Why does hydrogen create a placement problem?

Hydrogen has one valence electron like alkali metals but also needs one to complete its shell like halogens. It forms both H+ and H-, resembling both groups partially but neither fully.

How do you find the group from electronic configuration?

For s-block: group = valence electrons. For p-block: group = 10 + valence electrons. For d-block: group = electrons in (n-1)d + ns. For f-block: always Group 3.

Classification Of Elements

Q: How many elements are in each period?

Period 1 has 2 elements, periods 2 and 3 have 8 each, periods 4 and 5 have 18 each, and periods 6 and 7 have 32 each including 14 lanthanides and 14 actinides.

Q: What are s, p, d, f blocks?

Blocks are defined by which sub-shell receives the last electron: s-block (groups 1-2), p-block (groups 13-18), d-block or transition metals (groups 3-12), and f-block or inner transition metals (lanthanides and actinides).

ChemistryClassification Of Elements And Periodicity In PropertiesFor NEET aspirants

The classification of elements is the systematic arrangement of all known elements based on the periodicity of their properties. From Dobereiner's triads (1817) to Mendeleev's atomic-weight table (1869) and finally to Moseley's modern periodic law based on atomic number (1913), the periodic table groups elements into 7 periods and 18 groups. Elements are further classified by the sub-shell of the last differentiating electron into four blocks: s, p, d, and f. This arrangement predicts an element's chemical behaviour, valency, and position at a glance.

Key Formulas & Rules - Quick Reference

Modern Periodic Law (Moseley): Physical and chemical properties of elements are periodic functions of their atomic numbers, $Z$ .
Period number = principal quantum number $n$ of the valence (outermost) shell.
Block = sub-shell into which the last differentiating electron enters (s, p, d, or f).
Group (s-block): group number = number of valence electrons.
Group (p-block): group number $= 10 +$ number of valence electrons.
Group (d-block): group number $=$ electrons in $(n - 1) d + n s$ .
Group (f-block): always Group 3 (IIIB).
Number of elements per period $= 2 n^{2}$ maximum, but limited by the last stable shell filled.

1. History of the Development of the Periodic Table

All early attempts to classify the elements were based on atomic weight. Four milestones led to the modern table.

1817

Dobereiner

Law of Triads

Groups of 3 with similar properties

1866

Newlands

Law of Octaves

Every 8th element resembles the 1st

1869

Mendeleev

Periodic Table

By atomic weight; predicted gaps

1913

Moseley

Modern Periodic Law

Properties depend on atomic number Z

Figure: Historical milestones in the development of the periodic table from Dobereiner's triads (1817) to Moseley's modern periodic law (1913).

1.1 Dobereiner's Triads (1817)

Dobereiner classified elements with similar properties into groups of three called triads. In each triad, the atomic weight of the middle element was approximately the arithmetic mean of the other two.

Example triads:

Element	Li	Na	K
Atomic weight	7	23	39

Mean of Li and K $= (7 + 39) /2 = 23 =$ atomic weight of Na. Similar triads: (Ca, Sr, Ba) and (Cl, Br, I).

Drawback: The triad concept could be applied only to a very limited number of elements.

1.2 Newlands' Law of Octaves (1866)

Newlands noticed that when elements are arranged in increasing order of atomic weight, every eighth element shows properties similar to the first - like musical notes repeating in an octave (sa, re, ga, ma, pa, dha, ni, sa).

sa	re	ga	ma	pa	dha	ni	sa
Li	Be	B	C	N	O	F	Na
Na	Mg	Al	Si	P	S	Cl	K

Drawback: The law failed beyond calcium (atomic weight $> 40$ ), because after calcium the pattern of eight breaks down.

1.3 Lothar Meyer's Atomic Volume Curve (1869)

The German chemist Lothar Meyer plotted atomic volume (atomic weight / density in the solid state) against atomic weight and found that elements with similar properties occupy similar positions on the curve - showing that properties are a periodic function of atomic weight.

1.4 Mendeleev's Periodic Law (1869)

Dmitri Mendeleev, a Russian chemist, used a much broader range of physical and chemical properties. His Periodic Law stated:

"The physical and chemical properties of elements are a periodic function of their atomic weights."

