The Periodic Table: Organisation
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Four printable worksheets that build from the foundations up to exam-style questions — start at whatever level suits you.
Mendeleev arranged the elements in order of increasing atomic mass and left gaps for elements that had not yet been discovered. The modern periodic table arranges elements in order of increasing atomic number (number of protons). Why is atomic number a better organising principle than atomic mass? Consider the elements argon (Ar, Z = 18, A = 39.95) and potassium (K, Z = 19, A = 39.10).
Key facts
- How the periodic table is organised (periods = rows, groups = columns)
- Names and properties of key groups — alkali metals (1), alkaline earths (2), halogens (17), noble gases (18), transition metals
- Mendeleev ordered by atomic mass; the modern table orders by atomic number (Z)
Concepts
- Why elements in the same group share similar chemistry — same number of valence electrons
- Why Z ordering fixes Mendeleev's Te/I and Co/Ni anomalies
- How block structure (s, p, d, f) reflects which subshell is being filled
Skills
- From Z, determine an element's group, period and block
- Predict valence electrons and most likely ion charge from group position
- Compare reactivity of two elements in the same group using shell distance and shielding
The modern periodic table arranges elements in order of increasing atomic number (Z), not atomic mass (as Mendeleev used). This resolves a few inconsistencies Mendeleev faced (e.g. Te/I and Co/Ni are out of order by mass but in the correct order by Z).
The modern periodic table orders elements by increasing atomic number (Z) — not atomic mass — so chemical properties repeat because elements with the same valence-electron count align in vertical groups. Period number = number of occupied electron shells; main-group number = number of valence electrons. 118 elements confirmed; Mendeleev's predicted gaps (Ga, Sc, Ge) all matched.
Pause — copy the highlighted periodic table rule into your book before moving on.
Odd one out: three of these describe the modern periodic table; one does not. Which is the odd one out?
Periods = horizontal rows (period number = number of shells). Groups = vertical columns (main-group number = valence electrons). Metals: left + d-block (lose electrons → cations); non-metals: upper right (gain electrons → anions); metalloids: staircase boundary. Metal reactivity increases down a group; non-metal reactivity decreases down a group.
Add the highlighted period/group rules to your notes before the check below.
True or false: two elements in the same period have the same number of valence electrons.
| Group | Name | Valence e⁻ | Key properties | Examples |
|---|---|---|---|---|
| 1 | Alkali metals | 1 | Soft, low MP, very reactive metals; react vigorously with water to form H₂ + metal hydroxide; reactivity increases down the group | Li, Na, K, Rb, Cs, Fr |
| 2 | Alkaline earth metals | 2 | Harder/denser than Group 1; still reactive; form 2+ ions; less reactive than Group 1 | Be, Mg, Ca, Sr, Ba, Ra |
| 17 | Halogens | 7 | Non-metals; exist as diatomic molecules (F₂, Cl₂, Br₂, I₂); very reactive (one electron from full shell); reactivity decreases down the group; form −1 ions (halides) | F, Cl, Br, I, At |
| 18 | Noble gases | 8 (He: 2) | Full valence shell → extremely unreactive; monatomic gases; no tendency to form bonds; very low BP (weak dispersion only) | He, Ne, Ar, Kr, Xe, Rn |
| 3–12 | Transition metals | 1–2 (d-block) | Hard, high MP, multiple oxidation states; form coloured compounds; good conductors; catalytic properties | Fe, Cu, Zn, Ti, Cr, Mn |
We just saw the metal/non-metal divide and how period/group positions encode electron structure. That raises a question: what are the defining chemical properties of each main group, and why do those properties change predictably? This card answers it → each group's chemical behaviour is determined by its valence electron count and how easily those electrons are lost or gained.
Key groups: Group 1 — alkali metals (1 valence e⁻; react vigorously with water → H₂ + hydroxide). Group 2 — alkaline earth metals (2 valence e⁻; less reactive). Group 17 — halogens (7 valence e⁻; diatomic; reactivity decreases down group: F₂ > Cl₂ > Br₂ > I₂). Group 18 — noble gases (full outer shell; monatomic; almost completely unreactive).
Pause — write the highlighted key group properties into your book.
Cloze: drag the group labels into the right slots.
Group 1 elements are the ___, which have 1 valence electron and react vigorously with water. Group 17 elements are the ___, which exist as diatomic molecules and form −1 ions. Group 18 elements are the ___, which have full valence shells and are extremely unreactive. The d-block (Groups 3–12) holds the ___, which form coloured compounds and have multiple oxidation states.
