This MCQ module is based on: Periodic Trends Chemical
TOPIC 12 OF 28
Periodic Trends Chemical
🎓 Class 11
Chemistry
CBSE
Theory
Ch 3 – Classification of Elements and Periodicity in Properties
⏱ ~14 min
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Periodic Trends in Chemical Properties
3.8 Periodic Trends in Chemical Properties
Most of the trends in chemical properties of elements, such as diagonal relationships, inert pair effect, effects of lanthanoid contraction etc., will be dealt with along the discussion of each group in later units. In this section we shall study the periodicity of the valence state shown by elements and the anomalous properties of the second period elements (from lithium to fluorine).
3.8.1 Periodicity of Valence or Oxidation States
The valence is the most characteristic property of the elements and can be understood in terms of their electronic configurations. The valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals and/or equal to eight minus the number of outermost electrons as shown below:
| Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|---|
| Number of valence electrons | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 |
| Valence (max) | 1 | 2 | 3 | 4 | 3, 5 | 2, 6 | 1, 7 | 0, 2, 4, 6, 8 |
Variable valence: Many elements (especially groups 14–17) show multiple valences (e.g., S: 2, 4, 6 in H₂S, SO₂, SO₃ respectively). Transition metals show even more variable valence (Fe²⁺/Fe³⁺, Mn²⁺ to Mn⁷⁺). This is because their d-electrons can also participate in bonding.
Periodicity in Hydride and Oxide Formulae
The valence concept becomes clear when we look at the formulae of hydrides and oxides of period 2 and period 3 elements:
| Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 |
|---|---|---|---|---|---|---|---|
| Hydride (Period 2) | LiH | BeH₂ | BH₃ | CH₄ | NH₃ | H₂O | HF |
| Hydride (Period 3) | NaH | MgH₂ | AlH₃ | SiH₄ | PH₃ | H₂S | HCl |
| Highest Oxide | M₂O | MO | M₂O₃ | MO₂ | M₂O₅ | MO₃ | M₂O₇ |
3.8.2 Anomalous Properties of Second Period Elements
The first element of each of the groups 1 (lithium), 2 (beryllium) and groups 13–17 (boron to fluorine) differs in many respects from the other members of their respective group. For example, lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals, form compounds with pronounced covalent character; the other members of these groups predominantly form ionic compounds.
The anomalous behaviour is attributed to their:
- Small size
- Large charge/radius ratio (high charge density)
- High electronegativity
- Non-availability of d-orbitals in their valence shell
Diagonal Relationship
The first member of a group of elements in s- and p-blocks shows similarities with the second member of the next group. This similarity is due to the comparable charge/size ratios. This is termed as the diagonal relationship in the periodic table.
3.8.3 Periodic Trends and Chemical Reactivity
The atomic and ionic radii, generally, decrease in a period from left to right. As a consequence, the ionisation enthalpies generally increase (with some exceptions as outlined in section 3.7.1) and electron gain enthalpies become more negative across a period. In other words, the ionisation enthalpy of the extreme left element in a period is the least and the electron gain enthalpy of the element on the extreme right in the period is the highest.
This results in high chemical reactivity at the two extremes and the lowest in the centre. Thus, the maximum chemical reactivity at the extreme left (among alkali metals) is exhibited by the loss of an electron leading to the formation of a cation; at the extreme right (among halogens) it is shown by the gain of an electron forming an anion. This property can be related with the reducing and oxidising behaviour of the elements which you will learn later. However, here it can be directly related to the metallic and non-metallic character of the elements.
Reactivity Trends in Two Halves
| Property | Group 1 (alkali metals) | Group 17 (halogens) |
|---|---|---|
| Reactivity down group | INCREASES (Li → Cs) | DECREASES (F → I) |
| Reason | Lower IE makes electron loss easier | Less negative ΔₑgH; less attraction to e⁻ |
| Most reactive | Cs (or Fr, but radioactive) | F (gas, etches glass) |
| Least reactive | Li (within group) | I, At (within group) |
Periodic Trends in Oxide Character
The oxides of metals are generally basic, while oxides of non-metals are acidic. Some elements that lie at the metalloid border (Al, Sb, Bi, etc.) form amphoteric oxides that react with both acids and bases. As we move across a period, oxides become progressively less basic, then amphoteric, then acidic.