2. Mendeleev's Periodic Table

Mendeleev arranged the 63 elements known at the time into 7 horizontal rows (periods) and 8 vertical columns (groups) numbered I to VIII. Groups I to VII were further split into sub-groups A and B, and Group VIII was reserved for three sets of transition metals.

2.1 What Mendeleev did that was genius

Placed elements in order of atomic weight but excluded some that didn't fit and gave them independent positions.
Left gaps for undiscovered elements.
When properties didn't match the group, he named them by prefixing "Eka" to the preceding element. His predictions for Eka-aluminium (later Gallium) and Eka-silicon (later Germanium) matched real properties almost exactly.

2.2 Uses of Mendeleev's Periodic Table

Systematic study: Provided the first framework to study 63 elements together.
Atomic weight determination: Using $Atomic weight = Valency \times Equivalent weight$ , and valency $=$ group number, atomic weights of new elements were fixed.
Atomic weight correction: Beryllium's atomic weight had been calculated as $3 \times 4.5 = 13.5$ (assuming valency 3). Mendeleev showed this didn't fit; correcting valency to 2 gave $2 \times 4.5 = 9$ .
Discovery of new elements: Predicted Ga, Ge, Sc from gaps in the table.

2.3 Defects of Mendeleev's Periodic Table

Position of hydrogen is uncertain - it resembles both Group IA (alkali metals) and Group VIIA (halogens).
No separate positions given to isotopes.
Unclear placement of lanthanides and actinides.
Sub-groups A and B of a group have no resemblance except valency but were placed together.
Order of increasing atomic weight is not strictly followed. For example, cobalt (Co, at.wt. 58.9) was placed before nickel (Ni, at.wt. 58.7), and Ar (39.9) before K (39.1) - a violation of the strict weight order.

3. Long Form of the Periodic Table (Moseley's Table)

The table used today is the Long Form, based on Henry Moseley's discovery (1913) that the true periodic property is atomic number, not atomic weight. Moseley showed that the frequency of characteristic X-rays emitted by an element is proportional to $Z^{2}$ , fixing atomic number as the fundamental identity of an element.

Modern Periodic Law (Moseley): The physical and chemical properties of elements are a periodic function of their atomic numbers. When elements are arranged in the increasing order of atomic number, elements with similar properties recur at regular intervals.

Figure: Modern long-form periodic table with elements color-coded by block: s-block (groups 1-2), p-block (groups 13-18), d-block (transition metals, groups 3-12), and f-block (lanthanides and actinides).

3.1 Characteristics of Periods

The 1st period is the shortest - only 2 elements (H, He).
The 2nd and 3rd periods are short periods - 8 elements each.
The 4th and 5th periods are long periods - 18 elements each.
The 6th period contains 32 elements (including 14 lanthanides placed separately below).
The 7th period is incomplete and also contains 14 actinides placed separately.
Lanthanides and actinides are related to Group 3 (IIIB).
Third-period elements from Na to Cl are called typical elements; second-period elements are called bridge elements.
In a period, valency with respect to oxygen increases from 1 to 7: Na₂O, MgO, Al₂O₃, SiO₂, P₂O₅, SO₃, Cl₂O₇.

3.2 Characteristics of Groups

Moseley's table has 16 groups: I to VIII plus zero (noble gases). In modern IUPAC notation these become groups 1-18.
Inert (noble) gases occupy the zero group (Group 18).
Valency of an element in a group equals the group number.
Groups whose elements resemble typical elements are called normal elements: IA, IIA, IIIA, IVA, VA, VIA, VIIA and zero.
Groups whose elements do not resemble typical elements are called transition elements: IB, IIB, IIIB, IVB, VB, VIB, VIIB and VIII.

Solved Example 1

There are 2, 8 and 8 elements in the first, second and third periods of the periodic table respectively. Explain.

Solution:

In the 1st period, only the $1 s$ orbital fills; it holds a maximum of 2 electrons, so the period has 2 elements. In the 2nd period, $2 s$ and $2 p$ fill: capacity $2 + 6 = 8$ electrons, hence 8 elements. In the 3rd period, $3 s$ and $3 p$ fill: capacity $2 + 6 = 8$ electrons, hence 8 elements. The $3 d$ orbital is filled only after $4 s$ , so it appears in the 4th period, not the 3rd.

