JC1 H1 Chemistry Tuition Hougang | Atomic and Ionic Radius Trends Across Period 3

JC1 H1 Chemistry tuition for Hougang students should make the Period 3 atomic-radius trend easier to explain than a memorised downward arrow. Sodium and chlorine are both in the third period, yet the neutral sodium atom is much larger in the familiar atomic-radius comparison. Across the period, more protons are added while electrons continue filling the same principal shell. The increasing effective nuclear attraction generally pulls the outer electron distribution closer to the nucleus.

This H1 Chemistry 8873 guide explains atomic radii and ionic radii, why cations shrink and anions expand, and how to compare isoelectronic ions such as Na⁺, Mg²⁺ and Al³⁺. It also shows why P³⁻, S²⁻ and Cl⁻ have a different electron count, why one continuous ionic-radius line across the entire period can be misleading, and how shielding and nuclear charge influence size. The lesson is useful for answering unfamiliar 2027 H1 periodicity questions using first principles rather than numbers remembered from an earlier worksheet.

eduKate Singapore uses a premium three-student small-group teaching approach in which every learner’s structural explanation can be checked. The teaching centre for enquiries is 8 Fourth Avenue, Singapore 268674, near Sixth Avenue MRT. Hougang identifies families who may travel for lessons; it is not a claim of a classroom in Hougang, and actual class availability, fees and schedule should be confirmed.

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The Immediate Answer: Neutral Atoms and Ions Are Two Different Comparisons

Neutral atoms across Period 3: the atomic radius generally decreases from sodium towards chlorine as nuclear attraction increases within the broadly same occupied principal shell.
Forming a positive ion: removing electrons can remove an entire outer shell and reduce electron repulsion, so a cation is typically smaller than its parent neutral atom.
Forming a negative ion: adding electrons increases repulsion within the occupied shell, so an anion is typically larger than its neutral parent.
Na⁺, Mg²⁺ and Al³⁺: all contain ten electrons; greater proton number makes the species smaller across this isoelectronic group.
P³⁻, S²⁻ and Cl⁻: all contain eighteen electrons; greater proton number makes the species smaller across this second isoelectronic group.
Do not compare all ions by one simple line: the cations and anions do not share the same electron configurations, and the species represented must be specified.

The underlying question is what determines the extent of an electron cloud. The protons attract electrons, other electrons partly shield that attraction, and the number and arrangement of occupied shells determine where electrons reside. Ion charge is a useful clue, but it must be interpreted with electron count and proton number.

Exactly Where the 2027 H1 Chemistry Syllabus Places This Topic

The official 2027 Singapore-Cambridge H1 Chemistry 8873 syllabus, section 4 The Periodic Table, names atomic radius and ionic radius across Period 3 from sodium to chlorine and down Group 17 from chlorine to iodine. It requires qualitative explanations in terms of nuclear charge, shielding and changes in occupied electron shells. No exact table of fixed radius values is needed to explain the prescribed trends.

Other physical properties, including first ionisation energy, electronegativity, melting point and conductivity, appear alongside periodic size in the same syllabus section. Those are related but are not the same measurable property. This guide stays focused on the atom-versus-ion size decision; the related Hougang guides develop ionisation and Period 3 structural-property graphs.

The Meaning of Atomic Radius

An atom does not have a sharp solid surface like a tiny glass bead. Electron density spreads over regions of space, becoming smaller farther from the nucleus rather than ending at one perfectly defined boundary. Published atomic radii use conventions, such as covalent, metallic or van der Waals radius, appropriate to a particular comparison.

That means two reference tables can report different numerical atomic-radius values while representing the same qualitative chemical trends. H1 students should avoid combining one element’s covalent radius with another’s van der Waals radius and treating the difference as a precise like-for-like experimental result.

For the normal Period 3 teaching comparison, the key trend is the progressive reduction in the size of the occupied electron distribution under increasing nuclear attraction. The physics of the electrons supplies the reason, not a rule that atoms literally shrink in a mechanical vise.

Why Nuclear Charge Increases Across a Period

Sodium has 11 protons, magnesium 12, aluminium 13, silicon 14, phosphorus 15, sulfur 16 and chlorine 17. Each step adds a proton to the nucleus. The corresponding neutral atom also gains an electron, but the added electron occupies the same principal n = 3 shell across this period in the basic electronic-configuration model.