| Element (Period 3) | Oxide | Character | Example reaction |
|---|---|---|---|
| Na | Na₂O | Strongly basic | Na₂O + H₂O → 2 NaOH |
| Mg | MgO | Basic | MgO + 2 HCl → MgCl₂ + H₂O |
| Al | Al₂O₃ | Amphoteric | Reacts with HCl AND NaOH |
| Si | SiO₂ | Acidic (weakly) | SiO₂ + 2 NaOH → Na₂SiO₃ + H₂O |
| P | P₄O₁₀ | Acidic | P₄O₁₀ + 6 H₂O → 4 H₃PO₄ |
| S | SO₃ | Strongly acidic | SO₃ + H₂O → H₂SO₄ |
| Cl | Cl₂O₇ | Strongly acidic | Cl₂O₇ + H₂O → 2 HClO₄ |
🎯 Interactive: Element Property Comparator
Choose two elements to compare their key periodic properties side by side.
| Property | O | F |
|---|---|---|
| Atomic radius (pm) | 66 | 64 |
| Ionisation enthalpy (kJ/mol) | 1314 | 1681 |
| Electron gain enthalpy (kJ/mol) | -141 | -328 |
| Electronegativity | 3.5 | 4.0 |
| Common oxidation states | −2 | −1 |
| Highest oxide | — | OF₂ |
| Hydride | H₂O | HF |
Tip: Compare same-group elements (down trend) or same-period elements (across trend).
Worked Examples
Worked Example 1 (NCERT Problem 3.9): Oxide Acid-Base Character
Are the oxidation state and covalency of Al in [AlCl(H₂O)₅]²⁺ same?
Oxidation state of Al: Cl is −1, H₂O is neutral. The complex has +2 net charge. So:
+2 = (Al ox state) + (−1 from Cl) + (0 from H₂O) → Al is +3.
Covalency of Al: Number of bonds Al forms with surrounding atoms. Al is bonded to:
- 1 Cl (1 bond)
- 5 H₂O (each is a coordinate bond from O lone pair) = 5 bonds
Total bonds = 6.
So oxidation state (+3) ≠ covalency (6). They are different concepts: oxidation state is a charge bookkeeping; covalency is the count of bonds (regardless of bond type).
+2 = (Al ox state) + (−1 from Cl) + (0 from H₂O) → Al is +3.
Covalency of Al: Number of bonds Al forms with surrounding atoms. Al is bonded to:
- 1 Cl (1 bond)
- 5 H₂O (each is a coordinate bond from O lone pair) = 5 bonds
Total bonds = 6.
So oxidation state (+3) ≠ covalency (6). They are different concepts: oxidation state is a charge bookkeeping; covalency is the count of bonds (regardless of bond type).
Worked Example 2 (NCERT Problem 3.10): Predicting Oxide Behaviour
Show by chemical reaction with water that Na₂O is a basic oxide and Cl₂O₇ is an acidic oxide.
Na₂O is BASIC: reacts with water to form a base.
\[\text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH (a base)}\]
NaOH dissociates in water: NaOH → Na⁺ + OH⁻ (raises pH above 7).
Cl₂O₇ is ACIDIC: reacts with water to form an acid. \[\text{Cl}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{HClO}_4 \text{ (perchloric acid)}\] HClO₄ is one of the strongest known acids (pH well below 7).
Trend confirmed: Group 1 metals → basic oxides; Group 17 non-metals → acidic oxides.
Cl₂O₇ is ACIDIC: reacts with water to form an acid. \[\text{Cl}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{HClO}_4 \text{ (perchloric acid)}\] HClO₄ is one of the strongest known acids (pH well below 7).
Trend confirmed: Group 1 metals → basic oxides; Group 17 non-metals → acidic oxides.
Worked Example 3: Anomalous Behaviour of Period 2
List two reasons why Li differs from Na, K, and other alkali metals.
(1) Forms covalent compounds: LiCl, LiH have significant covalent character (unlike highly ionic NaCl, NaH). This is because Li⁺ is small with high charge density, polarising anions strongly (Fajans' rules).
(2) Reacts with N₂ to form nitride: 6 Li + N₂ → 2 Li₃N. None of the other alkali metals react with N₂ at room temperature. This is the diagonal similarity with Mg (3 Mg + N₂ → Mg₃N₂).
Other Li peculiarities: hydroxide LiOH is less basic than NaOH; LiNO₃ decomposes to Li₂O on heating (others give NO₂ + nitrite); Li₂CO₃ decomposes more easily than other alkali carbonates.
(2) Reacts with N₂ to form nitride: 6 Li + N₂ → 2 Li₃N. None of the other alkali metals react with N₂ at room temperature. This is the diagonal similarity with Mg (3 Mg + N₂ → Mg₃N₂).
Other Li peculiarities: hydroxide LiOH is less basic than NaOH; LiNO₃ decomposes to Li₂O on heating (others give NO₂ + nitrite); Li₂CO₃ decomposes more easily than other alkali carbonates.
📐 Activity 3.4 — Predict Oxide Reactions
Setup: Without consulting reference, predict the products and balanced reactions when:
- K₂O is dissolved in water.
