This MCQ module is based on: Periodic Trends Physical
TOPIC 11 OF 28
Periodic Trends Physical
🎓 Class 11
Chemistry
CBSE
Theory
Ch 3 – Classification of Elements and Periodicity in Properties
⏱ ~14 min
🌐 Language: [gtranslate]
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Periodic Trends in Physical Properties
3.7 Periodic Trends in Physical Properties
There are many observable patterns in the physical and chemical properties of elements as we descend in a group or move across a period in the Periodic Table. For example, within a period, chemical reactivity tends to be high in Group 1 metals, lowers towards the middle of the table, and increases to a maximum in the Group 17 non-metals. Likewise, within a group of representative metals (say alkali metals, Group 1), chemical reactivity increases on going down the group, whereas within a group of non-metals (say halogens), chemical reactivity decreases down the group. But why do these trends exist? And how can we explain them in terms of these properties?
3.7.1 Atomic Radius
You can well imagine that finding the size of an atom is a lot more complicated than measuring the radius of a ball. Do you know why? Firstly, because the size of an atom (~1.2 Å, i.e., 1.2 × 10⁻¹⁰ m) is incredibly small to be measured directly. Secondly, since the electron cloud surrounding the atom does not have a sharp boundary, the determination of the atomic size cannot be precise. In other words, there is no way in which we can isolate an atom and measure its diameter as we do for a ball.
One practical approach to estimate the size of an atom of a non-metallic element is to measure the distance between two atoms when they are bound together by a single bond in a covalent molecule and from this value, the "covalent radius" of the element can be calculated.
Types of Atomic Radii:
- Covalent radius: Half the internuclear distance between two atoms in a single covalent bond.
- Metallic radius: Half the internuclear distance between two adjacent atoms in a metallic crystal.
- Van der Waals radius: Half the distance between two non-bonded atoms in close contact (in solid state).
Periodic Trends in Atomic Radius
| Element (Period 2) | Atomic radius (pm) |
|---|---|
| Li | 152 |
| Be | 111 |
| B | 88 |
| C | 77 |
| N | 74 |
| O | 66 |
| F | 64 |
| Element (Group 1) | Atomic radius (pm) |
|---|---|
| Li | 152 |
| Na | 186 |
| K | 231 |
| Rb | 244 |
| Cs | 262 |
Trends:
- → Across a period (left to right): atomic radius DECREASES.
Reason: Nuclear charge (Z) increases by 1 per element, but electrons add to the same shell. Higher nuclear charge pulls the same shell of electrons closer, shrinking the atom. - ↓ Down a group: atomic radius INCREASES.
Reason: A new shell is added at each step. Even though Z increases, the increasing distance and shielding of inner electrons offsets the nuclear pull.
3.7.2 Ionic Radius
The removal of an electron from an atom results in the formation of a cation, whereas gain of an electron leads to an anion. The ionic radii can be estimated by measuring the distances between cations and anions in ionic crystals.
- Cation is always SMALLER than the parent atom (one electron lost; effective nuclear charge over fewer electrons increases). E.g., Na (186 pm) → Na⁺ (95 pm).
- Anion is always LARGER than the parent atom (one electron gained; same Z over more electrons; greater electron-electron repulsion). E.g., Cl (99 pm) → Cl⁻ (181 pm).
Isoelectronic Species
Atoms and ions which contain the same number of electrons are called isoelectronic species. For example, O²⁻, F⁻, Na⁺ and Mg²⁺ all have 10 electrons. Their radii decrease with increasing nuclear charge (Z):
| Species | Z (protons) | Electrons | Ionic radius (pm) |
|---|---|---|---|
| O²⁻ | 8 | 10 | 140 |
| F⁻ | 9 | 10 | 136 |
| Na⁺ | 11 | 10 | 95 |
| Mg²⁺ | 12 | 10 | 65 |
| Al³⁺ | 13 | 10 | 50 |
3.7.3 Ionisation Enthalpy
A quantitative measure of the tendency of an element to lose electron is given by its ionisation enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom (X) in its ground state:
X(g) → X⁺(g) + e⁻ ΔᵢH₁ = 1st Ionisation Enthalpy
Successive ionisations always require more energy:
ΔᵢH₁ < ΔᵢH₂ < ΔᵢH₃ < ...
