Lithium, Sodium and Potassium Behave Alike Because Each Has One Outer Electron
See why chemical properties repeat as atomic number rises, read an element's period, group and valence electrons straight from its electronic configuration, compare the alkali metals, alkaline earth metals, halogens and noble gases, and predict valency and formulae from position alone.
Why do elements far apart in atomic number behave almost identically?
Drop a small piece of sodium into water and it fizzes across the surface, giving off hydrogen. Drop in potassium and the same thing happens, only more violently. Lithium does it too, more gently. Three elements with atomic numbers , and — separated by eight each time — and the same chemistry.
The reason is not their mass or their size. It is their outermost shell.
- Lithium is
- Sodium is
- Potassium is
Each has exactly one electron in its outer shell, and chemistry is almost entirely a matter of what happens to outer electrons. So elements with the same number of outer electrons react in the same way.
As atomic number rises, electrons fill shells one at a time, and every time a new shell begins, the pattern of outer electrons starts over: one, two, three and so on up to a full shell. That regular repetition of similar outer configurations, and therefore of similar properties, is periodicity — and it is the whole reason the periodic table can be drawn as a grid of rows and columns.
The table then becomes a reading tool. From an element's atomic number you can write its configuration, and from the configuration you can read:
- Its period, from the number of shells
- Its group, from the number of outer electrons
- Its valency, and whether it is a metal or a non-metal
- The formula of its compounds, and roughly how it will behave
The four families at the edges of the table show the idea most clearly — the alkali metals and alkaline earth metals on the left, the halogens and noble gases on the right — so this part studies each of them and the trends within them.
This page covers the first part of the ICSE Class 10 Chemistry chapter on periodic properties: periodicity, locating elements up to calcium, the main families, and relating position to valency and behaviour.
The reason is not their mass or their size. It is their outermost shell.
- Lithium is
- Sodium is
- Potassium is
Each has exactly one electron in its outer shell, and chemistry is almost entirely a matter of what happens to outer electrons. So elements with the same number of outer electrons react in the same way.
As atomic number rises, electrons fill shells one at a time, and every time a new shell begins, the pattern of outer electrons starts over: one, two, three and so on up to a full shell. That regular repetition of similar outer configurations, and therefore of similar properties, is periodicity — and it is the whole reason the periodic table can be drawn as a grid of rows and columns.
The table then becomes a reading tool. From an element's atomic number you can write its configuration, and from the configuration you can read:
- Its period, from the number of shells
- Its group, from the number of outer electrons
- Its valency, and whether it is a metal or a non-metal
- The formula of its compounds, and roughly how it will behave
The four families at the edges of the table show the idea most clearly — the alkali metals and alkaline earth metals on the left, the halogens and noble gases on the right — so this part studies each of them and the trends within them.
This page covers the first part of the ICSE Class 10 Chemistry chapter on periodic properties: periodicity, locating elements up to calcium, the main families, and relating position to valency and behaviour.
What causes periodicity in the properties of elements?
Periodicity is the repetition of similar properties at regular intervals of atomic number, and it happens because similar outer electronic configurations recur.
The modern periodic law states that the properties of elements are a periodic function of their atomic numbers. The atomic number, not the atomic mass, is what the table is arranged by.
How the shells fill, in the simple shell picture used at this level. The shells are called , , and , and they can hold at most , , and electrons — given by for shell number — with the rule that the outermost shell never holds more than eight.
- **The shell fills with electrons, at helium
- The shell fills with , from lithium to neon
- The shell takes from sodium to argon before the next shell begins
- The shell starts with potassium and calcium
So the outer configuration runs one, two, three up to eight, and then starts again at one — which is exactly the repeating pattern in the properties.
Why the outer electrons decide the chemistry. Two things act on an outer electron:
- The pull of the nucleus, which grows with the nuclear charge — the number of protons
- The shielding by the inner shells, which reduces how much of that pull the outer electron feels
Atoms react to reach a stable arrangement — a full outer shell of eight, called an octet, or two for the first shell, called a duplet. An atom with one outer electron finds it easiest to lose that electron; an atom with seven finds it easiest to gain one. That tendency to lose or gain decides whether the element is a metal or a non-metal, what its valency is, and how vigorously it reacts.
Worked example — seeing the repetition.** Write the configurations of fluorine (), chlorine () and neon (), argon (), and say which pairs should behave alike.
- Fluorine: and chlorine: — seven outer electrons each, so both readily gain one electron. They behave alike
- Neon: and argon: — a complete outer octet each, so neither has any tendency to react. They behave alike
And the gap in each pair is eight, which is why the old observation that properties repeat after eight elements held so well for the lighter elements.
