Welcome to this deep dive into the fascinating world of the Modern Periodic Table and the periodic trends that govern the properties of elements! This is a core concept for your JEE preparation, as a solid understanding here will unlock many topics in inorganic and even physical chemistry. We'll start from the foundational principles and build up to the intricate details and exceptions vital for competitive exams.
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### The Modern Periodic Table: A Deeper Look
Our journey begins with understanding the backbone of modern chemistry: the Modern Periodic Table. While earlier attempts by Mendeleev were groundbreaking, they had limitations, primarily due to the dependence on atomic mass. The true revolution came with
Henry Moseley's work in 1913.
#### 1. Moseley's Law and the Fundamental Property
Moseley studied the X-ray spectra of various elements and discovered a remarkable regularity. He found that the square root of the frequency of X-rays emitted by an element was directly proportional to its atomic number (Z), not its atomic mass.
Mathematically, this is expressed as:
√ν = a(Z - b)
Where ν is the frequency of the X-rays, Z is the atomic number, and 'a' and 'b' are constants.
This groundbreaking discovery established that
atomic number (Z) is a more fundamental property of an element than atomic mass. It led to the modification of the periodic law, which now states:
"The physical and chemical properties of elements are a periodic function of their atomic numbers."
This revised law elegantly resolved anomalies in Mendeleev's table (e.g., the position of Ar and K, Co and Ni, Te and I), paving the way for the Modern Long Form of the Periodic Table, organized strictly by increasing atomic number.
#### 2. Organization of the Modern Periodic Table
The Modern Periodic Table is an arrangement of elements in increasing order of their atomic numbers, leading to a natural grouping based on similar electronic configurations and, consequently, similar chemical properties.
*
Periods (Horizontal Rows):
* There are
seven periods in total (1 to 7).
* A period number corresponds to the
principal quantum number (n) of the outermost electron shell (valence shell).
* As we move across a period, electrons are added to the same valence shell.
* The number of elements in each period is determined by the maximum number of electrons that can be accommodated in the respective shell (2n² rule, though slightly modified due to filling order).
*
Period No. |
n (Valence Shell) |
No. of Elements |
Description |
|---|
1 |
1 |
2 |
Shortest period (H, He) |
2 |
2 |
8 |
Short period |
3 |
3 |
8 |
Short period |
4 |
4 |
18 |
Long period |
5 |
5 |
18 |
Long period |
6 |
6 |
32 |
Longest period (includes Lanthanides) |
7 |
7 |
32 |
Incomplete period (includes Actinides) |
*
Finding Period: The highest principal quantum number (n) in the element's electronic configuration gives its period number. E.g., for Na (1s²2s²2p⁶3s¹), n=3, so it's in Period 3.
*
Groups (Vertical Columns):
* There are
18 groups (1 to 18) according to the IUPAC system.
* Elements in the same group have
similar valence shell electronic configurations and thus exhibit similar chemical properties.
*
Finding Group:
*
s-block elements: Group number = Number of valence electrons. (E.g., Na: [Ne]3s¹, 1 valence electron, Group 1).
*
p-block elements: Group number = 10 + Number of valence electrons (s and p electrons). (E.g., Cl: [Ne]3s²3p⁵, 7 valence electrons, Group 17).
*
d-block elements (Transition Elements): Group number = Number of electrons in (n-1)d subshell + Number of electrons in ns subshell. (E.g., Fe: [Ar]3d⁶4s², 6+2=8, Group 8).
*
f-block elements: These are generally placed below the main table and belong to Group 3 in the main block.
*
Blocks (s, p, d, f):
* This classification is based on the subshell to which the last differentiating electron enters.
*
s-block: Groups 1 and 2. Last electron enters 's' subshell. General electronic configuration:
ns¹⁻². Highly reactive metals.
*
p-block: Groups 13 to 18. Last electron enters 'p' subshell. General electronic configuration:
ns²np¹⁻⁶. Contains metals, non-metals, and metalloids. Group 18 are noble gases (np⁶).