4. Cause of Periodicity

The reason properties repeat at regular intervals is the repetition of the outermost electronic configuration. All elements in a group have the same outer-shell configuration - only the principal quantum number $n$ changes.

Figure: The cause of periodicity - alkali metals (Group 1) all end in

n s^{1}

, giving them nearly identical chemical behaviour.

Because chemical behaviour depends mainly on valence electrons, elements with the same valence configuration behave similarly - this is the microscopic explanation of periodicity.

5. IUPAC Nomenclature for Elements with Z > 100

To avoid disputes over naming rights for superheavy synthetic elements, IUPAC adopted a systematic naming scheme.

The name is derived directly from the digits of the atomic number using numerical roots for 0-9 and adding the suffix -ium.
In some cases, "bi-ium" and "tri-ium" are shortened to -bium and -trium.

5.1 Numerical Roots

Digit	0	1	2	3	4	5	6	7	8	9
Root	nil	un	bi	tri	quad	pent	hex	sept	oct	enn
Symbol	n	u	b	t	q	p	h	s	o	e

5.2 Recommended Names for Some Superheavy Elements

Z	IUPAC Systematic Name	Symbol	Official Name
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	Hassium (Hs)
109	Unnilennium	Une	Meitnerium (Mt)
110	Ununnilium	Uun	Darmstadtium (Ds)

Solved Example 2

Element with atomic number 119 has not yet been discovered. What would be its IUPAC systematic name and symbol? On the basis of the periodic table, predict its outer electronic configuration and expected oxide/chloride formulas.

Solution:

Digits of 119 are 1, 1, 9 → un + un + enn + ium = Ununennium, symbol Uue.

Predicted outer configuration: $8 s^{1}$ (it should start Period 8, in Group 1, alkali metal family).

Expected formulas: oxide Uue₂O, chloride UueCl - like other alkali metals.

6. Classification into Blocks (s, p, d, f)

Based on the sub-shell into which the last differentiating electron enters, the elements are divided into four blocks.

Figure: Four blocks of the periodic table classified by the sub-shell (s, p, d, or f) into which the last electron enters.

6.1 s-Block Elements

The last electron enters the outermost s-orbital. These occupy the extreme left of the periodic table (Groups 1 and 2 = IA and IIA).

Soft metals with low melting and boiling points.
Highly metallic; reactivity increases down the group.
Highly electropositive; low ionization enthalpies.
Valency $+ 1$ (alkali metals) or $+ 2$ (alkaline earth metals).
Most impart characteristic flame colours (Li = crimson, Na = yellow, K = violet, Ca = brick-red).
Strong reducing agents; good conductors of heat and electricity.

6.2 p-Block Elements

The last electron enters the outermost p-orbital. These occupy the right side of the periodic table (Groups 13 to 18 = IIIA to VIIA and zero).

Contain both metals (e.g. Al, Sn, Pb), non-metals (e.g. C, N, O), and metalloids (e.g. B, Si, As).
Metallic character decreases across a period and increases down a group.
Mostly form covalent compounds.
Ionization energy is higher than s-block elements.
Reducing character increases down a group; oxidizing character increases across a period.

6.3 d-Block Elements (Transition Metals)

The last electron enters the $(n - 1) d$ sub-shell. Located in the middle of the periodic table between s- and p-blocks (Groups 3 to 12 = IIIB to VIIB, VIII, IB, IIB).

Hard, ductile, malleable metals with high melting and boiling points.
Ionization energy between s- and p-block values.
Show variable oxidation states.
Good conductors of heat and electricity.
Form both ionic and covalent compounds; compounds are usually coloured and paramagnetic.
Most transition metals form alloys (e.g. bronze, brass, steel).

6.4 f-Block Elements (Inner Transition Metals)

The last electron enters the $(n - 2) f$ sub-shell. Placed separately below the main table; related to Group 3 (IIIB).

Two series: Lanthanides (Ce-Lu, $4 f$ filling) and Actinides (Th-Lr, $5 f$ filling).
Heavy metals with high melting and boiling points.
Show variable oxidation states; compounds are generally coloured.
Most actinides are radioactive; elements beyond U (Z > 92) are synthetic transuranics.