The growing proton count strengthens the attraction between the nucleus and the electron distribution. Inner electrons still shield the outer shell, but the shielding does not increase by a full compensating amount with each electron added to the same shell. Effective nuclear attraction therefore tends to increase.

Worked Example 1: Sodium Compared With Magnesium

Write sodium’s electron configuration as 1s²2s²2p⁶3s¹.
Write magnesium’s electron configuration as 1s²2s²2p⁶3s².
Count protons: Na has 11, Mg has 12.
Both neutral atoms have outer electrons in the third principal shell.
The extra proton increases nuclear attraction more than the added same-shell electron increases shielding in the broad comparison.
Predict that neutral magnesium has a smaller atomic radius than neutral sodium.

The wrong answer ‘magnesium has more electrons, so it must be larger’ ignores where those electrons are and the attraction of the additional proton. Atomic size does not follow electron count alone.

Worked Example 2: Aluminium Compared With Sulfur

Aluminium has 13 protons and three outer-shell electrons, conventionally written [Ne]3s²3p¹. Sulfur has 16 protons and [Ne]3s²3p⁴. Although sulfur has more electrons, its increased nuclear charge tends to hold the electrons within the third shell more strongly.

Identify Al and S as neutral Period 3 atoms with the same inner [Ne] electron core.
Count protons: 13 for Al and 16 for S.
Recognise that both have valence electrons within n = 3.
Use the greater effective nuclear attraction at sulfur in the broad trend.
Conclude that neutral sulfur is generally smaller than neutral aluminium in a consistent atomic-radius comparison.

This is an explanatory trend, not permission to assign a fabricated exact radius. A graph question may supply measured or convention-based values, which should take precedence over a hand-drawn line.

What Electron Shielding Actually Does

Inner electrons repel outer electrons and partially offset the nuclear attraction experienced by those outer electrons. This is the familiar concept of shielding. Across a period, the inner core largely remains the same while the proton number increases. Down a group, additional principal shells introduce substantial extra shielding and distance from the nucleus.

Students sometimes say shielding remains perfectly constant across a period. That is an oversimplification: electron–electron repulsion and shielding can change. The relevant H1 statement is that nuclear attraction generally increases sufficiently across the same outer shell to produce the decreasing-radius trend.

The Difference Between a Neutral Atom and Its Cation

A positive ion forms when an atom or species loses one or more electrons. If the outermost shell is emptied, the cation’s highest occupied principal shell can be smaller than the parent atom’s. Fewer electrons also change the balance of electron–electron repulsion and nuclear attraction.

Sodium is a clear example. Neutral Na has the arrangement [Ne]3s¹, whereas Na⁺ has [Ne] with no occupied third shell. The cation is substantially smaller than the neutral atom. That is a stronger explanation than saying ‘positive charges are automatically small’.

Worked Example 3: Sodium Atom Versus Sodium Ion

Start with neutral sodium, proton number 11 and 11 electrons.
Use its configuration [Ne]3s¹.
Remove the outer 3s electron to make Na⁺ with 10 electrons.
The occupied n = 3 shell disappears from the simple configuration description.
Predict that Na⁺ is smaller than Na.
Explain the reduction in outer electron-shell extent rather than simply pointing to the plus sign.

The mass of the sodium ion is almost unchanged compared with the neutral sodium atom, because the lost electron has tiny mass relative to the nucleus. A change in electronic size is not a proportional change in atomic mass.

Worked Example 4: Magnesium Atom Versus Mg²⁺

Neutral Mg has 12 protons and 12 electrons, with configuration [Ne]3s².
Losing two 3s electrons makes Mg²⁺, with 10 electrons.
Mg²⁺ now has the [Ne] electron configuration.
Both removed electrons came from the outer n = 3 shell.
Predict that Mg²⁺ is smaller than neutral Mg.
Do not conclude that two positive charges mean exactly half the neutral radius.

The number of electrons lost matters, but it cannot be used as a direct linear formula for radius. The correct reasoning depends on the orbital arrangement before and after ion formation.

Why an Anion Usually Expands Relative to Its Atom

When an atom gains electrons to form a negative ion, it may increase electron–electron repulsion within the occupied outer shell. The same nucleus now attracts more electrons, and the resulting electronic distribution is typically more spread out than for the neutral atom in a relevant ionic-radius comparison.

A chlorine atom with 17 protons and 17 electrons gains one electron to make Cl⁻ with 18 electrons. The highest occupied shell remains n = 3, but the extra electron changes the repulsion and attraction balance. The anion is larger than its neutral parent in the familiar comparison.

