend in 1st Ionization Energy
Group Trend:
As size of the atom increases
the nuclear charge has a
smaller effect on the outer most
electrons.
Period Trend:
Nuclear charge increases
as we move across a
period, thus nuclear
attraction increases.
III.) Trends in Ionic Size
IV.) Electronegativity
The ability of an atom of an element to
attract electrons when the atom is in a
compound.
In general, electronegativity values increase
from bottom to top within a group. For
representative elements, the values tend to
increase from left to right across a period.
The Periodic
Table
Chapter
6
The
End