Is Chemical Bonding and Molecular Structure part of the EAPCET Chemistry syllabus?

Yes. Chemical Bonding and Molecular Structure is taught in Intermediate first year Chemistry, and the EAPCET Engineering syllabus is based on the Intermediate first- and second-year syllabus. Check the official notification each year for any changes.

How many Chemistry questions are there in EAPCET?

40 of the 160 questions in the Engineering paper are Chemistry, spread across 24 chapters.

What mistakes do students make in Chemical Bonding and Molecular Structure?

Counting all valence electron pairs for hybridisation including π bonds — count only sigma bonds plus lone pairs on the central atom. Calling a molecule non-polar just because its bonds are polar or vice versa — symmetric shapes like CO₂, BF₃ and CCl₄ have zero dipole moment, while NH₃ and H₂O do not. Using the O₂ orbital order for N₂ or B₂ — for molecules up to N₂, π2p lies below σ2p, which is why B₂ is paramagnetic. Naming the electron-pair geometry instead of the molecular shape — NH₃ is sp³ but its shape is trigonal pyramidal, and H₂O is bent.

EAPCET Chemistry · Intermediate first year

Chemical Bonding and Molecular Structure

Q: How should I prepare Chemical Bonding and Molecular Structure for EAPCET?

Count bond pairs and lone pairs first; the hybridisation and shape follow directly.

Chemical Bonding covers ionic and covalent bonds, VSEPR theory, hybridisation, molecular orbital theory, dipole moment and hydrogen bonding.

Subject: Chemistry
Taught in: Intermediate first year
Chemistry in the paper: 40 of 160 questions

Key formulas and results

sp: linear; sp²: trigonal planar; sp³: tetrahedral; sp³d: trigonal bipyramidal; sp³d²: octahedral
Lone pairs distort shapes: $\ce{SF4}$ is see-saw, $\ce{ClF3}$ T-shaped, $\ce{XeF4}$ square planar
Bond order $= \frac{N_b - N_a}{2}$
Molecular orbital theory predicts O₂ to be paramagnetic
Dipole moment $\mu = q \times d$
Hydrogen bonding raises boiling points (H₂O vs H₂S)
Formal charge $= (\text{valence electrons}) - (\text{non-bonding electrons}) - \tfrac{1}{2}(\text{bonding electrons})$
Hybridisation number: $H = \tfrac{1}{2}[\text{valence electrons of central atom} + \text{number of monovalent atoms} - \text{charge of cation} + \text{charge of anion}]$ ; $H = 2, 3, 4, 5, 6$ means sp, sp², sp³, sp³d, sp³d²
Percentage ionic character $= \frac{\text{observed dipole moment}}{\text{dipole moment for complete charge transfer}} \times 100$
1 debye (D) $= 1 \times 10^{-18}\ \text{esu cm} = 3.336 \times 10^{-30}\ \text{C m}$
Schomaker–Stevenson bond length: $d(\text{A--B}) = r_A + r_B - 0.09\,|\chi_A - \chi_B|$ (Å)
Lattice energy $\propto \frac{z^+ \times z^-}{r^+ + r^-}$

What this chapter covers

Ionic and covalent bonding
Fajans’ rules
VSEPR theory and molecular shapes
Hybridisation
Molecular orbital theory
Dipole moment and hydrogen bonding

How it is usually asked

Hybridisation and shape of a molecule
Bond order and magnetic nature
Comparing dipole moments
Effects of hydrogen bonding

Understanding Chemical Bonding and Molecular Structure

Atoms bond to reach a stable, lower-energy arrangement. Ionic bonds form by electron transfer between a metal and a non-metal; covalent bonds form by sharing. Fajans’ rules say a small, highly charged cation and a large anion give more covalent character, which is why LiCl is more covalent than NaCl and AlCl₃ more covalent than NaCl. Real bonds lie between the two extremes, and polar covalent bonds give molecules a dipole moment.

VSEPR theory predicts shape by counting electron pairs around the central atom: two pairs are linear, three trigonal planar, four tetrahedral, five trigonal bipyramidal and six octahedral. Lone pairs repel more than bond pairs, so they squeeze bond angles (CH₄ 109.5°, NH₃ about 107°, H₂O about 104.5°) and change the shape name. Hybridisation is the matching description: the number of hybrid orbitals equals bond pairs (sigma) plus lone pairs on the central atom.

