History of the Periodic Table

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Transcript History of the Periodic Table

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Russian chemist
Looked for common
properties in
elements
Then arranged by
atomic mass
Noticed similar
properties
appeared at regular
intervals 
“periodic”
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English scientist
Elements fit into
patterns better if
arranged by atomic
number
e.g. Te and I
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Metals
Nonmetals
Metalloids
Transition metals  good conductors, shiny
Alkali metals  most reactive metals
Alkaline earth metals  reactive metals
Halogens  most reactive nonmetals
Noble gases  don’t react
Description of the Modern Periodic Table
• The vertical groups bring together elements with similar properties.
• The horizontal periods of the table are arranged in order of
increasing atomic number from left to right. The groups are numbered
at the top. and the periods at the extreme left in the periodic table on
the inside front cover.
• The first two groups--the s block- and the last six groups--the p
block-together constitute the main-group elements.
• The d-block elements are known as the transition elements. The fblock elements, sometimes called the innertransition elements,
would extend tile table to a widlh of 32 members if incorporated in the
main body of the table. The table would generally be too wide to fit on
a printed page, and so the f-block elements are extracted from the
table and placed at the bottom. The 14 elements following lanthanum
(Z = 57) are called the lanthanides, and the 14 following actinium (
Z = 89) are called the actinides.
Question
Which pair of elements would you expect to exhibit the most
similar physical and chemical properties?
 As and Se
Na and Cl
As and Sb
B and Be
10-2
Metals, Nonmetals, and Their Ions
Metals and nonmetals are distinguished by a collection of physical and
chemical properties.
Metalloids are elements that look like metals and in some ways behave
like metals, but also have some nonmetallic properties.
Atomic Properties and Periodic Table
The periodic trends of the following
properties will be studied here.
Atomic radius
Metallic and Non-metallic character
Ionization (energy) potential
Electron affinity
Electronegativity
The Size of Atoms
Atomic Radius
General definition
It is the distance from the centre of the nucleus to the outermost shell of an
atom.
The covalent radius is one-half the distance between
the nuclei oftwo identical atoms joined by a single
covalent bond.
The ionic radius is based on the distance between the
nuclei of ions joined by an ionic bond.
The metallic radius is one-half the distance between
the nuclei of two atoms in contact in the crystalline
solid metal.
157pm
Na
Covalent radius
Na
186pm
Na
Na
Metallic radius
99pm
Na-
Cl-
Ionic radius
Metallic and Non-metallic Character
Metallic character is the tendency of an
element to lose electrons and form positive
ions (cations). For e.g., alkali metals are the
most electropositive elements.
It is also known as electropositivity.
The tendency of an element to accept
electrons to form an anion is called its nonmetallic or electronegative character. For
e.g., chlorine, oxygen and phosphorous show
greater electronegative or non-metallic
character."
Trends across the period
Metallic character of elements
decreases as we move to the right.
Elements to the left have a
pronounced metallic character while
those to the right have a non-metallic
character.
Non-metallic character increases
from left to right.
Why does the metallic character decrease
from left to right across the period?
The elements to the left of the periodic table
have a tendency of losing electrons easily as
compared to those to the right.
As we move from left to right of the period,
the electrons of the outer shell experience
greater pull of the nucleus.
This greater force of attraction is because the
nuclear charge increases and the size of the
atom decreases from left to right.
Why does metallic character increase down
the group?
As we move down the group the number of
shells increases.
This causes the effective nuclear charge to
decrease due to the outer shells being further
away: in effect the atomic size increases.
The electrons of the outermost shell
experience less nuclear attraction and so can
lose electrons easily thus showing increased
metallic character.
Ionization energy (IE).
The amount of energy required to remove the most
loosely bound electron from an isolated gaseous atom is
called ionization energy (IE).
It is measured in the units of electron volts (eV) per atom
or kilo joules per mole of atoms (kJ mol-1).
Factors Governing Ionization Energy
•Size of the atom/ distance of the valence electron to the
nucleus
•the nuclear charge
•Screening effect of the inner electrons
Size of the atom
As the size of the atom increases the outermost
electrons are held less tightly by the nucleus
(attractive force between the electron and the
nucleus is inversely proportional to the
distance). As a result it becomes easier to
remove the electron and therefore the
ionization energy decreases with the increase in
atomic size.
Charge on the nucleus
•The attractive force between the
nucleus and the electron increases with
the increase in nuclear charge making it
more difficult to remove an electron.
•The ionization energy thus increases with
the increase in the nuclear charge.
Screening effect
In multielectron atoms, the outermost electrons are
shielded or screened from the nucleus by the inner
electrons.
The outer most electrons do not feel the complete
charge of the nucleus.
When the inner electrons are more, the screening
effect will be large, the nuclear attraction will be
less.
Thus when the inner electrons increase the ionization
energy will decrease.
Variation along a period
The ionization energy increases with increasing atomic
number in a period.
This is because
The nuclear charge increases on moving across a
period from left to right.
The atomic size decreases along a period though
the main energy level remains the same.
