5.1
Historical Background of the Periodic Table
Early Attempts at Classification
By the mid-19th century, over 60 elements had been discovered. Chemists noticed patterns in their properties and tried to organise them systematically.
| Year | Scientist | Contribution | Limitation |
|---|---|---|---|
| 1829 | J. W. Döbereiner | Law of Triads: groups of three elements with similar properties where the middle element's atomic mass was the average of the other two (e.g. Li, Na, K; Cl, Br, I; Ca, Sr, Ba). | Only worked for a few groups; could not classify all known elements. |
| 1864 | John Newlands | Law of Octaves: when elements were arranged by increasing atomic mass, every 8th element had similar properties to the first (like musical octaves). | Only worked for the first 20 elements. Failed for heavier elements; ridiculed by contemporaries. |
| 1869 | Dmitri Mendeleev | Periodic Table: arranged 63 known elements by increasing atomic mass and grouped by similar chemical properties. Left gaps for undiscovered elements and predicted their properties. | Had to reverse atomic mass order for some pairs (e.g. Te/I); based on mass not atomic number. |
| 1869 | Lothar Meyer | Independently proposed a similar arrangement, focusing on physical properties (atomic volumes). | Did not predict missing elements as effectively as Mendeleev. |
| 1913 | Henry Moseley | Used X-ray spectra to determine atomic numbers. Showed the periodic table should be arranged by atomic number (not mass), resolving anomalies. | — |
Why Mendeleev succeededMendeleev's key insight was leaving gaps for unknown elements rather than forcing elements into ill-fitting places. He predicted properties of eka-aluminium (gallium, discovered 1875), eka-boron (scandium, 1879), and eka-silicon (germanium, 1886) with remarkable accuracy — this validated his approach.
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Section 5.1 — Atomic Radius Trends
10 QuestionsQ1 of 10
Across a period (left to right), atomic radius:
Across period: Z increases, same shell → electrons added to same shell → same shielding → increasing Zeff → electrons pulled closer → decreasing radius.
Q2 of 10
Down a group, atomic radius:
Down group: new shells added → outer electrons further from nucleus → atomic radius increases despite increasing Z.
Q3 of 10
Which has the largest atomic radius?
Na is leftmost in Period 3 → largest radius (186 pm). Radius decreases across period: Na > Mg > Al > Si.
Q4 of 10
Ionic radius of Na⁺ compared to Na atom:
Na⁺: lost 1 electron from outer shell (3s¹ → gone). Fewer electrons, same protons → greater nuclear pull per electron → smaller ion. Na: 186 pm; Na⁺: 102 pm.
Q5 of 10
Ionic radius of Cl⁻ compared to Cl atom:
Cl⁻: gained 1 electron. More electrons (18 vs 17), same protons → more electron-electron repulsion → larger. Cl: 99 pm; Cl⁻: 181 pm.
Q6 of 10
Isoelectronic ions (same electrons): which is smallest?
All have 10 electrons. More protons → smaller. Na⁺ (Z=11) < F⁻ (Z=9) < O²⁻ (Z=8) < N³⁻ (Z=7). Na⁺ smallest in this isoelectronic series.
Q7 of 10
Effective nuclear charge (Zeff) increases across a period because:
Across period: protons increase (Z↑), electrons added to same shell (same shielding), so Zeff = Z − shielding increases.
Q8 of 10
Why does K have a larger atomic radius than Na despite more protons?
K: Period 4, outer electron in n=4 shell → further from nucleus → larger radius. The extra shell effect outweighs increased nuclear charge.
Q9 of 10
Transition metals (e.g. Fe, Co, Ni) have similar atomic radii because:
In transition metals: electrons fill 3d (inner). 3d electrons shield outer 4s electrons from increasing nuclear charge → Zeff on outer electrons changes little → approximately constant atomic radius across the row.
Q10 of 10
Which has a larger radius: Fe²⁺ or Fe³⁺?
Fe²⁺: 24 electrons. Fe³⁺: 23 electrons. Same protons, fewer electrons in Fe³⁺ → less repulsion → greater effective nuclear charge per electron → smaller radius. Fe²⁺ > Fe³⁺.
5.2
Mendeleev's Table vs the Modern Periodic Table
| Feature | Mendeleev's Table (1869) | Modern Periodic Table |
|---|---|---|
| Ordering basis | Increasing relative atomic mass | Increasing atomic number (proton number) |
| Number of elements | 63 elements known | 118 elements (as of current knowledge) |
| Groups | Vertical columns of similar chemical properties | 18 groups; elements in same group have same outer electron configuration |
| Periods | Horizontal rows | 7 periods; period number = number of electron shells |
| Noble gases | Not included (undiscovered) | Group 18 (He, Ne, Ar, Kr, Xe, Rn) |
| Gaps | Left gaps for undiscovered elements | No gaps; all positions filled |
| Anomalies | Some pairs out of order (e.g. Te before I) to preserve chemical groupings | Resolved by atomic number: correct order without anomalies |
| Blocks | Not recognised | Divided into s, p, d, f blocks based on electron configuration |
| Lanthanides/actinides | Not placed separately | Separated into f-block below main table |
The Te/I anomaly resolvedMendeleev placed Te (atomic mass 127.6) before I (atomic mass 126.9) despite I having a lower mass, to keep them in the correct groups chemically. Moseley showed Te has Z=52 and I has Z=53 — so the modern order (by atomic number) naturally places Te before I, resolving the anomaly.
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Section 5.2 — Electronegativity & Bond Polarity
10 QuestionsQ1 of 10
Electronegativity measures:
Pauling electronegativity: ability to attract shared electrons in a covalent bond. Increases across period and up a group. F is highest (4.0); Cs is lowest (0.79).