The block structure of the periodic table directly maps to which subshell the highest-energy (outermost) electrons occupy — a concept you'll explore fully in L16. For now, the key locations:
| Block | Groups | Subshell filling | Element type |
|---|---|---|---|
| s-block | 1, 2 | s subshell | Reactive metals (alkali/alkaline earth) + H and He |
| p-block | 13–18 | p subshell | Non-metals, metalloids, and some metals |
| d-block | 3–12 | d subshell | Transition metals |
| f-block | (separate rows) | f subshell | Lanthanides, actinides (rare earths) |
We just saw the key properties of main-group families. That raises a question: the periodic table is divided into s, p, and d blocks — what do these blocks tell you about electron configuration? This card answers it → the block indicates which type of subshell is being filled by the final electron, which directly connects table position to electron configuration.
s-block = Groups 1–2 (outermost electron in s subshell). p-block = Groups 13–18 (outermost in p subshell — includes non-metals, metalloids, some metals). d-block = Groups 3–12 (transition metals, d subshell fills). Quick rule: period = number of occupied shells; main-group number = valence electrons; block = the type of subshell being filled.
Add the highlighted block rules to your notes before the check below.
Odd one out: three of these elements are in the s-block. Which one is not?
6. Evaluate Mendeleev's contribution to chemistry, including: (a) the organising principle he used, (b) an example of a successful prediction he made about an undiscovered element, and (c) one limitation of his approach that the modern periodic table resolved. 4 MARKS
7. An unknown element W has Z = 20. (a) Determine its group, period, and block. (b) State the number of valence electrons and predict its ion charge. (c) Compare the reactivity of W to magnesium (Z=12) and explain using electron shell theory. 5 MARKS
We just saw the s, p, and d block organisation. That raises a question: how do you structure full-mark exam answers on the periodic table, including Mendeleev's contributions and reactivity trend explanations? This card answers it → for each question type, use the structured approach: state the trend, explain the electron-structure reason (Z, shielding, shell count).
For Mendeleev evaluation answers: name three features — ordering principle (atomic mass), a successful prediction (Ga, Sc, or Ge), and a key limitation (Te/I anomaly). For "predict position" questions: count shells → period; count valence electrons → group. For reactivity trend questions: down a group → more shells → more shielding → metals lose electrons more easily; non-metals gain electrons less easily.
Pause — copy the highlighted exam strategies into your book before moving on.
True or false: calcium (Ca, Z=20) is more reactive than magnesium (Mg, Z=12) because its valence electrons are further from the nucleus and more shielded.
Worked examples · reveal as you go
Element Y is in Group 2, Period 5. Predict Y's reactivity compared to calcium (Group 2, Period 4).
An element Q has Z = 119 (not yet discovered/confirmed). Using the periodic table, predict: (a) the group and period of Q, (b) the block, (c) likely physical properties (metal/non-metal), (d) number of valence electrons.
An unknown element X has atomic number Z = 38. Predict its group, period, block, and the charge of the ion it is most likely to form.
How close was your prediction?
Common errors · the 3 traps that cost marks
Misconception to fix
Wrong: Noble gases are inert because they have no electrons.
Misconception to fix
Right: Noble gases have full valence shells (8 electrons, or 2 for helium), which makes them extremely unreactive. They have the same number of electrons as any other element — it is the stable electron configuration, not the absence of electrons, that explains their low reactivity.
Confusing group number with valence electrons for the d-block
Group 6 is a transition-metal column (d-block), not the column of elements with 6 valence electrons. The "main-group number = valence electrons" rule only works for s-block (Groups 1–2) and p-block (Groups 13–18). An element with 6 valence electrons (s²p⁴) sits in Group 16, not Group 6.
Fix: For any p-block element, count valence electrons then convert with Group = 10 + (valence e⁻) so 6 valence e⁻ → Group 16.
Quick-fire practice · 5 reps +2 XP per reveal
Name the group: elements with 1 valence electron that react vigorously with water.
For sulfur (Z = 16), state the period, group and block.
Predict the most likely ion charge for an element in Group 16.
Why does atomic number, not atomic mass, fix the Te/I ordering problem?
Element X (Z = 38) and magnesium (Z = 12) are both in Group 2. Predict which is more reactive and justify in one sentence.
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Pick your answer, then rate your confidence — that tells the system what to drill next.
A student places an unknown element with Z = 16 on the periodic table. One line contains an error — click it.
- Z = 16 → electron configuration by shells = 2, 8, 6
- 3 occupied shells → Period 3
- 6 valence electrons → Group 6
- Outer electrons fill the p subshell → p-block element (sulfur, S)
Q1. 6. Evaluate Mendeleev's contribution to chemistry, including: (a) the organising principle he used, (b) an example of a successful prediction he made about an undiscovered element, and (c) one limitation of his approach that the modern periodic table resolved.
