Classification of elements and periodicity in properties

Aman Bathla (9034186096) Page 1
UNIT 3
CLASSIFATION OF ELEMENTS & PERIODICITY IN PROPERTIES
A well organized & tabulated classification of elements help to locate, identify & characterize
the elements & its properties.
VARIOUS ATTEMPTS OF THE CLASSIFICATION
DOBEREINER’S TRIADS
Element Atomic weight Element Atomic weight Element Atomic weight
Li 7 Ca 40 Cl 35.5
Na 23 Sr 88 Br 80
K 39 Ba 137 I 127
Main Features
 Dobereiner classified the elements into groups of three elements (called triads) with
similar properties.
 the atomic weight of the middle element was the arithmetic mean of the other two, e.g.
Mean of atomic weight of Na =
7 (𝑎𝑡𝑜𝑚𝑖𝑐 𝑤𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝐿𝑖)+39(𝑎𝑡𝑜𝑚𝑖𝑐 𝑤𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝐾)
2
= 23
This Dobereiner’s relationship is referred as “Law of triads”.
Limitations
 Dobereiner could not arrange all the known elements into triads.
 Dobereiner could identify only three such triads.
NEWLAND’S OCTAVES
Element Li Be B C N O F
Atomic weight 7 9 11 12 14 16 19
Element Na Mg Al Si P S Cl
Atomic weight 23 24 27 29 31 32 35.5
Main Features
 Newland arranged the elements in the increasing order of atomic weights.
 The properties of the every eighth element are similar to the first one, e.g.
Properties of Li are same as Na (the eighth element to Li)
This relationship is called as “Law of octaves”.
Limitations
 This classification was successful up to the element Ca.
 When Noble gas elements were discovered at a later stage, their inclusion in these
octaves disturbed the entire arrangement.
MENDELEEV’S PERIODIC TABLE
Main Features
 The Chemical & physical properties of elements are a periodic function of their atomic
masses.
This relationship is called as “Mendeleev’s Periodic Law”.
 Elements are arranged in tabular form in Rows & columns.

 The Horizontal Rows (7) present in the periodic table are called periods.
 The Vertical Columns present in it are called groups (8+1). Zero group was added later
on in the modified Mendeleev’s Periodic Table.
Merits
 Mendeleev’s periodic table was very helpful in remembering and studying the properties
of large number of elements.
 Mendeleev left space for the elements yet to be discovered. e.g.
He left spaces for Ga and Ge and named these elements as Eka-Aluminium (Ga) and
Eka-Silicon (Ge) respectively
 Mendeleev’s Periodic Table helped in correcting Atomic mass of doubtful elements on
the basis of their expected positions and properties.
Demerits
 Position of Hydrogen is not correctly defined in periodic table. It is placed in group 1
though it resembles both group 1 and 17.
 In certain pairs of elements increasing order of atomic masses was not obeyed, e.g.
Argon (Ar, atomic mass 39.9) is placed before Potassium (K, atomic mass 39.1)
 Isotopes were not given separate places in the periodic table although Mendeleev's
classification is based on the atomic masses.
 Some similar elements are separated and dissimilar elements are grouped together, e.g.
Copper and Mercury resembled in their properties but had been placed in different
groups. On the other hand Lithium and Copper were placed together although their
properties are quite different.
 Mendeleev did not explain the cause of periodicity among the elements.
 Lanthanoids and Actinoids were not given a separated position in the table.
Moseley performed experiments and studied the frequencies of the X-rays emitted from the
elements. With these experiments he concluded that atomic number is more fundamental
property of an element than its atomic mass.
MODERN PERIODIC TABLE
Main Features
 The Chemical & physical properties of elements are a periodic function of their atomic
numbers.
This relationship is called as “Modern Periodic Law”.
 Modern periodic table is also referred to as long form of periodic table.
 In the modern periodic table there are 7 Periods and 18 Groups.
Structural Features of the periodic Table
 Horizontal Rows in the periodic table are called Periods.
 Vertical columns in the periodic table are called Groups.
PERIODS
 The period number corresponds to principal quantum number (n).
 First Period contains 2 elements, and the subsequent periods consist of 8, 8, 18, 18 &
32 elements. The Numbers 2, 8, 8, 18, 18 & 32 are called magic numbers.
 In all the elements present in a period, the electrons are filled in the same valance
shell & are filled progressively in a period.