Q2 of 10
Across a period, electronegativity:
Across period: Zeff increases, same shell → stronger pull on bonding electrons → increasing electronegativity.
Q3 of 10
The most electronegative element is:
Fluorine (F): highest electronegativity = 4.0 (Pauling scale). F is small (strong nuclear pull) and very electronegative → forms very polar bonds.
Q4 of 10
A bond between atoms with electronegativity difference > 1.7 is typically:
Rough guide: ΔEN > 1.7 → ionic character dominates. 0.4–1.7 → polar covalent. < 0.4 → non-polar covalent. Not absolute — just a guide.
Q5 of 10
In the H-F bond (EN: H=2.2, F=4.0), the partial charges are:
F more EN → pulls electrons → F is δ⁻, H is δ⁺. The bond dipole points from H toward F.
Q6 of 10
Which molecule has the most polar bonds?
ΔEN: H-H = 0; H-C ≈ 0.4; H-N ≈ 0.9; H-F = 1.8. H-F most polar due to large electronegativity difference.
Q7 of 10
A molecule can have polar bonds but be non-polar overall if:
Symmetrical molecules: dipoles cancel. CO₂ (linear, 2 identical C=O dipoles cancel). CCl₄ (tetrahedral, 4 C-Cl dipoles cancel). Net dipole = 0 → non-polar molecule.
Q8 of 10
Down Group 17 (halogens), electronegativity:
Down group: atomic radius increases → bonding electrons further from nucleus → less pull → decreasing electronegativity. F > Cl > Br > I.
Q9 of 10
Bond polarity affects:
Polar bonds → polar molecules → stronger intermolecular forces → higher b.p., higher solubility in polar solvents. Also affects reactivity (nucleophilic/electrophilic sites).
Q10 of 10
The dipole moment of a molecule is:
Molecular dipole = vector sum of all bond dipoles. Accounts for both magnitude and direction of each bond's polarity. Symmetrical molecules: vector sum = 0 → μ = 0.
5.3
Location of Elements Based on Electronic Configuration
Period Number = Number of Electron Shells
The period an element is in equals the highest principal quantum number (n) in its electron configuration — i.e. the number of occupied electron shells.
Na: 1s²2s²2p⁶3s¹ → highest n = 3 → Period 3
Fe: [Ar]3d⁶4s² → highest n = 4 → Period 4
Br: [Ar]3d¹⁰4s²4p⁵ → highest n = 4 → Period 4
Group Number = Outer Electrons
For s- and p-block elements: Group number = number of outer shell (valence) electrons.
For d-block (transition metals): Group number = number of d electrons + s electrons (counted differently in different numbering systems; the modern IUPAC system numbers Groups 1–18).
C: [He]2s²2p² → 4 valence electrons → Group 14
Cl: [Ne]3s²3p⁵ → 7 valence electrons → Group 17
Fe: [Ar]3d⁶4s² → d-block, Group 8
Cu: [Ar]3d¹⁰4s¹ → d-block, Group 11
Finding period and group from configurationStep 1: Write the electron configuration. Step 2: Find the highest n value (= period). Step 3: Count valence electrons in the outer shell. Step 4: Use the block and valence count to assign the group. Remember: for p-block elements, group = 10 + (s + p valence electrons).
WORKED EXAMPLE
Locating Elements in the Periodic Table
1
Arsenic (As, Z=33): [Ar]3d¹⁰4s²4p³
Highest n = 4 → Period 4. Outer electrons: 4s²4p³ = 5 valence e− → Group 15 (p-block).
Highest n = 4 → Period 4. Outer electrons: 4s²4p³ = 5 valence e− → Group 15 (p-block).
2
Strontium (Sr, Z=38): [Kr]5s²
Highest n = 5 → Period 5. Outer electrons: 5s² = 2 → Group 2 (s-block).
Highest n = 5 → Period 5. Outer electrons: 5s² = 2 → Group 2 (s-block).
3
Tungsten (W, Z=74): [Xe]4f¹⁴5d⁴6s²
Highest n = 6 → Period 6. d-block (5d) → Group 6.
Highest n = 6 → Period 6. d-block (5d) → Group 6.
5.4
Classification into Blocks (s, p, d, f)
The Four Blocks
Elements are classified into blocks based on the subshell in which the last electron is placed (the differentiating electron).
| Block | Subshell filling | Groups | Periods | Key examples |
|---|---|---|---|---|
| s-block | 1s, 2s, 3s, 4s… filling | 1 & 2 (+ He) | All | H, He, Li, Na, K, Be, Mg, Ca |
| p-block | 2p, 3p, 4p… filling | 13–18 | 2–7 | C, N, O, F, Ne, Si, Cl, Ar, Br |
| d-block | 3d, 4d, 5d… filling | 3–12 | 4–7 | Sc, Ti, Cr, Mn, Fe, Co, Ni, Cu, Zn |
| f-block | 4f, 5f filling | Separate (below) | 6–7 | La–Lu (lanthanides), Ac–Lr (actinides) |
5.5
Characteristics of Different Blocks
s-Block — Groups 1 & 2
- Outer electrons in the s subshell: Group 1 has ns¹; Group 2 has ns².
- All are metals (except H — a non-metal).
- Low ionisation energies → form cations easily (Group 1: 1+; Group 2: 2+).
- Very reactive metals; reactivity increases down the group.
- Form basic oxides and hydroxides.
- Group 1: soft metals with low m.p.; Group 2: harder with higher m.p. than Group 1.
p-Block — Groups 13–18
- Outer electrons in the p subshell (after filling s subshell of same period).
