PERIODIC TABLE: DEVELOPMENT OF
THE PERIODIC TABLE
• In the late 1790s Lavoisier developed a list of the 23
known elements.
DEVELOPMENT OF THE PERIODIC
TABLE
• During the 1800s many more elements were discovered.
• Scientists used electricity to separate compounds.
• By 1870 there were 90 known elements.
DEVELOPMENT OF THE PERIODIC
TABLE
• Chemists were overwhelmed by the huge volume of
information.
• A new way of categorizing the information was needed.
• A breakthrough came in 1860 when chemists agreed
upon a method of accurately determining the atomic
masses of the elements.
DEVELOPMENT OF THE PERIODIC
TABLE
• The first attempt was made by John Newlands.
• He noticed the properties of the elements repeated every
eighth element.
• He called this the Law of Octaves.
DEVELOPMENT OF THE PERIODIC
TABLE
DEVELOPMENT OF THE PERIODIC
TABLE
• This was not widely accepted because it did not work for
all the elements.
• But, Newlands was correct in that the properties of
elements do repeat in a periodic way.
DEVELOPMENT OF THE PERIODIC
TABLE
• Next came Dmitri Mendeleev
DEVELOPMENT OF THE PERIODIC
TABLE
• Dmitri Mendeleev also arranged the elements by
increasing atomic mass.
• He did not limit the length of his rows
• He noticed the elements fell into columns with similar
properties.
DEVELOPMENT OF THE PERIODIC
TABLE
• Mendeleev also left empty spaces to account for
elements that had not yet been discovered.
DEVELOPMENT OF THE PERIODIC
TABLE
• There were problems with Mendeleev’s table because
new elements were discovered and atomic masses
better determined that placed them out of order.
• Henry Moseley had discovered that the atoms of each
element contain a unique number of protons in their
nuclei.
• He arranged the elements in order of increasing atomic
number – this adjustment fixed the periodic table.
DEVELOPMENT OF THE PERIODIC
TABLE
• The modern periodic table is based on the periodic law
which states that the physical and chemical properties of
the elements tend to change with increasing atomic
number in a periodic way.
MODERN PERIODIC TABLE
• Each element has its own box containing information
about the element.
• The elements are arranged in order of increasing atomic
number into a series of columns called groups, or
families, and rows called periods.
REGIONS OF THE PERIODIC TABLE
• There are five major regions of the periodic table:
metals
nonmetals
metalloids
noble gases
hydrogen
REGIONS OF THE PERIODIC TABLE
• The metals are the largest region of the periodic table.
Elements within this region are solid, generally lustrous,
ductile, and malleable. They are also excellent
conductors of heat and electricity.
REGIONS OF THE PERIODIC TABLE
• Nonmetals make up the second largest region of the
table. These elements have a variety of physical states
and properties. In general, they are poor conductors of
heat and electricity, and are brittle and non-lustrous.
REGIONS OF THE PERIODIC TABLE
• Metalloids are elements that have properties of both
metals and nonmetals. These elements are sandwiched
in between metal and nonmetal regions along the “stair
step”.
METALLOIDS
• Boron (B) – used in fiberglass, glassware, ceramics,
polymers, detergents, insecticides
• Silicon (Si) – semiconductors
• Germanium (Ge) – semiconductors
• Arsenic (As) – wood preservative, insecticide,
semiconductors
• Antimony (Sb) – batteries, low friction metals, flameproofing, ceramics, paints, glass, pottery
• Tellurium (Te) – blasting caps, added to copper and
stainless steel to improve machinability, reduce
corrosivity of lead
• Astatine (At) – radioactive tracer in cancer treatment
REGIONS OF THE PERIODIC TABLE
• The metalloids must be memorized since the periodic
tables we use are not color-coded.
• The noble gases are gases located in Group 18. These
gases do not normally react with anything.
FAMILY CHARACTERISTICS
Group 18 – Noble Gases
- Once called the inert gases
- No compounds of He, Ne or Ar
- Compounds of Xe, Kr and Rn
- Have full orbitals in the highest energy level; this is
called an octet.
