Read an Element's Address Straight Off Its Electron Count
Work out the group and period of any element in the first three periods from its electronic configuration, then use its position to predict its family, its size and whether it behaves as a metal.
Can you really work out where an element sits from its electrons alone?
Give a chemist the electronic configuration and nothing else — no name, no symbol, no atomic mass — and they will tell you three things immediately.
The element sits in period 3. It sits in group 17. And it is a reactive non-metal that will form salts with metals, because that is what group 17 does.
They have not looked anything up. The configuration contains the address, and the address contains the chemistry.
That is what the previous part of this chapter earned. Once the table is ordered by atomic number — and therefore by electron count — an element's position stops being a fact to memorise and becomes something you can derive.
The whole of this page is that one idea used four ways. Two numbers read off a configuration give you the position; the position gives you the family; the family gives you the reactivity; and the position within the grid gives you the size and the metallic character.
This page covers the second part of the ICSE Class 9 Chemistry chapter on the periodic table: finding the group and period from a configuration, the layout of the modern table, the four families you must be able to compare, and how atomic size and metallic character vary.
The element sits in period 3. It sits in group 17. And it is a reactive non-metal that will form salts with metals, because that is what group 17 does.
They have not looked anything up. The configuration contains the address, and the address contains the chemistry.
That is what the previous part of this chapter earned. Once the table is ordered by atomic number — and therefore by electron count — an element's position stops being a fact to memorise and becomes something you can derive.
The whole of this page is that one idea used four ways. Two numbers read off a configuration give you the position; the position gives you the family; the family gives you the reactivity; and the position within the grid gives you the size and the metallic character.
This page covers the second part of the ICSE Class 9 Chemistry chapter on the periodic table: finding the group and period from a configuration, the layout of the modern table, the four families you must be able to compare, and how atomic size and metallic character vary.
How do you find the group and period from an electronic configuration?
The period number is the number of shells that contain electrons, and the group number comes from the number of valence electrons.
The period rule is the simple one. Count the shells in the configuration. Hydrogen has one shell and sits in period 1. Carbon has two and sits in period 2. Sodium has three and sits in period 3.
The group rule has two cases, because the modern table numbers its columns 1 to 18 while the first three periods use only the outer two columns on the left and six on the right.
- Valence electrons 1 or 2: the group number is the number of valence electrons
- Valence electrons 3 to 8: the group number is the number of valence electrons plus 10
Worked examples, period by period.
- Sodium, — three shells, so period 3; one valence electron, so group 1
- Magnesium, — period 3, group 2
- Aluminium, — period 3, group 13
- Silicon, — period 3, group 14
- Phosphorus, — period 3, group 15
- Sulphur, — period 3, group 16
- Chlorine, — period 3, group 17
- Argon, — period 3, group 18
- Lithium, — period 2, group 1
- Carbon, — period 2, group 14
- Oxygen, — period 2, group 16
- Neon, — period 2, group 18
Notice that period 3 runs cleanly from group 1 to group 18 with a jump in the middle, from group 2 straight to group 13. The columns numbered 3 to 12 are empty in the first three periods, because the shell that fills there does not begin until period 4. That gap is not a mistake in the table — it is the space the transition elements occupy further down.
Helium is the exception you must remember. Its configuration is , which by the rule above would put it in group 2. But helium has a complete first shell — the duplet — so it is chemically inert and belongs with the other inert gases in group 18. It is placed there, above neon, and not with the reactive metals of group 2.
The reason the rule works at all is worth stating. Elements in the same group have the same number of valence electrons, and the valence electrons are what react. So the group number is not a label attached to an element from outside — it is a restatement of the element's outermost shell. Two elements in one group behave alike because their configurations end the same way, and that is the explanation the older periodic law could not supply.
The period rule is the simple one. Count the shells in the configuration. Hydrogen has one shell and sits in period 1. Carbon has two and sits in period 2. Sodium has three and sits in period 3.
The group rule has two cases, because the modern table numbers its columns 1 to 18 while the first three periods use only the outer two columns on the left and six on the right.
- Valence electrons 1 or 2: the group number is the number of valence electrons
- Valence electrons 3 to 8: the group number is the number of valence electrons plus 10
Worked examples, period by period.