One boundary case worth being exact about. Periodicity is a repetition of similar properties, not identical ones. Potassium reacts with water more violently than sodium, even though both have one outer electron, because potassium's outer electron is in a shell further from the nucleus and more shielded, so it is lost more easily. The outer configuration fixes the kind of behaviour; the shell number adjusts its strength — and that adjustment is the subject of the trends in the rest of this chapter.
The modern periodic law states that the properties of elements are a periodic function of their atomic numbers. The atomic number, not the atomic mass, is what the table is arranged by.
How the shells fill, in the simple shell picture used at this level. The shells are called , , and , and they can hold at most , , and electrons — given by for shell number — with the rule that the outermost shell never holds more than eight.
- **The shell fills with electrons, at helium
- The shell fills with , from lithium to neon
- The shell takes from sodium to argon before the next shell begins
- The shell starts with potassium and calcium
So the outer configuration runs one, two, three up to eight, and then starts again at one — which is exactly the repeating pattern in the properties.
Why the outer electrons decide the chemistry. Two things act on an outer electron:
- The pull of the nucleus, which grows with the nuclear charge — the number of protons
- The shielding by the inner shells, which reduces how much of that pull the outer electron feels
Atoms react to reach a stable arrangement — a full outer shell of eight, called an octet, or two for the first shell, called a duplet. An atom with one outer electron finds it easiest to lose that electron; an atom with seven finds it easiest to gain one. That tendency to lose or gain decides whether the element is a metal or a non-metal, what its valency is, and how vigorously it reacts.
Worked example — seeing the repetition.** Write the configurations of fluorine (), chlorine () and neon (), argon (), and say which pairs should behave alike.
- Fluorine: and chlorine: — seven outer electrons each, so both readily gain one electron. They behave alike
- Neon: and argon: — a complete outer octet each, so neither has any tendency to react. They behave alike
And the gap in each pair is eight, which is why the old observation that properties repeat after eight elements held so well for the lighter elements.
One boundary case worth being exact about. Periodicity is a repetition of similar properties, not identical ones. Potassium reacts with water more violently than sodium, even though both have one outer electron, because potassium's outer electron is in a shell further from the nucleus and more shielded, so it is lost more easily. The outer configuration fixes the kind of behaviour; the shell number adjusts its strength — and that adjustment is the subject of the trends in the rest of this chapter.
How do you find an element's period and group from its configuration?
The number of shells gives the period; the number of outer electrons gives the group.
The two rules:
- Period number = number of shells containing electrons
- Group number — for groups and it equals the number of valence electrons; **for groups to it equals the number of valence electrons plus
Valence electrons are the electrons in the outermost shell.
The elements up to calcium, with their positions:
- Period 1**: hydrogen in group ; helium in group
- Period 2: lithium group ; beryllium group ; boron group ; carbon group ; nitrogen group ; oxygen group ; fluorine group ; neon group
- Period 3: sodium group ; magnesium group ; aluminium group ; silicon group ; phosphorus group ; sulphur group ; chlorine group ; argon group
- Period 4, first two elements: potassium group ; calcium group
Worked example 1. An element has atomic number . State its configuration, period, group and number of valence electrons.
- Three shells, so Period 3
- **Five valence electrons, so group
- It is phosphorus
Worked example 2.** An element is in Period 3 and group . Find its atomic number.
**Period 3 means three shells; group means two outer electrons**, so the configuration is and
It is magnesium.
Worked example 3. An element has the configuration . Locate it.
**Four shells, so Period 4; two outer electrons, so group .** , calcium.
Two elements whose positions need explaining, because they break the simple rule.
Hydrogen has one electron, like the alkali metals, and is placed at the head of group . But it is a non-metal and a gas, and it also resembles the halogens: it can gain one electron to complete its shell, and it forms diatomic molecules. Its position is therefore described as anomalous — it fits in group by configuration and not by behaviour.
Helium has two electrons, which by the counting rule would put it in group . **It is placed in group instead, because its single shell is complete with two electrons — a duplet — so it is as unreactive as neon and argon. Position follows chemical behaviour here, and helium behaves as a noble gas.
One check that catches most errors. The atomic number must equal the sum of the numbers in the configuration, and no shell other than the first may be written with more than eight electrons at this level. is never correct for potassium** — the ninth electron goes into a new shell, giving , and that is exactly why potassium begins Period 4.
The two rules:
- Period number = number of shells containing electrons
- Group number — for groups and it equals the number of valence electrons; **for groups to it equals the number of valence electrons plus
Valence electrons are the electrons in the outermost shell.
