As compared to covalent compounds, electrovalent compounds, generally have

1. Atomic Stability and the Octet Rule (basic)

In the world of chemistry, stability is the ultimate goal. Just as we seek balance in our lives, atoms seek a state of minimum energy and maximum stability. This stability is defined by an atom's electronic configuration—specifically, how many electrons are buzzing around in its outermost shell, known as the valence shell. Through observations, scientists realized that a specific group of elements called the Noble Gases (like Helium, Neon, and Argon) rarely react with anything. Their secret? Their outermost shells are completely full Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47.

This observation led to the Octet Rule. Most atoms are considered stable when they have eight electrons in their valence shell. For example, Neon has an atomic number of 10, with 2 electrons in the inner (K) shell and 8 in the outer (L) shell. Elements that do not have this "magic number" are chemically reactive; they are constantly looking to gain, lose, or share electrons to mimic the configuration of the nearest noble gas Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59. There is one notable exception: Helium. Because its only shell (the K shell) can only hold 2 electrons, it is perfectly stable with a "duplet" rather than an octet Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60.

How an atom achieves this octet depends on its current count. As we see in the table below, atoms take the "path of least resistance" to reach stability:

Atom Type	Strategy	Example
Metals	Lose electrons (low valence count)	Sodium (Na) loses 1 electron to leave a stable shell of 8 underneath Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46.
Non-metals	Gain electrons (high valence count)	Chlorine (Cl) has 7 valence electrons and needs just 1 more to complete its octet Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47.
Carbon/Others	Share electrons	Carbon has 4 electrons; losing or gaining 4 is energy-intensive, so it shares instead 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.), Metals and Non-metals, p.47; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60

2. Electronegativity and Periodic Trends (basic)

At the heart of chemistry is the tendency of elements to achieve stability, usually by attaining a completely filled valence shell like that of noble gases Science, Chapter 3, p.46. Electronegativity is a measure of this "hunger" for electrons. Specifically, it is the ability of an atom in a chemical bond to attract the shared pair of electrons toward itself. Think of it as a tug-of-war: some atoms are very strong (electronegative) and pull the rope toward them, while others are weaker.

This "pull" is governed by two main factors: nuclear charge (the number of protons) and atomic radius (the distance of the valence electrons from the nucleus). As we move across a period (from left to right) in the periodic table, the number of protons increases, which increases the positive charge of the nucleus. Since the electrons are being added to the same shell, the nucleus exerts a stronger pull on them, causing the atomic radius to shrink and the electronegativity to increase. For instance, in Period 2, Fluorine is significantly more electronegative than Lithium.

Conversely, as we move down a group, new electron shells are added. This increases the distance between the nucleus and the outermost electrons, and the inner shells "shield" the outer electrons from the nuclear pull. Consequently, the atom's ability to attract shared electrons weakens, and electronegativity decreases. This explains why metals like Potassium (K) or Sodium (Na) at the bottom-left of the table are highly reactive but have very low electronegativity, often preferring to lose electrons rather than gain them Science, Chapter 3, p.45.

Understanding these trends is crucial because the difference in electronegativity between two atoms determines the nature of their bond. If the difference is very high (e.g., between a metal and a non-metal), one atom effectively "steals" the electron, leading to an ionic bond. If the difference is small, they share the electrons, forming a covalent bond Science, Chapter 4, p.60.

Direction	Electronegativity Trend	Reason
Across a Period (Left to Right)	Increases	Higher nuclear charge and smaller atomic radius.
Down a Group (Top to Bottom)	Decreases	Increased atomic size and electron shielding.

Sources: Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.45-46; Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.59-60

3. The Nature of Ionic and Covalent Bonds (basic)

At the heart of chemistry is the drive for stability. Atoms seek to achieve a noble gas configuration—a state where their outermost electron shell is full. To get there, they either trade electrons entirely or pool them together. This fundamental difference in how they reach stability dictates the physical personality of the resulting substance.

Ionic (Electrovalent) Bonding occurs when a metal transfers one or more electrons to a non-metal. This creates oppositely charged particles called ions. For example, in the formation of magnesium chloride (MgCl₂), magnesium gives up electrons to chlorine atoms Science, class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p.47. These ions don't just sit in pairs; they arrange themselves into a rigid, three-dimensional crystal lattice. Because positive and negative charges attract each other with immense electrostatic force, it takes a massive amount of thermal energy to pull them apart. This is why ionic compounds like NaCl (Table Salt) or CaO have exceptionally high melting points Science, class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p.48-49.

Covalent Bonding, on the other hand, is a story of cooperation. Instead of transferring electrons, atoms (like Carbon) share pairs of valence electrons to complete their shells Science, class X (NCERT 2025 ed.), Chapter 4: Carbon and its Compounds, p.59. While the bond within a molecule (intramolecular) is very strong, the forces between separate molecules (intermolecular) are relatively weak. Think of it like a group of strong individual bubbles: the bubble itself is tough, but the bubbles don't stick to each other very firmly. Because these intermolecular forces are weak, covalent compounds like methane (CH₄) or water generally have much lower melting and boiling points than ionic salts Science, class X (NCERT 2025 ed.), Chapter 4: Carbon and its Compounds, p.60.