- Sodium, — three shells, so period 3; one valence electron, so group 1
- Magnesium, — period 3, group 2
- Aluminium, — period 3, group 13
- Silicon, — period 3, group 14
- Phosphorus, — period 3, group 15
- Sulphur, — period 3, group 16
- Chlorine, — period 3, group 17
- Argon, — period 3, group 18
- Lithium, — period 2, group 1
- Carbon, — period 2, group 14
- Oxygen, — period 2, group 16
- Neon, — period 2, group 18
Notice that period 3 runs cleanly from group 1 to group 18 with a jump in the middle, from group 2 straight to group 13. The columns numbered 3 to 12 are empty in the first three periods, because the shell that fills there does not begin until period 4. That gap is not a mistake in the table — it is the space the transition elements occupy further down.
Helium is the exception you must remember. Its configuration is , which by the rule above would put it in group 2. But helium has a complete first shell — the duplet — so it is chemically inert and belongs with the other inert gases in group 18. It is placed there, above neon, and not with the reactive metals of group 2.
The reason the rule works at all is worth stating. Elements in the same group have the same number of valence electrons, and the valence electrons are what react. So the group number is not a label attached to an element from outside — it is a restatement of the element's outermost shell. Two elements in one group behave alike because their configurations end the same way, and that is the explanation the older periodic law could not supply.
What does the modern periodic table actually look like?
The modern periodic table has 18 vertical groups and 7 horizontal periods.
Elements are placed in order of increasing atomic number, reading left to right along a period and then starting the next period.
Each period corresponds to a shell being filled, and that is why the periods have the lengths they do.
- Period 1 holds only 2 elements — hydrogen and helium. The first shell takes two electrons and is finished
- Period 2 holds 8 elements, lithium to neon
- Period 3 holds 8 elements, sodium to argon
- Period 4 holds 18 elements
- Period 5 holds 18 elements
- Period 6 holds 32 elements
- Period 7 is incomplete
So a period ends where a shell is satisfied, and every complete period ends with a noble gas. That is the connection the previous page's fun fact pointed at: the zero group is the boundary between periods, not an appendix.
The groups fall into blocks according to which kind of shell is being filled.
- Groups 1 and 2 on the far left — the s-block, containing the most reactive metals
- Groups 13 to 18 on the right — the p-block, containing the non-metals, the metalloids and some metals
- Groups 3 to 12 in the middle — the d-block, whose members are called the transition elements. All of them are metals, and they are typically hard, dense, high-melting and able to show more than one valency
- Two separate rows placed below the main table — the f-block, called the inner transition elements, consisting of the lanthanoids and the actinoids
Why the f-block sits outside the grid. It has fourteen members in each row. Inserting them in their proper place inside period 6 and period 7 would make the table so wide that it would be unusable on a page or a wall. They belong inside the table and are printed below it for convenience — which is exactly what a question asking why they are shown separately is looking for. It is a printing decision, not a chemical one.
Reading the table as a map. Metals occupy the left and centre, non-metals the top right, and a staircase of metalloids — elements with properties of both — runs between them. So the left-to-right direction across any period is a journey from metal to non-metal, and the last section explains why.
Elements are placed in order of increasing atomic number, reading left to right along a period and then starting the next period.
Each period corresponds to a shell being filled, and that is why the periods have the lengths they do.
- Period 1 holds only 2 elements — hydrogen and helium. The first shell takes two electrons and is finished
- Period 2 holds 8 elements, lithium to neon
- Period 3 holds 8 elements, sodium to argon
- Period 4 holds 18 elements
- Period 5 holds 18 elements
- Period 6 holds 32 elements
- Period 7 is incomplete
So a period ends where a shell is satisfied, and every complete period ends with a noble gas. That is the connection the previous page's fun fact pointed at: the zero group is the boundary between periods, not an appendix.
The groups fall into blocks according to which kind of shell is being filled.
- Groups 1 and 2 on the far left — the s-block, containing the most reactive metals
- Groups 13 to 18 on the right — the p-block, containing the non-metals, the metalloids and some metals
- Groups 3 to 12 in the middle — the d-block, whose members are called the transition elements. All of them are metals, and they are typically hard, dense, high-melting and able to show more than one valency
- Two separate rows placed below the main table — the f-block, called the inner transition elements, consisting of the lanthanoids and the actinoids
Why the f-block sits outside the grid. It has fourteen members in each row. Inserting them in their proper place inside period 6 and period 7 would make the table so wide that it would be unusable on a page or a wall. They belong inside the table and are printed below it for convenience — which is exactly what a question asking why they are shown separately is looking for. It is a printing decision, not a chemical one.
