Three structures, labelled (A), (B) and (C), each show two glucose units joined by a glycosidic oxygen. In str

NCERT Exemplar · Q19

Q. Three structures, labelled (A), (B) and (C), each show two glucose units joined by a glycosidic oxygen. In structure (A) the two six-membered rings lie side by side and the linking oxygen runs from the anomeric carbon (C-1) of the first ring to the C-4 carbon of the second ring. In structure (C) the two rings are also joined C-1-to-C-4 through the linking oxygen. In structure (B) the second glucose unit is attached below the first through its terminal -CH2- group, so the linking oxygen runs from the anomeric carbon (C-1) of one unit to the C-6 (the CH2OH-derived carbon) of the other. Classify each linkage as C1-C4 or C1-C6. (A) (A) is between C1 and C4, (B) and (C) are between C1 and C6 (B) (A) and (B) are between C1 and C4, (C) is between C1 and C6 (C) (A) and (C) are between C1 and C4, (B) is between C1 and C6 (D) (A) and (C) are between C1 and C6, (B) is between C1 and C4 (A) (A) is between C1 and C4, (B) and (C) are between C1 and C6 (B) (A) and (B) are between C1 and C4, (C) is between C1 and C6 (C) (A) and (C) are between C1 and C4, (B) is between C1 and C6 (D) (A) and (C) are between C1 and C6, (B) is between C1 and C4

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Concept understanding — Biochemical Bonds

Biochemical Bonds: The Glue That Holds Life Together

Imagine you're building with LEGO bricks. Some bricks click together tightly and never come apart unless you really yank them. Others snap together lightly and can be pulled apart with a gentle tug. Some bricks don't even click — they just stick because of static cling or magnetism.

Biochemical bonds are exactly like that. They are the forces that hold atoms together inside the molecules of your body — your DNA, proteins, fats, and carbohydrates. Without these bonds, you'd literally fall apart into a pile of individual atoms.

The Core Idea

Atoms bond because being bonded is more stable (lower energy) than being alone. Think of it like this: a single atom is like a person standing alone in a cold room. Bonding is like huddling together for warmth — you lose some freedom of movement, but you gain stability.

In biochemistry, we care about four main types of bonds. They differ in strength, how they form, and what they do in living systems.

1. Covalent Bonds — The Strong, Permanent LEGO Clicks

This is the strongest bond in biology. Two atoms share electrons — like two people holding the same umbrella. Each atom contributes one or more electrons, and they both "own" the pair.

Key properties:

Very strong (100–400 kJ/mol)
Forms the backbone of all biomolecules
Takes a lot of energy (or enzymes) to break

Where you find it:

The carbon-carbon bonds in your DNA's sugar-phosphate backbone
The peptide bonds linking amino acids into proteins
The bonds within a glucose molecule

Note

A single covalent bond shares 2 electrons. A double bond shares 4. Triple bonds are rare in biology but exist (e.g., in cyanide).

2. Ionic Bonds — The Static Cling of Opposites

Some atoms steal electrons from others. When that happens, one atom becomes positively charged (lost an electron) and the other becomes negatively charged (gained one). Opposite charges attract — that's an ionic bond.

Key properties:

Moderate strength (5–100 kJ/mol in dry conditions)
Very weak in water (because water molecules get in between)
Easily broken by changes in pH or salt concentration

Where you find it:

In salt bridges that help proteins fold into their correct shape
Between the phosphate groups of DNA and positively charged proteins (histones)

Watch out

Ionic bonds are often called "bonds" but in water they behave more like attractions. Don't confuse them with covalent bonds — they're much weaker in biological fluids.

3. Hydrogen Bonds — The Gentle, Reversible Magnets

This is the most important weak bond in biology. A hydrogen atom that's already covalently bonded to an electronegative atom (like oxygen or nitrogen) gets a slight positive charge. It then gets attracted to another electronegative atom nearby.

Think of it like a weak magnet — it holds things together but can be easily undone.

Key properties:

Weak individually (5–30 kJ/mol)
But many together can be very strong
Easily broken by heat or changes in pH
Directional — they only work when atoms are properly aligned

Where you find it:

Between the two strands of DNA (this is what holds the double helix together)
In protein folding (between amino acids in the backbone)
Between water molecules (giving water its unique properties)

Important

Hydrogen bonds are the reason DNA can unzip for replication. If DNA used covalent bonds between strands, it would be impossible to separate without destroying the molecule.

4. Van der Waals Interactions — The Fleeting, Accidental Touches

Even neutral atoms have temporary, uneven distributions of electrons. These create tiny, momentary charges that attract nearby atoms. It's like two people accidentally brushing shoulders in a crowd — brief, weak, but real.

Key properties:

Extremely weak (0.5–5 kJ/mol per interaction)
Only work when atoms are very close (within 0.3–0.4 nm)
Add up significantly when many atoms are packed together

Where you find it:

In the hydrophobic core of proteins (where oily amino acids pack tightly)
Between lipid tails in cell membranes
In enzyme-substrate binding (helps "grip" the substrate)

Putting It All Together: A Biological Example

Consider a protein in your body. It's a long chain of amino acids held together by covalent peptide bonds. That chain then folds into a specific shape. The folding is guided by:

Hydrogen bonds between backbone atoms (forming alpha helices and beta sheets)
Ionic bonds between charged side chains …

Why this formula?

