Periodic Trends | O-Level Chemistry

Last updated: 8 October 2026

Diagnose → Learn → Master: Every study-note section includes 15 free MCQs — start with 5 Diagnostic MCQs to identify gaps in your understanding, then complete 10 Mastery MCQs near the end to check what you have learned.

The Periodic Table is not simply a chart for looking up element symbols. It is organised so that patterns in atomic structure become patterns in chemical behaviour. Once you learn how to read those patterns, you can make useful predictions even about elements you have not studied individually.

The key is to connect an element’s position to its electronic configuration. The number of occupied shells tells us its period, while the number of valence electrons helps us understand its group, the ions it may form and the way its reactivity changes.

This is what makes periodic trends so useful. Instead of memorising each element separately, we can use a small number of ideas to explain and predict the behaviour of whole groups of elements.

Diagnostic Quiz

What do you already know about the Periodic Table?

Try these five questions before reading the chapter. They will show whether the key ideas about protons, electron shells and valence electrons are already secure.

5 diagnostic questions
Important periodic-table misconceptions tested
Instant explanation for every answer option

Which quantity determines the identity of an element?

The number of protons

Correct. Every element has its own proton number. Changing the number of protons changes the identity of the element.

The number of neutrons

Isotopes of the same element can contain different numbers of neutrons, so neutron number does not determine the identity of the element.

The number of occupied electron shells

Several different elements can have the same number of occupied shells. The number of occupied shells instead tells us the Period number.

The number of valence electrons

Different elements in the same group can have the same number of valence electrons, so this does not uniquely identify an element.

The proton number uniquely identifies an element. This is why the modern Periodic Table is arranged in order of increasing proton number.

What does the Period number of an element tell you?

The number of valence electrons

Valence electrons are related to the Group number for the main-group elements, not the Period number.

The number of occupied electron shells

Correct. An atom in Period 3 has electrons occupying three electron shells.

The number of neutrons

Neutron number cannot be determined from the Period number.

The charge of the ion formed

Ionic charge depends on how many electrons are lost or gained, rather than directly on the Period number.

The Period number corresponds to the number of occupied electron shells in an atom.

Why do elements in the same group usually have similar chemical properties?

They contain the same total number of electrons.

Elements in the same group have different proton numbers and therefore different total numbers of electrons.

They contain the same number of occupied electron shells.

Elements in the same period have the same number of occupied shells. Moving down a group adds additional occupied shells.

They have the same number of valence electrons.

Correct. Valence electrons are involved in chemical bonding and reactions, so similar outer-shell configurations lead to similar chemical behaviour.

They have the same proton number.

Two atoms with the same proton number are atoms of the same element, not different elements within the same group.

Chemical behaviour depends strongly on the outer electrons. Main-group elements in the same group have similar outer-shell electronic configurations.

An atom has the electronic configuration 2,8,1. Which statement is correct?

It is in Period 3 and Group 1.

Correct. There are three occupied electron shells and one valence electron.

It is in Period 1 and Group 3.

This reverses the meanings of Period and Group. Three occupied shells mean Period 3.

It is in Period 2 and Group 8.

The atom contains three occupied electron shells, not two.

It is in Period 3 and Group 18.

Group 18 atoms have filled outer electron shells. This atom has only one valence electron.

In 2,8,1, the three occupied shells give Period 3 and the one valence electron places the element in Group 1.

A chlorine atom has the electronic configuration 2,8,7. What is it most likely to do when forming a simple ion?

Lose 1 electron to form Cl⁺.

Chlorine already has seven valence electrons. Losing one electron would move it further from a filled outer shell.

Gain 1 electron to form Cl⁻.

Correct. Gaining one electron gives chlorine the stable electronic configuration 2,8,8.

Gain 7 electrons to form Cl⁷⁻.

Chlorine needs only one additional electron to fill its outer shell.

Lose 7 electrons to form Cl⁷⁺.

Removing seven electrons would require far more energy than gaining the single electron needed to fill the outer shell.

Group 17 atoms have seven valence electrons and usually gain one electron when forming simple ions, producing 1− halide ions.