Reading the table as a map. Metals occupy the left and centre, non-metals the top right, and a staircase of metalloids — elements with properties of both — runs between them. So the left-to-right direction across any period is a journey from metal to non-metal, and the last section explains why.
How do the alkali metals, halogens and noble gases compare?
Four families are set for comparison, and each is defined by its valence electrons.
Alkali metals — group 1 (lithium, sodium, potassium and the rest; hydrogen is placed in this column but is not one of them).
- One valence electron, so valency 1, and they form ions with a single positive charge
- Very reactive metals — they react with cold water, giving hydrogen and a hydroxide
- Soft enough to cut with a knife, with low densities and low melting points for metals
- Their oxides and hydroxides are strongly basic
- Reactivity increases down the group, so potassium is more reactive than sodium, which is more reactive than lithium
Because they react with the air and with moisture, sodium and potassium are kept submerged in kerosene in every school laboratory — a bottle of sodium sitting in oil on the shelf is the group's reactivity made visible.
Alkaline earth metals — group 2 (beryllium, magnesium, calcium and the rest).
- Two valence electrons, so valency 2, forming doubly charged positive ions
- Reactive metals, but less reactive than the group 1 metal of the same period — magnesium does not react with cold water, while sodium does
- Harder, denser and higher melting than their group 1 neighbours
- Their oxides and hydroxides are basic
- Reactivity increases down the group
Halogens — group 17 (fluorine, chlorine, bromine, iodine).
- Seven valence electrons, so they need one more and have valency 1, forming singly charged negative ions
- Very reactive non-metals, existing as diatomic molecules such as
- They combine with metals to form salts — the name means salt-former
- They are coloured, and the colour deepens down the group
- Reactivity decreases down the group, which is the opposite of the metals
The iodine in iodised table salt and the chlorine used to treat drinking water are both halogens doing what group 17 does.
Noble gases — group 18 (helium, neon, argon, krypton, xenon).
- A complete outermost shell — an octet, or the duplet in helium's case
- Valency zero, and they are chemically inert
- They exist as single atoms — they are monatomic, unlike every other gas in this list
- Used wherever an unreactive atmosphere is needed, and in lighting
Look at why the two reactivity trends run in opposite directions. A metal reacts by losing its valence electron, and going down a group the atom gets bigger so the outer electron is held less tightly and is lost more easily — reactivity rises. A halogen reacts by gaining an electron, and going down a group the bigger atom attracts an incoming electron less strongly — reactivity falls.
So the same cause, increasing atomic size, produces opposite effects on the two families, because one family needs to give an electron away and the other needs to take one. That is the single most useful sentence in this topic, and a question asking why potassium is more reactive than lithium while iodine is less reactive than fluorine is asking for precisely it.
Alkali metals — group 1 (lithium, sodium, potassium and the rest; hydrogen is placed in this column but is not one of them).
- One valence electron, so valency 1, and they form ions with a single positive charge
- Very reactive metals — they react with cold water, giving hydrogen and a hydroxide
- Soft enough to cut with a knife, with low densities and low melting points for metals
- Their oxides and hydroxides are strongly basic
- Reactivity increases down the group, so potassium is more reactive than sodium, which is more reactive than lithium
Because they react with the air and with moisture, sodium and potassium are kept submerged in kerosene in every school laboratory — a bottle of sodium sitting in oil on the shelf is the group's reactivity made visible.
Alkaline earth metals — group 2 (beryllium, magnesium, calcium and the rest).
- Two valence electrons, so valency 2, forming doubly charged positive ions
- Reactive metals, but less reactive than the group 1 metal of the same period — magnesium does not react with cold water, while sodium does
- Harder, denser and higher melting than their group 1 neighbours
- Their oxides and hydroxides are basic
- Reactivity increases down the group
Halogens — group 17 (fluorine, chlorine, bromine, iodine).
- Seven valence electrons, so they need one more and have valency 1, forming singly charged negative ions
- Very reactive non-metals, existing as diatomic molecules such as
- They combine with metals to form salts — the name means salt-former
- They are coloured, and the colour deepens down the group
- Reactivity decreases down the group, which is the opposite of the metals
The iodine in iodised table salt and the chlorine used to treat drinking water are both halogens doing what group 17 does.