The First three periods are known as Short periods.
The Next three periods are known as Long Periods.
Seventh Period is an Incomplete Period.
GROUPS
 Elements in the same group have same number of valance electrons.
 Number of shell increases progressively in a group.
GROUP 1 -ALKALI METALS
हलीना ने की रब से फररयाद
H Li Na K Rb Cs Fr
GROUP 2 -ALKALINE EARTH METAL
बेटा माांगे कार सेंट्रो बाप रोये
Be Mg Ca Sr Ba Ra
GROUP 13 -BORON FAMILY
बेंगन आलू गाजर In थाली
B Al Ga In Tl
GROUP 14 -CARBON FAMILY
सस्सी गयी शाने पांजाब
C Si Ge Sn Pb
GROUP 15 –PINCOGENS (Cause suffocation)
नहीां पसांद ऐसे सब भाई
N P As Sb Bi
GROUP 16 –CHALCOGENS (Ore forming)
उस से टेपो
O S Se Te Po
GROUP 17 –HALOGENS (Salt forming)
फफर कल बाहर आयी आांटी
F Cl Br I At

GROUP 18 -NOBLE GASES
हीना नीना और करीना का X-Ray रांगीन
He Ne Ar Kr Xe Rn
BLOCKS
Elements are classified in blocks depending on the type of atomic orbital that are being filled
with Electrons.
s- BLOCK
s-block elements
The elements in which last electron enters into s- subshell are called s-block elements.
These elements are called Representative Elements.
Electronic Configuration- 𝒏𝒔 𝟏−𝟐
Group Included -1 & 2
p- BLOCK
p-block elements
The elements in which last electron enters into p- subshell are called p-block elements.
These elements are also called Representative Elements.
Electronic Configuration- 𝒏𝒔 𝟐
𝒏𝒑 𝟏−𝟔
Group Included- 13 – 18
d- BLOCK
d-block elements
The elements in which last electron enters into d- subshell are called d-block elements.
These elements are called Transition Elements.
Electronic Configuration- (𝒏 − 𝟏)𝒅 𝟏−𝟏𝟎
𝒏𝒔 𝟏−𝟐
Group Included- 3 – 12
f- BLOCK
f-block elements
The elements in which last electron enters into f- subshell are called f-block elements.
These elements are called Inner transition Elements. The elements after Uranium are called
Transuranic Element
Electronic Configuration- (𝒏 − 𝟐)𝒇 𝟏−𝟏𝟒(𝒏 − 𝟏)𝒅 𝟎−𝟏
𝒏𝒔 𝟐
Group Included- 3
Series included- 4f & 5f
Add to your knowledge:
1. Locating the position of an element in the periodic table by electronic configuration.
 The Principal quantum number of the valance shell represents the Period of the element.
 The sub-shell in which the last electron is filled corresponds to the block of the element.
 For s-block elements- group no. is equal to no. of ns electrons.
 For p-block elements- group no. is equal to 10 + no. of ns & np electrons.
 For d-block elements- group no. is equal to sum of the no. of (n-1) d & ns electrons.
 For f-block elements- group no. is 3.