The elements up to calcium, with their positions:
- Period 1**: hydrogen in group ; helium in group
- Period 2: lithium group ; beryllium group ; boron group ; carbon group ; nitrogen group ; oxygen group ; fluorine group ; neon group
- Period 3: sodium group ; magnesium group ; aluminium group ; silicon group ; phosphorus group ; sulphur group ; chlorine group ; argon group
- Period 4, first two elements: potassium group ; calcium group
Worked example 1. An element has atomic number . State its configuration, period, group and number of valence electrons.
- Three shells, so Period 3
- **Five valence electrons, so group
- It is phosphorus
Worked example 2.** An element is in Period 3 and group . Find its atomic number.
**Period 3 means three shells; group means two outer electrons**, so the configuration is and
It is magnesium.
Worked example 3. An element has the configuration . Locate it.
**Four shells, so Period 4; two outer electrons, so group .** , calcium.
Two elements whose positions need explaining, because they break the simple rule.
Hydrogen has one electron, like the alkali metals, and is placed at the head of group . But it is a non-metal and a gas, and it also resembles the halogens: it can gain one electron to complete its shell, and it forms diatomic molecules. Its position is therefore described as anomalous — it fits in group by configuration and not by behaviour.
Helium has two electrons, which by the counting rule would put it in group . **It is placed in group instead, because its single shell is complete with two electrons — a duplet — so it is as unreactive as neon and argon. Position follows chemical behaviour here, and helium behaves as a noble gas.
One check that catches most errors. The atomic number must equal the sum of the numbers in the configuration, and no shell other than the first may be written with more than eight electrons at this level. is never correct for potassium** — the ninth electron goes into a new shell, giving , and that is exactly why potassium begins Period 4.
How do the alkali metals, alkaline earth metals, halogens and noble gases compare?
Each family shares an outer configuration, so each shares a pattern of behaviour — and each shows a steady trend going down the group.
Group 1 — the alkali metals: lithium, sodium, potassium, rubidium, caesium.
- One valence electron, so valency , and they form ions with a charge of
- Soft enough to cut with a knife, with low densities — lithium, sodium and potassium float on water — and low melting points
- Highly reactive, so they are stored under kerosene to keep air and moisture away
- They react with water to give an alkali and hydrogen, which is how the group gets its name:
- Reactivity increases down the group — lithium reacts steadily, sodium vigorously, potassium violently enough for the hydrogen to catch fire
- Melting point decreases down the group
- They form basic oxides, and they are strong reducing agents because they lose electrons so easily
Group 2 — the alkaline earth metals: beryllium, magnesium, calcium, strontium, barium.
- Two valence electrons, so valency , forming ions with a charge of
- Harder, denser and higher melting than the alkali metals, because two electrons per atom take part in metallic bonding
- Less reactive than the alkali metals in the same period, since losing two electrons takes more energy than losing one
- Reactivity increases down the group: magnesium reacts only slowly with cold water but readily with steam, while calcium reacts with cold water:
- Their oxides are basic
Group 17 — the halogens: fluorine, chlorine, bromine, iodine.
- Seven valence electrons, so valency , gaining one electron to form ions with a charge of
- Non-metals existing as diatomic molecules — , , ,
- Their physical state changes down the group: fluorine is a pale yellow gas, chlorine a greenish-yellow gas, bromine a reddish-brown liquid, iodine a violet-black solid — melting and boiling points increase down the group
- Reactivity decreases down the group, the opposite of the metals, because a larger atom attracts an incoming electron less strongly
- A more reactive halogen displaces a less reactive one from its salt:
- They are strong oxidising agents, since they take electrons readily
Group 18 — the zero group or noble gases: helium, neon, argon, krypton, xenon.
- A complete outer shell — a duplet for helium and an octet for the rest — so valency zero
- Monatomic gases, existing as single atoms rather than molecules
- Chemically inert under ordinary conditions, having no tendency to lose, gain or share electrons
- Uses follow from their inertness: helium fills balloons, neon glows in lighted signs, argon fills electric bulbs so the hot filament does not react
The single most useful comparison in this section. Going down a group, metals become more reactive and non-metals become less reactive.
- For a metal, reacting means losing electrons, and a larger atom holds its outer electrons less tightly
- For a non-metal, reacting means gaining electrons, and a larger atom attracts an extra electron less strongly
One cause — increasing size — produces opposite trends in the two kinds of element, and stating that reason is worth more than listing either trend.
A boundary case about the displacement reaction. Bromine cannot displace chlorine from potassium chloride, because bromine is the less reactive of the two. Mixing bromine water with potassium chloride solution produces no change, and a question that asks you to complete such an equation expects the answer that no reaction occurs.