How the Periodic Table is arranged

Elements in the modern Periodic Table are arranged in order of increasing proton number, also called atomic number.

Hydrogen has proton number 1, helium has proton number 2, lithium has proton number 3, and so on. Moving from one element to the next increases the proton number by one.

Why proton number matters

The proton number determines the identity of an element. Every carbon atom contains 6 protons. Every oxygen atom contains 8 protons. If the number of protons changes, the element itself changes.

The Periodic Table is organised into horizontal rows called periods and vertical columns called groups.

Part of the Periodic Table	Direction	What it tells us
Period	Horizontal row	Number of occupied electron shells
Group	Vertical column	Outer-shell electronic pattern and many chemical properties

This organisation is not arbitrary. It reflects the way electrons are arranged around the nucleus, which is why an element’s position can be used to predict its behaviour.

Electronic configuration reveals Period and Group

An element’s electronic configuration tells us how its electrons are distributed among its electron shells. For the first 20 elements, this makes it straightforward to connect electronic configuration with position in the Periodic Table.

The Period number tells us the number of occupied shells

Consider sodium:

Na: 2,8,1

Sodium has electrons in three occupied shells, so sodium is found in Period 3.

Calcium has the electronic configuration:

Ca: 2,8,8,2

Calcium has four occupied shells, so it is found in Period 4.

Finding the Period

Count the number of occupied electron shells. That number gives the Period number.

The Group number is related to the valence electrons

The electrons in the outermost occupied shell are called valence electrons. These electrons are especially important because they are the electrons involved when atoms form ions and chemical bonds.

For the main-group elements considered at O-Level, the Group number is related to the number of valence electrons.

Group	Valence electrons	Example
Group 1	1	Na: 2,8,1
Group 2	2	Mg: 2,8,2
Group 13	3	Al: 2,8,3
Group 14	4	Si: 2,8,4
Group 15	5	P: 2,8,5
Group 16	6	S: 2,8,6
Group 17	7	Cl: 2,8,7
Group 18	Full outer shell	Ar: 2,8,8

For Groups 1 and 2, the Group number directly matches the number of valence electrons. For Groups 13 to 17, subtract 10 from the Group number. Group 18 contains elements with filled outer electron shells.

Helium is the important exception to the usual eight-electron pattern. Its electronic configuration is simply 2. The first electron shell can hold a maximum of two electrons, so helium already has a filled outer shell and belongs to Group 18.

Sulfur atom with electronic configuration 2,8,6 showing that three occupied electron shells place sulfur in Period 3 and six valence electrons place it in Group 16 — Sulfur has three occupied electron shells and six valence electrons, placing it in Period 3 and Group 16.

Worked reasoning: electronic configuration → position

An element has the electronic configuration 2,8,6.

It has three occupied shells, so it is in Period 3.

It has six valence electrons, so it is in Group 16.

The element is sulfur.

Worked reasoning: position → electronic configuration

Suppose element X is in Period 3, Group 17.

Period 3 tells us that X has three occupied electron shells. Group 17 tells us that it has seven valence electrons.

Its electronic configuration is therefore:

2,8,7

Why elements in the same group behave similarly

Lithium, sodium and potassium are different elements. Their atoms have different numbers of protons and different total numbers of electrons. Yet they react in similar ways.

The reason becomes clear when we compare their electronic configurations.

Element	Electronic configuration	Valence electrons
Lithium	2,1	1
Sodium	2,8,1	1
Potassium	2,8,8,1	1

Each atom has one valence electron. Because chemical reactions involve the outer electrons, these three elements tend to react in similar ways and form ions with the same charge.

The important link

Same group → similar outer-shell electronic configuration → similar chemical properties.

How properties change across a period

Moving from left to right across a period, the proton number increases by one from one element to the next. For neutral atoms, one electron is also added each time.

These additional electrons enter the same outer electron shell until that shell becomes full. The number of valence electrons therefore increases across the period.

Period 3 provides a useful example.