Periodic Trends in Ionisation Enthalpy
| Element (Period 2) | ΔᵢH₁ (kJ/mol) |
|---|---|
| Li | 520 |
| Be | 899 |
| B | 801 |
| C | 1086 |
| N | 1402 |
| O | 1314 |
| F | 1681 |
| Ne | 2081 |
Trends in IE:
- → Across a period: INCREASES (smaller atom + higher Z = electrons more tightly held).
- ↓ Down a group: DECREASES (larger atom + more shielding = electrons easier to remove).
- B (801) < Be (899): removing one 2p electron is easier than removing a 2s electron (s-orbital is more penetrating, more strongly bound).
- O (1314) < N (1402): half-filled 2p³ in N is unusually stable; removing electron from O's 2p⁴ requires breaking pair-up.
🎯 Interactive: Periodic Trend Visualizer
Pick a property and see its trend across Period 2 and down Group 1.
Insight: Atomic radius DECREASES across period; ionisation enthalpy and electronegativity INCREASE.
3.7.4 Electron Gain Enthalpy
When an electron is added to a neutral gaseous atom (X) to convert it into a negative ion, the enthalpy change accompanying the process is defined as the electron gain enthalpy:
X(g) + e⁻ → X⁻(g) ΔₑgH (electron gain enthalpy)
Convention: Negative ΔₑgH = energy released (atom "wants" the electron). Positive ΔₑgH = energy absorbed (atom resists adding electrons).
| Element | ΔₑgH (kJ/mol) | Notes |
|---|---|---|
| F | −328 | Highly negative (small atom but compactness reduces ΔₑgH) |
| Cl | −349 | MOST NEGATIVE (anomaly — F's compactness limits its ΔₑgH) |
| Br | −325 | Less negative than Cl |
| I | −295 | Even less negative (large atom) |
| N | +0 (slightly positive) | Half-filled 2p³ is stable; resists adding e⁻ |
| Noble gases (Ne, Ar...) | +ve | Already filled; would need to enter higher shell |
Key reason for Cl > F (more negative): Although F is smaller and more electronegative, its compactness causes greater inter-electronic repulsion when adding the new electron to the very small 2p subshell. Cl's larger size means the new electron experiences less repulsion → energy release is greater.
3.7.5 Electronegativity
Electronegativity is the relative tendency of an atom in a molecule to attract the shared pair of electrons towards itself. Unlike ionisation enthalpy and electron gain enthalpy, it is not a measurable physical property — it is a relative scale.
Pauling Electronegativity Values
| Element | EN (Pauling) |
|---|---|
| F (highest) | 4.0 |
| O | 3.5 |
| N, Cl | 3.0 |
| C, S, Br | 2.5 |
| H, P, I | 2.1 |
| Na | 0.9 |
| K, Cs (lowest stable) | 0.7 |
Trends in Electronegativity:
- → Across period: INCREASES (smaller atom holds bonding electrons more tightly).
- ↓ Down group: DECREASES (larger atom, lower attraction).
- Most electronegative: Fluorine (4.0). Least: Cs (0.7).
- Difference between two atoms' EN predicts ionic vs covalent bond character: ΔEN > 1.7 → predominantly ionic; ΔEN < 1.7 → covalent.
Worked Examples
Worked Example 1 (NCERT Problem 3.4): Smallest in a Period
Which of the following species will have the largest and the smallest size? Mg, Mg²⁺, Al, Al³⁺.
Mg: [Ne] 3s² (atom, period 3, ~160 pm).
Al: [Ne] 3s²3p¹ (atom, period 3, ~143 pm).
Mg²⁺: [Ne] (10 electrons over 12 protons; ionic radius ~65 pm).
Al³⁺: [Ne] (10 electrons over 13 protons; ionic radius ~50 pm).
Largest: Mg (atom, lowest effective nuclear charge per electron).
Smallest: Al³⁺ (10 electrons pulled by 13 protons).
Al: [Ne] 3s²3p¹ (atom, period 3, ~143 pm).
