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Periodic properties of elements
T.M.Sankaranarayanan
17.04.2010
Antoine-Laurent de Lavoisier
Born
Died
Influen
ces
26 August 1743(1743-08-26)
Paris, France
8 May 1794 (aged 50)
Paris, France
Guillaume-François Rouelle
Johann Wolfgang Döbereiner
In 1817, Johann Wolfgang
Döbereiner began to formulate one
of the earliest attempts to classify
the elements. He found that some
elements formed groups of three
with related properties. He termed
these groups "triads".
Alexandre-Emile Béguyer de Chancourtois
(20 January 1820 – 14 November 1886)
was a french geologist and mineralogist
who was the first to arrange the
chemical elements in order of atomic
weights, doing so in 1862.
De Chancourtois only published his
paper, but did not publish his actual
graph with the proposed arrangement.
Although his publication was significant,
it was ignored by chemists as it was
written in terms of geology. It was
Dmitri mendeleyev's table published in
1869 that became most recognized.
John Alexander Newlands
Born
Died
November 26, 1837 (183711-26)
London
July 29, 1898 (1898-07-30)
Nationality
English
Fields
analytical chemistry
Alma mater
Royal College of Chemistry
Known for
periodic table
john newlands was an English chemist who in 1865
classified the 56 elements that had been discovered at
the time into 11 groups which were based on similar
physical properties.
Newlands noted that many pairs of similar elements existed which differed by
some multiple of eight in atomic weight. However, his law of octaves, likening this
periodicity of eights to the musical scale, was ridiculed by his contemporaries. It
was not until the following century, with Gilbert N. lewis’ valence bond theory
(1916) and irving langmuir’s octet theory of chemical bonding[ (1919) that the
importance of the periodicity of eight would be accepted
Dmitri Mendeleev in 1897
Born
Died
8 February 1834(1834-02-08)
Verhnie Aremzyani, Russian
Empire
2 February 1907 (aged 72)
St. Petersburg, Russian
Empire
Nationality
Russian
Fields
Chemistry, physics and
adjacent fields
Alma mater
Saint Petersburg University
Notable
students
Dmitri Konovalov, Gemilian,
Valery, Baykov, Alexander
Known for
Inventing the Periodic table
of chemical elements
MENDELEEF’S PERIODIC TABLE
Mendeleef proposed following periodic law and arranged elements in the
increasing order of atomic masses.
Mendeleef’s Periodic law
"Physical and chemical properties of elements are periodic
functions of their atomic weights"
The elements are arranged in horizontal rows and vertical columns.
•The horizontal rows are called periods. The number of elements in a period may
vary. The first three periods of the Mendeleef table are called as short periods.
The other periods are known as long periods.
•The vertical columns in the table are known as groups or families. The groups are
sub divided into two subgroups ‘A’ and ‘B’. The elements arranged in a group possess
similar properties.
•The long period of the Mendeleef periodic table consists of two rows of elements
called as series.
* He considered the similarities in the formulae and the properties of the
compounds formed by the elements.
Mendeleef observed that elements with similar properties had
(i) Either almost the same atomic weights.
e.g. Fe(56), Co(59), Ni(59)
Os(191), Ir(193), Pt(195)
(ii) Or atomic weights which showed a constant increase
e.g. K(39), Rb(85), Cs(133)
Ca(40), Sr(88), Ba(137)
The elements with low atomic weights are called typical elements. These are
arranged in three short periods of the periodic table.
•Group VIII of the Mendeleef table contains three triads, namely, (Fe, Co, Ni
and Ru, Rh, Pd and Os, Ir, Pt). These triads are called transition elements
which include Sc(21) to Zn (30), lanthanides and actinides.
* From a study of adjacent elements and their compounds, Mendeleef was able
to predict the characteristics of certain elements which were found to be very
accurate. e.g.1) Eka Al - Gallium 2) Eka Si - Germanium 3) Eka B - Scandium
Comparison of properties predicted by Mendeleef with those observed.
Periodic Table
He corrected the atomic weights of some elements like Beryllium, Indium,
Uranium.
Limitations :
Zero group elements were not known at the time of Mendeleef.
•In every group lanthanides were placed though they have
dissimilar properties.
* At some places it violates the increasing order of atomic
weight rule.
Ar40 & K39
Co59 & Ni58
Te128 & I127
Th 232 & Pa 231
These four pairs are called anomalous pairs or inverted pairs.
* Elements with dissimilar properties were grouped together.
E.g. ' Th' is placed in III group and Ag is placed in I group.
* Atomic mass is taken as fundamental property.
MODERN LONG FORM OF PERIODIC TABLE
Modern periodic table was constructed by Neils Bohr based on modern
periodic law.
Modern periodic law
The chemical and physical properties of elements are the periodic
functions of their atomic numbers and electronic configurations.
Above law was proposed by Moseley. He found the relation between
atomic numbers (z) and the frequencies () of x- rays produced by
them.