Due to the increased nuclear charge and
simultaneous decrease in atomic size, the valence
electrons are more tightly held by the nucleus.
Therefore
more energy is needed to remove the electron and
hence ionization energy keeps increasing.
Variation down a group
The ionization energy gradually decreases in moving from top
to bottom in a group.
This is due to the fact that:
The nuclear charge increases in going from top to bottom in a
group.
An increase in the atomic size due to an additional energy
shell (level) 'n'.
Due to the increase in the number of inner electrons there is an
increase in the shielding effect on the outer most electron. The
effect of increase in atomic size and the shielding effect is
much more than the effect of increased nuclear charge.
As a result ,
the electron becomes less firmly held to the nucleus and so the
ionization energy decreases as we move down the group.
Successive ionization energies
The energies required to remove subsequent electrons from a
gaseous atom is called as successive ionization energies. They
are termed as first ,second, third …… ionization energy
depending on the removal of the first, second, third electron
respectively.
Successive ionization energies
The second ionization energies are higher than the
first due to the fact that after the removal of the
first electron the atom changes into a monovalent
positive ion.
In this ion, the number of electrons decreases but
the nuclear charge remains same and so the remaining
electrons are held more tightly by the nucleus and it
becomes difficult to remove the second electron.
Hence the value of the second ionization energy (IE2)
is higher than the first (IE1).
Electron affinity
Electron affinity is the amount of energy released when an
electron is added to an isolated gaseous atom.
Electron affinity is the ability of an atom to hold an additional
electron.
If the atom has more tendency to accept an electron then the
energy released will be large and consequently the electron affinity
will be high.
Just like a strong magnet
Factors affecting electron affinity
When the nuclear charge is high there is greater
attraction for the incoming electron. Therefore electron
affinity increases as the nuclear charge increases.
With the increase in the size of the atom the electron
affinity decreases because the distance between the
nucleus and the incoming electron increases.
Electron affinities are low or almost zero in elements
having stable electronic configurations (half filled and
completely filled valence subshells) because of the small
tendency to accept additional electron.
Variation along a period
The size of an atom decreases and the nuclear charge
increases on moving across a period. This results in
greater attraction for the incoming electron. Hence the
electron affinity increases in a period from left to right
Variation down a group
As we move down a group the atomic size and nuclear
size increases. As the effect of increase in atomic size
is more pronounced the additional electron feels less
attracted by the large atom. Consequently the electron
affinity decreases.
Electronegativity.
This the relative tendency of an atom in a
molecule to attract a shared pair of electrons
towards itself.
The value of electronegativity of an element
describes the ability of its atom to compete for
electrons with the other atom to which it is
bonded.
Variation along a period
As the nuclear charge increases from going left
to right in a period because the electrons enter
the same shell, the shielding is less effective.
Thus the increased nuclear charge attract the
shared pair of electrons more strongly resulting
in higher electronegativity from going left to
right in a period.
Variation down the group
Electronegativity decreases down the group
because the atomic size increases. The larger
the size of the atom the lesser the tendency to
attract the shared pair of electrons.
What happens if two atoms of equal electronegativity bond
together?
Consider a bond between two atoms, A and B.
If the atoms are equally electronegative, both have the same
tendency to attract the bonding pair of electrons, and so it will be
found half way between the two atoms.
To get a bond like this, A and B would usually have to be the same
atom.
Example,
H2, O2, Cl2 molecules.
This type of covalent bond is described as a non-polar
covalent bond.
Polar covalent bond
This is the bond between two unlike atoms, which
differ in their affinities for electrons.
When a covalent bond is formed between two atoms of
different elements, the bonding pair of electrons will lie
more towards the atom, which has more affinity for
electrons.
The electron pair do not lie exactly midway between the
two atoms.
The atom with higher affinity for electrons develops a
slight negative charge and the atom with lesser affinity
for electrons, a slight positive charge. Such molecules
are called 'polar molecules'.
In the hydrogen chloride (HCl) molecule, the bonding of
hydrogen and chlorine atoms lies more towards Cl atom
(because Cl is more electronegative) in the shared pair
of electrons. Therefore, Cl atom acquires a slight
negative charge, and H atom a slight positive charge.
This causes the covalent bond between H and Cl to have
an appreciable ionic character.
The compounds having polar bonds are termed polar
compounds. Polar substances in their pure forms, do not
conduct electricity, but give conducting solutions when
dissolved in polar
solvents.
The greater is the difference in the electronegativity
values of the combining atoms, greater is the polar
character in the bond so formed. For example, in the
series H - X (X=F, Cl, Br,I), the electronegativity
difference between H and X atom follows the order:
H- F > H - Cl > H - Br > H - I
Therefore, the polarity in the H - X bond follows the
order - F > H - Cl > H - Br > H I i.e., H - F bond is the
most polar and
H -I bond is the least polar in this series of
compounds.
Periodic Properties of the elements
Atomic Properties
Electron affinity
Atomic radius
Atomic properties and the periodic table—a summary
Electron affinity
Ionization energy
Atomic radius
metallic character
Ionization energy
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