Statement I : Water (H2O) is more polar than Hydrogen Sulphide . Statement II: Oxygen is more electro-negative than Sulphur.

1. Periodic Table Trends: Electronegativity (basic)

Electronegativity is a fundamental concept in chemistry that describes an atom's ability to attract and 'hog' shared electrons within a chemical bond. Think of it as a tug-of-war for electrons: the more electronegative an atom is, the more strongly it pulls electrons toward itself. This tendency is deeply rooted in an atom's desire to achieve a stable, completely filled valence shell, similar to the noble gases Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46. When two atoms with different electronegativity levels form a bond, the electron pair isn't shared equally; it shifts toward the stronger player, creating a partial negative charge (δ-) on the electronegative atom and a partial positive charge (δ+) on the other.

Across the Periodic Table, electronegativity follows very predictable trends. As you move from left to right across a period, electronegativity increases because the number of protons (nuclear charge) increases, pulling electrons more tightly. Conversely, as you move down a group, electronegativity decreases. This is because additional electron shells are added, increasing the distance between the nucleus and the bonding electrons, which weakens the 'tug.' For example, Oxygen (atomic number 8) is significantly more electronegative than Sulfur (atomic number 16) because Sulfur's valence electrons are further from the nucleus and more shielded.

This difference in 'pulling power' is what determines the polarity of a molecule. A classic example is comparing water (H₂O) and hydrogen sulfide (H₂S). Oxygen has a high electronegativity (about 3.5), while Sulfur is lower (about 2.5). Because the electronegativity difference between Oxygen and Hydrogen is much larger than that between Sulfur and Hydrogen, the O-H bonds in water are much more polar. This results in water being a highly polar molecule with a strong net dipole moment, which influences everything from its boiling point to how it acts as a solvent Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.59.

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

2. Chemical Bonding: Polar vs. Non-polar Covalent Bonds (basic)

In our previous step, we looked at how atoms bond to find stability. Now, let’s look closer at Covalent Bonding. At its simplest, a covalent bond is formed by the sharing of an electron pair between two atoms Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60. However, in the world of chemistry, sharing is not always equal. Think of it like a tug-of-war where one person is significantly stronger than the other.

Whether a bond is Polar or Non-polar depends on a property called Electronegativity—the "hunger" an atom has for electrons. When two identical atoms bond, like in a molecule of Nitrogen (N₂) or Hydrogen (H₂), they pull on the shared electrons with equal strength. This results in a Non-polar Covalent Bond because the electron cloud is distributed perfectly evenly Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60.

However, when atoms of different elements bond, one often pulls harder. In a water molecule (H₂O), the Oxygen atom is much more electronegative than the Hydrogen atoms. It pulls the shared electrons closer to itself. This creates a Polar Covalent Bond, where the Oxygen side becomes slightly negative (δ-) and the Hydrogen side becomes slightly positive (δ+). This separation of charges is what we call a dipole. Because these molecules don't fully transfer electrons (unlike ionic compounds), they don't form full ions, which is why they are generally poor conductors of electricity Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59.

Feature	Non-polar Covalent Bond	Polar Covalent Bond
Electron Sharing	Equal sharing between atoms.	Unequal sharing; electrons shift toward one atom.
Electronegativity Difference	Zero or very low.	Significant difference between atoms.
Example	H₂, O₂, CH₄ (C-H bonds)	H₂O, NH₃, HCl