- Contains metals (Al, Ga, Sn, Pb), metalloids/semimetals (B, Si, As, Ge, Te), and non-metals (C, N, O, F, Cl, Br) and noble gases (Group 18).
- Wide diversity of properties; includes most non-metals.
- Group 17 (halogens): highly reactive non-metals forming −1 ions.
- Group 18 (noble gases): full outer shell → chemically inert (with very few exceptions).
d-Block — Transition Metals (Groups 3–12)
- Inner d subshell fills. Outer electrons: (n−1)d and ns² (with exceptions: Cr, Cu).
- All are metals: high m.p., b.p., density, and tensile strength.
- Exhibit variable oxidation states (multiple stable ions: Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺).
- Form coloured compounds and ions (d–d electron transitions absorb visible light).
- Often good catalysts (Fe in Haber process; Ni in hydrogenation; Pt in catalytic converters).
- Form complex ions with ligands (e.g. [Fe(H₂O)⁶]³⁺, [Cu(NH₃)₄]²⁺).
f-Block — Lanthanides & Actinides
- Inner f subshell fills. Period 6 (4f): lanthanides (Ce–Lu); Period 7 (5f): actinides (Th–Lr).
- All are metals; many are radioactive (especially actinides).
- Lanthanides are used in powerful magnets, phosphors (screens, LEDs), and lasers.
- Actinides include uranium and plutonium (nuclear fuel/weapons).
- Similar chemical properties across the lanthanide series due to poor shielding by f electrons (“lanthanide contraction”).
5.6
Variation of Physical Properties Down Groups & Across Periods
Atomic Radius
Atomic RadiusHalf the distance between the nuclei of two adjacent atoms of the same element. (Van der Waals radius for non-bonded; covalent radius for bonded atoms.)
Across a period (left → right): DECREASES
Nuclear charge Z increases → same shell electrons are pulled in more strongly → electron cloud contracts → smaller radius. Shielding stays roughly constant (electrons added to same shell).
Down a group: INCREASES
Each successive element has an extra electron shell → outer electrons further from nucleus → larger radius. Despite increasing Z, the new shell adds more distance than the extra protons remove.
Ionisation Energy
Across a period: generally INCREASES (with dips at Group 13 and Group 16 — see Unit 2 for full explanation).
Down a group: DECREASES because outer electrons are in higher shells (further from nucleus, more shielding) → easier to remove.
Electronegativity
Across a period (left → right): INCREASES — nuclear charge increases, atomic radius decreases → nucleus attracts shared electrons more strongly.
Down a group: DECREASES — outer electrons further from nucleus, more shielding → nucleus attracts shared electrons less strongly.
Most electronegative: F (4.0) → O (3.5) → N,Cl (3.0) → Br (2.8). Least electronegative: Fr and Cs (down Group 1).
Melting Point
Across Period 3: Na → Cl
| Element | Na | Mg | Al | Si | P | S | Cl | Ar |
|---|---|---|---|---|---|---|---|---|
| m.p. (°C) | 98 | 650 | 660 | 1414 | 44 | 113 | −101 | −189 |
| Structure | Metal | Metal | Metal | Giant cov. | Simple mol. | Simple mol. | Simple mol. | Simple mol. |
- Na → Mg → Al: all metallic — m.p. increases because metallic bond strength increases (more delocalised e− per atom: 1, 2, 3).
- Si: giant covalent lattice → very high m.p. (1414°C): all bonds are strong covalent.
- P, S, Cl, Ar: simple molecular → low m.p.: only weak London dispersion forces between molecules. S (S₈ molecule, larger) > P (P₄) > Cl₂ > Ar because larger molecules have stronger LDF.
Down Group 1 — Physical Properties
| Property | Li | Na | K | Rb | Cs | Trend & Reason |
|---|---|---|---|---|---|---|
| Atomic radius (pm) | 152 | 186 | 231 | 244 | 262 | Increases: each element has one more shell. |
| m.p. (°C) | 181 | 98 | 64 | 39 | 29 | Decreases: larger ions + 1 delocalised e−/atom. Increased shielding → weaker metallic bond. |
| Density (g/cm³) | 0.53 | 0.97 | 0.86 | 1.53 | 1.87 | Generally increases (greater mass per atom), but irregularity at K (less dense than Na due to crystal structure). |
| IE₁ (kJ/mol) | 520 | 496 | 419 | 403 | 376 | Decreases: outer electron further and more shielded → easier to remove. |
| Reactivity | Least reactive | — | — | — | Most reactive | Increases: outer electron lost more easily (lower IE). |
Down Group 7 (Halogens) — Physical Properties
| Property | F | Cl | Br | I | Trend & Reason |
|---|---|---|---|---|---|
| Physical state (room temp) | Pale yellow gas | Green-yellow gas | Red-brown liquid | Grey-black solid | State increases: stronger LDF with increasing M. |
| Boiling point (°C) | −188 | −34 | 59 | 184 | Increases: more electrons → stronger LDF. |
| Atomic radius (pm) | 64 | 99 | 114 | 133 | Increases: each adds a shell. |
| Electronegativity | 4.0 | 3.0 | 2.8 | 2.5 | Decreases: outer electrons further away. |
| IE₁ (kJ/mol) | 1681 | 1251 | 1140 | 1008 | Decreases: outer electron in higher shell. |
| Colour | Pale yellow | Yellow-green | Red-brown | Dark grey/violet | Deepens: energy gap between HOMO and LUMO narrows → absorbs lower-energy (visible) light. |
Ionic Radius
Across a period: ionic radii are not simply compared since the ion type changes (from cations to anions). Within isoelectronic series: higher Z → smaller radius (more protons pulling same electrons).
