Detailed Concept Breakdown
8 concepts, approximately 16 minutes to master.
1. Atomic Structure and the Concept of Valency (basic)
To understand how the universe is built, we must look at the atom. At its core lies a nucleus (protons and neutrons), surrounded by electrons moving in specific energy levels called shells (K, L, M, N...). For an atom to be stable, it yearns for a "full" outermost shell—a state known as the Octet Rule. Most atoms seek eight electrons in their valence shell to achieve the stable configuration of a noble gas. Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46
Valency is essentially the "combining capacity" of an atom. It is determined by the number of valence electrons (electrons in the outermost shell). Atoms achieve stability in three primary ways:
- Losing electrons: Metals like Sodium (Atomic No. 11: 2, 8, 1) lose one electron to form a positive cation (Na⁺). Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46
- Gaining electrons: Non-metals like Chlorine (Atomic No. 17: 2, 8, 7) gain one electron to complete their octet.
- Sharing electrons: Elements like Oxygen (Atomic No. 8: 2, 6) share two electrons with another atom to form a double bond. Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60
A fascinating case is Nitrogen (Atomic No. 7). With an electronic configuration of 2, 5, it requires three more electrons to reach an octet. In a molecule of N₂, each nitrogen atom contributes three electrons, creating three shared pairs known as a triple bond. Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60 This drive to fill the outer shell is the fundamental reason why chemical reactions occur.
Remember If valence electrons are 1, 2, or 3, Valency = Valence electrons. If they are 5, 6, or 7, Valency = 8 minus Valence electrons.
Key Takeaway Valency is the number of electrons an atom must gain, lose, or share to achieve a stable, full outer shell (usually an octet).
Sources:
Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46; Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60
2. Chemical Bonding: Ionic vs. Covalent (basic)
At the heart of chemistry lies the quest for stability. Most atoms are naturally unstable and seek to achieve a noble gas configuration—a state where their outermost electron shell is completely full. To reach this "happy state," atoms engage in chemical bonding by either transferring or sharing their valence electrons. The two primary ways they do this are through Ionic and Covalent bonding.
Ionic bonding typically occurs between a metal and a non-metal. In this process, the metal atom loses one or more electrons to become a positively charged ion (cation), while the non-metal gains those electrons to become a negatively charged ion (anion). The bond itself is the strong electrostatic force of attraction between these oppositely charged ions. Because these forces are so powerful, ionic compounds like Sodium Chloride (NaCl) or Calcium Oxide (CaO) generally have very high melting and boiling points and can conduct electricity when dissolved in water or melted Science, class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p. 48.
In contrast, Covalent bonding involves the sharing of electron pairs between atoms, usually two non-metals. Neither atom is strong enough to completely steal an electron from the other, so they cooperate to fill their shells. For example, a Carbon atom shares its four outer electrons with four Hydrogen atoms to form Methane (CH₄). While the bonds within the molecule are strong, the intermolecular forces (the forces between separate molecules) are relatively weak. This explains why covalent compounds often have lower melting and boiling points compared to ionic ones Science, class X (NCERT 2025 ed.), Chapter 4: Carbon and its Compounds, p. 60.
Hydrogen is a fascinating case in this spectrum. Depending on its partner, it can act like a non-metal and share electrons (forming covalent bonds in H₂O) or, when paired with very active metals, it can actually gain an electron to become a hydride ion (H⁻), forming an ionic bond in compounds like Sodium Hydride (NaH) Science, class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p. 56.
| Feature |
Ionic Bond |
Covalent Bond |
| Mechanism |
Transfer of electrons |
Sharing of electron pairs |
| Participants |
Metal + Non-metal |
Non-metal + Non-metal |
| Physical State |
Usually crystalline solids |
Gases, liquids, or soft solids |
| Conductivity |
High (in molten/solution state) |
Generally poor conductors |
Key Takeaway Ionic bonds result from the total transfer of electrons creating charged ions, while covalent bonds arise from the mutual sharing of electron pairs to achieve stability.
Sources:
Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.48, 56; Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60
3. Basics of Redox Reactions (intermediate)
In chemistry, Redox reactions (short for Reduction-Oxidation) are processes where the chemical identities of atoms change through the transfer of electrons or the movement of oxygen and hydrogen. At its simplest level, oxidation is the gain of oxygen or the loss of hydrogen, while reduction is the loss of oxygen or the gain of hydrogen Science, class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12. These two processes always occur simultaneously; if one substance is oxidised, another must be reduced.
