Classification of Elements & Periodicity in Properties

Modern Periodic Table

Mendeleev made a successful effort in grouping elements in the form of his periodic
table. He had many achievements, but there were many limitations in his Periodic Table
as well.

Some limitations of Mendeleev’s periodic table are listed below:

• The position of hydrogen was not justified in Mendeleev’s periodic table.

• The discovery of isotopes revealed another limitation of Mendeleev’s periodic table.

• Although Mendeleev arranged the elements in the increasing order of their atomic
masses, there were instances where he had placed an element with a slightly higher
atomic mass before an element with a slightly lower atomic mass.

The limitations of Mendeleev’s periodic table forced scientists to believe that atomic
mass could not be the basis for the classification of elements.

In 1913, Henry Moseley demonstrated that atomic number (instead of atomic mass) is
a more fundamental property for classifying elements. The atomic number of an
element is equal to the number of protons present in an atom of that element. Since the
number of protons and electrons in an atom of an element is equal, the atomic number
of an element is equal to the number of electrons present in a neutral atom.

Atomic number = Number of protons = Number of electrons

The number of protons or electrons in an element is fixed. No two elements can have
the same atomic number. Hence, elements can be easily classified in the increasing
order of their atomic numbers. In the light of this fact, Mendeleev’s Periodic Law was
done away with. As a result, the modern periodic law came into the picture.

The modern periodic law states that the properties of elements are a periodic
function of their atomic numbers, not their atomic masses.

The table that is obtained when elements are arranged in the increasing order of their
atomic numbers is called the Modern Periodic Table or Long Form of the Periodic
Table as shown in the figure.
 The Modern periodic table

In the modern periodic table, the elements are arranged in rows and columns. These
rows and columns are known as periods and groups respectively. The table consists
of 7 periods and 18 groups.

Do You Know:

In the modern periodic table, hydrogen is placed above alkali metals because of
resemblance with their electronic configurations. However, it is never regarded as an
alkali metal. This makes hydrogen a unique element.

If you look at the modern periodic table, you will find that all elements in the same group
contain the same number of valence electrons. Let us see the following activity to
understand better.

Activity 1: Look at group two of the modern periodic table. Write the name of the first
three elements followed by their electronic configurations.
What similarity do you observe in their electronic configurations? How many
valence electrons are present in these elements?

The first three elements of group two are beryllium, magnesium, and calcium. All these
elements contain the same number of valence electrons. The number of valence
electrons present in these elements is 2. On the other hand, the number of shells
increases as we go down the group.

Again, if you look at periods in the modern periodic table, you will find that all elements
in the same period contain the same valence shell. Let us see the following activity to
understand better.

Activity 2: Look at the elements of the third period of the modern periodic table. Write
the electronic configuration of each element and calculate the number of valence
electrons present in these elements.

What do you observe from the given activity? Do these elements contain the
same number of shells? How many valence electrons are present in these
elements?

You will find that elements such as sodium, magnesium, aluminium, silicon,
phosphorus, sulphur, chlorine, and argon are present in that period. The valence shell in
all these elements is the same, but they do not have the same number of valence
electrons.

Electronic configuration
Name of the element
(K, L, M)

Sodium 2, 8, 1

Magnesium 2, 8, 2

Aluminium 2, 8, 3
 Silicon 2, 8, 4

Phosphorus 2, 8, 5

Sulphur 2, 8, 6

Chlorine 2, 8, 7

Argon 2, 8, 8

Thus, the number of electrons in the valence shell increases by one unit as the atomic
number increases by one unit on moving from left to right in a period.

Let us calculate the number of elements that are present in the first, second, third, and
fourth periods.

The maximum number of electrons that a shell can hold can be calculated using
the formula 2n2. Here, n represents the number of shells from the nucleus. For
example, n is equal to 1, 2, and 3 for K, L, and M shells respectively. Hence, the
maximum number of electrons that each of these shells can hold can be calculated by
substituting the value of n in the given formula.