Biochemical Bonds: Why the Key Formulas Hold

Biochemical bonds are the forces that hold atoms together in biomolecules. The key formulas come from electrostatics and quantum mechanics — not from biology itself. Let's break down the why behind the most important ones.

1. Ionic Bond Energy: Coulomb's Law

Formula:

E = \frac{k \cdot q_1 \cdot q_2}{r}

Why it holds:

Opposite charges attract — this is a fundamental law of physics (Coulomb's law).
In a biochemical context, consider a sodium ion ( $\text{Na}^+$ ) and a chloride ion ( $\text{Cl}^-$ ). The energy released when they come together is directly proportional to the product of their charges ( $q_1 q_2$ ) and inversely proportional to the distance ( $r$ ) between them.
The constant $k$ accounts for the medium (water vs. vacuum). In water, the effective force is weaker because water molecules partially shield the charges — this is why ionic bonds in biology are often weaker in aqueous environments.

Key insight: The formula is not arbitrary — it's derived from the inverse-square law of electrostatics, integrated over the distance the charges move toward each other.

2. Covalent Bond Energy: The Morse Potential (Approximation)

Formula (simplified):

E = D_e \left(1 - e^{-a(r - r_0)}\right)^2

Why it holds:

Covalent bonds arise from shared electrons between atoms. The energy is not a simple inverse-square law because electrons are delocalized.
The Morse potential is an empirical formula that captures two key observations:
1. At equilibrium distance ( $r_0$ ): Energy is minimum ( $E = 0$ in this form).
2. If atoms are pulled apart ( $r \to \infty$ ): Energy approaches $D_e$ (the bond dissociation energy).
3. If atoms are pushed too close ( $r \to 0$ ): Energy skyrockets due to Pauli repulsion (electrons can't occupy the same space).
The exponential term $e^{-a(r - r_0)}$ models the rapid drop in attractive force as distance increases — this comes from quantum mechanical overlap of electron clouds.

Key insight: The formula is a curve fit to quantum mechanical calculations, not a first-principles derivation. But it works because it respects the physics: attraction at long range, repulsion at short range, and a stable minimum.

3. Hydrogen Bond Energy: Dipole-Dipole Interaction

Formula (approximate):

E \approx -\frac{2 \mu_1 \mu_2}{4 \pi \epsilon_0 r^3} \cdot \cos\theta

Why it holds:

A hydrogen bond (e.g., between water molecules) is not a true bond — it's a strong dipole-dipole interaction.
The dipole moment ( $\mu$ ) arises because oxygen is more electronegative than hydrogen, creating partial charges ( $\delta^+$ and $\delta^-$ ).
The energy depends on:
- Strength of dipoles ( $\mu_1 \mu_2$ )
- Distance ( $r$ ) — falls off as $1/r^3$ , much faster than ionic bonds ( $1/r$ )
- Orientation ( $\cos\theta$ ) — strongest when dipoles are aligned head-to-tail

Key insight: The $1/r^3$ dependence comes from the derivative of the dipole field. Unlike point charges, dipoles have a field that decays faster — this is why hydrogen bonds are directional and weaker than covalent bonds.

4. Van der Waals Interaction: Lennard-Jones Potential

Formula:

E = 4\epsilon \left[ \left(\frac{\sigma}{r}\right)^{12} - \left(\frac{\sigma}{r}\right)^6 \right]

Why it holds:

Van der Waals forces arise from temporary fluctuations in electron distribution — even nonpolar molecules have instantaneous dipoles. …

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Tell the linkages apart by which carbon of the second glucose the oxygen reaches. A bridge through the terminal -CH2- (C-6) arm is a 1->6 linkage; a bridge to the ring carbon C-4 is a 1->4 linkage. (A) and (C) are 1->4; (B) is 1->6.

Concept

In a disaccharide/polysaccharide of glucose, the glycosidic bond always starts at the anomeric carbon (C-1) of one unit. What it connects TO defines the linkage:

C1-C4 (1->4): the oxygen joins C-1 to the ring carbon C-4 of the next unit; both connection points are ring carbons roughly at the same level (as in maltose/amylose).
C1-C6 (1->6): the oxygen joins C-1 to C-6, i.e. through the exocyclic -CH2- group that projects out of the ring (as at the branch points of glycogen/amylopectin).

Applying it …

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Concept understanding — Biochemical Bonds

Biochemical Bonds: The Glue That Holds Life Together

The Core Idea

1. Covalent Bonds — The Strong, Permanent LEGO Clicks

2. Ionic Bonds — The Static Cling of Opposites

3. Hydrogen Bonds — The Gentle, Reversible Magnets

4. Van der Waals Interactions — The Fleeting, Accidental Touches

Putting It All Together: A Biological Example

Why this formula?

Biochemical Bonds: Why the Key Formulas Hold

1. Ionic Bond Energy: Coulomb's Law

2. Covalent Bond Energy: The Morse Potential (Approximation)

3. Hydrogen Bond Energy: Dipole-Dipole Interaction

4. Van der Waals Interaction: Lennard-Jones Potential

Concept

Applying it …

Method: Distinguishing C1-C4 from C1-C6 Glycosidic Linkages

Core Concept

Steps

Applying it to this question

Concept

Applying it …

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