Down a group: ionic radius increases (each successive ion has one more shell of electrons).
Cations are smaller than their parent atoms (electrons removed → less repulsion → remaining electrons drawn in).
Anions are larger than their parent atoms (electrons gained → more repulsion → electron cloud expands).
Example isoelectronic series (all 10 electrons):
N³⁻ (Z=7) > O²⁻ (Z=8) > F⁻ (Z=9) > Na⁺ (Z=11) > Mg²⁺ (Z=12) > Al³⁺ (Z=13)
(same electrons, increasing nuclear charge → decreasing radius)
Summary of Period and Group Trends
| Property | Across a period (left → right) | Down a group |
|---|---|---|
| Atomic radius | Decreases (↓) — ↑Z, same shell | Increases (↑) — extra shell each period |
| Ionic radius | Decreases within isoelectronic series | Increases — extra shell each period |
| Ionisation energy | Generally increases (with dips) | Decreases — outer e− further, more shielded |
| Electronegativity | Increases — smaller atom, higher Z | Decreases — outer e− further away |
| Metallic character | Decreases — non-metals at right | Increases — outer e− easier to lose |
| Non-metallic character | Increases — towards right | Decreases — less electronegative down |
| Reactivity (metals) | Decreases (left has most reactive metals) | Increases — lose outer e− more easily |
| Reactivity (non-metals) | Increases towards Group 17 | Decreases — gain e− less readily |
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Section 5.3 — Ionisation Energy Trends
10 QuestionsQ1 of 10
First ionisation energy increases across a period because:
Across period: Z↑, same shell, same shielding → Zeff↑ → outer electrons held more tightly → more energy to remove.
Q2 of 10
Why is IE₁(O) < IE₁(N) despite O having more protons?
N: [He]2s²2p³ (half-filled 2p, extra stable). O: [He]2s²2p⁴ (one paired 2p electron — repulsion makes it easier to remove). O's IE₁ < N's despite higher Z.
Q3 of 10
Down Group 1, IE₁ decreases because:
Each period adds a new shell → outer electron further from nucleus → weaker attraction → less energy to remove → decreasing IE₁.
Q4 of 10
The large jump in successive ionisation energies identifies:
Large jump between IEₙ and IEₙ₊₁ means the n+1 th electron is in an inner shell. The number of electrons removed before the jump = number of outer electrons = group number.
Q5 of 10
Which element has IE₁ values (kJ/mol): 496, 4563? Jump between 1st and 2nd.
Jump after IE₁ → 1 outer electron → Group 1. IE₁ = 496 kJ/mol matches Na (496 kJ/mol). (K has IE₁ = 419, Li = 520)
Q6 of 10
IE₁ of noble gases is the highest in their period because:
Noble gases: full outer shell → maximum Zeff, maximum stability → highest IE₁ in period. He: 2372 kJ/mol; Ne: 2081 kJ/mol; Ar: 1521 kJ/mol.
Q7 of 10
Shielding by inner electrons reduces the experienced nuclear charge. Which provides the MOST shielding?
Inner shell electrons are between nucleus and outer electrons → most effective shielding. Same shell electrons: partial shielding (similar distance). Outer electrons: no shielding of inner.
Q8 of 10
Why is IE₁(Al) < IE₁(Mg)?
Mg: [Ne]3s². Al: [Ne]3s²3p¹. 3p is higher energy and more shielded than 3s → easier to remove → Al has lower IE₁ than Mg (despite higher Z).
Q9 of 10
The electron affinity of Cl is more negative than F. Why?
F: very small → adding e⁻ to 2p → strong repulsion with existing electrons → less energy released. Cl: larger → less repulsion when adding to 3p → more negative electron affinity despite lower EN.
Q10 of 10
Which statement about ionisation energy is incorrect?
All four statements are actually correct! Trick question — all are correct. Actually wait — statement A says IE always increases with each successive electron removed — this IS correct (each is harder). So actually all are correct. Let's reconsider... Actually the question should read: 'IE₁ always increases down a group' which would be INCORRECT. IE₁ DECREASES down a group.
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Exercises
- State the main contributions of Döbereiner, Newlands, and Mendeleev to the development of the periodic table. What was the fundamental difference between Mendeleev's arrangement and the modern periodic table?
Döbereiner (1829): Law of Triads — groups of 3 elements where middle element's mass was average of the other two. Newlands (1864): Law of Octaves — every 8th element repeated properties. Mendeleev (1869): arranged all known elements by atomic mass, grouped by similar chemical properties, and left gaps for undiscovered elements (predicted their properties). Fundamental difference: Mendeleev's table was ordered by atomic mass; the modern periodic table is ordered by atomic number (Moseley, 1913), which resolved anomalies like the Te/I inversion and gave a physically meaningful basis to periodicity.
- An element X has electron configuration [Kr]4d⁵5s¹. Identify: (a) the period, (b) the block, (c) the group of X. Identify the element.
(a) Highest n = 5 → Period 5. (b) Last electrons enter 4d subshell → d-block. (c) 4d⁵5s¹ = 6 outer electrons → Group 6. (Note the 5s¹ anomaly: half-filled d has stability.) Element: Molybdenum (Mo, Z=42).
- Explain why the atomic radius decreases across Period 3 from Na to Cl, but increases from Na to Cs down Group 1.
Across Period 3: nuclear charge Z increases from 11 (Na) to 17 (Cl), but all outer electrons are in the n=3 shell (shielding roughly constant). Greater nuclear charge pulls electrons closer to nucleus → electron cloud contracts → atomic radius decreases. Down Group 1: each element adds a new electron shell (n=2,3,4,5,6 for Li→Cs). The new shell places outer electrons much further from the nucleus; despite increasing Z, the extra shielding from inner shells means the net effect is a much larger atomic radius.