Noble gases — group 18 (helium, neon, argon, krypton, xenon).
- A complete outermost shell — an octet, or the duplet in helium's case
- Valency zero, and they are chemically inert
- They exist as single atoms — they are monatomic, unlike every other gas in this list
- Used wherever an unreactive atmosphere is needed, and in lighting
Look at why the two reactivity trends run in opposite directions. A metal reacts by losing its valence electron, and going down a group the atom gets bigger so the outer electron is held less tightly and is lost more easily — reactivity rises. A halogen reacts by gaining an electron, and going down a group the bigger atom attracts an incoming electron less strongly — reactivity falls.
So the same cause, increasing atomic size, produces opposite effects on the two families, because one family needs to give an electron away and the other needs to take one. That is the single most useful sentence in this topic, and a question asking why potassium is more reactive than lithium while iodine is less reactive than fluorine is asking for precisely it.
Why do atoms get smaller across a period and bigger down a group?
Atomic size decreases across a period from left to right, and increases down a group.
Across a period. Each step to the right adds one proton to the nucleus and one electron to the same outermost shell. The number of shells does not change, so the extra positive charge pulls the whole electron cloud in more tightly.
So sodium, at the left of period 3, is the largest atom in that period, and chlorine, near the right, is much smaller — even though chlorine has more electrons. More electrons in the same number of shells makes an atom smaller, not bigger, and that is the result students most often get backwards.
Down a group. Each step down adds a whole new shell, and the new shell lies further from the nucleus than the one before it. The added distance outweighs the added nuclear charge, so the atom gets bigger. Lithium is small, sodium larger, potassium larger still.
Metallic character decreases across a period and increases down a group.
A metal is an element that loses electrons readily. So metallic character follows whatever makes losing an electron easy — a large atom with its outer electron far from the nucleus, and a small nuclear pull on that electron.
- Across a period, atoms shrink and nuclear charge rises, so the outer electrons are held more tightly and are harder to lose. Metallic character falls. Period 3 begins with sodium, a violently reactive metal, passes through aluminium, a metal, silicon, a metalloid, and ends with sulphur and chlorine, which are non-metals
- Down a group, atoms grow and the outer electron is held more loosely, so it is lost more easily. Metallic character rises
Non-metallic character does exactly the opposite, because a non-metal gains electrons. It increases across a period and decreases down a group.
Every one of these four trends is the same fact stated four ways. Atomic size and nuclear pull are the only two things changing, and everything else follows from them — which is why it is far safer to reason from size than to memorise four separate arrows.
Worked reasoning — which is larger, magnesium or sulphur? Both are in period 3, with configurations and . Both have three shells, so the shell count cannot decide it. Sulphur has more protons — against magnesium's — so its nucleus pulls the same three shells in harder. Magnesium is the larger atom, and magnesium is the more metallic.
Worked reasoning — which is more reactive, lithium or potassium? Lithium is and potassium is . Both have one valence electron, but potassium's sits in the fourth shell, much further from the nucleus. It is lost more easily, so potassium is more reactive — as its group 1 position down the column requires.
One boundary case worth knowing. A positive ion is always smaller than its parent atom, because it has lost its outermost electrons entirely — sodium loses a whole shell to become . A negative ion is always larger than its parent atom, because the added electrons increase the repulsion within the same shell. So comparing an ion with an atom is not the same question as comparing two atoms, and a question that slips an ion into the list is testing whether you noticed.
Across a period. Each step to the right adds one proton to the nucleus and one electron to the same outermost shell. The number of shells does not change, so the extra positive charge pulls the whole electron cloud in more tightly.
So sodium, at the left of period 3, is the largest atom in that period, and chlorine, near the right, is much smaller — even though chlorine has more electrons. More electrons in the same number of shells makes an atom smaller, not bigger, and that is the result students most often get backwards.
Down a group. Each step down adds a whole new shell, and the new shell lies further from the nucleus than the one before it. The added distance outweighs the added nuclear charge, so the atom gets bigger. Lithium is small, sodium larger, potassium larger still.
Metallic character decreases across a period and increases down a group.
A metal is an element that loses electrons readily. So metallic character follows whatever makes losing an electron easy — a large atom with its outer electron far from the nucleus, and a small nuclear pull on that electron.
