7. 2. Antoine Lavoisier 1789:
• Made a list of 33 elements and divided
them into five classes:
• Light and heat: material substances
• Three gases: Oxygen, hydrogen, and
nitrogen.
• Seventeen metals: alphabetical order.
• Five earths: Now known to be oxides.
• Six non-metals.
8. • John Dalton 1808:
• Published the first list of
atomic weights
• In his ‘ Table of the relative atomic
weights of the ultimate particles
of gaseous and other bodies’.
• Changed chemists idea from a
Qualitative to a Quantitative basis.
9. • Dobereiner’s Triads 1829:
• Three chemically similar
elements, which he called
triads.
• Atomic weight of the middle
element in a triad is the
arithmetic mean of the other two.
10. • Dobereiner’s Triad:
• Drawbacks:
• Many elements could not be placed in a triad.
• For example: Iron, manganese, nickel,
cobalt, zinc and copper are six similar
elements, cannot be places in the triad.
13. • Newland’s Law of Octaves:
• Newland was the first person to give a
number to an element.
• He left places for elements still undiscovered.
• He altered positions of elements ( e.g
tellurium, iodine), if the atomic weights placed
them incorrectly.
• Newlands could classify elements upto
calcium only. The elements that were
discovered later could not be fitted into the
octave pattern.
14. • Mendeleev’s Periodic Table:
• (1884 – 1907)
• Dmitri Ivanovich Mendeleev
• First really successful arrangement
• “… the properties of elements are
in periodic dependence of their
atomic weights”.
• First to arrange the elements in the shape of a
table.
• This table showed how the properties of elements
were repeated, elements were related to each
other and chemical reactivities changed with
increasing atomic weight.
15.
16. •
•
• Arranged elements primarily in order of increasing
atomic weights.
• He left a number of gaps in his table.
• This table also helped in the correction of wrong atomic
weights.
• Eg. The atomic weight of beryllium was corrected.
• Eight out of ten gaps left by Mendeleev were filled.
• Predictive value and correct errors of atomic
weights.*
17. • Main defects of Mendeleev’s Periodic Table
• Position of Isotopes.
• Position of Rare Earth metals.
• Position of Hydrogen.
• Position of Elements having higher and lower
atomic weights.
• This led to further modification of the Periodic
Table.
with new developments into the structure of
the atom, it was realized that it is the ATOMIC
NUMBER that determines the fundamental
18. • The Modern Periodic Table.
• Henry Moseley ( 1913)
• The modern Periodic Law… “ The physical
and chemical properties of elements
are a periodic function of their
Atomic numbers”.
• Differs from Mendeleev table.
• The atomic number of an element is a more
fundamental property than atomic mass. Classifying
elements according to their atomic number lies in
the fact that, atomic numbers being equal to the
number of electrons, allows the elements to be
arranged according to their electronic configuration.
19. • Relation of Periodic Table and Electronic
Configuration:
• When elements are arranged in the increasing order of their
atomic numbers, there is periodic repetition of elements of
similar properties.
• The periodicity is of 2,8,8,18,18,32 elements. The cause of
periodicity is the repetition of similar electronic configurations in
the outermost shell.
•
20. Organization of the modern periodic table.
The main features of modern periodic table
• Elements are arranged in the increasing order of atomic
numbers.
21. • Elements in the modern periodic table are arranged in 7
periods and 18 groups.
• Horizontal rows are called periods.
• Vertical columns are called groups.
22. • The group number is assigned to an element
depending upon the number of valence electrons
present in the atom of the element.
24. • Going down the periodic table the number
of atomic orbital’s increases by one for
each row.
25. • In periods there are same number of shells, which
depicts the period number.
26.
27.
28.
29. • Group 1 Alkali metals:
• The elements such as lithium, sodium, potassium, rubidium,
cesium and francium have one electron in their outermost shell
and therefore show valency one. Hence placed in group 1.
