Mendeleev had a problem: 63 known elements, no organising principle, and chemistry was a confused mess. His insight was to arrange them by atomic mass and look for repeating patterns in properties. He was so confident in the pattern that he left blank spaces for undiscovered elements — and predicted their properties before anyone had seen them. When those elements were discovered (gallium in 1875, germanium in 1886), they matched his predictions almost exactly. The periodic table isn't just a chart — it's one of the greatest predictive tools in science.
Use the PDF for classwork, homework or revision. It includes key ideas, activities, questions, an extend task and success-criteria proof.
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).
Before reading on, write your best answer. What problem would Mendeleev's mass-based order create with these two elements?
Core Content
Wrong: Noble gases are inert because they have no electrons.
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.
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).
| 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 |
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) |
Worked Examples
Method A — Using position directly
Step 1 — Identify element Y
Group 2, Period 5 → alkaline earth metal in the 5th period
Counting down Group 2: Be (Period 2), Mg (Period 3), Ca (Period 4), Sr (Period 5)
Element Y = Strontium (Sr)
Step 2 — Valence electrons from group number
Group 2 → 2 valence electrons (all alkaline earth metals have 2 valence e⁻)
Step 3 — Predict reactivity vs Ca
Y (Sr) is below Ca in Group 2 (Period 5 vs Period 4).
Going down a group: atomic radius increases → valence electrons are further from the nucleus and shielded by more inner electron shells → weaker nuclear attraction on valence electrons → easier to lose them in reactions.
For metals: easier to lose valence electrons = more reactive.
Conclusion: Sr is MORE reactive than Ca.
Method B — Extrapolate from periodic table structure
Step 1 — Determine period
Current table: Period 7 ends at oganesson (Og, Z = 118).
Z = 119 would be in Period 8 (next row after Period 7 ends).
Period 8, first element → Group 1 (same column as H, Li, Na, K, Rb, Cs, Fr)
Step 2 — Determine group and block
Group 1, Period 8 → element is in s-block (Group 1 = s-block)
1 valence electron (Group 1 always has 1 valence e⁻)
Step 3 — Predict physical properties
All Group 1 elements are: soft, low melting point, highly reactive metals with 1 valence e⁻.
Reactivity trend: increases down the group (Fr > Cs > Rb > K > Na > Li).
Q (Period 8) would be the most reactive alkali metal known — even more reactive than Fr.
Likely very low density, very soft, reacts extremely violently with water.
Step 4 — Compare Methods A and B
Method A (known element): use Z and position directly.
Method B (undiscovered): extrapolate from group/period trends.
Both methods use the same principle: group = valence electrons = chemical character.
Activities
1 For each element, state the group, period, block, and number of valence electrons: (a) Lithium (Z=3), (b) Oxygen (Z=8), (c) Calcium (Z=20), (d) Bromine (Z=35).
2 An unknown element M is described as: "A soft, silvery metal that reacts violently with water, forming a strongly alkaline solution and releasing hydrogen gas. Its melting point is lower than sodium." Identify the most likely group for M. Which specific element is most likely M? Justify using periodic trends.
A Compare Mendeleev's periodic table and the modern periodic table by discussing: (i) the organising principle used, (ii) one advantage of the modern approach, (iii) one example where Mendeleev's approach was insufficient and how the modern approach resolved it.
B Element Z is in Period 3, Group 16. (a) Name and identify this element. (b) Predict its charge when it forms a simple ion. (c) Predict whether it would be more or less electronegative than the Group 16 element in Period 2. Justify.
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Multiple Choice
5 random questions from a replayable lesson bank — feedback shown immediately
Short Answer
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
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.)
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.
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).
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).
Return to your Think First response. You should now be able to explain why atomic number is superior to atomic mass for organising the periodic table:
The Periodic Table: Organisation
Tick when you've finished all activities and checked your answers.