e.g. The position of an element 1𝑠2
2𝑠2
2𝑝6
3𝑠2
3𝑝6
4𝑠2
3𝑑7
Principal quantum number for the valance shell is 4, so the element belong to Period No. 4
Last electron has been filled in 3d sub-shell, so the element belongs to d-block.
For d-block, group no. equal to 2 + 7=9
2. Locating the position of an element in the periodic table by Atomic number.
Period No.
No. of Elements
Present
Atomic Number
range
First 2 1-2
Second 8 3 - 10
Third 8 11 - 18
Fourth 18 19 - 36
Fifth 18 37 – 54
Sixth 32 55 – 86*
Seventh 32 87 – 118*
* The elements of atomic number range 58-71 belong to f-block have group no. 3rd
are called
LANTHANOIDS.
* The elements of atomic number range 90-103 belong to f-block have group no. 3rd
are called
ACTIANOIDS.
 For the given atomic number, find the Period number by the atomic number range.
 For group number, count the no. from the extreme left atomic number as first in the
atomic number range (you can count from the extreme right atomic number also). The no.
at which that atomic number exists is the group number of that element.
Remember: The element with atomic number 2 has group number 18.
The element with atomic number 5 to 10 or 13 to 18 has group number
equal to the number we have counted from extreme left atomic number plus 10.
3. Find the electronic configuration of an element in the periodic table by the position.
[ 𝑋] 𝑛𝑙 𝑥
 Firstly find the position of the element.
 The period to which the element belongs indicates the value of principal quantum number
(n) of the valance shell.
 The block to which the elements belong indicates the value of l.
 The noble gas before the element of which the electronic configuration to be find
indicates [X].
 The x can be identified by the group number as group has same outer most shell
configuration.
Group No. Elect. Conf. Group No. Elect. Conf. Group No. Elect. Conf.
1 ns1
7 ns2
nd5
13 ns2
np1
2 ns2
8 ns2
nd6
14 ns2
np2
3 ns2
nd1
9 ns2
nd7
15 ns2
np3
4 ns2
nd2
10 ns2
nd8
16 ns2
np4
5 ns2
nd3
11 ns1
nd10
17 ns2
np5
6 ns1
nd5
12 ns2
nd10
18 ns2
np6

IUPAC NOMENCLATURE OF ELEMENTS WITH ATOMIC NUMBER > 100
Digit 0 1 2 3 4 5 6 7 8 9
Name nil un bi tri quad pent hex sept oct enn
Abb. n u b t q p h s o n
e.g. The IUPAC name of the element having atomic number 109 is
Un + nil + enn + ium = Unnilennium (Une)
METALS, NON-METALS & METTLOIDS
 Elements which are electropositive in nature or have tendency to lose the electron(s) are
called Metals.
 Elements which are electronegative in nature or have tendency to gain the electron(s) are
called Non-Metals.
 Elements showing properties of both Metals & Non-metals are called Metalloids or
Semimetals.
 Most of the elements are metallic in nature while the non-metals are comparatively less.
 Metals are present on the left side as well as in the centre of the periodic table.
 Non-metals are present to the right top of the periodic table.
 Metalloids are present on the border line of metals & non-metals.
 Metallic character increases down the group & decreases along the period.
 Non-metallic character increases along the period & decreases down the group.
 All the noble gases are Non-metallic in nature.
 Hydrogen is the Non-metal.
PERIODIC TRENDS
This repetition of the same valance shell electronic configuration of the elements present in the
group & separated by the definite gaps of the atomic numbers is defined as periodicity.
ATOMIC RADII
 Covalent radius for a homo-nuclear molecule is defined as one half of the distance
between the centers of nuclei of two similar atoms bonded by single covalent bond.
 For hetero-nuclear molecule covalent radius may be defined as the distance between the
centre of nucleus of atom and mean position of the shared pair of electrons between the
bonded atoms.
 Metallic radius is defined as the one half of the inter-nuclear distance two neighbouring
atoms of a metal in a metallic lattice.
For simplicity term atomic radius is used for both covalent and metallic radius depending on
whether element is non-metal or a metal.
 Van der Waals radius is half of the distance between two similar atoms in separate
molecules in a solid.
Van der Waals radius > Metallic radius > Covalent radius
 Ionic radius may be defined as the effective distance from the nucleus of the ion up to
which it has an influence in the ionic bond.

