Consider the following statements with reference to the Periodic Table of chemical elements : I. lonisation potential gradually decreases along a period. II. In a group of elements, electron affinity decreases as the atomic weight increases. III. In a given period, electronegativity decreases as the atomic number increases. Which of these statement(s) is/are correct ?

1. Modern Periodic Law and Organization (basic)

The Modern Periodic Law is the fundamental principle that governs how we understand the building blocks of our universe. Historically, scientists like Mendeleev tried to organize elements based on their atomic mass, but this led to several confusing inconsistencies. The breakthrough came when Henry Moseley realized that the true identity of an element is defined by its atomic number—the number of protons in its nucleus. The Modern Periodic Law states: "The physical and chemical properties of the elements are a periodic function of their atomic numbers."

This organization creates a grid that acts as a predictive map for chemistry. The table is structured into two main directions:

• Periods (Horizontal Rows): Elements in a period have the same number of occupied electron shells (energy levels). As you move from left to right, the atomic number increases, and properties gradually shift.
• Groups (Vertical Columns): Elements in a group share the same number of valence electrons in their outermost shell. Because chemistry is all about how electrons interact, elements in the same group exhibit very similar chemical behaviors Science class X (NCERT 2025 ed.), Carbon and its Compounds, p.59.

The beauty of this arrangement is its periodicity. Similar to how a clock uses a repeating vibration to mark time Science-Class VII . NCERT(Revised ed 2025), Measurement of Time and Motion, p.111, the periodic table reflects a repeating pattern of chemical traits. This structure allows us to clearly categorize elements into metals, non-metals, and metalloids (like silicon and boron) which possess intermediate properties Science ,Class VIII . NCERT(Revised ed 2025), Nature of Matter: Elements, Compounds, and Mixtures, p.123.

Feature	Mendeleev's Periodic Law	Modern Periodic Law
Basis of Classification	Atomic Mass (Weight)	Atomic Number (Protons)
Logic	Mass determines properties	Electronic configuration determines properties
Reliability	Had anomalies (e.g., isotopes)	Universal and logically consistent

Sources: Science-Class VII . NCERT(Revised ed 2025), Measurement of Time and Motion, p.111; Science , class X (NCERT 2025 ed.), Carbon and its Compounds, p.59; Science ,Class VIII . NCERT(Revised ed 2025), Nature of Matter: Elements, Compounds, and Mixtures, p.123

2. Electronic Configuration and Valence Shells (basic)

To understand how elements behave, we must first look at how their electrons are organized. Electrons are not scattered randomly; they occupy specific energy levels called shells, labeled K, L, M, and N. The electronic configuration is simply a map of how many electrons live in each shell. For example, the K shell can hold a maximum of 2 electrons, while the L and M shells typically hold up to 8 in the basic representative elements we study first. Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47

The most important part of this map is the valence shell, which is the outermost shell of an atom. The electrons residing here are called valence electrons. These are the "social" electrons—they determine how an atom will react with others. Elements like Helium (2) or Neon (2, 8) have completely filled valence shells, making them incredibly stable and unreactive; we call these Noble Gases. Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46

Most other atoms have incomplete valence shells and are constantly "seeking" stability. They do this by trying to achieve a full octet (8 electrons in the outer shell, or 2 for Helium). This drive to reach a stable configuration is the fundamental reason behind all chemical reactions. For instance, Sodium (Atomic No. 11) has a configuration of 2, 8, 1. It has one lonely electron in its M shell, making it highly reactive as it seeks to lose that electron to reach the stable 2, 8 configuration of Neon. Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59

Element	Atomic Number	Configuration (K, L, M)	Valence Electrons
Magnesium (Mg)	12	2, 8, 2	2
Sulphur (S)	16	2, 8, 6	6
Argon (Ar)	18	2, 8, 8	8 (Stable)

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46-47; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59-60

3. Atomic Radius and Effective Nuclear Charge (intermediate)

To understand why atoms behave the way they do, we must first look at the Atomic Radius—the distance from the center of the nucleus to the outermost electron shell. This size isn't fixed like a hard ball; it is determined by a delicate balance of electrical forces. As we know from basic physics, opposite charges attract while like charges repel (Science, Class VIII. NCERT, Exploring Forces, p.71). In an atom, the positive protons in the nucleus are constantly trying to pull the negative electrons inward. However, the inner electrons act as a shield, pushing the outer electrons away. This net positive pull that an outer electron actually "feels" is called the Effective Nuclear Charge (Zeff).

As you move across a period (left to right) in the periodic table, the number of protons in the nucleus increases, but the electrons are added to the same principal energy level or shell. Because these new electrons don't provide much additional shielding, the Effective Nuclear Charge increases significantly. This stronger "tug" from the nucleus pulls the electron cloud tighter toward the center, causing the Atomic Radius to decrease. This explains why a Fluorine atom is much smaller than a Lithium atom, even though Fluorine has more subatomic particles.