Element K L M N O P Q
Li 2 1
Na 2 8 1
K 2 8 8 1
Rb 2 8 18 8 1
Cs 2 8 18 18 8 1
Fr 2 8 32 32 18 8 1
30. • Group 2 Alkaline earth metals:
• The elements such as beryllium, magnesium,
calcium, strontium, barium and radium have two
electrons in their outermost shell and therefore show
valency two. Hence placed in group 2.
Element K L M N O P Q
Be 2 2
Mg 2 8 2
Ca 2 8 8 2
Sr 2 8 18 8 2
Ba 2 8 18 18 8 2
Ra 2 8 18 32 18 8 2
31. • Group 17 Halogens:
• The elements such as fluorine, chlorine, bromine, iodine
and astatine have seven electrons in their outermost
shell and therefore show valency one. Hence placed in
group 17.
Element K L M N O P Q
F 2 7
Cl 2 8 7
Br 2 8 18 7
I 2 8 18 18 7
At 2 8 18 32 18 7
32. • Group 18 Noble gases:
• The elements such as helium, neon, argon, krypton,
xenon and radon have eight electrons in their outermost
shell and therefore show valency zero. Hence placed in
group 18.
Element K L M N O P Q
He 2
Ne 2 8
Ar 2 8 8
Kr 2 8 18 8
Xe 2 8 18 18 8
Rn 2 8 18 32 18 8
33. • Position of Hydrogen
• Has one valence electron
• In some reactions, this electron is given up
and the hydrogen ion undergoes reactions
similar to the alkali metals.
• Hydrogen differs from the alkali metals in
many ways, hydrogen is a diatomic non-metal.
• However, hydrogen, like halogens is a
monovalent, diatomic non-metal and
undergoes reactions similar to the halogens.
• Due to this dual nature, hydrogen is treated as
a special element and placed alone at the
head of the periodic table above the group I
alkali metals.
34. • H – similar to alkali metals and halogens.
• Similarities between hydrogen and alkali metals.
• H – 1
Li – 2,1
Na – 2,8,1
K – 2,8,8,1
• like all alkali metals hydrogen has 1 electron in the
valence shell.
• since it has only 1 valence electron it can lose that 1
electron to stabilize itself to satisfy the octet rule.
• When it combines with oxygen it forms the
corresponding metal oxides.
• Eg: 2Na + O2 2Na2O
35. • Similarities of hydrogen with halogens:
• Halogens are group 17 elements.
• All halogens are short of 1 valence electrons.
• F = 2,7
Cl = 2,8,7
• Outermost shell has 1 electron less than the nearest
noble gas.
• Similarly hydrogen has 1 electron less than the nearest
duplet configuration helium.
• Like all halogens gain 1 electron to complete the octet
hydrogen also has to gain 1 electron to form a duplet
structure. To stabilize itself.
• Halogens F2, Cl2, Br2 exists in its gaseous form
Hydrogen also exists in gaseous form H2.
• Atomicity of all halogens and hydrogen is 2.
36. Periodic Trends:
• Atomic Radius
• Metallic Character or Electropositivity
• Non-metallic Character or Electronegativity
• Ionization energy or Ionization Potential
• Electron Affinity
• Periodic trend is a regular variation of the
properties of an element with increasing
atomic number.
37. Gradation of properties in the
Periodic Table
•
• All periodic properties can be explained on the basis of two factors,
namely,
38. Atomic Radius/Size:
• Group: Atomic radius increases from top to bottom within a group.
• Reason: One shell is added and as a result, the atoms become larger.
Moreover, the higher number of sheilding electrons helps the valence shell to
spread out more, thus increasing the size of the atom. (Sheilding electrons: All
electrons other than valence electrons).
• Period: Atomic radius decreases from left to right within a period.
• Reason: This is caused by the increase in the number of protons and
electrons(increase in the atomic number) across a period which brings about an
increase in the effective nuclear charge. As a result the electrons are pulled
closer towards the nucleus, resulting in a smaller radius.
• Effective nuclear charge – the attractive positive charge of nuclear protons
acting on valence electrons.
39. Metallic Character or Electro positivity
• Group: Metallic character increases as you move down a group.
