Transcript periodic table of element

MODERN PERIODIC LAW
The physical and
chemical
properties of
elements are
periodic function
of their atomic
numbers.
PERIODIC LAW
The physical and
chemical
properties of
elements are
periodic function
of their atomic
weights.
In multielectron atoms, the electrons in the outermost shell
are screened from the nucleus by the inner electrons. As a
result, the electron in the outermost shell does not experience
the full charge of the nucleus.
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The actual nuclear charge felt by electron is called the
effective nuclear charge.
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The effective nuclear charge can be determined by
subtracting a screening constant for the inner electrons from
the actual nuclear charge( atomic number).
Z* =Z- S
where Z* is effective nuclear charge
Z is nuclear charge(atomic number)
S is screening constant.
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Write the electronic configuration of the element in the
following order and grouping: (1s), (2s,2p), (3s,3p), (3d),
(4s,4p), (4d), (4f), (5s,5p), (5d), (5f),…………
Electrons in any group higher in the sequence than the
electron under consideration contribute zero to screening
constant.
Electrons in the same group contribute S= 0.35 each. ( if
electron is present in 1s it will contribute S= 0.30).
For an electron in ns or np orbital, all electrons in (n-1) shell
contribute S= 0.85 each and all electrons in (n-2) shell
contribute S= 1.0 each.
For an electron in nd or nf orbital, all electrons in lower groups
contribute S= 1.0 each.
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The term atomic radius
means the distance from
centre of the nucleus to the
outermost shell of electrons.
An atom gets larger as the
number of electronic shells
increase; therefore the
radius of atoms increases as
you go down a certain group
in the periodic table of
elements.
In general, the size of an
atom will decrease as you
move from left to the right
of a certain period.
Ionization energy is the
quantity of energy that an
isolated, gaseous atom in the
ground electronic state must
absorb to discharge an
electron, resulting in a cation.
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H(g)→H+(g)+e−
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Ionization energies are
dependent upon the atomic
radius. Since going from
right to left on the periodic
table, the atomic radius
increases, and the ionization
energy increases from left
to right in the periods and up
the groups.
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Electron affinity is defined as
the change in energy (in
kJ/mole) of a neutral atom (in
the gaseous phase) when an
electron is added to the atom
to form a negative ion.
On moving down the group,
electron affinity decreases.
On moving across a period,
electron affinity increases.
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Electronegativity can be
understood as a chemical property
describing an atom's ability to
attract and bind with electrons.
From left to right across a
period of elements,
electronegativity increases. If
the valence shell of an atom is less
than half full, it requires less
energy to lose an electron than to
gain one. Conversely, if the valence
shell is more than half full, it is
easier to pull an electron into the
valence shell than to donate one.
From top to bottom down a
group, electronegativity
decreases. This is because atomic
number increases down a group,
and thus there is an increased
distance between the valence
electrons and nucleus, or a greater
atomic radius.
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Ionic character: ionic character depends upon
electronegativity difference between bonded atoms. The
greater the difference in the electronegativities of the
bonded atoms, the higher is the ionic character of the bond.
As we move down the group the degree of ionic character in
halides increases because of increase in electronegativity
difference.
HF > HCl > HBr > HI
As a result of ionic and covalent character , their properties
also change. For example, AlF3 is essentially ionic. AlCl3 has an
intermediate character and AlBr3 and AlI3 exist as covalent
dimers. Therefore, the melting point of these halides decrease
from AlF3 AlI3. AlF3 and AlCl3are conductors in fused sate
(ionic character) while AlBr3 and AlI3 are non-conductors
(covalent character).
• Oxidising and reducing character: Oxidation is a process in
which electrons are lost and reduction is a process in which
electrons are gained. Thus, an oxidising agent is a substance
which gains electrons while reducing agent is a substance which
loses electrons. Since the elements on the left of the periodic
table have low ionization energy values and therefore, have good
tendency to lose electrons. Consequently, they act as strong
reducing agents. On the other hand the elements on the right
hand side of the periodic table have great tendency to accept
electrons and therefore they act as good oxidising agents.
•Now , since ionization energy values decrease on moving down a
group therefore, reducing character of elements in a group
increases down a group. Conversely, the oxidising character of
elements decreases on moving down a group.
•As we move along a period from left to right, the ionization
energy increases and consequently, reducing character
decreases while oxidising character increases.
• LITHIUM IS THE STRONGEST REDUCING AGENT WHILE
FLOURINE IS THE STRONGEST OXIDISING AGENT.
• OXIDATION STATES: Like periodic properties, the
oxidation states of elements are closely related to the number
of electrons in the valence shell of their atoms. Oxidation
states of different elements correspond to the number of
electrons gained or lost by their atoms to acquire a complete
shell of eight electrons, ns2 np6.
Variation of oxidation state in a period: as we move along
a period , the positive oxidation state increases and negative
oxidation state decreases. For example, in second period, the
oxidation state increases from Li to C from +1 to +4 and then
decreases from nitrogen to fluorine from -3 to -1.
Variation in a group : In a group , the maximum oxidation
state shown by a p-block element is equal to the sum of its s
and p electrons, which is the same as its group number.
• acidic and basic character of hydroxides: The acidic or basic
character of hydroxides and oxides depend upon the position of the
element in the periodic table.
E------ O------H
There are two possibilities:
1. If the electronegativity of E is low, as in case of metals like Na, K, Mg
and Ca etc. the electrons of E-O bonds are drawn more closely to the
oxygen atom. This is because oxygen is highly electronegative element.
As a result bond between E-O breaks giving OH- ions. Therefore such
substance behaves as bases and readily react with acids like HCl to
get neutralized.
2. If the electronegativity of E is high as in case of non-metals like F, Cl,
N etc., the electrons of E-O bond are shared more equally between
non-metal and the oxygen atom. As a result, the oxygen –hydrogen
bond becomes weaker and cleaves forming a proton . Therefore, such
substances behaves as acids and are neutralized by bases.
•Reactivity of alkali and alkaline earth metals: As
the value of ionization energy decreases down the
group from Li to Cs, therefore, the reactivity of
alkali metal increases from Li to Cs. All the elements
are highly electropositive giving +1 ions. Because of
the very high second ionisation energies of these
elements, their oxidation state in compounds never
exceeds +1.
On the other hand , alkaline earth metals are in
general less reactive than alkali metals. This is
because of their relatively high ionization energies
and high heat of atomization in comparison to alkali
metals. The chemistry of this group is mainly
dominated by +2 oxidation state.
Some of the general important chemical trends are discussed below:
• The anomalous properties of elements in first short period (from Li to F)
are explained due to their peculiar atomic properties such as small size, low
ionisation energy and high electronegativity values.
•Diagonal relationship between elements of different groups (such as Li –
Mg, Be-Al, B-Si) are due to their similar atomic properties.
•Trends in bond type with the position of the element in the periodic table
and with oxidation state for a given element.
•The variable oxidation states of transition elements.
•The stability of compounds and their trends.
•Trends in stability of coordination compounds and the electron donor
power of various types of ligands.