Number of electrons that K shell can accommodate = 2n2

=2

Hence, K shell can accommodate only 2 electrons and only two elements are present in
the first period.

Similarly, the second and third shell (L and M respectively) can accommodate 8 and 18
electrons respectively. Since the outermost shell can contain only 8 electrons, there are
only 8 elements in both the periods.
 Important Note:

The position of an element in the Modern Periodic Table tells us about its chemical
reactivity. The valence electrons determine the kind and the number of bonds formed by
an element.

IUPAC Nomenclature for Elements with Atomic Number 100

• Latin word roots for various digits are listed in the given table.

Notation for IUPAC Nomenclature of Elements

Digit Name Abbreviation

0 nil n

1 un u

2 bi b

3 tri t

4 quad q

5 pent p

6 hex h
 7 sept s

8 oct o

9 enn e

• Latin words for various digits of the atomic number are written together in the order of
digits, which make up the atomic number, and at the end, ‘ium’ is added.

• Nomenclature of elements with the atomic number above 100 is listed below.

Nomenclature of Elements with Atomic Number Above 100

Atomic IUPAC Official IUPAC

Name Symbol
number Name Symbol

101 Unnilunium Unu Mendelevium Md

102 Unnilbium Unb Nobelium No

103 Unniltrium Unt Lawrencium Lr

104 Unnilquadium Unq Rutherfordium Rf

105 Unnilpentium Unp Dubnium Db

106 Unnilhexium Unh Seaborgium Sg

107 Unnilseptium Uns Bohrium Bh

108 Unniloctium Uno Hassnium Hs

109 Unnilennium Une Meitnerium Mt

110 Ununnilium Uun Darmstadtium Ds

111 Unununnium Uuu Rontgenium Rg

112 Ununbium Uub

113 Ununtrium Uut

114 Ununquadium Uuq

115 Ununpentium Uup

116 Ununhexium Uuh

 117 Ununseptium Uus

118 Ununoctium Uuo

Electronic Configuration and the Periodic Table

Electronic Configuration in Periods

• Period indicates the value of ‘n’ (principal quantum number) for the outermost or
valence shell.

• Successive periods in the periodic table are associated with the filling of the next higher
principal energy level (n = 2, n = 3, etc).

• First period (n = 1) → hydrogen (1s1) and helium (1s2) [2 elements]

• Second period (n = 2) → Li (1s2 2s1), Be (1s2 2s2), B (1s2 2s2 2p1) to Ne (2s2 2p6) [8
elements]

• Third period (n = 3) → filling to 3s and 3p orbitals gives rise to 8 elements (Na to Ar)

• Fourth period (n = 4) → 18 elements (K to Kr) − filling of the 4s and 4p orbitals

3d orbital is filled up before 4p orbitals (3d orbitals → energetically favourable)

• 3d-transition series → Sc (3d1 4s2) to Zn (3d10 4s2)

• Fifth period (n = 5) → 18 elements (Rb to Xe)

• 4d-transition series starts at Ytterbium and ends at Cadmium.

• Sixth period (n = 6) → 32 elements; electrons enter 6s, 4f, 5d, and 6p orbitals
successively. Elements from Z = 58 to Z = 71 are called 4f-inner transition series or
lanthanoid series (filling up of the 4f orbitals).

• Seventh period (n = 7) → electrons enter at 7s, 5f, 6d, and 7p orbitals successively.
Filling up of 5f orbitals after Ac (Z = 89) gives 5f-inner transition series or the actinoid
series.

Electronic Configuration in Groups

• Same number of electrons is present in the outer orbitals (that is, similar valence shell
electronic configuration).

• Electronic configuration of group 1 elements is given in the following table.