*
d-block (Transition Elements): Groups 3 to 12. Last electron enters (n-1)d subshell. General electronic configuration:
(n-1)d¹⁻¹⁰ns¹⁻². All are metals, typically colored compounds, variable valencies.
*
f-block (Inner Transition Elements): Lanthanides (4f series) and Actinides (5f series). Last electron enters (n-2)f subshell. General electronic configuration:
(n-2)f¹⁻¹⁴(n-1)d⁰⁻¹ns². Mostly metals, radioactivity in actinides.
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### Unveiling Periodic Trends – The Language of Elements
The term
periodicity refers to the recurrence of similar properties of elements at regular intervals when they are arranged in increasing order of their atomic numbers. This periodicity arises directly from the
periodic recurrence of similar valence shell electronic configurations. Elements with similar outer electronic configurations will naturally exhibit similar chemical behaviors.
Let's delve into the crucial periodic properties:
#### 1. Atomic Radius
The atomic radius is a measure of the size of an atom. However, defining the "boundary" of an atom is complex due to the probabilistic nature of electron clouds. Therefore, we use practical definitions:
*
Covalent Radius: Half the distance between the nuclei of two identical atoms bonded together by a single covalent bond (e.g., in Cl₂ molecule).
*
Metallic Radius: Half the inter-nuclear distance between two adjacent metal atoms in a metallic crystal lattice.
*
Van der Waals Radius: Half the inter-nuclear distance between two adjacent identical atoms belonging to different molecules in a solid state (non-bonded atoms). Typically, Van der Waals radius > Metallic radius > Covalent radius.
Factors Affecting Atomic Radius:
1.
Effective Nuclear Charge (Zeff): The net positive charge experienced by an electron in a multi-electron atom. Higher Zeff pulls electrons closer to the nucleus, reducing atomic size.
2.
Number of Electron Shells: As the number of shells increases, the atomic size increases.
3.
Shielding/Screening Effect: Inner shell electrons repel outer shell electrons, reducing the attraction of the nucleus for the valence electrons. A stronger shielding effect leads to a larger atomic size.
Trends in Atomic Radius:
*
Across a Period (Left to Right):
*
Trend: Atomic radius
decreases.
*
Reason: As we move across a period, electrons are added to the
same valence shell. However, the
nuclear charge (Z) increases by one unit for each successive element. The added electrons do not significantly shield each other, leading to an
increase in Zeff. This stronger attraction pulls the valence electrons closer to the nucleus, shrinking the atomic size.
*
Example: Li (152 pm) > Be (111 pm) > B (88 pm) > C (77 pm) > N (74 pm) > O (66 pm) > F (64 pm).
*
JEE Focus: The decrease is less pronounced towards the end of the period due to increased electron-electron repulsion. For noble gases, Van der Waals radius is quoted, which is larger, hence they appear larger, but if covalent radius could be measured, it would be very small.
*
Down a Group (Top to Bottom):
*
Trend: Atomic radius
increases.
*
Reason: As we move down a group, a
new principal quantum shell is added at each step. Although the nuclear charge increases, the addition of new shells (increasing 'n') and the
increased shielding effect from inner electrons outweigh the effect of increased nuclear charge. The valence electrons are farther from the nucleus, leading to a larger atomic size.
*
Example: Li (152 pm) < Na (186 pm) < K (227 pm) < Rb (248 pm) < Cs (265 pm).
*
JEE Focus - Lanthanoid Contraction: A notable exception and crucial concept! In the 6th period, after Lanthanum (La, Z=57), the 14 elements (Ce to Lu) involve the filling of the
4f subshell. The 4f electrons provide
very poor shielding effect due to their diffuse shapes. This poor shielding results in a significantly
higher Zeff for elements following the lanthanides. Consequently, the atomic radii of elements from the 6th period (e.g., Hf, Ta, W) are unexpectedly similar to their corresponding elements in the 5th period (e.g., Zr, Nb, Mo), instead of being much larger. This phenomenon is called
Lanthanoid Contraction. It significantly impacts the chemistry of post-lanthanide elements (e.g., similar radii and densities of Zr/Hf, Nb/Ta).