Group 1 — the alkali metals: lithium, sodium, potassium, rubidium, caesium.
- One valence electron, so valency , and they form ions with a charge of
- Soft enough to cut with a knife, with low densities — lithium, sodium and potassium float on water — and low melting points
- Highly reactive, so they are stored under kerosene to keep air and moisture away
- They react with water to give an alkali and hydrogen, which is how the group gets its name:
- Reactivity increases down the group — lithium reacts steadily, sodium vigorously, potassium violently enough for the hydrogen to catch fire
- Melting point decreases down the group
- They form basic oxides, and they are strong reducing agents because they lose electrons so easily
Group 2 — the alkaline earth metals: beryllium, magnesium, calcium, strontium, barium.
- Two valence electrons, so valency , forming ions with a charge of
- Harder, denser and higher melting than the alkali metals, because two electrons per atom take part in metallic bonding
- Less reactive than the alkali metals in the same period, since losing two electrons takes more energy than losing one
- Reactivity increases down the group: magnesium reacts only slowly with cold water but readily with steam, while calcium reacts with cold water:
- Their oxides are basic
Group 17 — the halogens: fluorine, chlorine, bromine, iodine.
- Seven valence electrons, so valency , gaining one electron to form ions with a charge of
- Non-metals existing as diatomic molecules — , , ,
- Their physical state changes down the group: fluorine is a pale yellow gas, chlorine a greenish-yellow gas, bromine a reddish-brown liquid, iodine a violet-black solid — melting and boiling points increase down the group
- Reactivity decreases down the group, the opposite of the metals, because a larger atom attracts an incoming electron less strongly
- A more reactive halogen displaces a less reactive one from its salt:
- They are strong oxidising agents, since they take electrons readily
Group 18 — the zero group or noble gases: helium, neon, argon, krypton, xenon.
- A complete outer shell — a duplet for helium and an octet for the rest — so valency zero
- Monatomic gases, existing as single atoms rather than molecules
- Chemically inert under ordinary conditions, having no tendency to lose, gain or share electrons
- Uses follow from their inertness: helium fills balloons, neon glows in lighted signs, argon fills electric bulbs so the hot filament does not react
The single most useful comparison in this section. Going down a group, metals become more reactive and non-metals become less reactive.
- For a metal, reacting means losing electrons, and a larger atom holds its outer electrons less tightly
- For a non-metal, reacting means gaining electrons, and a larger atom attracts an extra electron less strongly
One cause — increasing size — produces opposite trends in the two kinds of element, and stating that reason is worth more than listing either trend.
A boundary case about the displacement reaction. Bromine cannot displace chlorine from potassium chloride, because bromine is the less reactive of the two. Mixing bromine water with potassium chloride solution produces no change, and a question that asks you to complete such an equation expects the answer that no reaction occurs.
How can you predict valency, nature and formulae from an element's position?
The group tells you the valence electrons, the valence electrons tell you the valency and whether the element loses or gains electrons, and the valencies of two elements give the formula of their compound.
Valency from the group:
- Groups 1, 2 and 13 — valency equals the number of valence electrons: , and . These atoms lose electrons
- Group 14 — valency , usually by sharing electrons
- Groups 15, 16 and 17 — valency equals minus the valence electrons: , and . These atoms gain or share electrons
- Group 18 — valency
**So across a period the valency with respect to hydrogen or chlorine rises from to and then falls back to , and it repeats in the next period — another face of periodicity.
Metallic or non-metallic nature from position:
- Elements on the left of a period are metals, with one to three valence electrons that they tend to lose
- Elements on the right are non-metals, with five to seven valence electrons that they tend to gain
- Between them lie metalloids, such as boron and silicon, with intermediate properties
- Across a period, metallic character decreases; down a group, it increases
The nature of the oxide follows the nature of the element:
- Metals form basic oxides — sodium oxide, magnesium oxide, calcium oxide
- Aluminium oxide is amphoteric, reacting with both acids and bases
- Non-metals form acidic oxides — carbon dioxide, sulphur dioxide
Worked example 1 — a complete prediction.** Element has atomic number . Predict its position, valency, nature, the formula of its chloride and its oxide, and its reaction with water.
- Configuration , so Period 4, group 1
- **Valency **, losing one electron to form
- A highly reactive metal
- **Chloride and oxide , which is basic
- It reacts vigorously with cold water, giving an alkali and hydrogen
It is potassium, and every prediction is correct.
Worked example 2.** Element has atomic number . Predict the same.