Feature	Ionic Compounds	Covalent Compounds
Mechanism	Complete transfer of electrons.	Sharing of electron pairs.
Structure	Rigid crystalline lattice.	Discrete molecules.
Force	Strong electrostatic attraction.	Weak intermolecular forces.
Boiling Point	High (e.g., NaCl = 1686 K).	Relatively low.

Sources: Science , class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p.47-49; Science , class X (NCERT 2025 ed.), Chapter 4: Carbon and its Compounds, p.59-60

4. Solubility and Electrical Conductivity (intermediate)

When we explore why some substances dissolve in water or conduct electricity while others don't, we must look closely at the nature of their bonding. This behavior is fundamentally different for ionic (electrovalent) and covalent compounds. In the world of chemistry, solubility often follows the rule of "like dissolves like." Because ionic compounds are made of charged ions, they are generally soluble in polar solvents like water. The water molecules exert enough attraction to pull the ions out of their rigid lattice. However, these same compounds are usually insoluble in non-polar solvents such as kerosene or petrol Science, Class X, Metals and Non-metals, p.49.

Electrical conductivity requires the presence and movement of charged particles, such as ions or electrons. Ionic compounds are fascinating because their conductivity changes based on their physical state. In solid form, they do not conduct electricity because the ions are locked in a rigid crystal structure and cannot move Science, Class X, Metals and Non-metals, p.49. But once they are molten (melted) or dissolved in water, those strong electrostatic forces are overcome, allowing the ions to move freely toward electrodes to carry a current Science, Class X, Carbon and its Compounds, p.58.

In contrast, covalent compounds, like many carbon-based molecules, involve the sharing of electrons rather than a complete transfer. This means they do not typically form ions. Because their bonding does not give rise to any ions, most covalent compounds are poor conductors of electricity Science, Class X, Carbon and its Compounds, p.59. This is also why tools used by electricians, like screwdrivers, have plastic or rubber handles; these materials are covalent and act as insulators to prevent electric shocks Science-Class VII, The World of Metals and Non-metals, p.48.

Property	Ionic Compounds	Covalent Compounds
Solubility in Water	Generally Soluble	Generally Insoluble (exceptions exist)
Solubility in Petrol/Kerosene	Insoluble	Generally Soluble
Conductivity (Solid)	Non-conductor (Rigid)	Non-conductor
Conductivity (Solution/Molten)	Good conductor (Free ions)	Poor conductor (No ions)

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.49; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.58; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59; Science-Class VII (NCERT Revised ed 2025), The World of Metals and Non-metals, p.48

5. Allotropes of Carbon: Exceptions to the Rule (intermediate)

Carbon is a master of disguise. While it is a non-metal, its different physical forms—known as allotropes—often defy the general rules we associate with non-metallic elements. Typically, non-metals (and their covalent compounds) have relatively low melting points and are poor conductors of electricity Science, Class X, Carbon and its Compounds, p.59. However, carbon's allotropes like diamond and graphite are the "rebels" of the periodic table, showcasing properties usually reserved for metals or entirely unique materials.

For instance, while most non-metals are soft, brittle, or gaseous, diamond is the hardest natural substance known and possesses an incredibly high melting and boiling point Science, Class X, Metals and Non-metals, p.40. This extreme hardness stems from its internal architecture: each carbon atom is bonded to four other carbon atoms in a rigid, three-dimensional tetrahedral structure. Conversely, graphite challenges the rule that non-metals are electrical insulators. In graphite, carbon atoms are arranged in hexagonal layers where each atom is bonded to only three others. This configuration leaves one "free" valence electron per carbon atom, which is able to move through the structure, making graphite an excellent conductor of electricity Science, Class X, Carbon and its Compounds, p.61.

Feature	Diamond	Graphite
Structure	3D Rigid Tetrahedral	2D Hexagonal Layers
Hardness	Hardest natural substance	Soft and slippery (Lubricant)
Conductivity	Bad conductor (Insulator)	Very good conductor

Beyond these, we have fullerenes, such as C₆₀ (Buckminsterfullerene). These molecules are shaped like footballs and represent a distinct class of carbon allotropes Science, Class X, Carbon and its Compounds, p.61. Interestingly, the occurrence of these carbon forms is tied to specific geological formations. For example, in India, the Bhander and Bijwara series in Madhya Pradesh are famous for diamond mines, particularly in the Panna district, where the diamonds are prized for their brilliance and hardness Geography of India, Majid Husain, Resources, p.29.