Periodicity
 Across the periods- decreases
Explanation: The outer electrons are in the same valance shell & the effective nuclear
charge increases as the atomic number increases resulting in increased attraction of
electron to the nucleus.
 Down the group- increases
Explanation: No. of shell increases.
1. Among Iso-electronic species, the size of the anion is more than the parent atom & the size of
cation is less than the parent atom.
 Size of anion increases with increase in the magnitude of negative charge.
 Size of cation decreases with increase in the magnitude of positive charge.
2. The atomic radii of noble gases cannot be compared here.
3. Helium and francium are smallest and largest atoms respectively.
4. H–
and I–
ions are the smallest and largest anions respectively.
5. H+
and Cs+
ions are the smallest and largest cations respectively.
IONIZATION ENTHALPY
The energy required to remove an electron from an isolated gaseous atom in ground state is
called Ionization enthalpy.
M→𝑀+
+ 𝑒−
ΔiH1 (First Ionization Enthalpy)
The energy required to remove second most loosely bound electron is called second ionization
energy.
𝑀+
→𝑀2+
+ 𝑒−
ΔiH2 (Second Ionization Enthalpy)
Remember: Energy is always required to remove electrons from an atom and hence ionization
enthalpies are always positive.
Factor Affecting
 Size of atom or ion (I.E. increases with decreasing the size)
 Screening or shielding effect (I.E. increases with decreasing screening effect)
 The shielding of outermost electrons from the attractive forces of nucleus
by the inner shell electrons is called shielding effect or screening effect.
 Nuclear Charge (I.E. increases with increasing nuclear charge)
Remember: Second ionization enthalpy will be higher than the first ionization enthalpy (ΔiH2>
ΔiH1) because it is more difficult to remove an electron from a positively charged
ion than from a neutral atom as size of Cation is smaller than Parent atom. In the
same way the third ionization enthalpy will be higher than the second and so on.
Periodicity
 Across the periods- increases
Explanation: Increased nuclear charge & decrease in atomic size. Both the factors
increase the force of attraction towards nucleus and consequently, more and more energy
is required to remove the electrons.
Remember: The value of ionization enthalpy increases with breaks where the atoms
have somewhat stable configurations.
 Down the group- decreases
Explanation: Atomic Size Increases.

Unit kJ/mol or eV
 Electron volt is the energy acquired by an electron while moving under a potential
difference of one volt.
 Helium has the maximum ionization enthalpy.
 Cesium or francium has the lowest ionization enthalpy.
ELECTRON GAIN ENTHALPY
The amount of energy released when an electron is added in an isolated gaseous atom is called
Electron Gain enthalpy.
X(g) + 𝑒−
→𝑋−
(g) ΔegH (Electron Gain Enthalpy)
Remember: Electron gain enthalpies can be positive or negative. When energy is released
enthalpy will negative (in most of the cases) & when the energy is required
enthalpy will be positive.
Factor Affecting
 Size of atom or ion (EG.E. increases with increasing the size)
 Nuclear Charge (EG.E. increases with increasing nuclear charge)
Remember: First EG.E. is negative while the other successive EG.E. will be positive due to
repulsion between electrons already present in the anion & the electron being
added.
Periodicity
 Across the periods- increases (becomes more & more negative)
Explanation: Effective nuclear charge increases from left to right across a period and
consequently it will be easier to add an electron to a smaller atom since the added
electron on an average would be closer to the positively charged nucleus.
 Down the group- decreases (becomes less & less negative)
Explanation: Size of the atom increases and the added electron would be farther from the
nucleus.
Remember:
 Though Fluorine has smaller size than Chlorine, it’s EG.E. is less as compared to
Chlorine because due to small size the added electron experiences inter-electronic
repulsion & the incoming electron does not feel so much attraction.
 The elements with higher ionization enthalpy have higher negative electron gain
enthalpy.
 Chlorine has the maximum negative electron gain enthalpy.
 Noble gases have positive electron gain enthalpy.
 EG.E. is low for the element having half-filled & fully-filled orbitals.
ELECTRONEGATIVITY
The tendency of an atom to attract the shared pair electrons towards itself in covalent bond is
called electro-negativity.
Remember: Electro negativity is not a measurable quantity. However, a number of numerical
scales have been developed to find electro-negativity of elements. Pauling scale is widely used.
H X