Mg²⁺: [Ne] (10 electrons over 12 protons; ionic radius ~65 pm).
Al³⁺: [Ne] (10 electrons over 13 protons; ionic radius ~50 pm).
Largest: Mg (atom, lowest effective nuclear charge per electron).
Smallest: Al³⁺ (10 electrons pulled by 13 protons).
Worked Example 2 (NCERT Problem 3.5): Comparing IE
Which of the following pairs of elements would have a more negative electron gain enthalpy? (i) O or F, (ii) F or Cl.
(i) O vs F: Both are period 2, but F has higher Z and smaller radius.
F has more negative ΔₑgH (because adding e⁻ completes 2p⁵ → 2p⁶, fully filled, very stable).
F = -328 kJ/mol; O = -141 kJ/mol.
(ii) F vs Cl: Surprisingly, Cl has more negative ΔₑgH!
F = -328 kJ/mol; Cl = -349 kJ/mol.
Reason: F's small 2p shell has high inter-electronic repulsion when adding e⁻; Cl's larger 3p shell accommodates the new electron with less repulsion → more energy released.
F has more negative ΔₑgH (because adding e⁻ completes 2p⁵ → 2p⁶, fully filled, very stable).
F = -328 kJ/mol; O = -141 kJ/mol.
(ii) F vs Cl: Surprisingly, Cl has more negative ΔₑgH!
F = -328 kJ/mol; Cl = -349 kJ/mol.
Reason: F's small 2p shell has high inter-electronic repulsion when adding e⁻; Cl's larger 3p shell accommodates the new electron with less repulsion → more energy released.
Worked Example 3: Electronegativity Comparison
Arrange the following elements in increasing order of electronegativity: F, P, S, Si, Cl.
Looking up values: F = 4.0, Cl = 3.0, S = 2.5, P = 2.1, Si = 1.8.
Order: Si < P < S < Cl < F.
Reasoning: Period 3 elements (Si, P, S, Cl) — EN increases left-to-right. F (period 2, group 17) is even more electronegative because it's smaller and has higher effective nuclear charge per shell.
Order: Si < P < S < Cl < F.
Reasoning: Period 3 elements (Si, P, S, Cl) — EN increases left-to-right. F (period 2, group 17) is even more electronegative because it's smaller and has higher effective nuclear charge per shell.
📐 Activity 3.3 — Predicting Trends Without Tables
Setup: Without consulting any reference, predict the answers to these questions using only your knowledge of periodic trends:
- Which has larger atomic radius: Na or Mg?
- Which has higher first ionisation enthalpy: K or Cs?
- Which is more electronegative: O or S?
- Why does N have higher IE than O?
- Predict whether Cl⁻ or Cl is larger.
Predict your answers, then check:
1. Na > Mg. Across period 3, Z increases but same shell. Higher Z pulls electrons closer.
2. K > Cs. Down group 1, atom gets larger, easier to remove electron. K is smaller, holds electron more tightly.
3. O > S. Down group 16, electronegativity decreases. O (3.5) > S (2.5).
4. Half-filled stability of 2p³ in N. N has stable 2p³ (Hund's max multiplicity). Adding/removing destabilizes more than O's 2p⁴.
5. Cl⁻ > Cl. Anion is always larger than its parent atom. Adding electron increases electron-electron repulsion.
🎯 Competency-Based Questions
Q1. Among the following, which species has the largest size? Mg²⁺, Na, F⁻, O²⁻, Al³⁺. L3 Apply
Answer: All except Na are isoelectronic (10 electrons each); for those, size decreases with increasing Z. Na (atom) has 11 electrons. So among isoelectronic ions: O²⁻ (Z=8) is largest. But Na (atom, 186 pm) is even larger because its 11 electrons go into 3s shell (n=3), bigger than n=2 shell. Largest: Na.
Q2. Why is the second ionisation enthalpy of Na very high while the first is moderate? L4 Analyse
Answer: Na has [Ne] 3s¹. First electron removed is the loosely held 3s¹ → ΔᵢH₁ = 496 kJ/mol (moderate). After removal, Na⁺ has [Ne] config — fully filled stable inner shell. Removing a 2nd electron requires disrupting this stable shell from a 2p orbital. ΔᵢH₂ = 4562 kJ/mol (~9× larger). This is why Na is reliably +1 in compounds.