#### 2. Ionic Radius
The ionic radius is the effective distance from the center of the ion to its outermost electron shell.
*
Cations (Positive Ions):
* Formed by
loss of electrons.
*
Cations are always smaller than their parent neutral atoms.
*
Reason: Upon losing electrons, the remaining electrons experience a
greater effective nuclear charge (same positive nucleus, fewer negative electrons). Also, often an entire shell is removed.
*
Example: Na (186 pm) > Na⁺ (102 pm); Mg (160 pm) > Mg²⁺ (72 pm). Note that higher positive charge means smaller size (Mg²⁺ < Na⁺ for isoelectronic species).
*
Anions (Negative Ions):
* Formed by
gain of electrons.
*
Anions are always larger than their parent neutral atoms.
*
Reason: Upon gaining electrons, the number of electrons increases while the nuclear charge remains the same. This leads to
increased electron-electron repulsion and a
decreased effective nuclear charge experienced by each electron, causing the electron cloud to expand.
*
Example: Cl (99 pm) < Cl⁻ (181 pm); O (66 pm) < O²⁻ (140 pm). Note that higher negative charge means larger size (O²⁻ > F⁻ for isoelectronic species).
Trends in Ionic Radius:
*
Across a Period: For isoelectronic species (ions with the same number of electrons), ionic radius
decreases with increasing atomic number (Z) because Zeff increases.
*
Example of Isoelectronic Series (10 electrons): O²⁻ (140 pm) > F⁻ (133 pm) > Na⁺ (102 pm) > Mg²⁺ (72 pm) > Al³⁺ (53.5 pm).
*
Down a Group: Ionic radius
increases due to the addition of new electron shells, similar to atomic radius.
*
Example: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺.
#### 3. Ionization Enthalpy (IE) / Ionization Energy
Definition: The minimum amount of energy required to remove the most loosely bound electron from an
isolated gaseous atom in its ground state to form a gaseous cation.
M(g) + IE₁ → M⁺(g) + e⁻ (First Ionization Enthalpy, IE₁)
The energy required to remove the second electron from the unipositive ion is the
Second Ionization Enthalpy (IE₂), and so on.
M⁺(g) + IE₂ → M²⁺(g) + e⁻
Always, IE₁ < IE₂ < IE₃... because it's always harder to remove an electron from an already positively charged ion due to increased Zeff.
Factors Affecting Ionization Enthalpy:
1.
Atomic Size: Larger atoms have valence electrons farther from the nucleus, experiencing less attraction, hence lower IE.
2.
Effective Nuclear Charge (Zeff): Higher Zeff means stronger attraction for valence electrons, thus higher IE.
3.
Shielding Effect: Stronger shielding reduces Zeff, making it easier to remove electrons, thus lower IE.
4.
Penetration Effect: Electrons in orbitals closer to the nucleus (e.g., s-orbitals penetrate more than p, d, or f) are more strongly attracted. Thus, removing an electron from an s-orbital requires more energy than from a p-orbital in the same shell.
5.
Electronic Configuration: Half-filled (p³, d⁵, f⁷) and fully-filled (p⁶, d¹⁰, f¹⁴) subshells are extra stable. Removing an electron from such configurations requires exceptionally high energy.
Trends in Ionization Enthalpy:
*
Across a Period (Left to Right):
*
Trend: Generally
increases.
*
Reason: Atomic size decreases and Zeff increases across a period. Both factors make it harder to remove an electron.
*
Example: Li < Be < B < C < N < O < F < Ne (general trend).
*
JEE Focus - Exceptions to the Trend:
*
IE₁ of Group 13 < IE₁ of Group 2 (e.g., Boron < Beryllium):
* Be: [He]2s² (fully filled s-orbital, stable)
* B: [He]2s²2p¹ (electron in p-orbital, requires less energy)
* The 2p electron in Boron is further from the nucleus and experiences less penetration than the 2s electron in Beryllium. Therefore, removing a 2p electron from B is easier than removing a 2s electron from Be, despite Boron having a higher nuclear charge.