Worked Example 5: Chlorine Atom Versus Chloride Ion

Neutral chlorine has 17 protons and 17 electrons, written [Ne]3s²3p⁵.
Chloride has the same 17 protons but 18 electrons, written [Ar].
Adding the electron fills the n = 3 valence shell in the simple model.
With more electrons experiencing repulsion around the same nucleus, the anion’s electron distribution expands in the ordinary comparison.
Conclude that Cl⁻ is larger than Cl.
Do not attribute the change to chlorine acquiring a new proton or additional nuclear mass.

The crucial constant is proton number: becoming an ion does not change the identity of the element. Any explanation involving different proton numbers would instead describe a different chemical element.

Isoelectronic Species: The Best H1 Shortcut

Isoelectronic species have the same total number of electrons, and in the straightforward examples below also have the same electron configuration. Their proton counts can differ. When comparing such ions, the species with more protons generally holds the same number of electrons more tightly and has the smaller radius.

This removes one major source of confusion: electron count is fixed, so differences can be explained primarily through nuclear attraction. The model is particularly transparent for Na⁺, Mg²⁺ and Al³⁺, which all contain ten electrons.

Worked Example 6: Rank Na⁺, Mg²⁺ and Al³⁺

Count electrons in Na⁺: 11 − 1 = 10.
Count electrons in Mg²⁺: 12 − 2 = 10.
Count electrons in Al³⁺: 13 − 3 = 10.
Recognise the shared [Ne] configuration.
Compare proton numbers: Na 11, Mg 12, Al 13.
With the same electron count, increased nuclear charge contracts the ionic electron distribution.
The typical ionic-radius order is Na⁺ greater than Mg²⁺ greater than Al³⁺.

A common wrong answer ranks the radius solely from the displayed charge as though negative signs determine geometric size. The isoelectronic electron count and proton number provide a defensible first-principles comparison.

Worked Example 7: Check Two Unknown Ten-Electron Ions

A question gives species X⁺ and Y²⁺ with ten electrons each and states that X has proton number 11 while Y has 12. Without identifying the elements first, predict which ion is larger.

Note both ions contain ten electrons.
Identify X⁺ as having 11 protons.
Identify Y²⁺ as having 12 protons.
At fixed electron count, Y²⁺ attracts those electrons more strongly.
Predict X⁺ has the larger radius.
Only then recognise the named examples as Na⁺ and Mg²⁺.

This changed representation is important. A student who can only answer when familiar chemical symbols are printed may be recalling a table rather than applying the electron-count model.

Isoelectronic Anions: P³⁻, S²⁻ and Cl⁻

Phosphorus has proton number 15 and can be represented as P³⁻ in simple ionic compounds; sulfur has proton number 16 and can form S²⁻; chlorine has proton number 17 and forms Cl⁻. Each of these ions has eighteen electrons, corresponding to the [Ar] configuration.

Within this isoelectronic series, P³⁻ has the smallest proton number and the largest ionic radius, while Cl⁻ has the largest proton number and the smallest of the three ionic radii. The sequence is therefore P³⁻ greater than S²⁻ greater than Cl⁻ in the standard comparison.

Worked Example 8: Rank Three Eighteen-Electron Ions

Count P³⁻ electrons: 15 + 3 = 18.
Count S²⁻ electrons: 16 + 2 = 18.
Count Cl⁻ electrons: 17 + 1 = 18.
All three have the same electron configuration.
Compare proton counts: 15, 16 and 17.
Predict the radius order P³⁻ greater than S²⁻ greater than Cl⁻.
Explain why more protons attract the same electron population more strongly.

This comparison is different from neutral atomic radii even though the element names still belong to Period 3. The species have additional electrons, so their properties must not be inferred from neutral-atom tables without adjustment.

Why There Is No Single Smooth Ionic-Radius Line Across Every Period 3 Element

The simple cation group Na⁺, Mg²⁺ and Al³⁺ has ten electrons, while the anion group P³⁻, S²⁻ and Cl⁻ has eighteen. It is not sensible to draw one smooth ionic-radius curve through every neutral Period 3 element without naming the particular ions being compared.

Silicon and other elements may participate in predominantly covalent bonding rather than exist as the simple monatomic ions implied by such a line. The relevant ionic radius also depends on the coordination and convention used in crystallographic measurements.