• Reason: The atomic size is increasing due to an addition of a shell. When the
atomic size increases, the valence shell is farther away and the electrons of the
valence shell have less attraction to the nucleus and, as a result, can lose
electrons more readily. This causes an increase in metallic character.
• Period: From right to left across a period, metallic character increases.
• Reason: From right to left across a period, metallic character increases
because the atomic size increases and the attraction between valence electrons
and the nucleus is weaker, enabling an easier loss of electrons.
• Note: the opposite is true from left to right across a period. From left to
right the metallic character decreases.
Definition: The metallic character or
electropositive nature of an element
can be defined as a measure of the
ability of an atom to lose/donate
electrons to form positive ions.
40. Non-metallic character or Electronegativity
• Group: From top to bottom in a group, electro negativity decreases.
• Reason: From top to bottom down a group, electro negativity decreases.
This is because the number of shells increases down a group, and thus
there is an increases distance between the valence electrons and nucleus or
a greater atomic radius. As the pull of the nucleus becomes less, valence
electrons are more easily lost than gained.
• Period: From left to right across a period, the electronegativity increases.
• Reason: From left to right across a period, electronegativity increases
because the atomic size is decreasing and the effective nuclear charge is
increasing. So the nucleus is able to attract electrons more easily towards
itself to complete its valence shell.
Definition: Electronegativity or non-
metallic character is a measure of
the tendency of an atom to
gain/attract electrons towards itself
to form negative ions.
41. Ionization energy or Ionization Potential
• Group: The ionization energy of the elements decreases from top to bottom
within a group.
• Reason: The ionization energy of the elements within a group generally
decreases from top to bottom. This is due to increased atomic radius and
decreased pull of the nucleus on valence electrons. As a result, electrons
require very little energy to be removed.
• Period: The ionization energy of the elements increases from left to right
within a period.
• Reason: Atoms on the left side of the period are larger as their effective nuclear
charge is less. Hence they easily lose electrons and become cations. In other
words they have low ionization energies.
• Atoms on the right side of the period have smaller atomic radius and a greater
effective nuclear charge. Moreover, their valence shell is nearly filled and more
stable. Hence, it requires more energy to remove an electron. In other words,
they have higher ionization energies.
Definition: Ionization energy of an
atom is the energy required to
remove an electron from an
isolated, neutral gaseous atom to
form a positive ion.
42. Electron Affinity
• Group: Electron affinity decreases from top to bottom in a group.
• Reason: Electron affinity generally decreases down a group of elements
because the atomic size increases. A large atom has less nuclear pull on the
valence electrons hence less energy is released when an electron is added
to the valence shell. In other words, electron affinity decreases.
• Period: Electron affinity increases from left to right within a period.
• Reason: Moving from left to right across a period, atoms become smaller
as the nuclear charge increases. A small atom takes up electrons more
readily than a large atom, releasing higher amounts of energy. In other
words, the electron affinity increases from left to right across a period.
Definition: Electron affinity of an
atom is the amount of energy
released when an electron is
added to an isolated, neutral
gaseous atom to form a negative
ion.
44. Some points to note:
1. Periodicity: The properties that reappear at regular intervals, or in which
there is a gradual variation(i.e. increase or decrease) at regular intervals,
are called periodic properties and the phenomenon is called periodicity of
elements.
Reason for periodicity: After definite intervals of atomic number, similar
valence shell electronic configuration occurs in the same group. In the
same period, an increase or decrease in a particular property is seen due to
a gradual change in electronic configuration of the elements.
2. Modern periodic law: the physical and chemical properties of elements
are a periodic function of their atomic numbers.
3. The most non-metallic element in the periodic table is Fluorine.
4. the most metallic element in the periodic table is Francium/ Cesium.
5. Metals tend to lose electrons and are therefore, good reducing agents.
6. Nonmetals tend to gain electrons and are therefore, good oxidising
agents.
7. Transition metals have high boiling and melting points, are good
conductors of heat and electricity, display variable valencies, act as
