Present Form of the Periodic Table & Nomenclature of Super-heavy Elements

The Anatomy of the Long Form Periodic Table

The modern long form of the Periodic Table — the one hanging in every chemistry lab today — is the direct heir of Mendeleev's 1869 table, but reorganised on the basis of the Modern Periodic Law (Moseley, 1913): properties of elements are a periodic function of their atomic number $Z$.

The two dimensions

Horizontal rows — called periods: There are seven periods, numbered 1 through 7. Each period starts with an element whose outermost electron enters a new principal shell (quantum number $n$) and ends with a noble gas whose outermost shell is full. The length of each period is set by the number of orbitals that become available at that $n$:

Period	Starts with	Ends with	Length	Orbitals filled
1	H (Z = 1)	He (Z = 2)	2	$1s$
2	Li (Z = 3)	Ne (Z = 10)	8	$2s,\ 2p$
3	Na (Z = 11)	Ar (Z = 18)	8	$3s,\ 3p$
4	K (Z = 19)	Kr (Z = 36)	18	$4s,\ 3d,\ 4p$
5	Rb (Z = 37)	Xe (Z = 54)	18	$5s,\ 4d,\ 5p$
6	Cs (Z = 55)	Rn (Z = 86)	32	$6s,\ 4f,\ 5d,\ 6p$
7	Fr (Z = 87)	Og (Z = 118)	32	$7s,\ 5f,\ 6d,\ 7p$

The period number equals the highest principal quantum number ($n$) of the outermost electron. A memorable way to recall the lengths: $2, 8, 8, 18, 18, 32, 32$.

Vertical columns — called groups: There are eighteen groups, numbered 1 through 18. All members of a group have the same outer-shell (valence) electronic configuration (ignoring the minor quirks of Cr, Cu, Mo, Pd and Ag). That is why they show similar chemical behaviour — groups run "families" of elements.

Why 18 and not 8?

Mendeleev had only 8 vertical columns because he used atomic mass, did not yet understand orbitals, and lumped transition metals into sub-groups "A" and "B". The modern 18-column layout separates the s-, p-, d- and f-block elements cleanly so that chemically similar elements always line up under each other. A lithium atom has configuration $1s^{2}\,2s^{1}$; a sodium atom has $[\text{Ne}]\,3s^{1}$; a potassium atom has $[\text{Ar}]\,4s^{1}$. All three have a single $ns^{1}$ valence electron — so all three sit in Group 1 and all three behave like alkali metals.

Periods 6 and 7 & the f-block islands

Periods 6 and 7 are 32 elements long because the 4f and 5f orbitals open up. To keep the table a readable width, the 14 lanthanides (Z = 58–71) and 14 actinides (Z = 90–103) are pulled out and printed as two separate rows at the bottom, marked by the asterisks in the main table. Mentally, they belong in Period 6 and 7 respectively, inserted between the elements of Group 3.

The Four Blocks: s, p, d, f

The periodic table is divided into four blocks according to the sub-shell (orbital type) of the last-entering electron in the ground-state electronic configuration. This is the most powerful organising principle in modern chemistry — blocks decide the number of valence electrons, the typical oxidation states, and even the colour of compounds.

Which orbital is "last-entering"?

Electrons occupy orbitals in the order given by the Aufbau Principle:

$1s \to 2s \to 2p \to 3s \to 3p \to 4s \to 3d \to 4p \to 5s \to 4d \to 5p \to 6s \to 4f \to 5d \to 6p \to 7s \to 5f \to 6d \to 7p$

Whichever sub-shell type the last electron of a neutral atom occupies usually decides that atom's block.

The four blocks at a glance

Block	Groups	Valence sub-shells being filled	Typical valence config	# of elements
s	1, 2	outermost $ns$	$ns^{1-2}$	14
p	13–18	outermost $np$	$ns^{2}\,np^{1-6}$	36
d	3–12	inner $(n-1)d$	$(n-1)d^{1-10}\,ns^{0-2}$	40
f	(inside G3)	deeper inner $(n-2)f$	$(n-2)f^{1-14}\,(n-1)d^{0-1}\,ns^{2}$	28

There are 118 known elements: 14 s-block + 36 p-block + 40 d-block + 28 f-block = 118. (Hydrogen and helium belong to the s-block by electronic configuration; helium is placed in Group 18 because its chemistry is that of a noble gas.)