To understand these reactions more deeply, we look at oxidation states—a book-keeping system for electrons. Hydrogen is a fascinating case study because its oxidation state is not fixed. While we often think of it as +1, it adapts to its environment:
- Elemental form: In H₂, the oxidation state is 0.
- With non-metals: In compounds like H₂O or HCl, hydrogen is less electronegative than its partner and takes an oxidation state of +1.
- With metals: When reacting with highly reactive metals (like Sodium or Calcium) to form metal hydrides (e.g., NaH), hydrogen is more electronegative than the metal and takes an oxidation state of -1 Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.56.
| Process |
Oxygen Movement |
Hydrogen Movement |
Electron Movement (Modern) |
| Oxidation |
Gain of Oxygen |
Loss of Hydrogen |
Loss of Electrons |
| Reduction |
Loss of Oxygen |
Gain of Hydrogen |
Gain of Electrons |
These principles aren't just for the lab; they shape our planet. For instance, in geomorphic weathering, oxidation of iron minerals in the presence of water and air creates the distinct red colour of certain soils. Conversely, in oxygen-poor environments like waterlogged ground, reduction occurs, changing the soil's hue to a greenish or bluish-grey Physical Geography by PMF IAS, Geomorphic Movements, p.91.
Remember OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).
Key Takeaway Redox reactions involve the simultaneous transfer of chemical species or electrons, where hydrogen's oxidation state can vary between +1, 0, and -1 depending on whether it bonds with metals or non-metals.
Sources:
Science, class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12; Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.56; Physical Geography by PMF IAS, Geomorphic Movements, p.91
4. Electronegativity and Periodic Trends (intermediate)
In the world of chemistry, Electronegativity is best understood as a "tug-of-war" for electrons. While some atoms are happy to give away electrons completely to form ionic bonds (like Sodium giving an electron to Chlorine), most chemical relationships involve sharing. Electronegativity is a measure of an atom's ability to attract that shared pair of electrons toward itself within a chemical bond. This concept is fundamental to understanding why some substances are gases, why others are solid crystals, and how molecules like water behave. Science, Class X, Chapter 3: Metals and Non-metals, p.47.
The periodic table is not just a list; it is a map of these "pulling powers." We observe two distinct Periodic Trends in electronegativity:
- Across a Period (Left to Right): Electronegativity increases. This happens because the nuclear charge (number of protons) increases, pulling electrons more strongly, while the distance of the valence shell remains similar.
- Down a Group (Top to Bottom): Electronegativity decreases. As we move down, more electron shells are added. The valence electrons are further from the nucleus, and the inner shells "shield" the outer ones from the nucleus's pull.
| Trend Direction |
Electronegativity Change |
Primary Reason |
| Left to Right (Period) |
Increases |
Higher nuclear charge (more protons) |
| Top to Bottom (Group) |
Decreases |
Increased atomic radius and shielding effect |
Understanding these trends helps us predict oxidation states. For example, Hydrogen is a unique case because its electronegativity is intermediate. When Hydrogen bonds with more electronegative non-metals (like Oxygen in H₂O), it acts as if it has lost an electron (+1 oxidation state). However, when it reacts with less electronegative metals like Sodium (Na) to form Sodium Hydride (NaH), it actually pulls the electron toward itself, acting as a hydride ion (H⁻) with a -1 oxidation state. Science, Class X, Chapter 4: Carbon and its Compounds, p.59. This chemical flexibility is why Hydrogen holds a special, often debated position in the Activity Series. Science, Class X, Chapter 3: Metals and Non-metals, p.45.
Remember FON: Fluorine, Oxygen, and Nitrogen are the "bullies" of the periodic table—they are the most electronegative elements and will almost always win the electron tug-of-war.
Key Takeaway Electronegativity determines how atoms share electrons; it increases as you move toward the top-right of the periodic table (excluding noble gases), dictating the chemical character of the element.
Sources:
Science, Class X, Metals and Non-metals, p.45, 47; Science, Class X, Carbon and its Compounds, p.59
5. Properties of Metal Hydrides vs. Non-metal Hydrides (intermediate)
Hydrogen is often called the "chameleon" of the chemical world because of its unique ability to adapt to its surroundings. Depending on which element it bonds with, hydrogen can either give up, gain, or share its single electron. When hydrogen reacts with other elements, the resulting compounds are called hydrides. The nature of these hydrides changes drastically depending on whether the partner element is a metal or a non-metal.