- Their electron configuration is very stable
- Atoms from the other 17 groups either gain or lose
electrons to achieve a noble gas electron configuration
Group 1: Alkali Metals
• Have metallic properties – soft, shiny, highly reactive,
can be cut with a knife, react with oxygen, good
conductors of heat and electricity
• Single electron in the highest energy level
• By losing this electron the alkali metal achieves a noble
gas electron configuration and a positive charge of +1.
Group 2: Alkaline Earth Metals
• Also have metallic properties – harder, stronger, denser
than Group 1 metals.
• Less reactive than Group 1 metals.
• Need to lose 2 electrons to achieve a noble gas electron
configuration. Results in a positive charge of +2.
Groups 3 – 12: Transition Elements
• NOT A FAMILY – just the name of a section of the
periodic table.
• These are metals, but they are not as reactive as Group
1 and Group 2 metals.
• Electron configurations are unusual and sometimes do
not come out as predicted by the diagonal rule.
• Will lose electrons from both the s-orbitals and the dorbitals to form ions with various positive charges.
Bottom Rows: Inner Transition Elements
• NOT A FAMILY – just the name of a section of the
periodic table.
• Each row has its own name: Lanthanides and Actinides
• The lanthanides are reactive, shiny metals.
• The actinides are usually radioactive due to an unstable
arrangement of their protons and neutrons.
Groups 13 – 18: Main Block Elements
• NOT A FAMILY – just the name of a section of the
periodic table. These are also known as the
representative elements.
• They represent a wide variety of chemical and physical
properties.
• The metals will lose electrons and form positive ions
while the nonmetals will gain electrons to form negative
ions.
• The metalloids can both lose and gain electrons.
Group 17: Halogens
• The halogens combine easily with metals, especially the
alkali metals, to form compounds known as salts.
• The halogens are the most reactive nonmetals.
• Their electron configuration is one electron short of being
a noble gas electron configuration.
• Will gain an electron to form an ion with a charge of -1.
Hydrogen
• Most common element in the universe.
• It has only one electron and it reacts very rapidly with
other elements – behaving as a metal.
• It can gain an electron and attain a charge of -1.
EXAMPLES
• For each of the following elements, list the region of the
periodic table from which it comes and list the name of
its family (if applicable):
• Ba
• Metal, alkaline earth
• Sb
• Metalloid
• I
• Nonmetal, halogen
• Rn
• Noble gas, noble gas
EXAMPLES
• For each of the given elements, list two other elements
with similar chemical properties.
• Fe
• Ru, Hs, Os
• Se
• S, O
• Br
• I, Cl, F
• Rb
• K, Na
CLASSIFICATION OF THE ELEMENTS
Organizing by Electron Configuration
• Elements within the same group have the same electron
configuration in their outermost energy level.
• These electrons are called valence electrons.
• Atoms in the same group have similar chemical
properties because they have the same number and
arrangement of valence electrons.
Organizing by Electron Configuration
• The energy level of an element’s valence electron
indicates the period, or row, on the periodic table where
it can be found.
The s-, p-, d-, and f- Block Elements
• The s-block consists of Groups 1 & 2 along with
hydrogen and helium. In this block, the valence
electrons occupy only the “s” orbitals.
• The p-block elements consist of Groups 13 – 18. As one
progresses from left to right, one more electron is added
to the “p” orbital until it is filled with six electrons.
• The noble gases exhibit great stability because both their
“s” and “p” orbitals are filled.
The s-, p-, d-, and f- Block Elements
• The d-block contains the Transition Elements and it is
the largest of the electron blocks.
• These elements are characterized by a filled outermost
“s” orbital and either partially or completely filled “d”
orbitals.
• The f-block elements are found in both the lanthanide
and actinide series.
EXAMPLE
• Strontium has an abbreviated, noble gas, electron
configuration of [Kr] 5s2. Without using the periodic
table, determine:
• Group
• 2
• Period
• 5
• Block
• s
EXAMPLE
• Write the abbreviated, noble gas, electron configuration
of the element in Group 12, Period 4.
• Zn: [Ar] 3d104s2
PERIODIC TRENDS
• The periodic law states that the physical and chemical
properties of the elements are periodic functions of their
atomic numbers.
• Because the elements are arranged side by side in order
of increasing atomic number, one can view certain
important vertical and horizontal trends.
Atomic Radius
• Recall from Rutherford’s experiment that the nucleus
was found to occupy only a small fraction of the atom’s
entire volume.