Two key consequences

1. Group number from block & config: For s-block elements, group number = number of valence s-electrons. For p-block elements, group number = 10 + (valence s + valence p electrons). For the d-block, group number = (valence s electrons) + ((n−1)d electrons). This is a fast way to place an element without looking at a chart.

2. Type of chemistry follows block: s-block elements are reactive metals, p-block spans metals, metalloids, non-metals and noble gases, the d-block forms transition metals (variable valency, coloured ions, catalytic activity), and the f-block holds the inner transition metals (lanthanides + actinides).

Characteristic Features of Each Block

Knowing just the block of an element tells a trained chemist almost everything about how it will behave. Here is the minimum each Class 11 student should be able to rattle off.

s-block elements — Groups 1 & 2

• Outermost configuration $ns^{1}$ (Group 1, alkali metals) or $ns^{2}$ (Group 2, alkaline-earth metals).
• All are soft, silvery, reactive metals. They are so reactive that they are stored under kerosene (Na, K) or in sealed ampoules (Cs, Fr).
• They lose their $ns$ electron(s) easily, so ionization enthalpies are low and they form ionic compounds in the +1 (Group 1) or +2 (Group 2) oxidation state — and only those states.
• They impart characteristic colours to flames ($\text{Na}^{+}$ yellow, $\text{K}^{+}$ lilac, $\text{Ca}^{2+}$ brick-red, $\text{Sr}^{2+}$ crimson, $\text{Ba}^{2+}$ apple-green), which is the basis of the flame test.
• Hydrogen ($1s^{1}$) sits atop Group 1 by configuration, but its chemistry is special — we will discuss this briefly below.

p-block elements — Groups 13 to 18

• Valence configuration $ns^{2}\,np^{1}$ to $ns^{2}\,np^{6}$.
• The most chemically diverse block: it contains all non-metals except hydrogen, all metalloids ($\text{B, Si, Ge, As, Sb, Te, Po}$), the noble gases, and a handful of post-transition metals ($\text{Al, Ga, In, Tl, Sn, Pb, Bi}$).
• Exhibit multiple oxidation states that differ by 2 (inert-pair effect). For instance Tl shows +1 & +3, Pb shows +2 & +4, Sn shows +2 & +4 — with the lower state becoming more stable on going down the group.
• Noble gases (Group 18) have the stable $ns^{2}\,np^{6}$ octet and are largely unreactive, though Xe and Kr do form compounds (e.g. $\text{XeF}_{2},\ \text{XeF}_{4},\ \text{XeO}_{3}$).

d-block elements — Groups 3 to 12 (transition metals)

• Valence configuration $(n-1)d^{1-10}\,ns^{0-2}$. Defined formally by IUPAC as elements whose atoms or whose common ions have a partially filled $d$-subshell.
• All are metals: hard, shiny, high melting points, good conductors, mostly dense (tungsten, osmium, iridium are the densest).
• Show variable oxidation states (e.g. Mn: +2, +3, +4, +6, +7) because the $ns$ and $(n-1)d$ electrons have comparable energies and can both participate in bonding.
• Coloured ions due to $d{-}d$ electronic transitions (e.g. $\text{Cu}^{2+}$ blue, $\text{Ni}^{2+}$ green, $\text{MnO}_{4}^{-}$ purple).
• Form complex ions with ligands ($\text{[Fe(CN)}_{6}]^{4-}$, $\text{[Cu(NH}_{3})_{4}]^{2+}$, etc.).
• Catalytic activity: V₂O₅ (contact process), Fe (Haber), Pt/Pd (hydrogenation), Ni (vegetable-oil hydrogenation).
• Zn, Cd, Hg (Group 12) have completely filled $d^{10}$ shells in every common oxidation state, so many textbooks do not call them true transition metals — a subtle point JEE often tests.