- Configuration , so Period 3, group 16
- **Valency **, gaining two electrons to form
- A non-metal
- **Hydride **, and an oxide that is acidic
It is sulphur.
Worked example 3 — the formula of a compound between two elements. Predict the formula of the compound formed between and above.
** has valency and has valency .** Exchanging the valencies:
**Potassium sulphide, — two potassium atoms each losing one electron to supply the two electrons one sulphur atom gains.
Worked example 4.** Predict the formula of the oxide of an element with configuration .
**Group , valency **; oxygen has valency . Exchanging:
**The element is aluminium and the oxide is , which is amphoteric — a metal close enough to the non-metals for its oxide to show both kinds of behaviour.
Worked example 5.** Two elements and have configurations and . Give the formula and the kind of bond in their compound.
** is magnesium with valency ; is chlorine with valency .** The formula is , that is , and since a metal is transferring electrons to a non-metal, the bond is electrovalent — the subject of the chemical bonding chapter.
One trap in predictions of this kind. Valency is not the same as the number of valence electrons for the right-hand groups. Chlorine has seven valence electrons and a valency of one, and writing by mistake confuses the two. Always convert valence electrons to valency before exchanging.
Valency from the group:
- Groups 1, 2 and 13 — valency equals the number of valence electrons: , and . These atoms lose electrons
- Group 14 — valency , usually by sharing electrons
- Groups 15, 16 and 17 — valency equals minus the valence electrons: , and . These atoms gain or share electrons
- Group 18 — valency
**So across a period the valency with respect to hydrogen or chlorine rises from to and then falls back to , and it repeats in the next period — another face of periodicity.
Metallic or non-metallic nature from position:
- Elements on the left of a period are metals, with one to three valence electrons that they tend to lose
- Elements on the right are non-metals, with five to seven valence electrons that they tend to gain
- Between them lie metalloids, such as boron and silicon, with intermediate properties
- Across a period, metallic character decreases; down a group, it increases
The nature of the oxide follows the nature of the element:
- Metals form basic oxides — sodium oxide, magnesium oxide, calcium oxide
- Aluminium oxide is amphoteric, reacting with both acids and bases
- Non-metals form acidic oxides — carbon dioxide, sulphur dioxide
Worked example 1 — a complete prediction.** Element has atomic number . Predict its position, valency, nature, the formula of its chloride and its oxide, and its reaction with water.
- Configuration , so Period 4, group 1
- **Valency **, losing one electron to form
- A highly reactive metal
- **Chloride and oxide , which is basic
- It reacts vigorously with cold water, giving an alkali and hydrogen
It is potassium, and every prediction is correct.
Worked example 2.** Element has atomic number . Predict the same.
- Configuration , so Period 3, group 16
- **Valency **, gaining two electrons to form
- A non-metal
- **Hydride **, and an oxide that is acidic
It is sulphur.
Worked example 3 — the formula of a compound between two elements. Predict the formula of the compound formed between and above.
** has valency and has valency .** Exchanging the valencies:
**Potassium sulphide, — two potassium atoms each losing one electron to supply the two electrons one sulphur atom gains.
Worked example 4.** Predict the formula of the oxide of an element with configuration .
**Group , valency **; oxygen has valency . Exchanging:
**The element is aluminium and the oxide is , which is amphoteric — a metal close enough to the non-metals for its oxide to show both kinds of behaviour.
Worked example 5.** Two elements and have configurations and . Give the formula and the kind of bond in their compound.
** is magnesium with valency ; is chlorine with valency .** The formula is , that is , and since a metal is transferring electrons to a non-metal, the bond is electrovalent — the subject of the chemical bonding chapter.
One trap in predictions of this kind. Valency is not the same as the number of valence electrons for the right-hand groups. Chlorine has seven valence electrons and a valency of one, and writing by mistake confuses the two. Always convert valence electrons to valency before exchanging.
Exam tip
What layout keeps a periodic table answer free of errors?
Write the configuration first, then read off period, group, valency and nature in that order. Every later answer depends on the configuration, so a slip there spreads through the whole question.
- Check that the configuration adds up to the atomic number, and that no outer shell exceeds eight
- Period = number of shells; group = valence electrons for groups 1 and 2, and valence electrons plus 10 for groups 13 to 18
- Convert valence electrons to valency before writing a formula: minus the valence electrons for groups 15 to 17
- Exchange valencies to write a formula, and cancel any common factor
- Name the family when asked about properties: alkali metal, alkaline earth metal, halogen or noble gas
- Give the direction of every trend and the reason — metals more reactive down a group, non-metals less reactive, both because size increases
- Balance every equation you write, such as the reaction of sodium with water
- Explain hydrogen and helium separately if a question asks about their positions
- State the nature of the oxide from the nature of the element — basic, amphoteric or acidic
- Use the modern group numbers 1 to 18 consistently
The misconception to name. Elements in the same period are not chemically similar. Sodium and chlorine are both in Period 3 and could hardly be more different — one a soft reactive metal, the other a poisonous gas. Similar properties belong to elements in the same group, because it is the outer configuration that repeats down a column, not along a row.