Element	Electronic configuration	Broad classification
Sodium, Na	2,8,1	Metal
Magnesium, Mg	2,8,2	Metal
Aluminium, Al	2,8,3	Metal
Silicon, Si	2,8,4	Metalloid
Phosphorus, P	2,8,5	Non-metal
Sulfur, S	2,8,6	Non-metal
Chlorine, Cl	2,8,7	Non-metal
Argon, Ar	2,8,8	Noble gas

Metallic character decreases across a period

Metals generally form positive ions by losing electrons. Elements near the left-hand side of a period have relatively few valence electrons and show metallic behaviour.

Sodium, for example, loses one electron:

Na → Na⁺ + e⁻

Non-metals are found towards the right-hand side of a period. They have more valence electrons and may gain electrons to form negative ions or share electrons when forming covalent bonds.

Chlorine, for example, gains one electron:

Cl + e⁻ → Cl⁻

Across a period from left to right

Metallic character generally decreases.

Non-metallic character generally increases.

The change is gradual rather than an abrupt jump. Silicon lies near the boundary between metals and non-metals and is commonly described as a metalloid.

Using Group number to predict ionic charge

Atoms form ions by losing or gaining electrons. The number of protons in the nucleus does not change during ion formation.

Many of the simple ions encountered at O-Level form because atoms lose or gain electrons to obtain a filled outer electron shell, like that of a noble gas.

Group	Valence electrons	Usual electron change	Typical charge	Examples
Group 1	1	Loses 1 electron	+1	Li⁺, Na⁺, K⁺
Group 2	2	Loses 2 electrons	+2	Mg²⁺, Ca²⁺
Group 13	3	Loses 3 electrons	+3	Al³⁺
Group 14	4	Usually shares electrons instead	No simple general charge	Carbon and silicon usually form covalent bonds
Group 15	5	Gains 3 electrons	−3	N³⁻, P³⁻
Group 16	6	Gains 2 electrons	−2	O²⁻, S²⁻
Group 17	7	Gains 1 electron	−1	F⁻, Cl⁻
Group 18	Full outer shell	Does not normally lose or gain electrons	No common simple ion	He, Ne, Ar

Do not copy the Group number as the ionic charge

A chlorine atom is in Group 17 because it has seven valence electrons. It does not form Cl⁷⁻. It needs to gain only one electron to fill its outer shell, so it forms Cl⁻.

Worked reasoning: predicting the sulfide ion

A sulfur atom has the electronic configuration:

2,8,6

It needs two more electrons to obtain 2,8,8. It therefore gains two electrons:

S + 2e⁻ → S²⁻

The sulfur atom has gained two negatively charged electrons, so the resulting ion has a charge of 2−.

Trends down Group 1

Group 1 elements are called the alkali metals. Lithium, sodium and potassium are the familiar examples at O-Level.

They all have one valence electron and tend to lose this electron during chemical reactions.

Element	Electronic configuration	Ion formed
Lithium	2,1	Li⁺
Sodium	2,8,1	Na⁺
Potassium	2,8,8,1	K⁺

Reactivity increases down Group 1

As we move down Group 1, each element has one more occupied electron shell than the element above it.

The valence electron is therefore farther from the positively charged nucleus. There are also more inner electron shells between the nucleus and the valence electron. These inner shells reduce the attraction experienced by the valence electron through shielding.

The outer electron is therefore held less strongly and can be removed more easily.

Explaining the Group 1 reactivity trend

Down Group 1: more electron shells → greater distance and shielding → weaker attraction between the nucleus and valence electron → electron lost more easily → reactivity increases.

This is why potassium reacts more vigorously with water than sodium, while sodium reacts more vigorously than lithium.

2M + 2H₂O → 2MOH + H₂

Here, M represents a Group 1 metal.

Melting points generally decrease down Group 1

Group 1 elements are metals, so their structures contain positive metal ions attracted to delocalised electrons.

Down the group, the metal ions become larger. The attraction between the positive ions and the delocalised electrons becomes weaker, so less energy is needed to overcome the metallic bonding.

Their melting points therefore generally decrease down the group.

Keep the two Group 1 trends separate

Reactivity increases down Group 1 because the valence electron is lost more easily.