Atomic
Symbol Electronic configuration
number

3 Li 1s2 2s1(or) [He]2s1

11 Na 1s2 2s2 2p6 3s1(or) [Ne]3s1

19 K 1s2 2s2 2p6 3s2 3p6 4s1(or) [Ar]4s1

37 Rb 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1(or) [Kr]5s1

55 Cs 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s1(or)[Xe]6s1

87 Fr [Rn]7s1

Electronic Configurations and Types of Elements

• s- Block Elements

• Group 1 (alkali metals) − ns1 (outermost electronic configuration)

• Group 2 (alkaline earth metals) − ns2 (outermost electronic configuration)
• Alkali metals form +1 ion and alkaline earth metals form +2 ion.
• Reactivity increases as we move down the group.
• They are never found in the pure state in nature. (Reason − they are highly reactive)

• p - Block Elements
• Elements belonging to Groups 13 to 18
• Outermost electronic configuration varies from ns2np1 to ns2np6 in each period.
• Group 18 (ns2np6) − noble gases
• Group 17 − halogen
• Group 16 − chalcogens
• Non-metallic character increases from left to right across a period.

• d- Block Elements (Transition Elements)

• Elements of group 3 to group 12

• General electronic configuration is (n − 1) d1−10 ns0-2.
• Called transition elements
• Zn, Cd, and Hg with (n − 1) d10 ns2 configuration do not show properties of transition
elements.
• All are metals. They form coloured ions, exhibit variable oxidation states,
paramagnetism, and are used as catalysts.

• f- Block Elements

• Lanthanoids → Ce (Z = 58) to Lu (Z = 71)

• Actinoids → Th (Z = 90) to Lr (Z = 103)
• Outer electronic configuration → (n − 2) f1−14 (n −1) d0−1 ns2
• They are called inner-transition elements.
• All are metals.
• Actinoid elements are radioactive.
• Elements after uranium are called Transuranium elements.

Metals, Non-metals, and Metalloids

• Metals → Appear on the left side of the periodic table

• Non-metals → Located at the top right-hand side of the periodic table

• Elements change from metallic to non-metallic from left to right.

• Elements such as Si, Ge, As, Sb, Te show the characteristic properties of both metals
and non-metals. They are called semi-metals or metalloids.

Periodic Trends in Physical Properties

Atomic Radius

• Atomic radii decrease with the increase in the atomic number in a period.

• For example, atomic radii decrease from Li to F in the second period.

• Nuclear charge increases progressively by one unit on moving from left to right across
the period. As a result, the electron cloud is pulled closer to the nucleus by the
increased effective nuclear charge, which causes decrease in atomic size.

• Atomic radii increase from top to bottom within a group of the periodic table.

• Variation of atomic radii with atomic number among alkali metals and halogen:

Ionic Radius

• Cation is smaller than its parent atom.

• The size of the anion is larger than its parent atom.

Ionization Enthalpy
• Defined as the amount of energy required to remove the most loosely bound electron
from the isolated gaseous atom in its ground state

• Decreases with the increase in atomic size

• Increases with the increase in nuclear charge

• Decreases with the increase in the number of inner electrons

• Increases with the increase in penetration power of electrons

• Atom having a more stable configuration has high value of enthalpy.

• Variation across a period: Increases with the increase in atomic number across the
period.

• Variation in a group: Decreases regularly with the increase in atomic number within a
group.

Electron Gain Enthalpy

• Defined as the enthalpy change taking place when an isolated gaseous atom accepts
an electron to form a monovalent gaseous anion

• Larger the value of electron gain enthalpy, greater is the tendency of an atom to accept
electron.

• Greater the magnitude of nuclear charge, larger will be the negative value of electron
gain enthalpy.
• Larger the size of the atom, smaller will be the negative value of electron gain enthalpy.

• More stable the electronic configuration of the atom, more positive will be the value of
its electron gain enthalpy.

• Variation across a period − Tends to become more negative as we go from left to right
across a period

• Variation down a group − Becomes less negative on going down the group

Electronegativity

• Defined as the tendency of an atom in a molecule to attract the shared pair of electrons
towards itself

• Greater the effective nuclear charge, greater is the electronegativity.