7. Predicting Group, Period and Block from Electronic Configuration

Given only the atomic number, you can locate any element on the periodic table using three simple rules:

Period = principal quantum number $n$ of the valence shell (the shell with the highest $n$ that contains any electron).
Block = the sub-shell in which the last differentiating electron entered.
Group is computed from block-specific rules.

Group-number rules

Block	Group number rule
s-block	= number of electrons in the valence shell
p-block	$= 10 +$ number of electrons in the valence shell
d-block	= electrons in $(n - 1) d$ + electrons in $n s$
f-block	always Group 3 (IIIB)

Solved Example 3

Predict the period, group number and block of the following elements:
A (

Z = 8

), B (

Z = 11

), C (

Z = 28

), D (

Z = 54

Solution:

A (Z = 8): Configuration $1 s^{2} 2 s^{2} 2 p^{4}$ . Last electron in $2 p$ → p-block. Valence shell $n = 2$ , so Period 2. Valence electrons $= 2 + 4 = 6$ , so group $= 10 + 6 =$ 16 (VIA).

B (Z = 11): Configuration $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{1}$ . Last electron in $3 s$ → s-block. $n = 3$ → Period 3. Valence electrons $= 1$ , so group $=$ 1 (IA).

C (Z = 28): Configuration $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{8} 4 s^{2}$ . Last electron in $3 d$ → d-block. $n = 4$ → Period 4. Group $= 8 + 2 =$ 10.

D (Z = 54): Configuration ends $\dots 5 s^{2} 5 p^{6}$ . Last electron in $5 p$ → p-block. $n = 5$ → Period 5. Valence electrons $= 2 + 6 = 8$ , so group $= 10 + 8 =$ 18.

Solved Example 4

What is the atomic number of the element having the maximum number of unpaired

2 p

electrons? To which group does it belong?

Solution:

Maximum unpaired $2 p$ electrons $= 3$ (Hund's rule, half-filled $2 p^{3}$ ). This is Nitrogen, $Z = 7$ , configuration $1 s^{2} 2 s^{2} 2 p^{3}$ . It is a p-block element; group $= 10 + 5 =$ 15 (VA).

8. Bohr's Classification of Elements

Niels Bohr proposed an alternative classification based on the number of incomplete electron shells. Elements fall into four categories.

Figure: Bohr's classification of elements based on how many outer shells remain incomplete.

8.1 Inert (Noble) Gases

Both $s$ and $p$ orbitals of the outermost shell are completely filled.
Outermost configuration $= n s^{2} n p^{6}$ (except He, which is $1 s^{2}$ ).
Chemically inert; monoatomic gases at room temperature.

8.2 Representative (Normal) Elements

Only the outermost shell is incomplete; inner shells are complete.
Outermost shell has fewer than 8 electrons.
Comprise all s- and p-block elements except the noble gases.

8.3 Transition Elements

The last two shells (outermost and penultimate) are incomplete.
Last shell contains 1 or 2 electrons; penultimate shell contains 8-18 electrons.
Outer configuration $= (n - 1) d^{1 - 10} n s^{1 - 2}$ - the d-block.
Modern definition: elements with partly filled $d$ -orbitals in the ground state or in any stable oxidation state. By this definition Zn, Cd, Hg (with $d^{10}$ ) are not counted as true transition elements.

8.4 Inner Transition Elements

The last three shells (outermost, penultimate, pre-penultimate) are all incomplete.
Related to Group 3 (IIIB).
Outer configuration $= (n - 2) f^{1 - 14} (n - 1) d^{0 - 1} n s^{2}$ .
Elements of the 7th period beyond U (i.e. $Z > 92$ ) are called transuranics - all synthetic.

Solved Example 5

Boron is a metalloid. Comment.

Solution:

Boron sits at the boundary between metals (left) and non-metals (right) in the periodic table. It shows some metallic properties (e.g. lustrous surface, forms cations in fused salts) as well as clear non-metallic behaviour (poor conductivity, tendency to form covalent bonds, acidic oxide $B_{2} O_{3}$ ). This dual character defines a metalloid.

Solved Example 6

Why are inert gases monoatomic?

Solution:

Inert gases have a completely filled octet ( $n s^{2} n p^{6}$ ; He has $1 s^{2}$ ), giving them zero tendency to lose, gain, or share electrons. With no unpaired electrons and no available orbitals for bonding, they exist as isolated single atoms rather than as diatomic molecules.