Melting point decreases because metallic bonding becomes weaker.

Trends down Group 17

Group 17 elements are called the halogens. They are non-metals and exist as diatomic molecules.

This means their formulae are written as F₂, Cl₂, Br₂ and I₂ rather than F, Cl, Br and I.

Halogen	State at room temperature	Typical appearance
Fluorine, F₂	Gas	Pale yellow
Chlorine, Cl₂	Gas	Yellow-green
Bromine, Br₂	Liquid	Red-brown
Iodine, I₂	Solid	Grey-black

Reactivity decreases down Group 17

Each halogen atom has seven valence electrons. It tends to gain one electron when it forms a halide ion.

For example:

Cl + e⁻ → Cl⁻

Moving down Group 17, the number of occupied electron shells increases. An incoming electron would therefore be added farther from the nucleus and would experience more shielding from the inner electron shells.

The nucleus attracts the incoming electron less strongly, so gaining an electron becomes more difficult.

Explaining the Group 17 reactivity trend

Down Group 17: more electron shells → greater distance and shielding → weaker attraction for an incoming electron → electron gained less easily → reactivity decreases.

Comparison of Group 1 and Group 17 reactivity trends showing why Group 1 becomes more reactive while Group 17 becomes less reactive down the groups — Down both groups, atoms have more occupied electron shells. This makes the valence electron easier to lose in Group 1 but makes an incoming electron harder to gain in Group 17.

A more reactive halogen displaces a less reactive halogen

Because chlorine is above bromine in Group 17, chlorine is more reactive than bromine. Chlorine can therefore displace bromine from a solution containing bromide ions.

Cl₂ + 2KBr → 2KCl + Br₂

The ionic equation makes the electron transfer clearer:

Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

Bromine cannot displace chlorine from chloride ions because bromine is less reactive than chlorine.

How to predict a halogen displacement reaction

Compare the positions of the two halogens in Group 17. A halogen higher in the group is more reactive and can displace the halide ions of a halogen below it.

Melting and boiling points increase down Group 17

Halogen molecules become larger down the group and contain increasing numbers of electrons. The intermolecular attractions between the molecules become stronger.

More energy is therefore required to separate the molecules. This causes the melting and boiling points to increase down Group 17.

This trend helps explain why chlorine is a gas, bromine is a liquid and iodine is a solid at room temperature.

The colour becomes darker down Group 17

The halogens also show a visible colour trend. Their colours generally become darker down the group.

Do not mix up the Group 17 trends

Reactivity decreases down Group 17.

Melting and boiling points increase down Group 17.

Colour becomes darker down Group 17.

Group 18 and the noble gases

Group 18 elements are called the noble gases. Examples include helium, neon and argon.

Unlike the halogens, noble gases do not normally exist as molecules containing two atoms. They exist as individual atoms and are described as monatomic.

Why noble gases are very unreactive

The outer electron shell of a noble-gas atom is already filled.

Helium has a filled first shell containing two electrons. Neon and argon have eight electrons in their outermost occupied shells.

Because these atoms already have stable electronic configurations, they have little tendency to lose, gain or share electrons.

The reason for noble-gas stability

Noble gases are very unreactive because their atoms have filled outer electron shells.

Uses follow from their properties

The usefulness of a noble gas often comes directly from its properties.

Noble gas	Useful property	Example of use
Helium	Low density and non-flammable	Filling balloons
Neon	Produces a coloured glow during electrical discharge	Advertising signs
Argon	Very unreactive	Providing an inert atmosphere in lamps and during welding

Helium is especially useful in balloons because it has a low density but, unlike hydrogen, it is not flammable.