Sources: Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60

3. Molecular Geometry and VSEPR Theory (intermediate)

To understand why molecules take specific shapes, we must look at the Valence Shell Electron Pair Repulsion (VSEPR) Theory. At its core, this theory is based on a simple physical principle: electrons are negatively charged, and because like-charges repel, electron pairs surrounding a central atom will stay as far away from each other as possible. While we often draw molecules using electron dot structures on a flat page—as seen with Carbon Dioxide (CO₂) or Sulphur (S₈) in Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.61—these molecules actually exist in three-dimensional space to minimize this internal tension. There are two types of electron pairs to consider: Bond Pairs (shared between atoms) and Lone Pairs (unshared electrons). A crucial nuance of VSEPR is that lone pairs are more repulsive than bond pairs. Because a lone pair is attracted to only one nucleus, it spreads out more than a bond pair, which is pinned down between two nuclei. This 'extra space' taken up by lone pairs pushes the adjacent chemical bonds closer together, causing the bond angles to deviate from ideal geometric shapes. For example, while four electron pairs usually form a perfect Tetrahedral shape (109.5°), the presence of lone pairs in a water molecule (H₂O) 'squashes' the bonds, resulting in a Bent geometry with a smaller angle. In the UPSC context, understanding these shapes is vital because molecular geometry determines polarity. As noted in Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60, covalent molecules involve the sharing of electrons to achieve a noble gas configuration. However, if the geometry is asymmetrical (like the bent shape of H₂O), the electrical charges don't cancel out, making the molecule polar. Conversely, in a linear molecule like CO₂, even though the individual bonds are polar, their 180° alignment causes the 'pulls' to cancel each other out, resulting in a non-polar molecule. This relationship between 3D shape and physical properties explains everything from the boiling point of water to how enzymes recognize medicines in our body.

Sources: Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60; Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.61

4. Intermolecular Forces: Hydrogen Bonding (intermediate)

To understand Hydrogen Bonding, we must first distinguish between the forces inside a molecule and the forces between molecules. As we've seen, covalent bonds involve the sharing of electrons to achieve a stable configuration Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59. While these internal bonds are very strong, the intermolecular forces (the "glue" holding separate molecules together) are generally much weaker, which explains why many covalent compounds have low melting and boiling points Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60.

Hydrogen bonding is a special, stronger-than-average type of intermolecular force. It occurs when a Hydrogen atom is covalently bonded to a highly electronegative atom—specifically Nitrogen (N), Oxygen (O), or Fluorine (F). Because these atoms are "electron-hungry," they pull the shared electron pair toward themselves. This creates a dipole: the Hydrogen atom develops a partial positive charge (δ+), and the heteroatom (like Oxygen) develops a partial negative charge (δ-) Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.66. The electrostatic attraction between the δ+ Hydrogen of one molecule and the δ- atom of a neighboring molecule is the Hydrogen bond.

This force explains why water (H₂O) behaves so differently from similar molecules. Even though Oxygen and Sulfur are in the same chemical family, Oxygen is significantly more electronegative (approx. 3.5) than Sulfur (approx. 2.5). This makes the O-H bond much more polar than the S-H bond. Consequently, H₂O molecules stick together firmly via hydrogen bonds, remaining liquid at room temperature, while H₂S molecules lack this strong attraction and exist as a gas Science, Class VIII, NCERT (Revised ed 2025), Nature of Matter, p.123.

Feature	Van der Waals Forces	Hydrogen Bonding
Strength	Relatively Weak	Stronger (subset of dipole-dipole)
Requirements	Present in all molecules	H bonded to N, O, or F
Effect on Boiling Point	Lower	Significantly Higher

Sources: Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.59, 60, 66; Science, Class VIII, NCERT (Revised ed 2025), Nature of Matter: Elements, Compounds, and Mixtures, p.123

5. Physical Properties of Group 16 Hydrides (intermediate)

When we look at Group 16 of the periodic table (Oxygen, Sulfur, Selenium, Tellurium, and Polonium), we find that they all react with hydrogen to form hydrides with the general formula H₂E. While these molecules share a similar structural blueprint—specifically a bent molecular geometry caused by two lone pairs of electrons—their physical properties are dramatically different. For instance, while H₂O (water) is a liquid at room temperature and essential for life, H₂S (hydrogen sulfide) is a pungent gas often associated with industrial pollution or volcanic activity. Physical Geography by PMF IAS, Earths Atmosphere, p.270

The primary reason for these differences lies in molecular polarity. Polarity is determined by the electronegativity of the central atom. Oxygen is highly electronegative (approx. 3.5), meaning it has a very strong pull on the shared electrons in the O-H bond. Sulfur is much less electronegative (approx. 2.5). Because the difference in electronegativity between Oxygen and Hydrogen (ΔEN ≈ 1.4) is far greater than between Sulfur and Hydrogen (ΔEN ≈ 0.4), the O-H bond is significantly more polar. In a bent molecule like water, these strong individual bond dipoles do not cancel out; instead, they combine to create a large net dipole moment, making water a highly polar solvent.