Conversely, as you move down a group, the trend reverses. Although the nuclear charge (number of protons) is increasing, we are adding entirely new electron shells or principal quantum numbers. Each new shell is significantly further from the nucleus than the last. These additional layers of internal electrons provide a powerful "shielding effect," which outweighs the increase in nuclear charge. Consequently, the Atomic Radius increases as you move down a column.

Trend Direction	Atomic Radius	Primary Reason
Across a Period (→)	Decreases	Higher Zeff pulls electrons closer.
Down a Group (↓)	Increases	Addition of new principal energy levels (shells).

Sources: Science, Class VIII. NCERT, Exploring Forces, p.71

4. Chemical Bonding and the Octet Rule (intermediate)

In the vast world of chemistry, atoms are like people—they generally seek stability. This stability is achieved through Chemical Bonding. The driving force behind this is the Octet Rule: the tendency of atoms to prefer having eight electrons in their valence (outermost) shell. When an atom achieves this, it mirrors the electronic configuration of Noble Gases like Neon or Argon, which are naturally stable and unreactive Science, Class X, Metals and Non-metals, p.47.

Atoms reach this "happy state" of eight electrons primarily in two ways:

• Ionic Bonding: This involves the complete transfer of electrons. Metals (like Sodium) typically lose electrons to become positive cations, while non-metals (like Chlorine) gain them to become negative anions. Because these oppositely charged ions are held by strong electrostatic forces, ionic compounds usually have high melting and boiling points and conduct electricity when dissolved in water Science, Class X, Carbon and its Compounds, p.58.
• Covalent Bonding: This involves the sharing of electron pairs between atoms. This is the path chosen when losing or gaining electrons is energetically too difficult. For instance, Carbon has four valence electrons. Gaining four more to form a C⁴⁻ anion is difficult for a nucleus with only six protons to hold, and losing four to form a C⁴⁺ cation requires a massive amount of energy Science, Class X, Carbon and its Compounds, p.59. Instead, Carbon shares its electrons, forming covalent bonds that result in molecules with lower melting points and poor electrical conductivity Science, Class X, Carbon and its Compounds, p.59.

Conductivity

Feature	Ionic Bonding	Covalent Bonding
Mechanism	Transfer of electrons	Sharing of electrons
High (in molten/solution state)	Poor (no ions formed)
Forces	Strong electrostatic attraction	Weaker intermolecular forces

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.58; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59

5. Metallic and Non-Metallic Character (basic)

When we talk about Metallic Character, we are essentially discussing how easily an atom can lose electrons to form positive ions (cations). This property is also known as electropositivity. Conversely, Non-Metallic Character refers to an atom's tendency to gain electrons to form negative ions (anions), also known as electronegativity. In the periodic table, these characters are not distributed randomly; they follow very specific trends based on the atomic structure.

Metals are typically found on the left side of the periodic table. They tend to have 1, 2, or 3 electrons in their outermost shell, which they easily give away to achieve stability. This chemical behavior translates into physical properties like conductivity and malleability, but from a chemical standpoint, the most defining feature of metals is that they form basic oxides when they react with oxygen Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.41. For instance, when magnesium burns, it forms MgO, which reacts with water to form a base. However, some metals like Aluminium and Zinc form amphoteric oxides (like Al₂O₃ and ZnO), which can react with both acids and bases Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.41.

Feature	Metallic Character (Electropositivity)	Non-Metallic Character (Electronegativity)
Trend across a Period	Decreases (electrons held tighter by increasing nuclear charge)	Increases (stronger pull on electrons)
Trend down a Group	Increases (outer electrons are further from the nucleus)	Decreases (atomic size increases, reducing pull)
Nature of Oxide	Basic or Amphoteric Science, class X (NCERT 2025 ed.), Acids, Bases and Salts, p.22	Acidic

Non-metals are clustered on the upper right side of the table. Because they have more valence electrons and a smaller atomic radius compared to metals in the same period, they exert a stronger pull on electrons. This makes them highly effective at gaining electrons to complete their octet. The boundary between these two types is marked by "metalloids" like Silicon and Germanium, which show properties of both Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.39.

Sources: Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.39; Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.41; Science, class X (NCERT 2025 ed.), Acids, Bases and Salts, p.22

6. Ionization Potential and Electronegativity Trends (exam-level)

To master periodic trends, we must first understand the Nuclear Tug-of-War. Every atom consists of a positive nucleus pulling on negative electrons. Ionization Potential (IP) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom. Think of it as the 'cost' to kidnap an electron; the more tightly the nucleus holds it, the higher the cost. This is why ionizing radiations (like X-rays) are so potent—they possess the high penetration power necessary to provide this energy and break atomic bonds Environment, Shankar IAS Academy (ed 10th), Environmental Pollution, p.82. Electronegativity (EN), on the other hand, is a measure of an atom's 'greed'—its ability to attract a shared pair of electrons in a chemical bond.