- The melting points of Period 3 elements vary widely: Na 98°C, Si 1414°C, Cl₂ −101°C. Explain this variation in terms of structure and bonding.
Na: metallic solid. Metallic bonds (Na⁺ cations + 1 delocalised e−/atom). Moderate m.p. Si: giant covalent lattice. Each Si sp³-bonded to 4 others → all bonds must be broken → very high m.p. requires vast energy. Cl₂: simple molecular. Consists of Cl₂ molecules held together only by weak London dispersion forces (LDF). Small molecules, few electrons → very weak LDF → very low energy to separate molecules → very low m.p.
- Arrange the following in order of increasing ionic radius: Na⁺, Mg²⁺, Al³⁺, F⁻, O²⁻. Explain your reasoning.
All are isoelectronic: each has 10 electrons. Increasing nuclear charge pulls the same number of electrons in more strongly → smaller radius. Z values: O(8) < F(9) < Na(11) < Mg(12) < Al(13). Order of increasing radius (smallest Z → largest radius): Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻. O²⁻ has the lowest nuclear charge (Z=8) holding 10 electrons → largest radius. Al³⁺ has the highest Z (13) → smallest radius.
- Compare the physical properties of the halogens F₂, Cl₂, Br₂, and I₂ at room temperature (state, colour, boiling point). Explain the trend in boiling point.
F₂: pale yellow gas, b.p. −188°C. Cl₂: yellow-green gas, b.p. −34°C. Br₂: red-brown liquid, b.p. +59°C. I₂: grey-black solid, b.p. +184°C. Trend: b.p. increases down the group. All halogens are non-polar diatomic molecules → only London dispersion forces between molecules. Going down Group 17: molar mass increases (F₂=38, Cl₂=71, Br₂=160, I₂=254 g/mol) → more electrons per molecule → more polarisable → stronger instantaneous dipoles → stronger LDF → more energy needed to separate molecules → higher b.p.
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Multiple Choice Quiz — 25 Questions
Unit 5 Quiz
Select one answer per questionQ1
Mendeleev's key innovation over Newlands was:
Mendeleev's insight was leaving deliberate gaps where elements were missing and predicting their properties — validated when Ga, Sc, and Ge were discovered matching his predictions.
Q2
The modern periodic table is arranged by:
Moseley (1913): atomic number (proton number) is the correct basis for ordering, not atomic mass. This resolved anomalies like Te/I.
Q3
An element has configuration [Xe]4f¹⁴5d⁶6s². Which block is it in?
The differentiating electron enters 5d → d-block. (The 4f is completely filled; the 5d is partially filled.) Period 6, d-block, Group 8 — this is Osmium (Os).
Q4
The period number of an element equals:
Period = highest n in electron config. E.g., Na: [Ne]3s¹ → highest n=3 → Period 3.
Q5
Atomic radius across Period 3 decreases because:
Across a period: Z increases by 1 each step; all electrons added to same shell (same shielding). Net effect: greater nuclear attraction → electron cloud pulled in → smaller radius.
Q6
Which element has the highest electronegativity?
Fluorine: most electronegative element (EN=4.0, Pauling scale). Top-right of periodic table: highest Z, smallest atomic radius → greatest attraction for shared electrons.
Q7
Why does the melting point of metals increase from Na to Al across Period 3?
Na: 1, Mg: 2, Al: 3 delocalised e− per atom. More electrons in sea + higher ionic charge (Na⁺, Mg²⁺, Al³⁺) + smaller ionic radii → stronger metallic bond → higher m.p.
Q8
Silicon has the highest melting point in Period 3 because:
Si: giant covalent 3D lattice (like diamond, sp³). Every bond is a strong covalent bond. Melting requires breaking vast numbers of covalent bonds → very high m.p. (1414°C).
Q9
In which block is phosphorus (P, Z=15)?
P: [Ne]3s²3p³. The differentiating electron enters 3p → p-block. Period 3, Group 15.
Q10
Which of the following has the largest atomic radius?
Cs is at the bottom of Group 1 → 6 electron shells → largest atomic radius (262 pm). Radius increases down a group as each element adds a new shell.
Q11
The d-block elements are characterised by:
Transition (d-block) metals: variable oxidation states (due to similar energies of d and s orbitals), coloured compounds (d-d transitions), catalytic properties, complex ion formation.
Q12
The ionic radius order for the isoelectronic series N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ is:
All have 10 electrons. Higher Z → greater pull on same electrons → smaller radius. Z: N(7) < O(8) < F(9) < Na(11) < Mg(12). Largest radius: N³⁻; smallest: Mg²⁺. Order: Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻.
Q13
Which scientist determined that atomic number, not mass, is the basis for the periodic table?
Henry Moseley (1913) used X-ray diffraction to measure atomic numbers and showed the correct ordering principle. Killed in WWI at age 27.
Q14
First ionisation energy generally increases across Period 2 because:
Across a period: Z increases, atomic radius decreases, shielding stays roughly constant. Greater nuclear attraction on outer electrons → more energy needed to remove them → higher IE.
Q15
Why do noble gases (Group 18) have very low boiling points and are chemically inert?
Noble gases: complete outer shell (8 electrons, or 2 for He) → no tendency to form bonds → inert. Monoatomic, non-polar → only very weak LDF → low b.p. b.p. increases down the group (more electrons → stronger LDF).
Q16
Which element is in Period 4, Group 2?
Group 2, Period 4: [Ar]4s² → Calcium (Ca). Mg is Period 3, Sr is Period 5, Be is Period 2.