f-block elements — lanthanides and actinides

• Valence configuration $(n-2)f^{1-14}\,(n-1)d^{0-1}\,ns^{2}$; last electron enters the deep, inner $(n-2)f$ sub-shell.
• Placed inside Group 3 of Periods 6 and 7 but shown as two separate rows at the bottom to keep the main table compact.
• Also called inner transition elements.
• Lanthanides (Ce to Lu) — all show a dominant +3 state; they exhibit lanthanide contraction (discussed in Section 8).
• Actinides (Th to Lr) — variable oxidation states because 5f, 6d and 7s energies are close. All isotopes beyond U (Z = 92) are radioactive and man-made.

Old vs New Group Numbering: Why the Switch?

Open any Indian high-school chemistry book printed before 1990 and you will see group numbers written as IA, IIA, IIIA, IVA, VA, VIA, VIIA, VIII (three columns!), IB, IIB, IIIB, … 0. That is the old Mendeleev-style notation. After 1988 IUPAC officially recommended the unambiguous numbering 1, 2, 3, …, 18.

Three headaches with the old scheme

1. 'A' vs 'B' meant different things in different books. The American CAS convention used 'A' for main-group elements on the left and right (IA, IIA, IIIA, … VIIIA), with 'B' for the d-block in the middle (IIIB, IVB, …, VIIB, VIII, IB, IIB). The European CAS convention did almost the opposite — Group 11 (Cu, Ag, Au) was IB under one system and IA under another. Students reading imported books got a clashing set of labels for the same column.

2. Group VIII was three columns crammed into one. Iron, cobalt and nickel (the "iron triad") plus ruthenium, rhodium, palladium plus osmium, iridium, platinum were all dumped into a single Roman numeral VIII, even though they are clearly three distinct vertical families. IUPAC's 1–18 splits them into groups 8, 9 and 10.

3. The noble gases were called 'Group 0' — a hangover from the false belief that they had zero valency.

The IUPAC fix

IUPAC issued its final recommendation in 1988: number the groups simply 1, 2, 3, …, 18 from left to right. Each column now has one unambiguous number, and chemically similar elements line up under each other in the new scheme. NCERT Class 11 prints both numbering systems side by side; both are considered correct in examinations, but the 1–18 IUPAC system is preferred.

Exam tip: Always carry the old-to-new conversion in your head — questions frequently give group labels as IA or VIIA and expect you to know these are Group 1 (alkali metals) and Group 17 (halogens) respectively. The image above and the table below are your cheat-sheet.

Old	New	Family name
IA	1	Alkali metals
IIA	2	Alkaline-earth metals
IIIB	3	Scandium family
IVB	4	Titanium family
VB	5	Vanadium family
VIB	6	Chromium family
VIIB	7	Manganese family
VIII	8, 9, 10	Iron/platinum triads
IB	11	Coinage metals
IIB	12	Zinc family
IIIA	13	Boron family
IVA	14	Carbon family
VA	15	Pnictogens
VIA	16	Chalcogens
VIIA	17	Halogens
0	18	Noble gases

IUPAC Nomenclature for Elements with $Z > 100$

Super-heavy elements — those with atomic number greater than 100 — are almost all man-made. They are produced atom-by-atom in giant particle accelerators, often by bombarding a heavy target with a lighter ion. Many live only for milliseconds before decaying radioactively.

Because competing laboratories in the USA, Russia and Germany would synthesise the same element and each propose its own name, IUPAC was forced in 1977 to adopt a temporary systematic nomenclature that would give every new element a unique, agreed-upon name before the priority disputes were resolved.

The rules in full

Rule 1 — Numerical roots for each digit. Assign a short Latin/Greek root to each digit from 0 to 9:

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

Rule 2 — Write the digits of the atomic number left-to-right and concatenate the corresponding roots.

Rule 3 — Append the suffix -ium.