Common mistakes in periodic-trend questions

Common mistake	Why it is wrong	Better idea
Confusing the Period number with the Group number.	They describe different features of the electronic configuration.	Period tells you the number of occupied shells; Group is related to the number of valence electrons.
Using the Group number directly as the ionic charge.	A Group 17 atom does not gain seven electrons.	Work out how many electrons must actually be lost or gained to obtain a filled outer shell.
Saying elements in the same group have the same number of electron shells.	The number of occupied shells increases down a group.	Elements in the same group have similar numbers of valence electrons, not identical numbers of shells.
Using the same reactivity trend for Groups 1 and 17.	The two groups react by different electron changes.	Group 1 loses an electron more easily down the group; Group 17 gains an electron less easily down the group.
A less reactive halogen displaces a more reactive halogen.	Displacement occurs only when the free halogen is more reactive.	Chlorine can displace bromide ions, but bromine cannot displace chloride ions.
Every filled outer shell contains eight electrons.	The first electron shell can hold a maximum of two electrons.	Helium is stable with the electronic configuration 2.
Ion formation changes the proton number.	Ion formation involves electrons, not protons.	The proton number remains unchanged, so the identity of the element remains unchanged.
Reactivity and melting-point trends must always move in the same direction.	These properties depend on different factors.	Explain reactivity using electron loss or gain, and explain melting or boiling points using attractions between particles.

Periodic trends checklist

Period and Group

Can I identify horizontal periods and vertical groups?
Can I find the Period from the number of occupied electron shells?
Can I relate Group number to valence electrons?

Electronic configuration

Can I work from electronic configuration to Period and Group?
Can I work from Period and Group to an electronic configuration?
Do I remember the special case of helium?

Ion formation

Can I predict whether an atom loses or gains electrons?
Can I predict common simple ionic charges?
Do I remember that proton number does not change during ion formation?

Group 1

Can I state that reactivity increases down Group 1?
Can I explain the trend using distance and shielding?
Can I state that melting point generally decreases down the group?

Group 17

Can I state that reactivity decreases down Group 17?
Can I explain why gaining an electron becomes harder?
Can I state the colour and melting/boiling-point trends?

Groups 17 and 18

Can I predict a halogen displacement reaction?
Do I remember that halogens are diatomic?
Can I explain why noble gases are very unreactive?

Mastery Quiz

Can you apply periodic trends?

These ten questions go beyond simple recall. Use electronic configuration, position and trend reasoning to work out each answer.

10 mastery questions
Application and misconception testing
Instant explanation for every answer option

Element X has the electronic configuration 2,8,6. Which prediction is correct?

X is in Period 6 and Group 3.

The configuration contains three occupied shells, not six. The six is the number of valence electrons.

X is in Period 3, Group 16 and can form X²⁻.

Correct. Three occupied shells give Period 3. Six valence electrons give Group 16, and gaining two electrons produces a 2− ion.

X is in Period 3, Group 16 and can form X⁶⁻.

X needs only two additional electrons to obtain a filled outer shell, not six.

X is in Period 2, Group 6 and can form X²⁺.

X has three occupied shells and is therefore in Period 3. With six valence electrons, it is more likely to gain two electrons.

Read the electronic configuration in two stages: count occupied shells for the Period, then count valence electrons for the Group.

An unknown element is in Period 4, Group 1. Which electronic configuration is expected for an O-Level model of its atom?

2,8,1

This has only three occupied shells, so it represents a Period 3 element.

2,8,7

This has three occupied shells and seven valence electrons, corresponding to Period 3, Group 17.

2,8,8,1

Correct. Four occupied shells give Period 4 and one valence electron gives Group 1.

2,8,8

This has only three occupied shells and a filled outer shell.

Period 4 means four occupied electron shells. Group 1 means one valence electron, giving 2,8,8,1 for potassium.

Potassium reacts more vigorously with water than sodium. Which explanation best accounts for this?

Potassium has fewer occupied electron shells.

Potassium actually has more occupied electron shells than sodium.

Potassium's valence electron is farther from the nucleus and more shielded.

Correct. The attraction between the nucleus and the valence electron is weaker, so the electron is lost more easily.

Potassium has more valence electrons than sodium.

Both potassium and sodium are in Group 1 and have one valence electron.

Potassium gains an electron more easily than sodium.

Group 1 metals react by losing their single valence electron rather than gaining an electron.

Group 1 reactivity increases down the group because greater distance and shielding weaken the attraction between the nucleus and the valence electron.