Sources: Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59, 61; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.40; Geography of India, Majid Husain (McGrawHill 9th ed.), Resources, p.29

6. Crystal Lattice and Electrostatic Forces (exam-level)

To understand why certain materials withstand the heat of a furnace while others melt instantly, we must look at the invisible architecture of their atoms. In ionic compounds, atoms don't just 'share' electrons; they transfer them, creating cations (positive) and anions (negative). These oppositely charged particles exert a powerful electrostatic force on one another—a fundamental non-contact force where unlike charges attract Science, Class VIII, Exploring Forces, p.71. Instead of forming small, independent groups, these ions arrange themselves into a massive, repeating three-dimensional structure called a crystal lattice. This lattice is incredibly stable because every single ion is locked into place by the collective pull of all the surrounding ions of the opposite charge. Because this electrostatic attraction within the lattice is so strong, a 'considerable amount of energy' is required to overcome it Science, Class X, Metals and Non-metals, p.49. Imagine trying to pull apart two powerful magnets; now imagine millions of them interlocked in a grid. This is why ionic compounds like Sodium Chloride (NaCl) or Calcium Oxide (CaO) are solids at room temperature and possess exceptionally high melting and boiling points. In contrast, most carbon-based (covalent) compounds exist as discrete molecules. While the bonds holding the atoms together inside a molecule are strong, the forces between the molecules are quite weak Science, Class X, Carbon and its Compounds, p.59. Consequently, very little heat is needed to move these molecules apart, leading to the lower melting points we see in substances like wax or alcohol.

Feature	Ionic Compounds	Covalent (Molecular) Compounds
Nature of Force	Strong electrostatic attraction between ions.	Weak intermolecular forces (between molecules).
Physical Structure	Rigid, 3D Crystal Lattice.	Independent, discrete molecules.
Thermal Property	High melting and boiling points.	Low melting and boiling points.
State at Room Temp	Solid and generally brittle.	Liquid, gas, or soft solids.

Sources: Science, Class VIII (NCERT 2025), Exploring Forces, p.71; Science, Class X (NCERT 2025), Metals and Non-metals, p.49; Science, Class X (NCERT 2025), Carbon and its Compounds, p.59

7. Thermal Properties: Melting and Boiling Points (exam-level)

To understand why some substances melt only at extremely high temperatures while others evaporate quickly, we must look at the microscopic 'glue' holding their particles together. The melting point (the temperature at which a solid becomes a liquid) and boiling point (liquid to gas) are direct measures of the strength of these attractive forces. In chemical compounds, we see a massive divide between ionic and covalent structures regarding their thermal stability. Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.49 notes that ionic compounds generally require a tremendous amount of heat to change state, whereas carbon-based covalent compounds often have much lower thresholds.

Ionic (electrovalent) compounds, such as Sodium Chloride (NaCl), do not exist as isolated molecules. Instead, they form a rigid crystal lattice—a massive, three-dimensional grid where every positive ion is surrounded by negative ions and vice versa. These electrostatic forces of attraction are incredibly powerful. To melt such a solid, you must provide enough thermal energy to overcome these strong bonds throughout the entire lattice. This explains why ionic compounds are typically hard solids at room temperature with high melting points. Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59.

In contrast, covalent compounds (like Methane, CH₄, or Ethanol) usually consist of discrete, independent molecules. While the atoms inside a single molecule are held together by very strong covalent bonds, the forces between the neighboring molecules—known as intermolecular forces—are relatively weak. When you heat a covalent substance, you aren't breaking the atoms apart; you are simply overcoming these weak intermolecular tugs. Because these forces are easily disrupted, covalent compounds have significantly lower melting and boiling points and are frequently found as liquids or gases at room temperature. Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60.

Feature	Ionic Compounds	Covalent Compounds
Structure	Rigid 3D Crystal Lattice	Discrete individual molecules
Force to Overcome	Strong Electrostatic Bonds	Weak Intermolecular Forces
Thermal Energy Needed	Very High	Relatively Low
State at Room Temp	Solid	Often Liquid or Gas

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.49; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.60

8. Solving the Original PYQ (exam-level)

This question bridges your understanding of chemical bonding with the physical properties of matter. As you have learned, electrovalent (ionic) compounds are formed by the complete transfer of electrons, resulting in a rigid crystal lattice. According to Science, Class X (NCERT 2025 ed.), these structures are held together by powerful electrostatic forces of attraction between oppositely charged ions. In contrast, covalent compounds typically exist as discrete molecules held together by much weaker intermolecular forces, such as van der Waals forces, which require significantly less energy to overcome.

To arrive at the correct answer, you must apply the logic of thermal energy requirement. Because the ionic lattice is so stable, a massive amount of heat is needed to break these bonds to transition the substance from a solid to a liquid or a gas. This leads us directly to (D) high melting point and high boiling point. Remember: the physical state is a direct reflection of the bond strength; since ionic bonds are much stronger than the forces between covalent molecules, their resistance to heat is consistently higher across both phase changes.

UPSC often uses mixed combinations like those in options (B) and (C) as traps to see if you are guessing or if you truly understand the underlying principle. Since melting and boiling both involve overcoming the same attractive forces, it is illogical for a standard compound to have a "low" melting point but a "high" boiling point relative to another. Options (A), (B), and (C) are incorrect because they underestimate the thermal stability and the magnitude of energy required to disrupt an ionic assembly compared to a covalent one.