Rule 4 — Small spelling tweaks for readability:

• bi + ium becomes bium, tri + ium becomes trium — drop one i so we do not write biium or triium.
• The double-n of enn is kept: un + enn + ium → ununennium.
• The root nil keeps both l's.

Rule 5 — The element's symbol is the initial letter of each root, in order, capitalised appropriately (first letter upper-case, rest lower-case).

Worked example 1 — Element 104

• Digits: 1, 0, 4 → roots: un, nil, quad
• Systematic name: un + nil + quad + ium = Unnilquadium
• Symbol: Unq
• Permanent name: Rutherfordium, symbol Rf.

Worked example 2 — Element 118

• Digits: 1, 1, 8 → roots: un, un, oct
• Systematic name: un + un + oct + ium = Ununoctium
• Symbol: Uuo
• Permanent name: Oganesson, symbol Og.

Worked example 3 — Element 120 (not yet synthesised)

• Digits: 1, 2, 0 → roots: un, bi, nil
• Systematic name: un + bi + nil + ium = Unbinilium
• Symbol: Ubn

Status of super-heavy elements today

All elements up to Z = 118 (Oganesson, Og) have been officially named. The permanent names for Z = 113, 115, 117 and 118 are Nihonium (Nh), Moscovium (Mc), Tennessine (Ts) and Oganesson (Og). Elements Z = 119 and 120 are still targets of current experimental campaigns and continue to be referred to systematically as Ununennium (Uue) and Unbinilium (Ubn) until their discoveries are confirmed.

Why Systematic Names Matter — A Quick History of Priority Disputes

If you think priority disputes are an abstract worry, remember this: for three decades (~1964 to ~1997), element Z = 104 was called rutherfordium in America and kurchatovium in the USSR — at the height of the Cold War, each country refused to accept the other's name. Chemistry journals had to print papers with both names side by side.

Similar stand-offs occurred for:

• Element 105 — "hahnium" vs "nielsbohrium" vs "joliotium", before IUPAC finally chose dubnium in 1997.
• Element 106 — US vs Russian claims were resolved by naming it seaborgium (after Glenn T. Seaborg), controversial at the time because Seaborg was still alive — the first element named after a living person.
• Element 109 — named meitnerium after Lise Meitner, correcting her historical exclusion from the Nobel Prize for the discovery of nuclear fission.

The IUPAC systematic names prevent such disputes from paralysing the literature: everyone agrees on Unnilquadium (Z = 104) until the committees settle on a permanent name. In that sense the rule is not trivia — it is diplomacy.

What students should remember for exams

1. IUPAC systematic names apply only to elements with $Z > 100$, not to elements already known and named.
2. Know the ten numerical roots by heart — examiners test them by asking for the systematic name of an atomic number, or the atomic number of a systematic name.
3. Every systematic name ends in -ium.
4. The symbols are three-letter (Uuo, Unq, Uub) unlike the one- or two-letter symbols of permanent names.
5. Spelling tweaks: double-i becomes single-i when bi or tri meets ium; the root nil keeps both l's; enn keeps both n's.

Memory Capsule: Quick Revision

• Dimensions: 18 Groups (valence config) and 7 Periods (orbital filling).
• Period Lengths: $2, 8, 8, 18, 18, 32, 32$.
• Blocks:
  • s-block: Groups 1-2. Reactive metals, +1/+2 states.
  • p-block: Groups 13-18. Metals, non-metals, metalloids.
  • d-block: Groups 3-12. Transition metals, variable oxidation, coloured ions.
  • f-block: Lanthanides (4f) & Actinides (5f). Inner transition elements.
• Nomenclature ($Z>100$): Numerical roots + -ium. Symbols are 3 letters.
• IUPAC Grouping: 1-18 replaces Roman numerals to avoid ambiguity.

Solved Examples

Example 1: Length of Period 6

Why is Period 6 of the periodic table 32 elements long? Name the orbitals that fill during this period.