Chlorine water is added to aqueous potassium bromide. What is expected?

Bromine is formed because chlorine is more reactive than bromine.

Correct. Chlorine is higher than bromine in Group 17 and can displace bromide ions.

No reaction occurs because all halogens have the same reactivity.

Halogen reactivity decreases down Group 17, so the halogens do not have equal reactivities.

Bromide ions displace chlorine because bromine is below chlorine.

Being lower in Group 17 means bromine is less reactive, not more reactive.

Potassium metal is formed.

Potassium ions remain unchanged in this displacement reaction. The electron transfer occurs between chlorine and bromide ions.

The ionic equation is Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂. A more reactive halogen displaces a less reactive halogen from its halide.

Bromine is a liquid at room temperature whereas chlorine is a gas. Which explanation best accounts for this difference?

Bromine molecules contain ionic bonds whereas chlorine molecules do not.

Both Br₂ and Cl₂ are simple covalent molecules. Neither consists of ions.

Bromine is less reactive, so its molecules cannot move.

Reactivity does not determine whether particles can move. Liquid bromine molecules are continually moving.

Br₂ molecules are larger and have stronger intermolecular attractions.

Correct. More energy is required to overcome these stronger attractions, so bromine has a higher boiling point.

The covalent bond inside Br₂ is broken when bromine boils.

Boiling separates molecules from one another. The covalent Br–Br bonds within the molecules remain intact.

Down Group 17, the molecules become larger and intermolecular attractions become stronger. Melting and boiling points therefore increase.

An unfamiliar element Z is below potassium in Group 1. Which prediction is most reasonable?

Z should be less reactive because it has more electron shells.

More shells increase distance and shielding, making the valence electron easier rather than harder to remove.

Z should gain one electron during its reactions.

Group 1 atoms have one valence electron and normally lose it to form 1+ ions.

Z should have seven valence electrons.

Every Group 1 element has one valence electron.

Z should react with water more vigorously than potassium.

Correct. Group 1 reactivity increases down the group because the valence electron becomes easier to lose.

Periodic trends allow us to predict the behaviour of unfamiliar elements. An element below potassium should continue the trend of increasing Group 1 reactivity.

Which pair correctly shows the simple ions normally formed by magnesium and chlorine?

Mg⁺ and Cl⁺

Magnesium normally loses two electrons, while chlorine gains one electron.

Mg²⁺ and Cl⁻

Correct. Magnesium has two valence electrons and loses both. Chlorine has seven valence electrons and gains one.

Mg²⁻ and Cl⁺

This reverses the usual electron changes. Magnesium forms a positive ion, while chlorine forms a negative ion.

Mg⁸⁺ and Cl⁷⁻

Ionic charges are based on the number of electrons actually lost or gained, not simply copied from electron-shell numbers.

Mg: 2,8,2 → Mg²⁺ by losing two electrons. Cl: 2,8,7 → Cl⁻ by gaining one electron.

Which set of trends is correct as you move down Group 17?

Reactivity increases; boiling point decreases; colour becomes lighter.

All three trends are reversed. Reactivity decreases, boiling point increases and colour becomes darker.

Reactivity decreases; boiling point decreases; colour becomes darker.

The reactivity and colour trends are correct, but boiling point increases rather than decreases.

Reactivity decreases; boiling point increases; colour becomes darker.

Correct. These are the characteristic trends down Group 17.

Reactivity increases; boiling point increases; colour becomes darker.

The physical-property trends are correct, but reactivity decreases down Group 17.

Do not assume every property changes in the same direction. Reactivity decreases down Group 17, while melting and boiling points increase.

Why is argon suitable for providing an inert atmosphere around hot materials?

Argon has only one valence electron.

Argon is a Group 18 element and has a filled outer electron shell.

Argon reacts rapidly with oxygen and removes it.

Argon is useful because it is very unreactive, not because it reacts rapidly with oxygen.

Argon has a filled outer electron shell and is very unreactive.

Correct. Its stable electronic configuration means that it does not readily react with surrounding substances.

Argon is a reactive metal.

Argon is a noble gas and is a non-metal.