Factor Affecting
 Size of atom (electro-negativity increases with decreasing the size)
 Charge on the ion (electro-negativity increases with increasing the magnitude of charge)
Periodicity
 Across the periods- increases
Explanation: The attraction between the valance electron(s) & the nucleus increases as
atomic size decreases in period.
 Down the group- decreases
Explanation: Atomic size increases.
Remember: Metals have low electro-negativities and non-metals have high electro-negativities.
Electron Gain Enthalpy Electro-negativity
Electron gain enthalpy is the amount
of energy released when an electron
is added in an isolated gaseous atom.
Electro-negativity is the tendency of an
atom to attract the shared pair electrons
towards itself in covalent bond.
Measured in kJ/mol Unit-less & can be find by Pauling scale
Measures the amount of energy Measures the ability to gain electron
Can be either positive or negative Always a positive value
Add to your knowledge: Fluorine is the most electro-negative element.
PERIODIC TRENDS IN CHEMICAL PROPERTIES
VALENCY
The number of electrons which an atom loses or gains or shares with other atom to attain the
noble gas configuration is termed as its valency.
The electrons present in outermost shell are called as valence electron.
The oxidation state of an element in a given compound may be defined as the charge acquired
by its atom on the basis of electro-negativity of the other atoms in the molecule.
Remember:
 Nowadays the term Oxidation state is frequently used as valance.

 In case of representative elements, the valence (Oxidation State) of an atom is generally
equal to either the number of valence electrons (s- and p-block elements) or equal to eight
minus the number of valence electrons.
Group 1 2 13 14 15 16 17 18
No. of Valance electron 1 2 3 4 5 6 7 8
Valance 1 2 3 4 3,5 2,6 1,7 0,8
 Transition and inner-transition elements, exhibit variable valence due to involvement of
not only the valence electrons but d- or f-electrons as well. However, their most common
valence are 2 and 3.
Periodicity
 Across the periods- first increase & then decrease.
 Down the group- remain same
CHEMICAL REACTIVITY
Reactivity of Metals (Metallic Character)
Tendency of metal to lose electrons from their outermost shell is called its reactivity.
Periodicity
 Across the periods- decrease
 Down the group- increase
Reactivity of Non-metals (Non-metallic Character)
Tendency of Non-metal to gain electron is called its reactivity.
Periodicity
 Across the periods- increase
 Down the group- decrease
1. Solubility of alkali metals carbonates and bicarbonates
Periodicity
2. Solubility of alkaline earth metal hydroxides and sulphates
Periodicity
3. Basic strength of alkaline earth metal hydroxides
Periodicity
DIAGONAL RELATIONSHIP
Some elements of certain groups of 2nd
period resemble much in properties with the elements of
3rd
period of next group i.e. elements of 2nd
& 3rd
period are diagonally related in properties. This
relationship is called diagonal relationship.
2nd
period Li Be B C
3rd
period Na Mg Al Si

Cause
 Similar atomic size
 Similar Ionization Enthalpy
 Similar Electro-negativity
Consequence
Anomalous behaviour of second period element
Reasons
 Small size of these atoms
 High electro-negativity
 Large charge/radius ratio
 These elements also have only 4 valence orbitals available (2s and 2p) for
bonding as compared to the 9 available (3s, 3p, and 3d) to the other members of
the respective groups, so their maximum co-valency is 4.
INERT PAIR EFFECT
Down the group, the tendency of s-block electrons to participate in chemical bonding decreases.
This effect is known as inert pair effect.
Cause
Poor shielding of the ns2
electrons by the d & f electrons, as a result ns2
electrons are held tightly
by the nucleus & cannot participate in bonding.
Consequence
Lower period elements do not show ‘Group oxidation state - 2’ oxidation state.
LANTHANOID CONTRACTION
The gradual decrease in atomic & ionic size of lanthanoides with increase in atomic number is
called lanthanoid contraction.
Cause
The 4f electrons are ineffective in screening the outer electrons from the nucleus causing
imperfect shielding.
Consequence
The atomic radii of second row transition elements are almost similar to those of the third row
transition elements.
Reasons
The increase in size on moving down the group from second to third transition elements
is cancelled by the decrease in size due to the lanthanide contraction.

Classification of elements and periodicity in properties

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