Q17
The law of triads was proposed by:
Döbereiner (1829): Law of Triads — middle element's mass = average of outer two, and properties are intermediate. E.g. Li(6.9), Na(23), K(39.1): (6.9+39.1)/2 = 23 ✓.
Q18
Which property INCREASES down Group 17 (halogens)?
Boiling point increases down Group 17: F₂(−188) < Cl₂(−34) < Br₂(59) < I₂(184)°C. More electrons → stronger LDF. EN, IE, and oxidising power all decrease down the group.
Q19
The f-block elements (lanthanides and actinides) are separated below the main table because:
14 f-block elements per period would make the table impractically wide (32 columns). They involve filling inner 4f/5f subshells, so they are placed separately for clarity.
Q20
A cation is smaller than its parent atom because:
Cation: fewer electrons but same Z. Less electron-electron repulsion → remaining electrons drawn closer to nucleus → smaller radius. Also may lose an entire outer shell (e.g. Na → Na⁺: loses n=3 shell entirely → much smaller).
Q21
Which correctly describes p-block elements?
p-block: Groups 13–18, Periods 2–7. Contains great diversity: metals (Al, Pb), metalloids (Si, Ge, As), non-metals (C, N, O, Cl), noble gases (He*, Ne, Ar).
Q22
Why does metallic character increase down Group 1?
Metallic character = tendency to lose outer electrons. Down Group 1: outer electrons in higher shells, more shielded, further away → lower IE → lost more easily → more metallic character.
Q23
The anomaly in Mendeleev's table where Te appeared before I was resolved by:
Moseley showed atomic number is the correct ordering principle. Te (Z=52) and I (Z=53) are correctly placed by Z, even though Te has higher average atomic mass due to its isotope distribution.
Q24
Which element is in Period 5, Group 17?
Group 17, Period 5: [Kr]4d¹⁰5s²5p⁵ → Iodine (I, Z=53). Cl is Period 3, Br is Period 4, At is Period 6.
Q25
Which trend is correct for electronegativity?
EN increases across a period (higher Z, smaller atom, stronger attraction for shared electrons) and decreases down a group (outer electrons further away, more shielded). Maximum: F (top right); minimum: Cs/Fr (bottom left).
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Unit Test — 50 Marks
Section A — Short Answer
20 marksQ1 [4 marks]
Describe Mendeleev's contribution to the periodic table. State two ways in which the modern periodic table differs from Mendeleev's original table. [4]
Mendeleev (1869): arranged 63 known elements in order of increasing atomic mass, grouping elements with similar chemical properties in vertical columns. Crucially, he left gaps for undiscovered elements and predicted their atomic masses and properties (e.g. eka-silicon = Ge, eka-aluminium = Ga). [2] Two differences from modern table: (1) Mendeleev ordered by atomic mass; modern table ordered by atomic number (Moseley, 1913), resolving anomalies. [1] (2) Modern table includes noble gases (Group 18), discovered after Mendeleev's table and not originally present. Also valid: f-block separated, 118 elements vs 63, blocks recognised. [1]
Q2 [4 marks]
An element X has the electron configuration [Kr]4d¹⁰5s²5p⁴. (a) State the period and group of X. [2] (b) Identify the block. [1] (c) Name element X. [1]
(a) Highest n = 5 → Period 5. Valence electrons in 5p subshell: 5s²5p⁴ = 6 outer electrons → Group 16. [2] (b) Differentiating electron enters 5p → p-block. [1] (c) Period 5, Group 16 = Tellurium (Te, Z=52). [1]
Q3 [4 marks]
Explain the following trend in first ionisation energy down Group 1: Li (520) > Na (496) > K (419) > Rb (403) > Cs (376 kJ mol−¹). [4]
Going down Group 1, each element has one more electron shell [1]: Li: 2s¹ outer; Na: 3s¹; K: 4s¹; Rb: 5s¹; Cs: 6s¹. This means: (1) The outer electron is in a progressively higher shell → greater distance from the nucleus [1]. (2) More inner electron shells provide greater shielding of the outer electron from the nuclear charge [1]. Both factors reduce the effective nuclear charge felt by the outer electron → it is held less tightly → less energy required to remove it → IE₁ decreases down the group. Despite increasing Z (3 → 55), the shielding and distance effects dominate [1].
Q4 [4 marks]
Describe and explain the trend in melting points across Period 3 elements (Na → Ar). [4]
Na, Mg, Al: metallic structures. Melting point increases (98 → 650 → 660°C) because number of delocalised e− per atom increases (1, 2, 3) → stronger metallic bond → more energy to disrupt. [1] Si: giant covalent lattice (sp³, 3D); m.p. = 1414°C — all bonds are strong covalent → very high m.p. [1] P, S, Cl, Ar: simple molecular. Low m.p. because only weak London dispersion forces between molecules, not covalent bonds. S (S₈, larger molecule) > P (P₄) > Cl₂ > Ar because larger molecules/more electrons → stronger LDF. [2]
Q5 [4 marks]
Arrange these in order of increasing atomic radius: F, Na, Cl, K, Cs. Explain the two factors that determine the trend. [4]
Increasing atomic radius: F < Cl < Na < K < Cs. [1] F (64pm) and Cl (99pm) are both p-block; Cl has an extra shell → larger. Na (186pm) is in Period 3 Group 1 → larger than Cl (Group 17 Period 3) because less nuclear charge pulling electrons inward. K (231pm) and Cs (262pm) are Group 1 Periods 4 and 6 → increase down the group. [1] Factor 1: Shell number — more shells → outer electrons further from nucleus → larger radius. [1] Factor 2: Nuclear charge (Z) across a period — more protons attract electrons inward → smaller radius. Down a group, increased shielding dominates over increased Z → radius increases. [1]
Section B — Extended Answer
30 marksQ6 [8 marks]
Describe the historical development of the periodic table from Döbereiner to the modern table. Include: Law of Triads, Law of Octaves, Mendeleev's table (contributions and limitations), and Moseley's contribution. Explain why Mendeleev's arrangement is considered a greater achievement than Newlands'. [8]
Döbereiner (1829) [1]: Law of Triads: grouped 3 elements (e.g. Li/Na/K; Cl/Br/I) where middle element's mass ≈ average of other two, and properties intermediate. Limited: only worked for a few groups.