Q2. 7. An unknown element W has Z = 20. (a) Determine its group, period, and block. (b) State the number of valence electrons and predict its ion charge. (c) Compare the reactivity of W to magnesium (Z=12) and explain using electron shell theory.
📖 Comprehensive answers (click to reveal)
Activity 1
1. (a) Li (Z=3): Group 1, Period 2, s-block, 1 valence electron. (b) O (Z=8): Group 16, Period 2, p-block, 6 valence electrons. (c) Ca (Z=20): Group 2, Period 4, s-block, 2 valence electrons. (d) Br (Z=35): Group 17, Period 4, p-block, 7 valence electrons.
2. Group 1 (alkali metal) — reacts with water to form alkaline solution and H₂ is characteristic of Group 1. Lower MP than Na suggests it is further down Group 1 than Na (Period 3). MP decreases: Li > Na > K > Rb > Cs. Element M is most likely potassium (K) — one period below Na, MP = 63°C vs Na's 98°C. (Could also be Rb or Cs if even lower MP.)
Activity 2
A: (i) Mendeleev used increasing relative atomic mass as the ordering principle; the modern table uses increasing atomic number (Z). (ii) The modern approach correctly places Te before I — by Ar, Te (127.6) > I (126.9) which would put I before Te by mass; by Z, Te (52) < I (53), placing Te first as its chemical properties require (Te is in Group 16, I in Group 17). (iii) The same Te/I anomaly: Mendeleev was forced to override mass ordering to preserve chemical group relationships, acknowledging his principle was imperfect. The modern table resolved this completely because Z directly determines electron configuration and chemical properties, explaining why some elements are "out of order" by mass.
B: (a) Period 3, Group 16 = Sulfur (S). (b) Group 16 elements have 6 valence electrons; they need 2 more to reach a full shell of 8 → S forms S²⁻ ion (charge = 2−). (c) Less electronegative than the Period 2 element (oxygen, O). Electronegativity decreases down a group because the valence shell is further from the nucleus and more shielded → weaker nuclear attraction on the bonding electrons → less tendency to attract electrons → lower electronegativity.
❓ Multiple Choice
1. C — Period number = number of occupied electron shells. Group = valence electrons. Ar is the ordering principle of Mendeleev's (not modern) table.
2. B — Z ordering correctly places Te (Z=52) before I (Z=53). Their chemical properties (Te is Group 16, I is Group 17) are consistent with Z ordering. Ar ordering fails here.
3. D — Same valence electrons = same bonding tendencies = similar chemistry. Same shells = same period, not group.
4. A — The valence electron distance and shielding explanation. B is wrong — Group 1 always has 1 valence electron. MP (D) is a result of reactivity, not a cause.
5. C — Halogens gain an electron to complete their valence shell. Fluorine (Period 2) has its valence shell closest to the nucleus → strongest nuclear pull → greatest tendency to gain → most reactive. Physical state (D) affects rate, not reactivity (tendency).
Short Answer Model Answers
Q6 (4 marks): (a) Mendeleev organised elements in order of increasing relative atomic mass and identified that properties repeated at regular intervals — he arranged elements with similar properties into vertical groups (1 mark). (b) Example: Mendeleev predicted the existence of eka-aluminium (later discovered as gallium, 1875) and eka-silicon (germanium, 1886), leaving gaps in his table and predicting their properties (atomic mass, density, valence) from the surrounding elements — the actual properties of the discovered elements closely matched his predictions, validating the periodic pattern (1 mark). (c) Limitation: ordering by mass led to anomalies where elements would be misplaced — Te (Ar 127.6) would come after I (Ar 126.9) by mass, but their chemical properties required Te in Group 16 and I in Group 17 (opposite order). Mendeleev overrode mass ordering empirically. The modern table resolved this by using Z instead: Te (Z=52) correctly precedes I (Z=53) (1 mark). Overall evaluation: Mendeleev's table was a landmark achievement — it unified chemistry and had predictive power, though it was later refined when atomic structure was understood (1 mark).
Q7 (5 marks): (a) Z=20. Electron shells: 2,8,8,2 → 4 occupied shells → Period 4. Last 2 electrons in s subshell → s-block, Group 2. Element W = calcium (Ca) (1 mark). (b) Valence electrons = 2 (Group 2). Ca loses 2 electrons to achieve noble gas configuration → Ca²⁺ ion (charge = 2+) (1 mark). (c) W (Ca, Period 4) is more reactive than Mg (Period 3). Both are in Group 2 (s-block, 2 valence electrons). Going from Period 3 to Period 4: atomic radius increases (additional electron shell) → valence electrons are further from the nucleus (1 mark) → more inner electron shells provide greater shielding of nuclear charge (1 mark) → weaker effective nuclear attraction on the 2 valence electrons → they are more easily removed in reactions → Ca is more reactive than Mg (1 mark).
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