Q3. Why does O have lower first ionisation enthalpy than N, even though O has a higher atomic number? L5 Evaluate
Answer: N has electronic configuration [He] 2s² 2p³ — the 2p subshell is exactly half-filled. According to Hund's rule, half-filled subshells have extra stability (exchange energy). Removing one electron from N would disrupt this stable configuration, requiring extra energy.
O has [He] 2s² 2p⁴ — the additional 2p electron pairs up with another, causing electron-electron repulsion. This makes the paired electron easier to remove. So N (1402) > O (1314) kJ/mol, despite Z(O) > Z(N).
O has [He] 2s² 2p⁴ — the additional 2p electron pairs up with another, causing electron-electron repulsion. This makes the paired electron easier to remove. So N (1402) > O (1314) kJ/mol, despite Z(O) > Z(N).
Q4. Critically evaluate: "Cl has more negative electron gain enthalpy than F because Cl is more electronegative." L5 Evaluate
Answer: The statement is INCORRECT in reasoning. While Cl does have more negative ΔₑgH than F (-349 vs -328 kJ/mol), F is MORE electronegative (4.0 vs 3.0). Electronegativity and electron gain enthalpy are related but distinct: electronegativity reflects the atom's pull on shared electrons in a bond, while ΔₑgH is for adding an electron to an isolated atom.
The reason Cl's ΔₑgH is more negative is the SMALL SIZE of F: when F gains an electron, the new electron enters its compact 2p subshell where there is high inter-electronic repulsion, partially offsetting the energy released. Cl's 3p subshell is more spacious, less repulsion → more energy released.
The reason Cl's ΔₑgH is more negative is the SMALL SIZE of F: when F gains an electron, the new electron enters its compact 2p subshell where there is high inter-electronic repulsion, partially offsetting the energy released. Cl's 3p subshell is more spacious, less repulsion → more energy released.
Q5. HOT (Create): A new "transition metal" of period 8 (Z = 124) is hypothesized. Predict its atomic radius, first ionisation enthalpy (compare with Au), and electronegativity. L6 Create
Sample Solution:
Atomic radius: Period 8 means a new shell (n=8). Compared to Au (period 6, ~144 pm covalent), Z=124 should have larger atomic radius (~ 200+ pm) due to extra shell, despite higher Z (relativistic effects matter for super-heavy elements!).
Ionisation enthalpy: With more shielding from filled inner shells, IE should be LOWER than Au (890 kJ/mol). Predicted: ~ 700-800 kJ/mol. (Caveat: relativistic contraction of s-orbitals can sharply increase IE for some super-heavy elements like 112-118.)
Electronegativity: Lower than Au (2.5). Predicted: ~ 1.7-2.0.
Caveats: Relativistic effects in super-heavy elements can drastically alter trends! Element 124's actual properties cannot be reliably predicted by simple periodic trends alone — quantum-mechanical calculations are required.
Atomic radius: Period 8 means a new shell (n=8). Compared to Au (period 6, ~144 pm covalent), Z=124 should have larger atomic radius (~ 200+ pm) due to extra shell, despite higher Z (relativistic effects matter for super-heavy elements!).
Ionisation enthalpy: With more shielding from filled inner shells, IE should be LOWER than Au (890 kJ/mol). Predicted: ~ 700-800 kJ/mol. (Caveat: relativistic contraction of s-orbitals can sharply increase IE for some super-heavy elements like 112-118.)
Electronegativity: Lower than Au (2.5). Predicted: ~ 1.7-2.0.
Caveats: Relativistic effects in super-heavy elements can drastically alter trends! Element 124's actual properties cannot be reliably predicted by simple periodic trends alone — quantum-mechanical calculations are required.
🧠 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: Atomic radius of Na (186 pm) is greater than Mg (160 pm).
R: Across a period, atomic radius decreases due to increased nuclear charge with the same shell.
Answer: (A). Both true; R correctly explains A. From Na (Z=11) to Mg (Z=12), nuclear charge increases by 1 but electrons go into the same 3s shell, increasing effective nuclear pull and contracting the atom.