• It is the electron cloud surrounding the nucleus that
determines the boundaries of the atom.
Atomic Radius
• Because the electrons travel in this “cloud region”, it is
difficult to measure the size of an atom.
• The atomic radius is determined in two primary ways.
One is to measure the distance between centers of like
atoms in a diatomic molecule. The other is to measure
the bond lengths of atoms in compounds.
Atomic Radius
• Looking at the periodic table one can notice things that
affect the size of an atom.
• As one moves down through a group a new principal
energy level is added. Each new level is physically
farther from the nucleus.
• As each energy level is added, the energy levels and
electrons closer to the nucleus shield electrons in the
new energy level from the pull of the nucleus. This
shielding effect allows the outer electrons to drift farther
away from the nucleus.
Atomic Radius
• When crossing a period, row, from left to right each atom
gains both a proton and an electron.
• But, no new principal energy level is added so the new
electrons enter the same energy level as all the others.
• There is no shielding.
• The additional protons in the nucleus provide more pull
on the electron cloud and bring it closer to the nucleus.
Atomic Radius
• THE TREND:
Atomic radius increases from top to bottom in a group
and decreases from left to right across a period.
Atomic Radius
Ionic Radius
• Atoms can gain or lose electrons to form charged
particles.
• These particles are called ions.
• When atoms lose an electron they form a positively
charged ion and become smaller in size.
• Positive ions are smaller than their parent atom.
Ionic Radius
• When atoms gain an electron they form a negatively
charged ion and become larger in size.
• Negative ions are larger than their parent atom.
Ionization Energy
• To form a positive ion, an electron must be removed from
a neutral atom.
• The electron is removed from the outermost energy level
of the atom, known as the valence shell.
• The electrons in that “shell” are known as valence
electrons.
• This requires energy to overcome the attraction between
the positive nucleus and the negatively charged electron.
Ionization Energy
• This energy is known as the ionization energy.
• Ionization energy is defined as the energy required to
remove an electron from a gaseous atom.
• The energy required to remove the first electron from an
atom is called the first ionization energy.
• The process is:
• ATOM + IONIZATION ENERGY → ION+ + e-
Ionization Energy
• As one moves down a group (column) it becomes easier
to remove electrons because they are farther from the
nucleus and the shielding effect increases.
• As one moves left to right across a period (row) it
becomes more difficult to remove electrons because
there is no increase in the shielding effect, the positive
charge in the nucleus increases, and the electrons are
closer to the nucleus.
• Trend:
• Ionization energy decreases from top to bottom down a
group and increases from left to right across a period.
Ionization Energy
Ionization Energy
Electronegativity
• The electronegativity of an element indicates the relative
ability of an atom to attract electrons to itself when the
atom is involved in a chemical bond.
• These values cannot be measured. In fact, they were
developed by a chemist, Linus Pauling, as a way to
explain chemical bonding in molecules. They are based
on fluorine being the element with the strongest
attraction for electrons, so it is a relative scale.
Electronegativity
• The noble gases are ignored since they are basically
inert.
• Fluorine is the most electronegative element and is
assigned a value of 3.98 while cesium and francium are
the least electronegative elements.
• In a chemical bond, the atom with the greater
electronegativity more strongly attracts the electrons in
that bond.
Electronegativity
• Trend:
• Electronegativity decreases from top to bottom in a
group (column) and increase from left to right across a
period (row).
Electronegativity
Example:
• In each of the following sets of elements, which element
would be expected to have the highest ionization
energy?
• a. Cs, K, Li
Li
• b. Ba, Sr, Ca
Ca
• c. I, Br, Cl
Cl
• d. Mg, Si, S
S
Example:
• Arrange the following sets of elements in order of
increasing atomic size.
• a. Sn, Xe, Rb, Sr
Xe < Sn < Sr < Rb
• b. Rn, He, Xe, Kr
He < Kr < Xe < Rn
• c. Pb, Ba, Cs, At
At < Pb < Ba < Cs
Example:
• In each of the following sets of elements, indicate which
element has the smallest electronegativity value.
• a. Na, K, Rb
Rb
• b. S, Na, Si
Na
• c. P, N, As
As
• d. O, N, F
N