Newlands (1864) [1]: Law of Octaves: elements arranged by atomic mass; every 8th element had similar properties. Worked well for first 20 elements; broke down for heavier elements. Ridiculed for forcing dissimilar elements together.
Mendeleev (1869) [3]: Arranged all 63 known elements by atomic mass in groups of similar chemical properties. Key innovations: (1) Left deliberate gaps for undiscovered elements; (2) Predicted properties of eka-aluminium (Ga, discovered 1875), eka-boron (Sc, 1879), eka-silicon (Ge, 1886) — all confirmed. Limitation: had to reverse some pairs (e.g. Te before I by chemical properties, even though I has lower mass) and could not explain why periodicity occurs.
Moseley (1913) [2]: Used X-ray spectra to measure atomic numbers accurately. Showed that the correct ordering basis is atomic number (proton number), not atomic mass. This resolved all anomalies (Te Z=52 before I Z=53 is correct by atomic number) and gave a physical explanation for periodicity (period = electron shell; group = outer electron configuration).
Mendeleev > Newlands [1]: Mendeleev included all elements, left gaps, and made quantitative predictions that were later verified. Newlands forced elements into octaves without gaps, failed for heavier elements, and made no predictions. The predictive power and systematic approach of Mendeleev were far more scientifically valuable.
Newlands (1864) [1]: Law of Octaves: elements arranged by atomic mass; every 8th element had similar properties. Worked well for first 20 elements; broke down for heavier elements. Ridiculed for forcing dissimilar elements together.
Mendeleev (1869) [3]: Arranged all 63 known elements by atomic mass in groups of similar chemical properties. Key innovations: (1) Left deliberate gaps for undiscovered elements; (2) Predicted properties of eka-aluminium (Ga, discovered 1875), eka-boron (Sc, 1879), eka-silicon (Ge, 1886) — all confirmed. Limitation: had to reverse some pairs (e.g. Te before I by chemical properties, even though I has lower mass) and could not explain why periodicity occurs.
Moseley (1913) [2]: Used X-ray spectra to measure atomic numbers accurately. Showed that the correct ordering basis is atomic number (proton number), not atomic mass. This resolved all anomalies (Te Z=52 before I Z=53 is correct by atomic number) and gave a physical explanation for periodicity (period = electron shell; group = outer electron configuration).
Mendeleev > Newlands [1]: Mendeleev included all elements, left gaps, and made quantitative predictions that were later verified. Newlands forced elements into octaves without gaps, failed for heavier elements, and made no predictions. The predictive power and systematic approach of Mendeleev were far more scientifically valuable.
Q7 [8 marks]
Describe and explain the variation of the following properties across Period 3 and down Group 1: (a) atomic radius [2]; (b) ionisation energy [2]; (c) electronegativity [2]; (d) reactivity of the metal with water [2]. [8]
(a) Atomic radius [2]: Across Period 3: decreases Na(186)→Cl(99pm). Z increases 11→17; all outer electrons in n=3; shielding roughly constant. Higher nuclear charge attracts electrons inward → radius decreases. Down Group 1: increases Li(152)→Cs(262pm). Each element adds a new electron shell; despite increasing Z, the greater distance and shielding of the new shell dominate → radius increases.
(b) Ionisation energy [2]: Across Period 3: generally increases (with dips at Al and S). Higher Z, smaller atomic radius, constant shielding → outer electrons held more tightly. Down Group 1: decreases Li(520)→Cs(376 kJ/mol). Outer electron in higher shell (greater distance + more shielding) → weaker nuclear hold → lower IE.
(c) Electronegativity [2]: Across Period 3: increases Na(0.9)→Cl(3.0, Pauling). Increasing Z + decreasing radius → greater attraction for shared electrons. Down Group 1: decreases Li(1.0)→Cs(0.7). Outer electrons further and more shielded → less ability to attract shared electrons.
(d) Reactivity with water [2]: Across Period 3: Na reacts vigorously; Mg reacts slowly (with cold water/steam); Al reacts only with steam; Si, P, S, Cl not reactive as metals (already non-metals or metalloids). Reactivity as metals decreases across the period. Down Group 1: reactivity increases Li(slow, fizzes) → Na(vigorous) → K(very vigorous, ignites) → Rb/Cs(explosive). Lower IE down the group → outer electron lost more easily → more vigorous reaction with water.
(b) Ionisation energy [2]: Across Period 3: generally increases (with dips at Al and S). Higher Z, smaller atomic radius, constant shielding → outer electrons held more tightly. Down Group 1: decreases Li(520)→Cs(376 kJ/mol). Outer electron in higher shell (greater distance + more shielding) → weaker nuclear hold → lower IE.
(c) Electronegativity [2]: Across Period 3: increases Na(0.9)→Cl(3.0, Pauling). Increasing Z + decreasing radius → greater attraction for shared electrons. Down Group 1: decreases Li(1.0)→Cs(0.7). Outer electrons further and more shielded → less ability to attract shared electrons.