A scientific answer should specify the ionic species, not merely say ‘the ionic radius of Period 3 falls continuously’. The change between the familiar cation and anion sets is not one common isoelectronic progression.

Worked Example 9: Why Comparing Na⁺ With Cl⁻ Needs Care

Na⁺ contains ten electrons with a [Ne] configuration. Cl⁻ contains eighteen electrons with an [Ar] configuration. The two species have different occupied principal shells, so one cannot rank them through the simple ‘more protons wins with equal electrons’ rule.

Count electrons for Na⁺: ten.
Count electrons for Cl⁻: eighteen.
Recognise that the ion configurations differ.
Use structural or consistently defined ionic-radius evidence for a quantitative comparison.
Explain why the isoelectronic rule is not directly applicable.
Do not pretend both ions have the same shielding and electronic-shell count.

The general size of the outer electron cloud and the ionic environment still matter. The learning point is to choose the right comparison model before declaring a size order.

Down Group 17: More Shells Matter

The H1 syllabus also expects radius trends down Group 17 from chlorine through bromine to iodine. Here the main difference is that the outer electrons occupy increasingly distant principal shells, with greater shielding by inner electrons. Despite increasing proton number, the overall atomic radius increases down the group.

The same broad trend applies to the monatomic halide anions Cl⁻, Br⁻ and I⁻. They have different electron counts and occupied shells, with larger ionic radii down the group. This is not the same experiment as an isoelectronic ten-electron series.

Worked Example 10: Chlorine Compared With Bromine

Write chlorine’s outer electron shell as n = 3.
Write bromine’s outer shell as n = 4 in the familiar shell model.
Recognise an additional occupied principal shell for bromine.
Identify the increased electron distance and shielding down the group.
Explain why these factors generally outweigh the added nuclear charge for atomic size.
Predict that bromine has a larger atomic radius than chlorine.

The explanation should not say bromine is larger because it is heavier alone. Mass and size may correlate across a group, but the causal H1 account concerns electron shells, shielding and attraction.

Worked Example 11: Halide Ion Sizes

Compare chloride, bromide and iodide anions.
Recognise their outer occupied shells as n = 3, 4 and 5 respectively.
Describe increased shielding and extent of the outer electronic distribution.
Predict the familiar ionic size order Cl⁻ smaller than Br⁻ smaller than I⁻.
State that each has charge −1, so ionic charge alone cannot explain their differing sizes.

The point transfers to other groups. A student who uses proton number alone would predict the opposite down-group trend and miss the effect of new shells.

Atomic Radius and First Ionisation Energy Are Related, Not Identical

The tendency for neutral atomic radius to decrease across Period 3 often accompanies a broad rise in first ionisation energy: outer electrons experience stronger attraction and can require more energy to remove. But subshell differences and electron pairing cause local deviations in ionisation energies that are not simply the inverse of radius.

For example, magnesium and aluminium show a first-ionisation exception because aluminium’s outer 3p electron is in a different subshell environment from magnesium’s 3s electron. The correct student response distinguishes a broad size trend from a precise electronic-energy comparison.

Worked Example 12: Why a Shorter Radius Does Not Always Give the Higher First IE

A student sees neutral aluminium smaller than neutral magnesium in the broad Period 3 atomic-radius comparison and therefore predicts that aluminium’s first ionisation energy must be higher. The actual familiar first-ionisation comparison has Mg above Al.

Recognise both neutral atoms have n = 3 outer electrons.
Mg’s first electron removed is from 3s.
Al’s first electron removed is from 3p.
The 3p electron is in a higher-energy, less tightly bound subshell environment in the relevant comparison.
Therefore Al’s first ionisation energy can be lower despite its generally smaller atomic radius.
State that radius alone is not a complete predictor of ionisation-energy differences.

This is a particularly valuable extension for learners who have memorised two trends as though they were mathematically identical.

Electronegativity and Radius

As effective attraction for shared electrons increases across a period, electronegativity also tends to rise. A smaller atom often attracts bonding electrons more strongly, but bond polarity depends on the two atoms involved and their chemical environments. Electronegativity is not simply an ionic radius measured in another unit.

A tutor should ask a student whether the question concerns a free neutral atom, an ion in a crystal or an atom in a covalent bond. All three involve electrons and nuclei, yet the most appropriate model and reported quantity differ.