Where a & b are constants characteristic of elements
Periodic Table Summary
•The subdivisions of the periodic table are periods, groups, and classes.
The horizontal rows are called periods. The vertical columns are called
groups.
•The entire table consists of three classes: metals, non-metals, and
semimetals.The subdivisions of the periodic chart have been explained
such that the student should be able to identify them if given a periodic
table.
•Elements of the same group share certain physical and chemical
characteristics. Examples of the characteristics of several groups are
listed below.
•Group 0 contains elements that are unreactive gases.
•Group IA contains elements that are chemically active soft metals.
•Group VIIA contains elements that are chemically active nonmetals.
•Groups IB through VIIB and VIII are called transition groups and
•elements found in them display properties of metals.
•The valence of an atom is defined as the number of electrons an element
•gains or loses, or the number of pairs of electrons it shares when it
interacts with other elements.
Number of elements in each period
Development of the Periodic Table
• Mendeleev: The First Periodic Table
– Arranged elements by increasing atomic mass
– Noticed a regular pattern in chemical and
physical properties
– Blank spaces – Predicted the existence of
elements not yet discovered based on
properties
• Moseley
– Determined atomic number of elements by
determining nuclear charge
– Arranged table by increasing atomic number –
remains this way today
The Modern Periodic Table
• Periods – Horizontal Rows
• Groups/Families – Vertical Columns
• Elements in groups have similar physical
and chemical properties
• Periodic Law: When elements are
arranged in order of increasing atomic
number, there is a regular pattern in
their chemical and physical properties
Electron Configurations & Periodicity
• Electrons – play a large part in
determining
physical
&
chemical
properties of elements
• Arrangements related to position on the
table
– Groups have the same outer orbital
configuration in consecutively higher
energy levels
– Periods have outer electrons in the same
energy level
Classifications
• Noble Gases: elements in which outermost s and p
sublevels are filled [full octet of valence e-]
• Representative (Main Group 1A-7A) elements:
outermost s or p sublevels are only partially filled.
The number of valence electrons can be determined
by the group number
• Transition Metals – elements in which s and nearby
d sublevels contain electrons
• Inner Transition Metals – elements in which s and
nearby f sublevels contain electrons
Reading Configurations from the Periodic Table
• Read the table like sentences in a paragraph from
left to right & top to bottom until a particular
element of interest is reached.
• Each period corresponds to a principle energy
level, each group corresponds to a particular
sublevel and orbital position.
– Subtract 1 from the principle energy level for
transition metals (d-block)
– Subtract 2 from the principle energy level for
inner transition metals (f-block)
Trends in Atomic Size
• Atomic Radius: half of the distance between
the nuclei of two atoms of the same element
bonded together.
• Group Trends: Generally increase as one goes
down a group in the periodic table
– Electrons occupy consecutively higher energy levels,
farther from the nucleus
– Nuclear charge increases, but outermost orbitals are
larger due to distance
– Shielding effect increases with additional number of
occupied orbitals between the outermost electrons
and the nucleus
Atomic Radius Cont.
• Period Trends: Decreases left to right
across the periodic table
• Electrons are added to the same
principle energy level
• Increased nuclear charge “pulls” them
closer to the nucleus
Trends in Atomic Radius
• Atomic Size increases
from top to bottom
within a group.
• This is due to shielding
effect.
The
inner
electrons
shield
the
outer electrons from the
nucleus. So, the nucleus
has less pull. Therefore,
the atom gets bigger
down a group.
Trends in Atomic Radius
• Atomic size usually
decreases across a period
from left to right.
• This is due to the nuclear
charge (which is positive)
pulling on the electrons
(which are negative.)
• The nuclear charge
increases causing it to
have more pull. It pulls
the energy levels in making
the atom smaller.
Trends in Ionic Radius
• An Ion is an atom or a group
of atoms with a charge.
• Positive ions form when
electrons are lost. (cations)
• Negative ions form when
electrons are gained.
(anions)
Sodium Ion
• Cations are always smaller
than the atoms from which
they form.
• Anions are always larger
than the atoms form which
they form.
Chlorine Ion
Trends in Ionic Radius
180
• In general, ionic
radius increases
down a group due
to the electron
shielding.
160
140
120
Lithium
Sodium
Potassium
Rubidium
Cesium
100
80
60
40
20
0
Group Trend
Trends in Ionic Radius
• Ionic radius
decreases across a
period for the
cations, and
decreases across a
period for the
anions due to
nuclear charge.
Lithium
160
140
Berylliu
m
Boron
120
100
Carbon
80
60
Nitrogen
40
Oxygen
20
0
Period Trend
Fluorine
Trends in Ionic Size
• Positive ions are smaller than the neutral atoms
from which they are formed
– Electrons are lost from energy levels farthest
from the nucleus
– Remaining electrons pulled closer to the nucleus
• Negative ions are larger than the neutral atoms
from which they are formed
– Electrons are gained, resulting in smaller
effective nuclear charge for the greater number
of electrons
– Repulsive forces between electrons increases
across the periods
Ionic Size Trends Cont.