A second trap. Writing the group number as the number of valence electrons for groups 13 to 18. Chlorine has seven valence electrons and is in group 17, not group 7. The older numbering used numbers such as seven for the halogens, and mixing the two systems within one answer is marked wrong.
- Check that the configuration adds up to the atomic number, and that no outer shell exceeds eight
- Period = number of shells; group = valence electrons for groups 1 and 2, and valence electrons plus 10 for groups 13 to 18
- Convert valence electrons to valency before writing a formula: minus the valence electrons for groups 15 to 17
- Exchange valencies to write a formula, and cancel any common factor
- Name the family when asked about properties: alkali metal, alkaline earth metal, halogen or noble gas
- Give the direction of every trend and the reason — metals more reactive down a group, non-metals less reactive, both because size increases
- Balance every equation you write, such as the reaction of sodium with water
- Explain hydrogen and helium separately if a question asks about their positions
- State the nature of the oxide from the nature of the element — basic, amphoteric or acidic
- Use the modern group numbers 1 to 18 consistently
The misconception to name. Elements in the same period are not chemically similar. Sodium and chlorine are both in Period 3 and could hardly be more different — one a soft reactive metal, the other a poisonous gas. Similar properties belong to elements in the same group, because it is the outer configuration that repeats down a column, not along a row.
A second trap. Writing the group number as the number of valence electrons for groups 13 to 18. Chlorine has seven valence electrons and is in group 17, not group 7. The older numbering used numbers such as seven for the halogens, and mixing the two systems within one answer is marked wrong.
Did you know
Why is argon, not air, sealed inside an ordinary light bulb?
The filament of a traditional bulb glows because it is extremely hot — hot enough that, in air, it would react with oxygen almost at once and burn through. A vacuum would stop that, but the hot metal slowly evaporates in a vacuum, blackening the glass and thinning the filament. The answer is to fill the bulb with a gas that presses on the filament without reacting with it.
Argon does exactly that, and its position in the table explains why.
- It is in group 18, with the configuration — a complete octet
- It has no tendency to lose, gain or share electrons, so it does not react with the white-hot metal
- It is heavy enough to slow the evaporation of the filament by crowding the space around it
The same inertness is the whole value of the noble gases, and each is used where not reacting is the point.
- Helium fills balloons because it is light and cannot catch fire, unlike hydrogen, which is lighter still but burns
- Neon glows with a bright red-orange light when electricity passes through it at low pressure, without being chemically changed — which is why it is used in lighted signs
- Argon also shields metals during welding, so the hot metal is not attacked by the air
And the contrast with the neighbours makes the periodicity vivid. Move one step to the left of argon and you reach chlorine, — one electron short of an octet and so desperate to gain one that it is among the most reactive elements known. Move one step to the right and you reach potassium, — one electron over, and so eager to lose it that it reacts violently with water.
**Three consecutive atomic numbers, , and , and the chemistry swings from extreme reactivity to none at all and back to extreme reactivity. Nothing about the atoms' masses changes much across that step; only the outer shell does — and that is the most striking single demonstration of why the table is arranged by electronic configuration.
One more consequence of that arrangement. Hydrogen fills balloons less often than helium even though it is lighter, because hydrogen has one electron and reacts readily, while helium's complete duplet makes it safe. The position of each element in the table predicts which gas is safe to put near a flame**, without any experiment at all.
Argon does exactly that, and its position in the table explains why.
- It is in group 18, with the configuration — a complete octet
- It has no tendency to lose, gain or share electrons, so it does not react with the white-hot metal
- It is heavy enough to slow the evaporation of the filament by crowding the space around it
The same inertness is the whole value of the noble gases, and each is used where not reacting is the point.
- Helium fills balloons because it is light and cannot catch fire, unlike hydrogen, which is lighter still but burns
- Neon glows with a bright red-orange light when electricity passes through it at low pressure, without being chemically changed — which is why it is used in lighted signs
- Argon also shields metals during welding, so the hot metal is not attacked by the air
And the contrast with the neighbours makes the periodicity vivid. Move one step to the left of argon and you reach chlorine, — one electron short of an octet and so desperate to gain one that it is among the most reactive elements known. Move one step to the right and you reach potassium, — one electron over, and so eager to lose it that it reacts violently with water.