Worked Example 13: Explain a Sodium–Chlorine Bond Without Calling It Atomic-Size Ranking

Sodium and chlorine are different neutral atoms before reaction. In the ordinary ionic model, sodium loses an electron and chlorine gains one, forming Na⁺ and Cl⁻. The resulting ionic sizes and strong lattice attractions must be analysed using those ionic species.

Identify the neutral atom configurations Na: [Ne]3s¹ and Cl: [Ne]3s²3p⁵.
Describe transfer of one electron in the simplified ionic formation picture.
Write the ions Na⁺ and Cl⁻.
Recognise Na⁺ has ten electrons, while Cl⁻ has eighteen.
Explain why their ionic radii cannot be copied unchanged from neutral Na and Cl atomic radii.
Use the extended ionic-lattice model, not a pair of isolated neutral atoms, to discuss melting and conduction.

This connects periodicity with the H1 ionic-lattice topic without repeating the entire NaCl-versus-MgO melting-point guide.

Measurements: Coordination Environment Can Affect Reported Ionic Radii

Ionic radius is often inferred from separations in crystal structures, rather than measured by placing a ruler between an ion’s visible edge and its nucleus. The effective size of an ion in a particular coordination environment can depend on the adopted radius convention and how many neighbours surround it.

H1 students need not calculate crystallographic radii from advanced diffraction data to understand the prescribed general trend. They should, however, avoid combining inconsistent definitions into one quantitative graph without considering how the values were determined.

Worked Example 14: Evaluate an Inconsistent Data Table

A question quotes a metallic radius for neutral sodium and a six-coordinate crystal ionic radius for Mg²⁺, then asks for a precise percentage change ‘from sodium to magnesium’ as though the two numbers were the same kind of measurement.

Identify one reported quantity as a neutral metallic atomic radius.
Identify the other as an ionic radius of Mg²⁺ in a specified crystal environment.
Recognise that electron count, charge and measurement convention are different.
State that a calculated numerical difference would not be a direct like-for-like atomic-size trend.
Request consistent atomic or ionic radius definitions, or focus on an appropriate qualitative explanation.

The result tests scientific literacy: numbers are not automatically comparable because they share the unit picometres.

A Reliable Six-Step Radius Question Method

Write down the exact species: neutral atom, cation or anion.
Count the number of protons from atomic number.
Count the number of electrons after any ion charge is applied.
Write or recognise the occupied principal shells.
Check whether the species are isoelectronic and choose the appropriate nuclear-attraction comparison.
Explain the size trend using shielding, occupied shells and nuclear charge rather than charge symbols alone.

That sequence works across the named Period 3 and Group 17 examples. The short diagnostic should precede all calculations or memorised radius rankings.

Seven Common H1 Periodicity Errors

Claiming more electrons always make a neutral atom larger across a period.
Claiming more protons always make an atom smaller down a group without accounting for added shells.
Treating Na⁺ and Cl⁻ as an isoelectronic pair.
Ranking Na⁺, Mg²⁺ and Al³⁺ without counting their ten electrons.
Assuming Cl⁻ is smaller than its parent neutral Cl because it contains a full shell.
Comparing published radii drawn from different definitions as though they were equivalent measurements.
Calling an atom’s mass the sole determinant of its electronic radius.
Inferring every individual first-ionisation exception from atomic radius alone.
Drawing a smooth ionic-radius curve across species that do not exist as equivalent simple monatomic ions.

What a Premium Three-Student Tutorial Makes Visible

Imagine three students attempting a radius ranking. One knows Na⁺, Mg²⁺ and Al³⁺ are ten-electron ions but mistakenly says Al³⁺ is largest because its charge is greatest. Another applies the same ten-electron rule to Cl⁻ and Na⁺. The third knows both rankings but incorrectly explains down-group size by nuclear charge rather than occupied shells.

All three need different corrections. A tutor working with three students can inspect the first written explanation before giving the model, ask an individual repair question and then change the symbols. Each student must independently complete the new ranking.

An Illustrative 90-Minute JC1 H1 Session

0–10 minutes: retrieve proton number, electron count and Period 3 shell configurations.
10–25 minutes: compare neutral Na, Mg, Al and S atomic sizes with reasoning.
25–40 minutes: form Na⁺, Mg²⁺, Al³⁺ and rank the ten-electron ions.
40–55 minutes: rank the eighteen-electron P³⁻, S²⁻ and Cl⁻ ions.
55–68 minutes: explain neutral versus cation/anion size and down-Group-17 trends.
68–80 minutes: critique inconsistent radius and ionisation-energy explanations.
80–90 minutes: complete an unfamiliar isoelectronic set and plan delayed retrieval.