- Across a period, atoms shrink and nuclear charge rises, so the outer electrons are held more tightly and are harder to lose. Metallic character falls. Period 3 begins with sodium, a violently reactive metal, passes through aluminium, a metal, silicon, a metalloid, and ends with sulphur and chlorine, which are non-metals
- Down a group, atoms grow and the outer electron is held more loosely, so it is lost more easily. Metallic character rises
Non-metallic character does exactly the opposite, because a non-metal gains electrons. It increases across a period and decreases down a group.
Every one of these four trends is the same fact stated four ways. Atomic size and nuclear pull are the only two things changing, and everything else follows from them — which is why it is far safer to reason from size than to memorise four separate arrows.
Worked reasoning — which is larger, magnesium or sulphur? Both are in period 3, with configurations and . Both have three shells, so the shell count cannot decide it. Sulphur has more protons — against magnesium's — so its nucleus pulls the same three shells in harder. Magnesium is the larger atom, and magnesium is the more metallic.
Worked reasoning — which is more reactive, lithium or potassium? Lithium is and potassium is . Both have one valence electron, but potassium's sits in the fourth shell, much further from the nucleus. It is lost more easily, so potassium is more reactive — as its group 1 position down the column requires.
One boundary case worth knowing. A positive ion is always smaller than its parent atom, because it has lost its outermost electrons entirely — sodium loses a whole shell to become . A negative ion is always larger than its parent atom, because the added electrons increase the repulsion within the same shell. So comparing an ion with an atom is not the same question as comparing two atoms, and a question that slips an ion into the list is testing whether you noticed.
Exam tip
Exam tip: count the shells, then count the valence electrons
Period = number of shells. Group = valence electrons, plus 10 if there are 3 or more. Write both steps out; do not guess.
Helium is in group 18, not group 2 — it has a complete duplet and is inert. Learn it as the exception.
Groups 3 to 12 are empty in the first three periods. Period 3 jumps from group 2 straight to group 13.
Learn the period lengths: 2, 8, 8, 18, 18, 32, and period 7 incomplete. Every complete period ends with a noble gas.
Name the blocks with their groups: s-block 1–2, p-block 13–18, d-block 3–12 (transition elements), f-block below the table (inner transition — lanthanoids and actinoids).
The f-block is printed separately to keep the table narrow enough to use — a printing decision, not a chemical one.
For each family, lead with the valence electrons: group 1 has 1, group 2 has 2, group 17 has 7, group 18 has a complete shell and valency zero.
Halogens are diatomic; noble gases are monatomic. That one contrast is often worth a mark on its own.
Reactivity increases down group 1 and DECREASES down group 17. Explain it by size: a bigger atom loses an electron more easily and attracts one less strongly.
Atomic size decreases across a period (same shells, more nuclear charge) and increases down a group (a new shell). Always give the reason, not just the direction.
And remember that a positive ion is smaller than its atom and a negative ion is larger — a size comparison that includes an ion is a different question.
Helium is in group 18, not group 2 — it has a complete duplet and is inert. Learn it as the exception.
Groups 3 to 12 are empty in the first three periods. Period 3 jumps from group 2 straight to group 13.
Learn the period lengths: 2, 8, 8, 18, 18, 32, and period 7 incomplete. Every complete period ends with a noble gas.
Name the blocks with their groups: s-block 1–2, p-block 13–18, d-block 3–12 (transition elements), f-block below the table (inner transition — lanthanoids and actinoids).
The f-block is printed separately to keep the table narrow enough to use — a printing decision, not a chemical one.
For each family, lead with the valence electrons: group 1 has 1, group 2 has 2, group 17 has 7, group 18 has a complete shell and valency zero.
Halogens are diatomic; noble gases are monatomic. That one contrast is often worth a mark on its own.
Reactivity increases down group 1 and DECREASES down group 17. Explain it by size: a bigger atom loses an electron more easily and attracts one less strongly.
Atomic size decreases across a period (same shells, more nuclear charge) and increases down a group (a new shell). Always give the reason, not just the direction.
And remember that a positive ion is smaller than its atom and a negative ion is larger — a size comparison that includes an ion is a different question.