**Three consecutive atomic numbers, , and , and the chemistry swings from extreme reactivity to none at all and back to extreme reactivity. Nothing about the atoms' masses changes much across that step; only the outer shell does — and that is the most striking single demonstration of why the table is arranged by electronic configuration.
One more consequence of that arrangement. Hydrogen fills balloons less often than helium even though it is lighter, because hydrogen has one electron and reacts readily, while helium's complete duplet makes it safe. The position of each element in the table predicts which gas is safe to put near a flame**, without any experiment at all.
Exam relevance
How does the periodic table prepare you for JEE and NEET Chemistry?
This is foundation work for Class 11 Classification of Elements and Periodicity in Properties, and for the s-block and p-block chapters, which are examined in both JEE Main and NEET Chemistry.
Where the shell picture leads. Class 11 replaces shells with orbitals — , , and — and divides the table into four blocks by the orbital that receives the last electron. Groups 1 and 2 become the s-block, groups 13 to 18 the p-block, and the transition elements between them the d-block. Your rule that groups 13 to 18 have ten added to the valence count is the d-block's ten electrons, which is why it works.
Where the group rules lead. Locating an element from its atomic number is asked directly in both JEE Main and NEET, usually for elements beyond calcium where the electrons must be counted. The method — configuration first, then period and group — is identical, and a candidate fluent at it here needs only the orbital filling order added.
Where the family properties lead. Class 11 The s-Block Elements treats the alkali and alkaline earth metals in depth — their reactions with water, oxygen and hydrogen, the solubility of their compounds, and the anomalous behaviour of lithium and beryllium. The trends you learn here are the backbone of that chapter, and NEET frequently asks which alkali metal is most reactive or which hydroxide is most basic.
Where the halogens and noble gases lead. Class 12 The p-Block Elements treats groups 15 to 18, including the displacement reactions of the halogens, their oxidising power and the compounds of xenon. The observation here that a more reactive halogen displaces a less reactive one is examined directly, often as an assertion-reason question.
Where the valency and oxide predictions lead. The acidic, basic and amphoteric character of oxides across a period is a recurring question in both exams, and it is the same pattern as worked example 4. Aluminium oxide being amphoteric is asked by name.
Question types to expect. At this level: configuration to period and group, family properties, trends down a group, and formula prediction. In competitive papers: block identification, s-block and p-block reactions, nature of oxides, and anomalous properties of the first element of a group.
The single trap that costs marks. Treating the trend in reactivity as the same for metals and non-metals. Metals grow more reactive down a group and non-metals less reactive, for one reason — increasing size — acting on opposite processes. NEET sets questions where only that distinction separates the options.
A second trap. Placing helium in group 2 by counting electrons. Its complete duplet puts it in group 18, and the same principle — position decided by behaviour rooted in configuration — explains why hydrogen's position is called anomalous. Both are recurring objective questions.
Board versus competitive emphasis. The ICSE paper marks the configuration, the stated period and group, the balanced equation and the reason behind a trend; a competitive paper marks a single identification or ordering, often for heavier elements. The transferable habit is reading every property back to the outer configuration — because that is what the Class 11 orbital treatment does too, only in finer detail.
Where the shell picture leads. Class 11 replaces shells with orbitals — , , and — and divides the table into four blocks by the orbital that receives the last electron. Groups 1 and 2 become the s-block, groups 13 to 18 the p-block, and the transition elements between them the d-block. Your rule that groups 13 to 18 have ten added to the valence count is the d-block's ten electrons, which is why it works.
Where the group rules lead. Locating an element from its atomic number is asked directly in both JEE Main and NEET, usually for elements beyond calcium where the electrons must be counted. The method — configuration first, then period and group — is identical, and a candidate fluent at it here needs only the orbital filling order added.
Where the family properties lead. Class 11 The s-Block Elements treats the alkali and alkaline earth metals in depth — their reactions with water, oxygen and hydrogen, the solubility of their compounds, and the anomalous behaviour of lithium and beryllium. The trends you learn here are the backbone of that chapter, and NEET frequently asks which alkali metal is most reactive or which hydroxide is most basic.
Where the halogens and noble gases lead. Class 12 The p-Block Elements treats groups 15 to 18, including the displacement reactions of the halogens, their oxidising power and the compounds of xenon. The observation here that a more reactive halogen displaces a less reactive one is examined directly, often as an assertion-reason question.
Where the valency and oxide predictions lead. The acidic, basic and amphoteric character of oxides across a period is a recurring question in both exams, and it is the same pattern as worked example 4. Aluminium oxide being amphoteric is asked by name.