• The trends follow the same pattern as
atomic size
– Larger from top to bottom
– Smaller from left to right
Trends in Ionization Energy
• Ionization Energy: energy needed to remove
an electron from a gaseous atom
• Group 1: easily loses its 1 valence electron
– Low first ionization energy
– Second ionization energy will be very high
since it is “happy” with losing 1 electron.
• Group 2: easily lose 2 valence electrons
– Low first and second ionization energies
– High third ionization energy
Trends in Ionization Energy
• The energy required to remove an
electron from an atom is called
ionization energy.
• First Ionization energies are the
energies required to remove the first
electron.
• Second Ionization energy is the energy
required to remove the second electron,
and so on…
Trends in Ionization Energy
Atomic Symbol
Number
Name
1st
2nd
3rd
1
H
Hydrogen
1312
2
He
Helium
2372
5250
3
Li
Lithium
520
7298
11815
4
Be
Beryllium
899
1757
14848
21006
5
B
Boron
800
2427
3659
25025
4th
• Lithium only has 1 electron to lose in order to bond.
Notice that between 1st and 2nd ionization energies
that the number goes up. Because once lithium has
lost it’s one electron it becomes stable. Elements
like being stable.
So, the ionization energy
increases.
Trends in Ionization Energy
• First Ionization energies tend to decrease from top
to bottom within a group due to electron shielding.
Trends in Ionization Energy
• First Ionization energies tend to increase from
left to right across a period due to increased
nuclear charge.
Ionization Energy Trends
• Group Trends: decreases from top to bottom
on the periodic table
– Outermost electrons are farther from the
effect of the nuclear charge and therefore
easier to remove
– Shielding effect increases down the table
• Period Trends: increases from left to right
– Nuclear charge is increasing with no
increase in shielding effect
– Outermost electrons are closer to the
nucleus
Electron Affinity Trends
• Electron Affinity: energy change that
accompanies the addition of an
electron to a gaseous atom
– Most electron affinities are negative
because most elements release energy
when they become negative ions
• Group: Decreases from top to bottom
• Period: Increases left to right
Electron Affinity Trends
(Same as for Ionization Energy)
• Group Trends: decreases from top to bottom
on the periodic table
• Outermost electrons are farther from the
effect of the nuclear charge and therefore
easier to remove
• Shielding effect increases down the table
• Period Trends: increases from left to right
• Nuclear charge is increasing with no
increase in shielding effect
• Outermost electrons are closer to the
nucleus
Electronegativity
• Electronegativity: tendency for an atom
to attract electrons to itself when it is
chemically combined with another
element
• Group: decreases top to bottom
• Period: increases left to right
Electronegativity Trends
(Similar to Ionization Energy)
• Group Trends: decreases from top to bottom
on the periodic table
• Outermost electrons are farther from the
effect of the nuclear charge and therefore
more difficult to attract
• Shielding effect increases down the table
• Period Trends: increases from left to right
• Nuclear charge is increasing with no increase
in shielding effect
• Outermost electrons are closer to the
nucleus and attracted more easily
Trends in Electronegativity
• Is the ability of an
atom of an element to
attract electrons when
the atom is in a
compound.
• Notice that noble
gases have no values.
This is due to the fact
that most do not form
compounds.
Electronegativity Values
Li Be B C N O F
1.0 1.5 2.0 2.5 3.0 3.5 4.0
Na Mg Al Si P S Cl
0.9 1.2 1.5 1.8 2.1 2.5 3.0
K Ca Ga Ge As Se Br
0.8 1.0 1.6 1.8 2.0 2.4 2.8
Rb Sr In Sn Sb Te I
0.8 1.0 1.7 1.8 1.9 2.1 2.5
Trends in Electronegativity
• In general, electronegativity values decrease from
top to bottom within a group due to the shielding
effect. The inner electrons shield the nucleus
from being able to attract electrons.
Trends in Electronegativity
• In general, electronegativity values tend to
increase from left to right across a period due to
an increase in nuclear charge.
General Explanations
• Shielding effect plays a major role in
how strongly the nucleus can pull on
its outermost electrons – This mainly
effects group trends.
• Nuclear charge is a measure of the
strength of a nucleus’ pull - This
greatly
effects
period
trends
because shielding effect is not an
issue across a period.
Effective Nuclear charge
• The effective nuclear charge, is the
net positive charge experienced by an
electron in a multi-electron atom. The
term "effective" is used because the
shielding effect of negative electrons
prevents higher orbital electrons from
experiencing the full nuclear charge. It
is possible to determine the strength of
the nuclear charge by looking at the
oxidation number of the atom.
Summary of Trends
Thank you