This schedule is a teaching framework, not a claim that every actual group follows the same times. The tutor should adapt to the JC1 student’s current college syllabus, conceptual readiness and available class placement.

Three Learner Pathways

Repair: atom and ion electron counts are unstable

Begin with three simple species and count their protons and electrons. Do not jump into radius trends until the learner can explain that forming Na⁺ removes an electron but keeps its eleven protons.

Stabilisation: correct trends, weak explanations

Require the learner to name the outer occupied shell and decide whether the species are isoelectronic before writing a radius order. Replace vague ‘more charge’ answers with the actual proton–electron attraction argument.

Extension: mixed radius and periodic-property evidence

Present an unfamiliar chart containing ionic radius, first ionisation energy and electronegativity. Ask which trends are broadly related and where a local exception needs an orbital explanation. Require an appropriately qualified conclusion.

Four-Week Independent Retrieval

Week 1: neutral Period 3 atomic-radius trends and electronic configurations.
Week 2: cation formation and the ten-electron isoelectronic series.
Week 3: anion formation, the eighteen-electron series and Group 17 down-group trends.
Week 4: mixed unfamiliar-radius comparisons and short explanations of model limits.

The learner may need fewer or more sessions. The point is delayed independent recall rather than repeated copying of one ranking table.

What Hougang Parents Can Check After a Lesson

Ask the student why neutral magnesium is smaller than neutral sodium despite having more electrons, why Na⁺ is smaller than Na, and why Cl⁻ is larger than Cl. Then give two unfamiliar ions with the same electron count but different nuclear charges and ask which is smaller.

A learner who can answer all four without notes is demonstrating transferable understanding of the atomic model. Some students can achieve the same progress through college consultations and deliberate self-study; tuition is most useful where a specific gap remains visible.

Class Details and Venue for Families From Hougang

Level: JC1 H1 Chemistry 8873. Topic: Period 3 atomic and ionic radii, isoelectronic species, shielding, nuclear attraction and Group 17 size trends. Method: premium three-student small-group tutorials with independent first attempts, worked comparisons and targeted correction. Typical session to enquire about: a compatible 90-minute weekly group where available. Teaching centre: eduKate Singapore, 8 Fourth Avenue, Singapore 268674, near Sixth Avenue MRT.

Hougang describes the family search audience, not the class venue. Current fees, timetable, available places and travelling arrangements require confirmation. No particular examination result can be guaranteed.

Frequently Asked Questions

Why does atomic radius decrease across Period 3?

Nuclear charge increases while the outer electrons remain in the same principal shell, so effective attraction generally contracts the neutral electron distribution.

Is Na⁺ smaller than neutral Na?

Yes. Sodium loses its third-shell 3s electron to form a ten-electron Na⁺ ion with the [Ne] configuration.

Which is smallest: Na⁺, Mg²⁺ or Al³⁺?

Al³⁺. All three have ten electrons, and aluminium has the greatest proton number, giving the strongest attraction within this isoelectronic series.

Which is largest: P³⁻, S²⁻ or Cl⁻?

P³⁻ in the standard comparison. All have eighteen electrons, and phosphorus has the fewest protons.

Why is Cl⁻ larger than Cl?

The chloride anion has an additional electron around the same nucleus, increasing repulsion and expanding the electron distribution in the familiar comparison.

Why do halogen atomic radii increase down Group 17?

Additional occupied principal shells and shielding outweigh the increasing nuclear charge in the broad size trend.

Is atomic radius measured as one fixed visible edge?

No. Reported radii follow conventions and bonding environments; electron density has no sharp solid boundary.

Are Chemistry classes located in Hougang?

No. The stated centre is Fourth Avenue near Sixth Avenue MRT. Hougang families may enquire about suitable current groups.

Decide Which Electrons Are Being Compared

The strongest periodic-radius answer identifies the species, counts the electrons and protons, recognises the relevant occupied shell, and then explains the attraction and shielding. When that sequence is secure, a new Period 3 or Group 17 comparison becomes a reasoning task rather than a ranking to memorise.

Arrange a parent–student consultation with eduKate Singapore to review the student’s current H1 periodicity work and suitable premium three-student teaching provision.