Metal Hydrides (also known as Ionic or Saline hydrides) are formed when hydrogen reacts with highly electropositive metals, such as the alkali metals (e.g., Sodium) or alkaline earth metals (e.g., Calcium). In these compounds, hydrogen is more electronegative than the metal. Therefore, it pulls an electron toward itself to form a hydride ion (H⁻), achieving a stable configuration like Helium. Here, the oxidation state of hydrogen is -1. These compounds are typically solids with high melting points and conduct electricity when molten, similar to the ionic compounds we see in Science, class X (NCERT 2025 ed.), Metals and Non-metals, p. 58. A classic reaction is Sodium reacting with hydrogen to form Sodium Hydride (2Na + H₂ → 2NaH).
Non-metal Hydrides (Covalent or Molecular hydrides) are formed when hydrogen bonds with non-metals like Oxygen, Nitrogen, or Carbon. Since non-metals are more electronegative than hydrogen, they do not "give" their electrons away. Instead, they share electrons through covalent bonds to fill their outer shells Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p. 59. In these cases, hydrogen usually assumes an oxidation state of +1. Common examples include water (H₂O), methane (CH₄), and ammonia (NH₃). Unlike metal hydrides, these are often gases or liquids at room temperature and do not conduct electricity because they do not form ions in their pure state.
To keep these straight, think of the "pull" on the electron. If the partner is a weak metal, Hydrogen takes the electron (H⁻). If the partner is a strong non-metal, Hydrogen effectively loses control of it (H⁺).
| Feature |
Metal Hydrides |
Non-metal Hydrides |
| Bond Type |
Ionic (Saline) |
Covalent (Molecular) |
| Oxidation State of H |
-1 |
+1 |
| Physical State |
Crystalline Solids |
Gases, Liquids, or Soft Solids |
| Example |
NaH, CaH₂ |
H₂O, NH₃, CH₄ |
Key Takeaway Hydrogen acts as an anion (H⁻) with a -1 oxidation state in metal hydrides, but typically shares electrons with a +1 oxidation state in covalent non-metal hydrides.
Sources:
Science , class X (NCERT 2025 ed.), Metals and Non-metals, p.58; Science , class X (NCERT 2025 ed.), Carbon and its Compounds, p.59
6. Rules for Assigning Oxidation Numbers (exam-level)
To master redox reactions, we must first understand
Oxidation Numbers—a formal 'bookkeeping' system where we assign a hypothetical charge to an atom to track electron movement. While some elements have fixed values, others are chameleons.
Hydrogen is the perfect example of this versatility because of its unique electronic configuration. As we've seen in our study of atomic structure, elements react to attain a stable, completely filled valence shell
Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p. 46. For hydrogen, which has only one electron in its 1s shell, stability means either losing that electron or gaining one more to look like Helium.
The oxidation state of hydrogen depends entirely on its 'partner' in a compound. When hydrogen exists in its elemental form as a diatomic molecule (H₂), its oxidation number is 0 because the electrons are shared equally between identical atoms Science, Class VIII (Revised ed 2025), Nature of Matter, p. 123. However, when it forms compounds, we follow two primary rules based on electronegativity (the 'hunger' for electrons):
- With Non-metals (+1): In most common compounds like H₂O, HCl, or CH₄, hydrogen is bonded to more electronegative non-metals. Here, hydrogen effectively 'loses' its share of electrons, resulting in an oxidation state of +1.
- With Active Metals (-1): Metals at the top of the activity series, like Sodium (Na) or Calcium (Ca), are extremely electropositive—they want to give away electrons Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p. 50. When hydrogen reacts with these metals to form metal hydrides (e.g., NaH, CaH₂), the metal forces its electron onto the hydrogen. In these cases, hydrogen acts as a hydride ion (H⁻) with an oxidation state of -1.
| Chemical Environment |
Oxidation Number |
Example |
| Elemental Form |
0 |
H₂ gas |
| Combined with Non-metals |
+1 |
H₂O, NH₃, HCl |
| Combined with Active Metals |
-1 |
NaH, MgH₂, LiH |
Remember: "Non-plus, Met-minus." Hydrogen is Plus when with a Non-metal, and Minus when with a Metal.