Did you know
Why the table is wide in the middle and not at the top
Look at a printed periodic table and the shape is odd. The top two rows are stubby — two elements, then eight, then eight. Then the table suddenly widens to eighteen, stays there for two rows, and widens again to thirty-two.
A table designed by someone tidying a list would be a rectangle. This one is not, and the shape is not decoration.
Every step in the width is a shell capacity showing through. The first shell holds two electrons, so period 1 has two elements and stops. The second and third shells effectively take eight before the next shell begins, so those periods have eight members each. From period 4 onwards a further set of positions becomes available, and eighteen elements are needed to fill the row. Further down, thirty-two.
The numbers 2, 8, 8, 18, 18, 32 are not arbitrary lengths chosen by a designer. They are electron counts, and they are the same counts the previous chapter derived from the shell-filling rules.
So the periodic table has the outline it has because atoms are built the way they are. The bulge in the middle of the table is the transition-metal block, and that block exists because a shell that holds eighteen electrons starts filling late. The two rows stranded at the bottom exist because a set of positions holding fourteen electrons starts filling later still.
There is a nice consequence for a student. If you know the period lengths you can work out where an element must be without seeing a table at all. Element number 11 must start period 3, because periods 1 and 2 together hold elements. Element number 19 must start period 4, because .
The shape of the chart on the laboratory wall is a picture of electron shells, drawn by nature rather than by a printer — which is why every version of the table you will ever see has the same strange silhouette.
A table designed by someone tidying a list would be a rectangle. This one is not, and the shape is not decoration.
Every step in the width is a shell capacity showing through. The first shell holds two electrons, so period 1 has two elements and stops. The second and third shells effectively take eight before the next shell begins, so those periods have eight members each. From period 4 onwards a further set of positions becomes available, and eighteen elements are needed to fill the row. Further down, thirty-two.
The numbers 2, 8, 8, 18, 18, 32 are not arbitrary lengths chosen by a designer. They are electron counts, and they are the same counts the previous chapter derived from the shell-filling rules.
So the periodic table has the outline it has because atoms are built the way they are. The bulge in the middle of the table is the transition-metal block, and that block exists because a shell that holds eighteen electrons starts filling late. The two rows stranded at the bottom exist because a set of positions holding fourteen electrons starts filling later still.
There is a nice consequence for a student. If you know the period lengths you can work out where an element must be without seeing a table at all. Element number 11 must start period 3, because periods 1 and 2 together hold elements. Element number 19 must start period 4, because .
The shape of the chart on the laboratory wall is a picture of electron shells, drawn by nature rather than by a printer — which is why every version of the table you will ever see has the same strange silhouette.
Exam relevance
Why do JEE Main and NEET keep coming back to periodic trends?
Because a trend question can be set on any four elements in the table, so the supply of questions is effectively unlimited.
This is the foundation for Class 11 Chemistry Classification of Elements and Periodicity in Properties, examined in both JEE Main and NEET. The group-and-period rule used here is replaced by a more powerful one: the block is read from the subshell being filled, and the group is read from the s- and p-electron count or, for the d-block, from the sum of the d- and s-electrons. The method generalises to every element in the table rather than just the first three periods.
The four qualitative trends become four named quantities. Class 11 examines atomic and ionic radii, ionisation enthalpy, electron gain enthalpy and electronegativity, each with its own trend and its own exceptions to account for. Ranking four given elements on one of these is among the most frequent JEE Main inorganic question types, and the reasoning is exactly the size-versus-nuclear-charge argument used on this page.
The exceptions are where the marks are. Class 11 asks why the ionisation enthalpy of beryllium exceeds that of boron, why nitrogen's exceeds oxygen's, and why the electron gain enthalpy of chlorine exceeds that of fluorine. Each exception is a departure from the smooth trend learnt here, so the trend has to be secure before the exception can make sense.
Ionic radii are asked as isoelectronic comparisons. Given a set of ions with the same number of electrons — such as the ten-electron set from the previous chapter — you are asked to order them by size. The answer turns on nuclear charge alone, and the rule that a cation is smaller and an anion larger than the parent atom is the starting point. This is a recurring NEET and JEE Main item.
The families organise the rest of inorganic chemistry. Class 11 The s-Block Elements studies groups 1 and 2 exactly as compared here, and The p-Block Elements continues into groups 13 to 18; Class 12 covers the d- and f-block. The anomalous behaviour of the first member of each group, the diagonal relationship and the inert-pair effect are recurring JEE Advanced themes, and all three are trend arguments.