• Smaller the atomic radius, greater is the electronegativity.

• In a period − Increases on moving from left to right

• In a group − Decreases on moving down a group

Valency

• It is defined as the number of univalent atoms which can combine with an atom of the
given element.

• Valency is given by the number of electrons in outermost shell.

• If the number of valence electrons ≤4: valency = number of valence electrons

• If the number of valence electrons >4: valency = (8 - number of valence electrons)

• In a period − Increases from 1 to 4 and then decreases from 4 to zero on moving from
left to right

• In a group − No change in the valency of elements on moving down a group. All

elements belonging to a particular group exhibit same valency.

Non −Metallic (and Metallic Character) of an Element

• Non-metallic elements have strong tendency to gain electrons.

• Non-metallic character is directly related to electronegativity and metallic character is
inversely related to electronegativity.

• Across a period, electronegativity increases. Hence, non-metallic character increases

(and metallic character decreases).

• Down a group, electronegativity decreases. Hence, non-metallic character decreases

(and metallic character increases).

The periodic trends of various properties of elements in the periodic table are shown in
figure.

Periodic Trends in Chemical Properties

Periodicity of Valence or Oxidation States

• Valence of the elements = Number of electrons in the outermost orbitals (if valence
electrons ≤ 4)

• Or, valency of the element = 8 − Number of outermost electrons (if valence electrons >
4)

Group 1 2 13 14 15 16 17 18

Number of valence electrons 1 2 3 4 5 6 7 8

 Valence 1 2 3 4 3,5 2,6 1,7 0,8

• The given table shows the periodic trends observed in the valence of elements
(hydrides and oxides).

Group 1 2 13 14 15 16 17

Formula of hydride LiH B2H6 CH4 NH3 H2O HF

NaH CaH2 AlH3 SiH4 PH3 H2S HCl

KH GeH4 AsH3 H2Se HBr

SnH4 SbH3 H2Te HI

Formula of oxide Li2O MgO B2O3 CO2 N2O3, N2O5

Na2O CaO Al2O3 SiO2 P4O6, P4O10 SO3 Cl2O7

K2O SrO Ga2O3 GeO2 As2O3, As2O5 SeO3

BaO In2O3 SnO2 Sb2O3,Sb2O5 TeO3

 PbO2 Bi2O3

• Many elements exhibit variable valence (particularly transition elements and actinoids).

Anomalous Properties of Second Period Elements

• First member of each group (the element in the second period from lithium to fluorine)
differs in many respects from the rest of the members of the same group.

• For example, the behaviour of Li and Be is more similar with the second element of the
following group i.e., Mg and Al respectively.

• Such sort of similarity is commonly known as diagonal relationship in periodic

properties.

• Reasons for the different chemical behaviour of the first member of a group of elements
in the s-and p-blocks as compared to the other members in the same group are as
follows.

• Small atomic size of the first element

• Large charge/radius ratio
• High electronegativity
• Absence of d-orbitals in the valence shell
• Ability of form pπ − pπ multiple bonds

• First member of each group of p-block element has the tendency to

form pπ − pπ multiple bonds to itself and to the other second period elements. For
example, C = C, C ≡ C, C = O, C = N

• Reason − This property of the elements is due to their small size.

• Higher members of the group have little tendency to form pπ − pπ bonds.

Periodic Trends and Chemical Reactivity

• High chemical reactivity at the two extremes of a period and the lowest in the centre

• Maximum chemical reactivity is at the extreme left of a period because of the ease of
electron loss (or low ionization enthalpy).

• Elements at the extreme left exhibit strong reducing behaviour and elements at the
extreme right exhibit strong oxidizing behaviour.
• Oxides formed by the elements on the left are basic and by the elements on the right
are acidic in nature.

• Oxides of elements in the centre are amphoteric or neutral.

• The electron gain enthalpy increases across a period and decreases down the group.