Key Takeaway Hydrogen does not have a single fixed oxidation state; it varies between 0, +1, and -1 depending on whether it is alone, bonded to a non-metal, or bonded to an active metal.
Sources:
Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.50; Science, Class VIII (Revised ed 2025), Nature of Matter: Elements, Compounds, and Mixtures, p.123
7. The Unique Position and Behavior of Hydrogen (exam-level)
Hydrogen occupies a truly unique position in the periodic table, often described as a 'rogue' element because it doesn't fit perfectly into any single group. With an atomic number of 1, it possesses a single electron in its innermost
K shell. To achieve the stable electronic configuration of the nearest noble gas (Helium), it requires exactly one more electron
Science, Class X, Chapter 4, p.59. This 'half-full' status allows hydrogen to exhibit a chameleon-like chemical behavior, adapting its oxidation state based on the electronegativity of its partner.
In its elemental form as a
diatomic molecule (H₂), two hydrogen atoms share their electrons equally to fill their shells, resulting in an
oxidation state of 0 Science, Class X, Chapter 4, p.59. However, when hydrogen reacts with non-metals like Oxygen or Chlorine (which are more electronegative), it behaves somewhat like an alkali metal, tending to lose its electron density. This gives it an
oxidation state of +1. For instance, in water (H₂O) or Hydrochloric acid (HCl), hydrogen exists as H⁺ ions, though in aqueous solutions, these ions are so reactive they must combine with water molecules to form
hydronium ions (H₃O⁺) Science, Class X, Chapter 2, p.23.
Conversely, when hydrogen encounters highly reactive metals such as Sodium (Na) or Calcium (Ca), the roles are reversed. Because these metals are less electronegative than hydrogen, they give up their electrons to hydrogen. In these
metal hydrides (e.g., NaH or CaH₂), hydrogen acts as a
hydride ion (H⁻) with an
oxidation state of -1. This dual ability to either lose, gain, or share electrons makes the common assumption that hydrogen 'always' has a +1 charge a significant misconception in chemistry.
| Chemical Environment | Bonding Type | Oxidation State |
|---|
| Elemental Form (H₂) | Covalent (Equal sharing) | 0 |
| With Non-metals (H₂O, HCl) | Covalent/Ionic (Polar) | +1 |
| With Active Metals (NaH, CaH₂) | Ionic (Hydride) | -1 |
Remember Hydrogen is "High-Stakes": It wants 1 electron to reach 2 (Helium). If it gives one away, it's +1; if it steals one, it's -1; if it shakes hands, it's 0.
Key Takeaway Hydrogen’s oxidation state is not fixed; it varies between 0, +1, and -1 depending on whether it is alone, paired with a more electronegative non-metal, or paired with a less electronegative metal.
Sources:
Science, Class X, Chapter 4: Carbon and its Compounds, p.59; Science, Class X, Chapter 2: Acids, Bases and Salts, p.23; Science, Class X, Chapter 1: Chemical Reactions and Equations, p.12
8. Solving the Original PYQ (exam-level)
Now that you have mastered the fundamental concepts of atomic bonding and electronegativity, this question tests your ability to apply those building blocks to the versatile nature of hydrogen. Oxidation numbers are not fixed properties; they represent the charge distribution within a molecule based on how atoms share or transfer electrons. As you have learned, hydrogen occupies a unique position in the periodic table, acting sometimes like an alkali metal and sometimes like a halogen. This chemical "dual personality" is the key to unlocking the correct answer.
Let’s walk through the logic: in its elemental form (H2), hydrogen shares electrons equally, resulting in an oxidation state of 0. When it bonds with more electronegative non-metals, such as in water (H2O), it effectively "loses" its electron to become +1. However, when it reacts with highly electropositive metals to form metal hydrides (like NaH or CaH2), it "gains" an electron, taking on a -1 state. Because it can exist as 0, +1, or -1 depending on its partner, we conclude that (D) Hydrogen can have more than one oxidation number is the correct statement.
In the context of the UPSC Civil Services Examination, this question highlights a classic examiner's trap: the use of extreme qualifiers. Options A, B, and C all use the word "always," which is a major red flag in science-based questions. Nature rarely follows absolute rules without exceptions. By identifying the specific instances—such as those detailed in Science, class X (NCERT 2025 ed.)—where hydrogen deviates from its common +1 state, you can eliminate the rigid options and choose the one that reflects scientific reality.