(d) Reactivity with water [2]: Across Period 3: Na reacts vigorously; Mg reacts slowly (with cold water/steam); Al reacts only with steam; Si, P, S, Cl not reactive as metals (already non-metals or metalloids). Reactivity as metals decreases across the period. Down Group 1: reactivity increases Li(slow, fizzes) → Na(vigorous) → K(very vigorous, ignites) → Rb/Cs(explosive). Lower IE down the group → outer electron lost more easily → more vigorous reaction with water.
Q8 [6 marks]
Classify each element below into the correct block, period, and group, and give a brief reason. (a) Selenium (Se, Z=34); (b) Barium (Ba, Z=56); (c) Nickel (Ni, Z=28). [6]
(a) Se (Z=34) [2]: Configuration: [Ar]3d¹⁰4s²4p⁴. Highest n=4 → Period 4. Last electron enters 4p → p-block. Valence electrons: 4s²4p⁴ = 6 → Group 16. Se is a non-metal/metalloid (chalcogen), below sulfur in Group 16.
(b) Ba (Z=56) [2]: Configuration: [Xe]6s². Highest n=6 → Period 6. Last electron enters 6s → s-block. Valence electrons: 6s² = 2 → Group 2. Ba is an alkaline earth metal; reactive, forms Ba²⁺ ions.
(c) Ni (Z=28) [2]: Configuration: [Ar]3d⁸4s². Highest n=4 → Period 4. Last electron enters 3d → d-block. Outer electrons: 3d⁸4s² → Group 10 (transition metal). Ni is a silvery-white transition metal; forms Ni²⁺ ions; used in alloys and catalysts.
(b) Ba (Z=56) [2]: Configuration: [Xe]6s². Highest n=6 → Period 6. Last electron enters 6s → s-block. Valence electrons: 6s² = 2 → Group 2. Ba is an alkaline earth metal; reactive, forms Ba²⁺ ions.
(c) Ni (Z=28) [2]: Configuration: [Ar]3d⁸4s². Highest n=4 → Period 4. Last electron enters 3d → d-block. Outer electrons: 3d⁸4s² → Group 10 (transition metal). Ni is a silvery-white transition metal; forms Ni²⁺ ions; used in alloys and catalysts.
Q9 [8 marks]
Compare the characteristics of s-, p-, d-, and f-block elements in terms of: (a) which subshell is filling; (b) typical physical properties (state, conductivity, melting point); (c) typical chemical properties (ion formation, reactivity); (d) one important use or application of each block. Give at least one example element per block. [8]
s-block [2]: (a) s subshell filling (Groups 1–2). (b) Metals (except H); low-moderate m.p. (Na: 98°C); good conductors (metallic bonding). (c) Lose outer electrons easily to form cations (Na⁺, Ca²⁺); low IE → very reactive metals; react with water/air. React more vigorously down the group. (d) Na used in sodium-vapour lamps; Ca in cement/construction; Mg in alloys (aircraft).
p-block [2]: (a) p subshell filling (Groups 13–18). (b) Diverse — metals (Al, Pb), metalloids (Si, Ge), non-metals (C, N, Cl), noble gases (Ar). Wide range of m.p. and states. Non-metals are non-conductors; Al is a good conductor. (c) Non-metals tend to gain electrons or share; halogens form −1 ions; Group 14–16 form various ions and covalent bonds. Noble gases are inert (full outer shell). (d) Si in semiconductors/electronics; Cl in water treatment (Cl₂); N₂ as inert atmosphere in food packaging.
d-block [2]: (a) (n−1)d subshell filling (Groups 3–12). (b) All metals; high m.p. (Fe: 1538°C), high density, high tensile strength; excellent conductors. (c) Variable oxidation states (Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺); form coloured compounds (d–d transitions); form complex ions. Good catalysts. (d) Fe in steelmaking; Cu in electrical wiring; Pt as catalytic converter; Ni in hydrogenation of fats.
f-block [2]: (a) 4f (lanthanides) or 5f (actinides) filling. (b) All metals; often radioactive (actinides); dense, high m.p. Similar properties across each series (poor f-shielding → similar Zeff). (c) Lanthanides usually form Ln³⁺ ions; actinides can form multiple oxidation states. (d) Nd in powerful neodymium magnets (NdFeB); U/Pu in nuclear fuel and weapons; La in catalytic converters and lighting.
p-block [2]: (a) p subshell filling (Groups 13–18). (b) Diverse — metals (Al, Pb), metalloids (Si, Ge), non-metals (C, N, Cl), noble gases (Ar). Wide range of m.p. and states. Non-metals are non-conductors; Al is a good conductor. (c) Non-metals tend to gain electrons or share; halogens form −1 ions; Group 14–16 form various ions and covalent bonds. Noble gases are inert (full outer shell). (d) Si in semiconductors/electronics; Cl in water treatment (Cl₂); N₂ as inert atmosphere in food packaging.
d-block [2]: (a) (n−1)d subshell filling (Groups 3–12). (b) All metals; high m.p. (Fe: 1538°C), high density, high tensile strength; excellent conductors. (c) Variable oxidation states (Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺); form coloured compounds (d–d transitions); form complex ions. Good catalysts. (d) Fe in steelmaking; Cu in electrical wiring; Pt as catalytic converter; Ni in hydrogenation of fats.
f-block [2]: (a) 4f (lanthanides) or 5f (actinides) filling. (b) All metals; often radioactive (actinides); dense, high m.p. Similar properties across each series (poor f-shielding → similar Zeff). (c) Lanthanides usually form Ln³⁺ ions; actinides can form multiple oxidation states. (d) Nd in powerful neodymium magnets (NdFeB); U/Pu in nuclear fuel and weapons; La in catalytic converters and lighting.