For NEET, periodicity appears as recall plus one-step reasoning — identify the block or group from a configuration, pick the largest or most electronegative species from a list, match a family to a property. Biology leans on the table as well: the essential and trace mineral elements of Mineral Nutrition and the ion movements of Transport in Plants assume you can read metals and non-metals off it.
What the questions look like. For board work, expect give the group and period from a configuration, state the number of elements in each period, name the four families with their valence electrons and valency, compare two named families, and explain a trend in atomic size, metallic character or reactivity with a reason. Trend answers with a direction but no reason lose most of the marks. For JEE Main and NEET, expect block identification, trend rankings, isoelectronic size ordering and assertion-reason items.
How board and competitive emphasis differ. A board paper rewards the derivation of the position and a reasoned trend. A competitive paper assumes both and tests the exceptions to the trend instead.
The single trap that costs the most marks. Assuming an atom with more electrons must be bigger. Across a period the electrons go into the same shells while the nuclear charge rises, so the atom gets smaller — chlorine has far more electrons than sodium and is much the smaller atom. The defence is to check the shell count first: if two species have the same number of shells, size is decided by nuclear charge alone, and only a change in shell count makes an atom bigger.
This is the foundation for Class 11 Chemistry Classification of Elements and Periodicity in Properties, examined in both JEE Main and NEET. The group-and-period rule used here is replaced by a more powerful one: the block is read from the subshell being filled, and the group is read from the s- and p-electron count or, for the d-block, from the sum of the d- and s-electrons. The method generalises to every element in the table rather than just the first three periods.
The four qualitative trends become four named quantities. Class 11 examines atomic and ionic radii, ionisation enthalpy, electron gain enthalpy and electronegativity, each with its own trend and its own exceptions to account for. Ranking four given elements on one of these is among the most frequent JEE Main inorganic question types, and the reasoning is exactly the size-versus-nuclear-charge argument used on this page.
The exceptions are where the marks are. Class 11 asks why the ionisation enthalpy of beryllium exceeds that of boron, why nitrogen's exceeds oxygen's, and why the electron gain enthalpy of chlorine exceeds that of fluorine. Each exception is a departure from the smooth trend learnt here, so the trend has to be secure before the exception can make sense.
Ionic radii are asked as isoelectronic comparisons. Given a set of ions with the same number of electrons — such as the ten-electron set from the previous chapter — you are asked to order them by size. The answer turns on nuclear charge alone, and the rule that a cation is smaller and an anion larger than the parent atom is the starting point. This is a recurring NEET and JEE Main item.
The families organise the rest of inorganic chemistry. Class 11 The s-Block Elements studies groups 1 and 2 exactly as compared here, and The p-Block Elements continues into groups 13 to 18; Class 12 covers the d- and f-block. The anomalous behaviour of the first member of each group, the diagonal relationship and the inert-pair effect are recurring JEE Advanced themes, and all three are trend arguments.
For NEET, periodicity appears as recall plus one-step reasoning — identify the block or group from a configuration, pick the largest or most electronegative species from a list, match a family to a property. Biology leans on the table as well: the essential and trace mineral elements of Mineral Nutrition and the ion movements of Transport in Plants assume you can read metals and non-metals off it.
What the questions look like. For board work, expect give the group and period from a configuration, state the number of elements in each period, name the four families with their valence electrons and valency, compare two named families, and explain a trend in atomic size, metallic character or reactivity with a reason. Trend answers with a direction but no reason lose most of the marks. For JEE Main and NEET, expect block identification, trend rankings, isoelectronic size ordering and assertion-reason items.
How board and competitive emphasis differ. A board paper rewards the derivation of the position and a reasoned trend. A competitive paper assumes both and tests the exceptions to the trend instead.
The single trap that costs the most marks. Assuming an atom with more electrons must be bigger. Across a period the electrons go into the same shells while the nuclear charge rises, so the atom gets smaller — chlorine has far more electrons than sodium and is much the smaller atom. The defence is to check the shell count first: if two species have the same number of shells, size is decided by nuclear charge alone, and only a change in shell count makes an atom bigger.
Key takeaways
Position, families and trends in the periodic table: quick revision
- Period number = the number of occupied shells.
- Group number = valence electrons if there are 1 or 2, and valence electrons + 10 if there are 3 or more.