As we move Left to Right across a Period, the number of protons in the nucleus increases (increasing the Effective Nuclear Charge), but the electrons are added to the same energy shell. This stronger positive charge pulls the electron cloud closer, reducing the atomic radius. Because the electrons are now closer to the nucleus and more strongly attracted, both Ionization Potential and Electronegativity increase. It becomes harder to remove an electron (higher IP) and easier for the atom to attract new ones (higher EN). This concept of 'potential' and 'energy' required to move charges is a fundamental principle of physics Science, Class X (NCERT 2025 ed.), Electricity, p.174.

Conversely, as we move Down a Group, the trends reverse. Each step down adds a new principal quantum shell, placing the outermost electrons significantly further from the nucleus. This distance, combined with the 'shielding effect' of inner electrons, weakens the nuclear pull. Consequently, both Ionization Potential and Electronegativity decrease. The atom becomes 'generous' with its electrons because the nucleus has a weaker grip on them. This is why elements at the bottom-left (like Cesium) are highly reactive metals, while those at the top-right (like Fluorine) are the most electronegative non-metals.

Trend Direction	Ionization Potential	Electronegativity	Reasoning
Across a Period (→)	Increases	Increases	Nuclear charge rises; radius falls.
Down a Group (↓)	Decreases	Decreases	New shells added; shielding increases.

Sources: Environment, Shankar IAS Academy (ed 10th), Environmental Pollution, p.82; Science, Class X (NCERT 2025 ed.), Electricity, p.174

7. Electron Affinity and Electron Gain Enthalpy (exam-level)

At its heart, chemistry is driven by the search for stability. Just as we seek balance, atoms strive to reach a stable octet or a noble gas configuration (Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.59). While Ionization Energy measures the effort needed to remove an electron, Electron Gain Enthalpy (Δₑ𝓰H) measures the energy change when an electron is added to an isolated gaseous atom to form a negative ion, or anion (Physical Geography by PMF IAS, Thunderstorm, p.348). If energy is released during this process, the value is negative; the more an atom 'wants' an electron, the more negative (and higher in magnitude) its electron gain enthalpy will be.

The behavior of this property follows distinct patterns across the periodic table. Across a period (left to right), the effective nuclear charge increases while the atomic radius decreases. This means the nucleus exerts a stronger pull on any incoming electron, making it easier to add one. Consequently, electron gain enthalpy becomes more negative as we move toward the halogens. For instance, Chlorine (atomic number 17) has seven valence electrons and is extremely 'eager' to gain one more to complete its outermost shell (Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.47). Conversely, noble gases like Neon or Argon already have stable configurations and require energy to force an extra electron in, resulting in positive gain enthalpy values.

Down a group, the situation changes. As the atomic weight and principal quantum number increase, new electron shells are added, making the atom larger. The nucleus is now further away from the outermost shell where the new electron would reside. Because of this increased distance and the 'shielding' effect of inner electrons, the nuclear attraction for an added electron weakens. Therefore, electron gain enthalpy generally becomes less negative (decreases in magnitude) as we move down a group.

Factor	Trend (Across Period)	Trend (Down Group)
Nuclear Charge	Increases (Stronger pull)	Increases (but distance matters more)
Atomic Radius	Decreases (Electron closer to nucleus)	Increases (Electron further from nucleus)
Δₑ𝓰H Magnitude	Becomes more negative	Becomes less negative

Sources: Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.59; Physical Geography by PMF IAS, Thunderstorm, p.348; Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.47

8. Solving the Original PYQ (exam-level)

Now that you have mastered the fundamental properties of atoms—such as effective nuclear charge and atomic radius—you can see how they act as the logic gates for this UPSC question. Periodic trends are not just facts to memorize; they are the direct results of the electromagnetic pull between the nucleus and the valence electrons. In this question, the examiner is testing your ability to predict how these forces shift as we move across a period or down a group. This question is a classic exercise in applying the principle that as the nucleus gets "stronger" (left to right) or the shells get "farther" (top to bottom), the energy required to manipulate electrons changes predictably.

Let’s walk through the logic like we’re in the exam hall. Statement II is the anchor here: as you move down a group, the atomic weight and the number of electron shells increase. Because the outermost electrons are now further from the nucleus, the atom's ability to attract an incoming electron—its electron affinity—generally weakens. This makes Statement II correct. Conversely, the UPSC often uses directional traps to catch students. In Statements I and III, the trends are described backwards. As you move left to right across a period, the atomic number increases and the nucleus exerts a stronger pull on electrons. This means both ionization potential (the energy to remove an electron) and electronegativity (the tendency to attract shared electrons) must increase, not decrease. Therefore, Statements I and III are false.

By eliminating the incorrect directions in I and III, we arrive confidently at (B) II only. The common trap here is the phrase "gradually decreases" or "decreases as atomic number increases"; always pause to visualize whether the nucleus is becoming more or less effective at grabbing electrons. If the effective nuclear charge is rising (across a period), most energy-related trends will rise with it. Use the logic of atomic tug-of-war to verify these trends rather than relying on pure rote memory. Mastering Periodic Trends Infographic, ACS