- SO₃ is mixed with water.
- Al₂O₃ is mixed with HCl.
- Al₂O₃ is mixed with NaOH (showing amphoteric nature).
1. K₂O is basic (Group 1): K₂O + H₂O → 2 KOH.
2. SO₃ is acidic (non-metal oxide): SO₃ + H₂O → H₂SO₄ (sulphuric acid).
3. Al₂O₃ + 6 HCl → 2 AlCl₃ + 3 H₂O (acts as base, neutralizes acid).
4. Al₂O₃ + 2 NaOH + 3 H₂O → 2 Na[Al(OH)₄] (sodium aluminate; acts as acid, neutralizes base).
Insight: Al₂O₃ shows AMPHOTERIC behaviour — reacting with both acids and bases. This is characteristic of the metalloid/metal-non-metal border region.
🎯 Competency-Based Questions
Q1. Which is the most basic oxide among Na₂O, Al₂O₃, P₄O₁₀? L3 Apply
Answer: Na₂O. Group 1 oxide → strongly basic. Al₂O₃ is amphoteric. P₄O₁₀ is acidic. Trend: across period 3, oxide character changes basic → amphoteric → acidic.
Q2. Which group of elements would you expect to form predominantly ionic compounds? Why? L4 Analyse
Answer: Group 1 (alkali metals) and Group 2 (alkaline earth metals) typically form ionic compounds with non-metals (especially Group 17 halogens).
Reason: Their low ionisation enthalpies make electron loss easy (form +1 or +2 cations). When combined with electronegative non-metals (high ΔₑgH, large EN), large electron-transfer (EN difference > 1.7) leads to ionic bonding.
Reason: Their low ionisation enthalpies make electron loss easy (form +1 or +2 cations). When combined with electronegative non-metals (high ΔₑgH, large EN), large electron-transfer (EN difference > 1.7) leads to ionic bonding.
Q3. The first element of each main group differs in many properties from other members. List 3 reasons. L4 Analyse
Answer:
- Small atomic size — leads to high charge density and high polarising power.
- High electronegativity — they hold bonding electrons more tightly.
- Absence of d-orbitals in valence shell — limits maximum coordination number to 4 (e.g., F has max 4 bonds; Cl can have 7 because of 3d availability).
Q4. Evaluate: A student claims "All alkali metals are equally reactive because they all have 1 valence electron." Critique. L5 Evaluate
Answer: The claim is INCORRECT. While all alkali metals have ns¹ configuration, their REACTIVITY varies significantly:
- Reactivity INCREASES down the group (Li < Na < K < Rb < Cs). - Reason: ionisation enthalpy DECREASES down group (Li 520 → Cs 376 kJ/mol). Lower IE → easier to lose electron → more reactive. - Cs reacts EXPLOSIVELY with water; Li reacts moderately.
This shows that valence electron count alone doesn't determine reactivity; how easily that electron is lost is governed by IE, atomic radius, and shielding.
- Reactivity INCREASES down the group (Li < Na < K < Rb < Cs). - Reason: ionisation enthalpy DECREASES down group (Li 520 → Cs 376 kJ/mol). Lower IE → easier to lose electron → more reactive. - Cs reacts EXPLOSIVELY with water; Li reacts moderately.
This shows that valence electron count alone doesn't determine reactivity; how easily that electron is lost is governed by IE, atomic radius, and shielding.
Q5. HOT (Create): Design an experiment to verify the diagonal relationship between Li and Mg. List specific observations you would expect. L6 Create
Sample Experimental Design:
- Reaction with N₂: Burn small samples of Li and Na in pure N₂ atmosphere.
Expected: Li and Mg both form nitrides (Li₃N, Mg₃N₂). Na, K, Rb, Cs do NOT react with N₂. → Confirms Li ↔ Mg diagonal. - Carbonate decomposition: Heat Li₂CO₃, Na₂CO₃, K₂CO₃, MgCO₃.
Expected: Li₂CO₃ and MgCO₃ decompose easily to oxide + CO₂. Na₂CO₃, K₂CO₃ are heat-stable. → Diagonal similarity. - Solubility of fluorides: Test solubility in water.
Expected: LiF and MgF₂ are sparingly soluble (covalent character + lattice energy). NaF, KF are soluble. → Diagonal similarity. - Hydroxide basicity: Compare LiOH, NaOH, Mg(OH)₂ in dissolving in water.
Expected: LiOH and Mg(OH)₂ are weak bases (less soluble) compared to NaOH (very strong, very soluble). → Diagonal similarity.
🧠 Assertion–Reason Questions
Choose: (A) Both true, R explains A. (B) Both true, R doesn't explain A. (C) A true, R false. (D) A false, R true.