Group 18 atoms have filled outer electron shells and are therefore very unreactive under ordinary conditions.

Four elements lie consecutively in the same period. Their numbers of valence electrons are 1, 2, 3 and 4. What general change would you expect as you move from the first element to the fourth?

Metallic character generally decreases.

Correct. Moving from left to right across a period, the elements generally become less metallic.

The number of occupied electron shells increases by one each time.

Elements in the same period have the same number of occupied electron shells.

The proton number decreases.

Proton number increases from left to right across the Periodic Table.

All four elements should have identical chemical properties.

Similar chemical properties are associated mainly with elements in the same group, not consecutive elements across a period.

Across a period, the number of valence electrons increases while metallic character generally decreases as the elements change towards non-metallic behaviour.

Frequently asked questions

Questions students often ask about periodic trends

These questions connect the main trends and clarify a few ideas that can easily become oversimplified when the Periodic Table is treated as a list of facts.

Why do chemical properties repeat in the Periodic Table?

Chemical properties depend strongly on valence electrons. As proton number increases, electron shells fill in a repeating pattern. Eventually, the same type of outer-shell electronic configuration appears again in a new electron shell.

Lithium has the electronic configuration 2,1 while sodium has 2,8,1. Both have one valence electron, so both belong to Group 1 and show similar chemical behaviour.

Why do Group 1 and Group 17 have opposite reactivity trends?

Group 1 atoms react mainly by losing their single valence electron. Down the group, that electron is farther from the nucleus and more shielded, so it is lost more easily. Reactivity therefore increases.

Group 17 atoms react mainly by gaining an electron. Down the group, an incoming electron is farther from the nucleus and more shielded, so it is attracted less strongly. Reactivity therefore decreases.

Does every stable outer electron shell contain eight electrons?

No. The first electron shell can hold a maximum of only two electrons. Helium is therefore stable with the electronic configuration 2.

For many of the other elements commonly studied at O-Level, a filled outer shell contains eight electrons.

Why do Group 14 elements not simply form 4+ or 4− ions?

Group 14 atoms have four valence electrons. Losing four electrons or gaining four electrons is generally much less favourable than sharing electrons.

This is why carbon and silicon commonly form covalent bonds rather than simple 4+ or 4− ions.

Are all elements in the same group equally reactive?

No. Elements in the same group have similar types of reactions because they have similar outer-shell electronic configurations, but their reactivities can change down the group.

Sodium and potassium both react with water, for example, but potassium reacts more vigorously.

Why can reactivity decrease while boiling point increases?

These properties depend on different factors. Reactivity concerns how easily an atom undergoes an electron change during a chemical reaction.

Boiling point depends on the strength of attractions between particles. In Group 17, reactivity decreases down the group while intermolecular attractions become stronger, causing boiling point to increase.

Can I predict properties of an element that I have never studied?

Often, yes. If you know its position in the Periodic Table, you can make predictions based on trends.

For example, an unfamiliar Group 1 element below potassium would be expected to have one valence electron, form a 1+ ion and be more reactive than potassium.

Why does a halogen displacement reaction involve electron transfer?

A more reactive halogen gains electrons to form halide ions, while the halide ions of the less reactive halogen lose electrons and form halogen molecules.

For example:

Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

Chlorine gains electrons more readily than bromine, so chlorine can displace bromide ions.

Why are noble gases monatomic while halogens are diatomic?

Noble-gas atoms already have filled outer electron shells, so they do not normally need to bond to other atoms to obtain a stable electronic configuration.

Halogen atoms have seven valence electrons. Two halogen atoms can share a pair of electrons to form a covalent bond, producing molecules such as Cl₂ and Br₂.

Periodic trends become much easier when you stop treating them as separate facts to memorise. The position of an element tells you about its electronic configuration, and that electronic configuration helps explain its chemical behaviour. Periods reveal the number of occupied shells, groups reveal the outer-electron pattern, and changes in distance and shielding explain why reactivity changes down a group. Once these connections are clear, the Periodic Table becomes a tool for reasoning rather than a chart to memorise.