*
IE₁ of Group 16 < IE₁ of Group 15 (e.g., Oxygen < Nitrogen):
* N: [He]2s²2p³ (half-filled p-orbital, stable)
* O: [He]2s²2p⁴ (electron in paired p-orbital, experiences repulsion)
* Removing an electron from the stable half-filled 2p subshell of Nitrogen requires more energy than removing an electron from the 2p⁴ configuration of Oxygen, where the paired electron experiences repulsion, making it easier to remove.
*
Down a Group (Top to Bottom):
*
Trend: Generally
decreases.
*
Reason: Atomic size increases and shielding effect increases down a group. The valence electrons are farther from the nucleus and less strongly attracted, requiring less energy to remove.
*
Example: Li > Na > K > Rb > Cs.
*
JEE Focus - Exceptions for d-block and f-block:
* Due to
Lanthanoid Contraction, elements in the 5th period (e.g., Zr) and 6th period (e.g., Hf) have very similar atomic radii and consequently, their ionization enthalpies are also very similar, sometimes even reversing the trend (e.g., IE₁ of Hf is slightly higher than Zr). This is because the increased Zeff due to poor 4f shielding makes it harder to remove electrons.
#### 4. Electron Gain Enthalpy (EGE) / Electron Affinity (EA)
Definition: The enthalpy change that occurs when an electron is added to an
isolated gaseous atom in its ground state to form a gaseous anion.
X(g) + e⁻ → X⁻(g) + EGE
* If energy is
released (exothermic process), EGE is
negative. This indicates the atom has a tendency to accept electrons.
* If energy is
absorbed (endothermic process), EGE is
positive. This indicates the atom resists accepting electrons (or requires energy input to force an electron onto it).
Factors Affecting Electron Gain Enthalpy:
1.
Atomic Size: Smaller atoms generally have more negative EGE (stronger attraction for added electron).
2.
Effective Nuclear Charge (Zeff): Higher Zeff means stronger attraction for the incoming electron, leading to a more negative EGE.
3.
Electronic Configuration: Elements with stable (half-filled or fully-filled) electronic configurations have very low or even positive EGE values (e.g., noble gases, alkaline earth metals, nitrogen), as adding an electron would destabilize them.
4.
Electron-Electron Repulsion: For very small atoms, the added electron experiences significant repulsion from existing electrons in the small valence shell, which can make EGE less negative or even positive.
Trends in Electron Gain Enthalpy:
*
Across a Period (Left to Right):
*
Trend: Generally becomes
more negative (i.e., elements have a greater tendency to accept electrons).
*
Reason: Atomic size decreases and Zeff increases, making it easier for the nucleus to attract an additional electron. Non-metals (right side) generally have high negative EGE.
*
Example: Li (–60 kJ/mol) → F (–328 kJ/mol) → Ne (+116 kJ/mol). Noble gases have positive EGE due to stable octet.
*
Down a Group (Top to Bottom):
*
Trend: Generally becomes
less negative (i.e., tendency to accept electrons decreases).
*
Reason: Atomic size increases, and Zeff slightly decreases, making the attraction for an incoming electron weaker.
*
Example: F (–328 kJ/mol) > Cl (–349 kJ/mol) > Br (–325 kJ/mol) > I (–295 kJ/mol).
*
JEE Focus - Exceptions and Anomalies:
*
Halogens (Group 17): Chlorine has more negative EGE than Fluorine!
*
Reason: Fluorine is a very small atom. The incoming electron experiences significant
electron-electron repulsion from the existing electrons in the very compact 2p subshell of fluorine. In contrast, chlorine is larger, and its 3p subshell is more diffuse, allowing the incoming electron to be accommodated with less repulsion, leading to a more stable anion and thus a more negative EGE.
*
Group 15 (Nitrogen, Phosphorus): Nitrogen has a positive EGE, while phosphorus has a negative EGE.