Common Mistakes to Avoid

Watch out

Confusing atomic weight and atomic number as the periodic property. Mendeleev used atomic weight; the modern periodic law (Moseley) uses atomic number. Small confusion here changes your entire reasoning about the pair reversals (Co/Ni, Ar/K, Te/I).
Placing helium in Group 2 instead of Group 18. He has $1 s^{2}$ (looks like s-block) but chemically behaves as a noble gas - always place it with Group 18.
Treating "block" and "group" as the same thing. Block is the sub-shell filled last; group is a column number. An s-block element is in Group 1 or 2, but not every element in Group 1 or 2 is only s-block (e.g. H is a special case).
Forgetting that Group 3 for d-block starts at Sc, not at B. Group 3 in the modern IUPAC (long form) contains Sc, Y, La, Ac - not boron.
Adding " $10 +$ " to s-block group numbers. The $+ 10$ correction is only for p-block; for s-block, group = valence electrons directly.
Skipping the noble-gas core in IUPAC name derivation. Numerical roots come from the actual digits of $Z$ , so 119 = un-un-enn-ium, not "one-hundred-nineteen-ium".

Frequently Asked Questions

Q1. What is the modern periodic law?

The modern periodic law, given by Henry Moseley in 1913, states that the physical and chemical properties of elements are a periodic function of their atomic numbers. Atomic number, not atomic weight, is the fundamental property that determines an element's identity and chemical behaviour.

Q2. How many elements are in each period of the modern periodic table?

Period 1 has 2 elements, periods 2 and 3 have 8 each (short periods), periods 4 and 5 have 18 each (long periods), period 6 has 32 elements including 14 lanthanides, and period 7 also has 32 elements including 14 actinides. The number equals the electrons needed to complete each shell's filling sequence.

Q3. What are s, p, d, f blocks?

The four blocks are defined by which sub-shell receives the last differentiating electron: s-block (Groups 1 and 2), p-block (Groups 13 to 18), d-block or transition metals (Groups 3 to 12), and f-block or inner transition metals (lanthanides and actinides placed below the main table).

Q4. Why is Mendeleev's periodic table considered a great achievement despite its defects?

Mendeleev's genius was in leaving gaps for undiscovered elements and predicting their properties with remarkable accuracy - Eka-aluminium (Gallium), Eka-silicon (Germanium), and Eka-boron (Scandium) were all later discovered with properties matching his predictions. He also corrected several atomic weights based on periodic trends.

Q5. Why does the position of hydrogen create a problem?

Hydrogen has one valence electron like Group 1 alkali metals, but it also needs one electron to complete its shell like Group 17 halogens. It forms $H^{+}$ (like alkali metals) as well as $H^{-}$ (hydride, like halogens). Because it fits partially into both groups but fully into neither, its position remains ambiguous.

Q6. How do you find the group of an element from its electronic configuration?

If the last electron enters $n s$ : group = number of valence electrons (s-block). If it enters $n p$ : group = 10 + number of valence electrons (p-block). If it enters $(n - 1) d$ : group = number of electrons in $(n - 1) d + n s$ (d-block). All f-block elements are placed in Group 3.

Q7. What is the IUPAC name of element 118?

By the systematic method, 118 gives un-un-oct-ium = Ununoctium (Uuo). The official IUPAC name of element 118 is Oganesson (Og), named after Yuri Oganessian. It is the heaviest element currently discovered and completes the 7th period.

Q8. Why are lanthanides and actinides shown separately below the main periodic table?

Both series contain 14 elements in which the last electron enters an f-orbital. If they were placed inside the main table, periods 6 and 7 would become 32 columns wide - impractical for printing and reading. Placing them below preserves clarity while showing that they belong to Group 3.

Q9. What is the difference between representative and transition elements?

Representative (normal) elements are s- and p-block elements (except noble gases) - only the outermost shell is incomplete. Transition elements are d-block metals - the outermost and penultimate shells are both incomplete, giving them variable oxidation states and coloured compounds.

Previous year questions on Classification Of Elements

5 questions from past papers, each with a step-by-step solution.

Ready to master Classification Of Elements And Periodicity In Properties?

Take a full mock test, practice concept-by-concept, and get an AI-powered rank prediction — all on Fundamenthol.

Practice questions →Take a mock test