Solution:

Period 6 begins when the electron enters the $6s$ orbital (Cs, Z = 55) and ends when the $6p$ orbital is completely filled (Rn, Z = 86). The number of orbitals available between these points is:

Sub-shell	6s	4f	5d	6p	Total
Orbitals	1	7	5	3	16
Electrons	2	14	10	6	32

Sixteen orbitals × 2 electrons each = 32 electrons, so Period 6 contains 32 elements (Cs 55 → Rn 86). The $4f$ block (14 elements Ce → Lu) is the "stretched" part that makes Periods 6 and 7 longer than Periods 4 and 5.

Example 2: Block identification from configuration

Identify the block of the following elements: (a) $[\text{Kr}]\,4d^{10}\,5s^{1}$, (b) $[\text{Xe}]\,4f^{7}\,6s^{2}$, (c) $[\text{Ar}]\,3d^{5}\,4s^{1}$, (d) $[\text{Ne}]\,3s^{2}\,3p^{4}$.

Solution:

• (a) $[\text{Kr}]\,4d^{10}\,5s^{1}$ corresponds to Ag. Although its ground-state configuration is exceptional, Ag belongs to Group 11 of the d-block and is treated with the coinage metals Cu and Au.
• (b) $[\text{Xe}]\,4f^{7}\,6s^{2}$ is Europium, a typical f-block lanthanide.
• (c) $[\text{Ar}]\,3d^{5}\,4s^{1}$ is Chromium, an exceptional but standard d-block element.
• (d) $[\text{Ne}]\,3s^{2}\,3p^{4}$ is Sulphur, a p-block element.

Answer: (a) d-block, (b) f-block, (c) d-block, (d) p-block.

Example 3: Predicting block, period and group

An element has atomic number $Z = 35$. Without consulting a periodic table predict its (i) period, (ii) block, (iii) group and (iv) common oxidation states.

Solution:

Write the electronic configuration by applying the Aufbau order:

$$1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,4s^{2}\,3d^{10}\,4p^{5}$$

• (i) Period: highest $n$ of any occupied orbital = 4, so Period 4.
• (ii) Block: last-entering electron is in $4p$, so p-block.
• (iii) Group: for p-block, group = 10 + (valence s + valence p) = 10 + (2 + 5) = 17 (halogens).
• (iv) Oxidation states: halogens show $-1$ (most common), and $+1, +3, +5, +7$ in their oxoacids and oxoanions.

The element is bromine, which matches Group 17, Period 4.

Example 4: Block of La (NEET 2019 style)

The element with atomic number 57 belongs to (a) s-block (b) p-block (c) d-block (d) f-block.

Solution:

Z = 57 is Lanthanum, electronic configuration $[\text{Xe}]\,5d^{1}\,6s^{2}$. The last electron enters the $5d$ orbital, so strictly Lanthanum is a d-block element (Group 3, Period 6) — option (c).

Caveat: NCERT sometimes classifies La as the "first lanthanide" — a relic of older textbooks that grouped the f-block by chemistry rather than by the last-entering electron. JEE/NEET questions follow the IUPAC d-block rule, so select (c) d-block.

Example 5: IUPAC systematic name

Give the IUPAC systematic name and symbol of the element with atomic number 105.

Solution:

Split Z = 105 into digits: 1, 0, 5.

• 1 → un (u)
• 0 → nil (n)
• 5 → pent (p)

Concatenate and add -ium:

$$\text{un} + \text{nil} + \text{pent} + \text{ium} \;=\; \mathbf{Unnilpentium}$$

Symbol = Unp.

(Permanent name: Dubnium, Db.)

Example 6: From systematic name to atomic number

The IUPAC systematic symbol of a newly-synthesised element is Uuq. What is its atomic number and the likely permanent name (if any)?

Solution:

Read the three letters as initial letters of the roots:

• U → un → 1
• u → un → 1
• q → quad → 4

So Z = 114. The systematic name is Ununquadium. The permanent name is Flerovium (Fl).

Example 7: Counting elements in each block

How many elements are there in (i) the s-block, (ii) the p-block, (iii) the d-block and (iv) the f-block? Verify that the four blocks together account for all 118 known elements.