*
Reason: Nitrogen has a stable half-filled 2p³ configuration and a very small size, leading to strong electron-electron repulsion and requiring energy to add an electron. Phosphorus is larger, and the 3p³ configuration is less compact, allowing for easier electron addition.
*
Group 2 (Alkaline Earth Metals) and Group 18 (Noble Gases): Have positive (or near zero) EGE values due to stable s² and s²p⁶ configurations, respectively.
#### 5. Electronegativity (EN)
Definition: The
tendency of an atom in a chemical compound to attract the shared pair of electrons towards itself. It is a
relative property, not an absolute energy value.
It differs from Electron Gain Enthalpy: EGE is about an isolated gaseous atom gaining an electron, while EN is about an atom in a *bond* pulling shared electrons.
Common Scales for Electronegativity:
*
Pauling Scale: Most widely used, based on bond dissociation energies. Fluorine is assigned 4.0 (highest).
*
Mulliken Scale: Average of ionization enthalpy and electron affinity.
Electronegativity = (IE + EGE) / 2 (Both in electron volts)
Factors Affecting Electronegativity:
1.
Effective Nuclear Charge (Zeff): Higher Zeff, higher electronegativity.
2.
Atomic Size: Smaller atoms hold valence electrons more tightly and attract shared electrons more strongly, hence higher electronegativity.
3.
Hybridization: For carbon, sp-hybridized carbon is more electronegative than sp² which is more electronegative than sp³ (due to higher s-character, electrons are closer to the nucleus).
Trends in Electronegativity:
*
Across a Period (Left to Right):
*
Trend: Increases.
*
Reason: Zeff increases and atomic size decreases, leading to a stronger attraction for shared electrons.
*
Example: Li (0.98) < Be (1.57) < B (2.04) < C (2.55) < N (3.04) < O (3.44) < F (3.98).
* Noble gases are generally not assigned electronegativity values as they don't readily form bonds.
*
Down a Group (Top to Bottom):
*
Trend: Decreases.
*
Reason: Atomic size increases and Zeff slightly decreases, reducing the attraction for shared electrons as they are farther from the nucleus.
*
Example: F (3.98) > Cl (3.16) > Br (2.96) > I (2.66).
Applications of Electronegativity (JEE Focus):
*
Bond Polarity: Larger the electronegativity difference between two bonded atoms, the more polar the bond.
*
Acidic/Basic Character of Oxides and Hydroxides:
*
Across a period: Acidic character of oxides increases (e.g., Na₂O (basic) → SiO₂ (amphoteric/weakly acidic) → P₄O₁₀ (acidic) → SO₃ (acidic) → Cl₂O₇ (strongly acidic)). Non-metals form acidic oxides.
*
Down a group: Basic character of oxides increases (e.g., Li₂O < Na₂O < K₂O < Rb₂O < Cs₂O, all basic). Metallic character increases.
*
Hydroxides (M-O-H): If M is highly electronegative (non-metal), it attracts electrons from O-H bond, releasing H⁺ (acidic). If M is less electronegative (metal), it releases OH⁻ (basic).
#### 6. Metallic and Non-metallic Character
*
Metallic Character: The tendency of an element to lose electrons and form positive ions.
* High metallic character means low IE, low EGE, low EN.
*
Trend: Decreases across a period (metals on left, non-metals on right) and increases down a group.
*
Non-metallic Character: The tendency of an element to gain electrons and form negative ions.
* High non-metallic character means high IE, high (more negative) EGE, high EN.
*
Trend: Increases across a period and decreases down a group.
*
Metalloids: Elements exhibiting properties intermediate between metals and non-metals (e.g., B, Si, Ge, As, Sb, Te, At).
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### Conclusion
A thorough understanding of the Modern Periodic Table and its underlying periodic trends is fundamental to excelling in Chemistry. Remember to not just memorize the trends but understand the
reasons behind them, especially the exceptions and anomalies, as these are frequently tested in JEE. Factors like Zeff, atomic size, shielding, and electronic configuration are the keys to unlocking the periodic behavior of elements. Keep practicing with examples, and you'll master this crucial topic!