This high polarity in water leads to strong intermolecular attractions, specifically hydrogen bonding, which explains why water has a much higher boiling point compared to other hydrides in the group. As we move down the group from Oxygen to Polonium, the electronegativity of the central atom decreases, causing the polarity of the H-E bond to diminish. This is a classic example of how fundamental atomic properties like electronegativity dictate the macroscopic physical behavior of substances we encounter in the environment. Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.56

Property	Water (H₂O)	Hydrogen Sulfide (H₂S)
Electronegativity of Central Atom	High (~3.5)	Moderate (~2.5)
Bond Polarity	Very High	Low
Physical State (Room Temp)	Liquid	Gas

Sources: Physical Geography by PMF IAS, Earths Atmosphere, p.270; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.56

6. Net Dipole Moment and Molecular Polarity (exam-level)

To understand why some substances behave differently than others, we must look at Molecular Polarity. This is determined by two critical factors: the difference in electronegativity between atoms and the spatial geometry of the molecule. Electronegativity is essentially the 'greediness' of an atom for electrons. When two different atoms form a bond, the more electronegative atom pulls the shared electrons closer to itself, creating a bond dipole (a tiny separation of charge). As we see in the bonding structure of water Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60, oxygen and hydrogen share electrons, but because oxygen is significantly more electronegative, the O-H bonds are highly polar. However, a molecule is only 'polar' overall if it has a Net Dipole Moment. Think of this as a tug-of-war in three dimensions. The net dipole moment is the vector sum of all individual bond dipoles. If a molecule is perfectly symmetrical, the 'pulls' cancel each other out, leaving the molecule non-polar. But if the molecule is asymmetrical—like the 'bent' shape of a water molecule—the dipoles do not cancel out, resulting in a strong net dipole. This net polarity is what allows water to dissolve many substances and gives it a high boiling point compared to molecules of similar mass Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.67.

To illustrate how electronegativity and geometry work together, consider this comparison:

Molecule	Electronegativity Difference	Geometry	Net Dipole Moment
H₂O (Water)	High (O is very electronegative)	Bent	High (Polar)
H₂S (Hydrogen Sulphide)	Low (S is less electronegative)	Bent	Low (Weakly Polar)
CO₂ (Carbon Dioxide)	Moderate	Linear	Zero (Non-polar)

Sources: Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60; Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.67

7. Solving the Original PYQ (exam-level)

You have just mastered the building blocks of electronegativity and molecular geometry, and this PYQ is the perfect application of those concepts. To determine a molecule's polarity, we must evaluate how unevenly electrons are shared between atoms. In General Science (Chemistry), we learn that the "pull" an atom exerts on shared electrons creates a dipole moment. This question tests your ability to link a specific atomic property (electronegativity) to a resultant molecular characteristic (polarity). It bridges the gap between atomic structure and chemical behavior.

Let's walk through the reasoning as a coach: First, identify the periodic trend. Oxygen sits above Sulphur in Group 16, meaning Oxygen is more electronegative and holds onto electrons more tightly. Second, consider the bond. This higher electronegativity creates a much stronger partial negative charge on Oxygen compared to Sulphur. Third, look at the shape. Both H2O and H2S have a bent geometry due to lone pairs, so their bond polarities do not cancel out. Because the individual O-H bonds are more polar than S-H bonds, the net molecular polarity of water is significantly higher. Thus, Statement II is the direct scientific cause of Statement I, making (A) the correct answer.

UPSC frequently uses (B) as a trap by providing two factually correct statements that lack a causal link. For instance, if Statement II mentioned water’s high boiling point, it would be true, but it wouldn't explain why water is more polar. Options (C) and (D) are designed to catch students who are shaky on Periodic Table trends. Always ask yourself: "Does the second statement provide the 'Why' for the first?" In this case, the electronegativity difference is the fundamental reason the dipole moment is larger, confirming the causal relationship.