Solution:

• s-block: Groups 1 and 2, across 7 periods → $2 \times 7 = 14$. Including H and He gives 14 elements.
• p-block: Groups 13–18, across Periods 2–7 → $6 \times 6 = 36$. ⇒ 36 elements.
• d-block: Groups 3–12, across Periods 4–7 → $10 \times 4 = 40$. ⇒ 40 elements.
• f-block: 14 lanthanides + 14 actinides = 28 elements.

$$14 + 36 + 40 + 28 \;=\; 118 \;\checkmark$$

This matches the 118 elements currently on the IUPAC table.

Example 8: Group from old to new numbering (Board Level)

An element was described in an older textbook as "Group IIIA, Period 4". Give its position in the modern IUPAC notation, identify the element and state its common oxidation state.

Solution:

Old "IIIA" means the main-group column with 3 valence electrons → New Group 13.

Period 4, Group 13 ⇒ electronic configuration $[\text{Ar}]\,3d^{10}\,4s^{2}\,4p^{1}$ ⇒ Gallium (Ga), Z = 31.

Common oxidation state = +3.

Example 9: Why noble gases are in Group 18, not Group 0

Justify the IUPAC decision to call the noble-gas column "Group 18" rather than retaining Mendeleev's label "Group 0".

Solution:

Mendeleev's zero label reflected the outdated idea that noble gases were chemically inert — they had "zero" valency. However:

• Bartlett prepared the first noble-gas compound in 1962.
• Xenon forms several compounds such as $\text{XeF}_{2}, \text{XeF}_{4}, \text{XeF}_{6}, \text{XeO}_{3}$.
• Krypton and radon also form compounds.

Since noble gases are not zero-valent in all circumstances, the label "0" is misleading. IUPAC retired it and called the column simply Group 18 — the 18th vertical stripe of the table.

Example 10: Where does hydrogen really belong?

Hydrogen ($1s^{1}$) is sometimes placed in Group 1 (alkali metals) and sometimes in Group 17 (halogens). Give the argument for each placement and state what NCERT recommends.

Solution:

Arguments for Group 1:

• Electronic configuration $ns^{1}$ like alkali metals.
• Forms $\text{H}^{+}$ in many compounds.
• Shows oxidation state +1 in most compounds.

Arguments for Group 17:

• One electron short of a noble-gas configuration.
• Can form $\text{H}^{-}$ in metal hydrides.
• Exists as diatomic $\text{H}_{2}$ like halogens.

NCERT view: Hydrogen is usually shown at the top of Group 1, but its chemistry is unique and its position is considered special or ambiguous.

Example 11: Assigning a future element (JEE Main style)

Predict the electronic configuration, block and group of the undiscovered element with $Z = 121$.

Solution:

Using the extended Aufbau order for a theoretical Period 8 element:

$$[\text{Og}]\,8s^{2}\,5g^{1}$$

Here $[\text{Og}]$ stands for the closed-shell configuration of oganesson (Z = 118). The next two electrons enter $8s$ (Z = 119 and Z = 120), and the 121st electron is predicted to enter the $5g$ sub-shell.

• Block: theoretical g-block.
• Systematic name: digits 1, 2, 1 → un-bi-un-ium → Unbiunium, Ubu.

This is a theoretical extension of periodic-table logic beyond the currently confirmed elements.

Example 12: School Exam short-answer

State the modern periodic law. Write the IUPAC systematic name and symbol of elements with atomic numbers 108 and 111.

Solution:

Modern Periodic Law: "The physical and chemical properties of elements are a periodic function of their atomic numbers."

Z = 108: digits 1, 0, 8 → un + nil + oct + ium = Unniloctium (Uno). (Permanent name: Hassium, Hs.)

Z = 111: digits 1, 1, 1 → un + un + un + ium = Unununium (Uuu). (Permanent name: Roentgenium, Rg.)

A one-mark bonus is often given for writing the permanent name as well.