A: Al₂O₃ is amphoteric — it reacts with both acids and bases.
R: Al lies on the metal/non-metal borderline of the periodic table.
Answer: (A). Both true; R explains A. Al's intermediate metallic/non-metallic character shows up in its oxide reacting both with acids (as a base, making AlCl₃) and bases (as an acid, making aluminate).
A: Lithium shows similarity with magnesium rather than with potassium.
R: Li and Mg have similar charge/radius ratios (diagonal relationship).
Answer: (A). Both true; R explains A. The unusual smallness of Li gives it Mg-like polarising power, leading to similar chemistry. Yet Li is also genuinely an alkali metal with Na, K — it just bridges these two groups.
A: Down group 1, reactivity decreases.
R: Atomic size increases down a group.
Answer: (D). A is FALSE — reactivity of alkali metals INCREASES down the group (Cs is most reactive). R is TRUE — atomic size does increase. The reason is: as size grows, IE decreases, valence electron is held more loosely, easier to lose → more reactive.
Frequently Asked Questions — Periodic Trends in Chemical Properties
How does valency vary across the periodic table?
Valency is the combining capacity of an element, generally equal to the number of electrons gained, lost or shared. For s- and p-block elements: Group 1 = +1, Group 2 = +2, Group 13 = +3, Group 14 = +4, Group 15 = −3 or +5, Group 16 = −2 or +6, Group 17 = −1 or +7, Group 18 = 0 (mostly). Across a period, valency typically increases from 1 to 4 then decreases back to 0. Down a group, valency stays the same. NCERT Class 11 Chemistry presents both the maximum valency (group number) and minimum valency (8 − group number for non-metals).
What are oxidation states and how do they vary periodically?
Oxidation state is the apparent charge an atom would have if all its bonds were ionic. For s-block elements, the oxidation states are fixed (+1 for Group 1, +2 for Group 2). p-block elements show variable oxidation states with the inert pair effect making heavier elements prefer lower states (Tl⁺ > Tl³⁺). Transition metals exhibit multiple oxidation states due to participation of (n−1)d electrons. Highest oxidation state across a period typically increases up to the central transition metal, then decreases. NCERT Class 11 Chemistry uses these trends to predict bonding and redox behaviour.
Why does metallic character decrease across a period?
Metallic character — the tendency of an atom to lose electrons and form positive ions — depends on ionisation enthalpy. Across a period, ionisation enthalpy increases due to higher effective nuclear charge and smaller atomic radius, making it harder to lose electrons. Therefore metallic character decreases left to right: Na (highly metallic) → Mg → Al (still metallic) → Si (metalloid) → P, S, Cl (non-metals). Down a group, ionisation enthalpy decreases, so metallic character increases: C (non-metal) → Si (metalloid) → Ge → Sn → Pb (metal). NCERT Class 11 Chemistry uses this trend extensively.
What is the anomalous behaviour of Period-2 elements?
Period-2 elements (Li, Be, B, C, N, O, F) show several properties different from their heavier group members because of: (1) very small atomic and ionic size, (2) high ionisation enthalpy, (3) absence of d-orbitals, (4) high electronegativity. For example, Li shows more covalent character than other alkali metals; Be is amphoteric while heavier Group 2 metals are basic; B forms electron-deficient compounds; N₂ is unreactive due to triple bond while P₄ is reactive; F is the most electronegative with anomalously high electron gain enthalpy of Cl. NCERT Class 11 Chemistry highlights these anomalies in detail.
What is the diagonal relationship?
Diagonal relationship is the similarity in chemical properties between an element in Period 2 and the element diagonally placed in Period 3 — specifically: Li & Mg, Be & Al, B & Si. These pairs show similar size, electronegativity and polarising power. For example, Li (like Mg) forms ionic nitride Li₃N and burns in oxygen to form normal oxide Li₂O. Be and Al both form amphoteric oxides and covalent halides. B and Si both form covalent network solids and acidic oxides. NCERT Class 11 Chemistry uses diagonal relationship as a teaching tool for predicting properties.
How do reactivity and acid-base nature of oxides vary?
Across a period, oxides transition from basic (alkali metals) → amphoteric (Al₂O₃, BeO) → acidic (CO₂, SO₃, Cl₂O₇). Down a group, basicity of oxides increases. Reactivity of metals increases down Groups 1 and 2 because of decreasing ionisation enthalpy. Reactivity of non-metals (halogens) decreases down the group because of decreasing electron-gain enthalpy. NCERT Class 11 Chemistry presents these trends as foundational concepts for understanding inorganic reactions in Class 12 and for solving acid-base equilibrium problems in Chapter 6.
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