Question types to expect. At this level: configuration to period and group, family properties, trends down a group, and formula prediction. In competitive papers: block identification, s-block and p-block reactions, nature of oxides, and anomalous properties of the first element of a group.
The single trap that costs marks. Treating the trend in reactivity as the same for metals and non-metals. Metals grow more reactive down a group and non-metals less reactive, for one reason — increasing size — acting on opposite processes. NEET sets questions where only that distinction separates the options.
A second trap. Placing helium in group 2 by counting electrons. Its complete duplet puts it in group 18, and the same principle — position decided by behaviour rooted in configuration — explains why hydrogen's position is called anomalous. Both are recurring objective questions.
Board versus competitive emphasis. The ICSE paper marks the configuration, the stated period and group, the balanced equation and the reason behind a trend; a competitive paper marks a single identification or ordering, often for heavier elements. The transferable habit is reading every property back to the outer configuration — because that is what the Class 11 orbital treatment does too, only in finer detail.
Key takeaways
What must you be able to do from this part?
One law, two reading rules and four families.
- Modern periodic law: properties are a periodic function of atomic number
- Periodicity is the repetition of similar properties because similar outer configurations recur as shells fill
- Shell capacities are , , , , but the outermost shell never holds more than eight
- Atoms react to reach an octet, or a duplet for the first shell
- Period = number of shells; group = valence electrons for groups 1 and 2, valence electrons plus 10 for groups 13 to 18
- **** is , Period 3, group 15, phosphorus; Period 3 group 2 is , , magnesium; ** is calcium
- Hydrogen's position is anomalous — group 1 by configuration, but a non-metal resembling the halogens
- Helium is in group 18, because its single shell is complete
- Alkali metals** have one valence electron, form ions, are soft and stored under kerosene, react with water to give alkali and hydrogen, and grow more reactive down the group
- Alkaline earth metals have two valence electrons, form ions, are harder and less reactive than alkali metals, and also grow more reactive down the group
- Halogens have seven valence electrons, form ions, exist as diatomic molecules, change from gas to liquid to solid down the group, grow less reactive down the group, and a more reactive one displaces a less reactive one:
- Noble gases have complete shells, valency zero, are monatomic and inert, and are used in balloons, lighted signs and bulbs
- Valency equals the valence electrons for groups 1, 2 and 13, is for group 14, and is minus the valence electrons for groups 15 to 17
- Metals lie on the left and form basic oxides; non-metals lie on the right and form acidic oxides; aluminium oxide is amphoteric
- Metallic character decreases across a period and increases down a group
- Exchange valencies for a formula: with gives ; with oxygen gives ; with gives
The quickest self-test needs only the numbers to . Pick any five at random, write each configuration, and give the period, group, valency and one compound formula for each — then check whether any two of your five should behave alike.
- Modern periodic law: properties are a periodic function of atomic number
- Periodicity is the repetition of similar properties because similar outer configurations recur as shells fill
- Shell capacities are , , , , but the outermost shell never holds more than eight
- Atoms react to reach an octet, or a duplet for the first shell
- Period = number of shells; group = valence electrons for groups 1 and 2, valence electrons plus 10 for groups 13 to 18
- **** is , Period 3, group 15, phosphorus; Period 3 group 2 is , , magnesium; ** is calcium
- Hydrogen's position is anomalous — group 1 by configuration, but a non-metal resembling the halogens
- Helium is in group 18, because its single shell is complete
- Alkali metals** have one valence electron, form ions, are soft and stored under kerosene, react with water to give alkali and hydrogen, and grow more reactive down the group
- Alkaline earth metals have two valence electrons, form ions, are harder and less reactive than alkali metals, and also grow more reactive down the group
- Halogens have seven valence electrons, form ions, exist as diatomic molecules, change from gas to liquid to solid down the group, grow less reactive down the group, and a more reactive one displaces a less reactive one:
- Noble gases have complete shells, valency zero, are monatomic and inert, and are used in balloons, lighted signs and bulbs
- Valency equals the valence electrons for groups 1, 2 and 13, is for group 14, and is minus the valence electrons for groups 15 to 17
- Metals lie on the left and form basic oxides; non-metals lie on the right and form acidic oxides; aluminium oxide is amphoteric
- Metallic character decreases across a period and increases down a group
- Exchange valencies for a formula: with gives ; with oxygen gives ; with gives
The quickest self-test needs only the numbers to . Pick any five at random, write each configuration, and give the period, group, valency and one compound formula for each — then check whether any two of your five should behave alike.