Q.How many hydrogen-bonded water molecule(s) are associated in CuSO4.5H2O?
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Hydrogen Bonding: From Intuition to Precision
Imagine you're holding two magnets. If you bring the north pole of one close to the south pole of another, they snap together. Now imagine a much weaker version of that — a tiny tug, not a full lock. That's the spirit of hydrogen bonding.
In chemistry, atoms in a molecule share electrons through covalent bonds. But electrons aren't shared equally in all cases. Some atoms are greedy — they pull the shared electrons closer to themselves. Oxygen, nitrogen, and fluorine are the biggest electron-hoarders. When one of these atoms bonds with hydrogen, the hydrogen ends up with a slight positive charge (because its electron has been pulled away), and the other atom gets a slight negative charge.
Now here's the key: that slightly positive hydrogen is attracted to any nearby slightly negative atom (like oxygen, nitrogen, or fluorine) on another molecule. This attraction is a hydrogen bond.
A hydrogen bond is not a true chemical bond like a covalent or ionic bond. It's an intermolecular force — a strong dipole-dipole attraction — but weaker than covalent bonds (about 1/10th to 1/20th the strength).
The Precise Definition
A hydrogen bond is an attractive interaction between a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) and another electronegative atom (N, O, or F) that has a lone pair of electrons.
We can write it as:
X—H⋯Y
where X and Y are N, O, or F. The dotted line (⋯) represents the hydrogen bond. X—H is the donor (the molecule that provides the hydrogen), and Y is the acceptor (the molecule that provides the lone pair).
Why Only N, O, and F?
Three things make these three elements special:
- High electronegativity — They pull electrons hard, creating a large partial positive charge on hydrogen.
- Small size — The lone pair on Y is compact, allowing the hydrogen to get very close. Closer distance means stronger attraction.
- Lone pairs — They have unshared electron pairs that can act as the acceptor.
Chlorine is electronegative, but it's too large — the hydrogen can't get close enough for a strong bond. Carbon is not electronegative enough.
What Makes Hydrogen Bonding Special?
Unlike other dipole-dipole interactions, hydrogen bonds are directional and stronger. They're about 5–30 kJ/mol, compared to 0.5–2 kJ/mol for ordinary van der Waals forces. This strength has dramatic consequences.
Real-World Consequences
Water's high boiling point — Water (H2O) boils at 100∘C, while hydrogen sulfide (H2S) boils at −60∘C. Both are similar molecules, but water forms hydrogen bonds; H2S does not (sulfur is not electronegative enough). Those bonds must be broken to boil water, requiring much more energy.
Ice floats — In liquid water, molecules jostle and form temporary hydrogen bonds. When water freezes, the molecules arrange into a hexagonal lattice held open by hydrogen bonds. This structure is less dense than liquid water — hence ice floats. Without hydrogen bonding, ice would sink, and lakes would freeze from the bottom up, killing aquatic life.
DNA double helix — The two strands of DNA are held together by hydrogen bonds between base pairs (adenine-thymine and guanine-cytosine). These bonds are strong enough to keep the strands together, but weak enough to be unzipped during replication. …
Why this formula?
Hydrogen Bonding: Why It Happens — The Reasoning, Not Just the Rule
Hydrogen bonding is not a full covalent bond — it's a special type of intermolecular attraction. To understand why it occurs, we must look at the electronic structure of the atoms involved.
1. The Core Requirement: A "Naked" Proton
A hydrogen bond forms when a hydrogen atom is covalently bonded to a highly electronegative atom (like F, O, or N). Why?
- Electronegativity difference pulls the bonding electron pair strongly toward the electronegative atom.
- The hydrogen atom is left with almost no electron cloud — it becomes a partially positive proton (δ+).
Key idea: The hydrogen is now a small, dense positive charge — it can get very close to a lone pair on another electronegative atom.
2. The Electrostatic Attraction (The "Why")
The partially positive hydrogen (δ+) is attracted to a lone pair of electrons on another electronegative atom (the acceptor).
This is electrostatic — Coulomb's law governs it:
F=4πε01⋅r2q1q2
- q1 = partial positive charge on H
- q2 = partial negative charge on lone pair
- r = distance between them
Because the hydrogen is so small, r is very small → force is strong (stronger than van der Waals, weaker than covalent).
3. Why Only F, O, N?
Not all electronegative atoms work. The atom must have:
| Property | Why it matters |
|---|---|
| High electronegativity | Pulls electron density away from H, creating δ+ |
| Small atomic size | Allows close approach of the H to the lone pair |
| At least one lone pair | Provides the negative site for attraction |
F, O, and N satisfy all three. Cl is electronegative but too large — the H cannot get close enough for a strong bond.
4. The "Formula" for Hydrogen Bond Strength
There is no single formula for hydrogen bond energy, but the strength depends on:
EH-bond∝r2δ+⋅δ−
Where:
- δ+ = partial charge on H (depends on electronegativity of donor atom)
- δ− = partial charge on acceptor lone pair
- r = distance between H and acceptor atom
Typical strengths (for context):
- Covalent bond: ~400 kJ/mol
- Hydrogen bond: 10–40 kJ/mol
- van der Waals: ~1–5 kJ/mol
5. Directionality — The "Linear" Preference
Hydrogen bonds are directional: the strongest interaction occurs when the donor H–X bond and the acceptor lone pair are collinear (180° angle).
Why? Because:
- The positive charge on H is concentrated along the bond axis …
Out of the five water molecules in CuSO4·5H2O, four are directly coordinated to Cu²⁺; only ONE water molecule is linked purely by hydrogen bonds (to the sulphate ion and to a coordinated water molecule). …
In blue vitriol, CuSO4⋅5H2O, the five water molecules are not all bonded the same way:
- Four of the water molecules are bonded directly to the Cu²⁺ ion through coordinate (dative) bonds, forming the square-planar complex [Cu(H2O)4]2+.
- The fifth water molecule is not coordinated to copper at all. It is held in the crystal purely by hydrogen bonds — bridging between the sulphate (SO42−) ion and one of the four coordinated water molecules of the copper complex. …
Recall the structure of blue vitriol: 4 waters coordinate Cu²⁺ as a complex ion, and the 5th water sits outside the coordination sphere, held only …
- Saying all 5 water molecules are hydrogen-bonded, or all 5 are coordinated — only 4 are coordinated and 1 is hydrogen-bonded. …
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