- Sodium — period 3, group 1. Aluminium — period 3, group 13. Chlorine — period 3, group 17. Argon — period 3, group 18. Carbon — period 2, group 14.
- Helium is the exception — configuration , but a complete duplet, so group 18, not group 2.
- Groups 3 to 12 are empty in periods 1 to 3; period 3 jumps from group 2 to group 13.
- 18 groups and 7 periods. Period lengths 2, 8, 8, 18, 18, 32, with period 7 incomplete. Every complete period ends with a noble gas.
- Blocks: s groups 1–2; p groups 13–18; d groups 3–12, the transition elements, all metals, often with more than one valency; f below the table, the inner transition elements — lanthanoids and actinoids — printed separately only to keep the table narrow.
- Alkali metals (group 1): 1 valence electron, valency 1, very reactive, soft, low melting, basic oxides, stored under kerosene, reactivity increases down the group.
- Alkaline earth metals (group 2): 2 valence electrons, valency 2, reactive but less so than group 1 in the same period, harder and denser, basic oxides.
- Halogens (group 17): 7 valence electrons, valency 1, very reactive non-metals, diatomic, salt-formers, coloured, reactivity decreases down the group.
- Noble gases (group 18): complete outermost shell, valency zero, chemically inert, monatomic.
- The opposite reactivity trends have one cause: a bigger atom loses an electron more easily (metals more reactive down a group) and attracts an incoming electron less strongly (halogens less reactive down a group).
- Atomic size decreases across a period — same number of shells, more nuclear charge pulling them in.
- Atomic size increases down a group — a new shell is added each step.
- Metallic character decreases across a period and increases down a group; non-metallic character does the reverse.
- Magnesium is larger and more metallic than sulphur — same three shells, fewer protons.
- Potassium is more reactive than lithium — the valence electron is further out.
- A positive ion is smaller than its atom; a negative ion is larger.
- Metals on the left, non-metals at the top right, metalloids on the staircase between them.
Write down any configuration from to and give yourself ten seconds to name the period, the group and the family — then check whether the chemistry you predicted is the chemistry the element actually has.
- Group number = valence electrons if there are 1 or 2, and valence electrons + 10 if there are 3 or more.
- Sodium — period 3, group 1. Aluminium — period 3, group 13. Chlorine — period 3, group 17. Argon — period 3, group 18. Carbon — period 2, group 14.
- Helium is the exception — configuration , but a complete duplet, so group 18, not group 2.
- Groups 3 to 12 are empty in periods 1 to 3; period 3 jumps from group 2 to group 13.
- 18 groups and 7 periods. Period lengths 2, 8, 8, 18, 18, 32, with period 7 incomplete. Every complete period ends with a noble gas.
- Blocks: s groups 1–2; p groups 13–18; d groups 3–12, the transition elements, all metals, often with more than one valency; f below the table, the inner transition elements — lanthanoids and actinoids — printed separately only to keep the table narrow.
- Alkali metals (group 1): 1 valence electron, valency 1, very reactive, soft, low melting, basic oxides, stored under kerosene, reactivity increases down the group.
- Alkaline earth metals (group 2): 2 valence electrons, valency 2, reactive but less so than group 1 in the same period, harder and denser, basic oxides.
- Halogens (group 17): 7 valence electrons, valency 1, very reactive non-metals, diatomic, salt-formers, coloured, reactivity decreases down the group.
- Noble gases (group 18): complete outermost shell, valency zero, chemically inert, monatomic.
- The opposite reactivity trends have one cause: a bigger atom loses an electron more easily (metals more reactive down a group) and attracts an incoming electron less strongly (halogens less reactive down a group).
- Atomic size decreases across a period — same number of shells, more nuclear charge pulling them in.
- Atomic size increases down a group — a new shell is added each step.
- Metallic character decreases across a period and increases down a group; non-metallic character does the reverse.
- Magnesium is larger and more metallic than sulphur — same three shells, fewer protons.
- Potassium is more reactive than lithium — the valence electron is further out.
- A positive ion is smaller than its atom; a negative ion is larger.
- Metals on the left, non-metals at the top right, metalloids on the staircase between them.
Write down any configuration from to and give yourself ten seconds to name the period, the group and the family — then check whether the chemistry you predicted is the chemistry the element actually has.