A: Ionisation enthalpy of N (1402) is higher than O (1314).
R: N has electronic configuration 2p³ (half-filled, extra stable).
Answer: (A). Both true; R correctly explains A. The half-filled 2p³ in N has exchange-energy stability. Removing electron from N requires disrupting this stable configuration; from O's paired 2p⁴, the paired electron is easier to remove (less stable to begin with).
A: All cations are smaller than their parent atoms.
R: Removing electrons reduces electron-electron repulsion and increases effective nuclear charge per remaining electron.
Answer: (A). Both true; R correctly explains A. Fewer electrons sharing the same nuclear charge → each electron is more tightly held → ion shrinks. (Often outermost electron(s) and even outermost shell are lost.)
Frequently Asked Questions — Periodic Trends in Physical Properties
How does atomic radius vary in the periodic table?
Atomic radius generally decreases across a period from left to right because the effective nuclear charge increases while electrons are added to the same shell, pulling the electron cloud closer. Atomic radius increases down a group because new shells are added, increasing the distance of the outermost electrons from the nucleus. For example, in Period 2: Li (152 pm) > Be (112 pm) > B (88 pm) > C (77 pm); in Group 1: Li (152) < Na (186) < K (231). NCERT Class 11 Chemistry uses these trends to explain bond lengths and reactivity.
What is ionisation enthalpy and how does it vary periodically?
Ionisation enthalpy (IE) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom: M(g) → M⁺(g) + e⁻, ΔH = IE. It generally increases across a period (electrons more tightly held due to increasing Z_eff) and decreases down a group (outer electrons farther and shielded). Some exceptions exist due to electronic configuration stability — IE of B < Be (loss of single 2p vs paired 2s) and IE of O < N (loss from paired 2p vs half-filled stable 2p³). NCERT Class 11 Chemistry emphasises these trends and anomalies.
What is electron gain enthalpy?
Electron gain enthalpy is the enthalpy change when an isolated gaseous atom gains an electron to form an anion: X(g) + e⁻ → X⁻(g), ΔH = Δ_eg H. A negative value means energy is released and the atom readily accepts an electron. Electron gain enthalpy generally becomes more negative across a period (atom becomes more eager to gain an electron to complete its octet) and less negative down a group. Halogens have the most negative electron gain enthalpies. Noble gases have positive values because adding an electron disrupts their stable configurations. NCERT Class 11 Chemistry covers this trend.
What is electronegativity and why is it important?
Electronegativity is the relative tendency of an atom in a covalent bond to attract the shared pair of electrons toward itself. Unlike electron gain enthalpy (a measured property), electronegativity is a relative scale developed by Linus Pauling. Fluorine is the most electronegative element (4.0), and caesium/francium are the least. Electronegativity increases across a period and decreases down a group. The difference in electronegativity between bonded atoms predicts bond polarity — small differences give covalent bonds, large differences give ionic bonds. NCERT Class 11 Chemistry uses electronegativity widely.
Why does ionic radius differ from atomic radius?
When an atom loses electrons to form a cation, the ionic radius is smaller than the atomic radius because the same nuclear charge attracts fewer electrons, contracting the electron cloud. For example, Na (186 pm) → Na⁺ (102 pm). When an atom gains electrons to form an anion, the ionic radius is larger than the atomic radius because the increased electron-electron repulsion expands the cloud. For example, Cl (99 pm) → Cl⁻ (181 pm). For isoelectronic species (same number of electrons), the ionic radius decreases as nuclear charge increases. NCERT Class 11 Chemistry explores these patterns.
What is effective nuclear charge and shielding?
Effective nuclear charge (Z_eff) is the net positive charge experienced by an outer electron after the inner electrons partially shield it from the full nuclear charge: Z_eff = Z − S where Z is atomic number and S is the shielding constant. Inner shell electrons shield outer electrons most effectively, while electrons in the same shell shield each other less. NCERT Class 11 Chemistry uses Z_eff to explain why atomic radius decreases across a period (Z increases but S stays roughly constant) and why ionisation energy increases across a period. The concept is foundational for all periodic trends.
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