Molecular orbital theory combines atomic orbitals into bonding and antibonding orbitals. Bond order = (Nb − Na)/2; a higher bond order means a shorter, stronger bond, and any unpaired electron makes the species paramagnetic. For O₂ and heavier, σ2p lies below π2p; for N₂ and lighter molecules, π2p lies below σ2p. EAPCET often asks for the bond order or magnetic nature of species such as O₂⁺, O₂⁻ or N₂⁺, and for shapes of molecules with lone pairs.

Worked example

EAPCET-style question

The bond order and magnetic nature of the $\ce{O2+}$ ion are, respectively:

(1)2.5, paramagnetic
(2)2.5, diamagnetic
(3)1.5, paramagnetic
(4)3, diamagnetic

Solution

$\ce{O2}$ has 16 electrons, so $\ce{O2+}$ has 15.
Fill the MOs (σ2p below π2p for oxygen): $\sigma 1s^2\ \sigma^* 1s^2\ \sigma 2s^2\ \sigma^* 2s^2\ \sigma 2p^2\ \pi 2p^4\ \pi^* 2p^1$ .
Bonding electrons $N_b = 2 + 2 + 2 + 4 = 10$ ; antibonding electrons $N_a = 2 + 2 + 1 = 5$ .
Bond order $= \frac{10 - 5}{2} = 2.5$ .
One electron in $\pi^* 2p$ is unpaired, so $\ce{O2+}$ is paramagnetic.

Answer: 2.5, paramagnetic

Shortcuts and results to remember

Bond order from the total electron count of a diatomic species: 6 → 1, 8 → 0, 10 → 1, 11 → 1.5, 12 → 2, 13 → 2.5, 14 → 3, 15 → 2.5, 16 → 2, 17 → 1.5, 18 → 1, 19 → 0.5, 20 → 0.
VSEPR repulsion order is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair; in trigonal bipyramidal molecules the axial bonds are longer and weaker than the equatorial ones.
Fajans’ rules: covalent character increases with a smaller cation, a larger anion, a higher charge on either ion, and an 18-electron (pseudo noble gas) cation such as Cu⁺ or Ag⁺.
Dipole moment exceptions: $\ce{CH3Cl} > \ce{CH3F} > \ce{CH3Br} > \ce{CH3I}$ , and $\ce{NH3}\ (1.47\ \text{D}) > \ce{NF3}\ (0.23\ \text{D})$ because the lone pair moment adds in $\ce{NH3}$ but opposes the bond moments in $\ce{NF3}$ .
Hydrogen bonding makes the first member boil highest or out of order: $\ce{HCl} < \ce{HBr} < \ce{HI} < \ce{HF}$ , $\ce{H2S} < \ce{H2Se} < \ce{H2Te} < \ce{H2O}$ , and $\ce{PH3} < \ce{AsH3} < \ce{NH3} < \ce{SbH3}$ .

Common mistakes to avoid

Counting all valence electron pairs for hybridisation including π bonds — count only sigma bonds plus lone pairs on the central atom.
Calling a molecule non-polar just because its bonds are polar or vice versa — symmetric shapes like CO₂, BF₃ and CCl₄ have zero dipole moment, while NH₃ and H₂O do not.
Using the O₂ orbital order for N₂ or B₂ — for molecules up to N₂, π2p lies below σ2p, which is why B₂ is paramagnetic.
Naming the electron-pair geometry instead of the molecular shape — NH₃ is sp³ but its shape is trigonal pyramidal, and H₂O is bent.

How to prepare it

Count bond pairs and lone pairs first; the hybridisation and shape follow directly.

Once the formulas feel familiar, test them the way the exam does: practise Chemical Bonding and Molecular Structure questions chapter-wise, then attempt a full previous year paper to see whether it holds up under a 3-hour timer.

Related Chemistry chapters

See every chapter in the EAPCET Chemistry syllabus.

Practise Chemical Bonding and Molecular Structure chapter-wise

Chapter-wise quizzes show exactly which topics in this chapter you still get wrong — before a full paper does.

Frequently asked questions

Is Chemical Bonding and Molecular Structure part of the EAPCET Chemistry syllabus?: Yes. Chemical Bonding and Molecular Structure is taught in Intermediate first year Chemistry, and the EAPCET Engineering syllabus is based on the Intermediate first- and second-year syllabus. Check the official notification each year for any changes.
How should I prepare Chemical Bonding and Molecular Structure for EAPCET?: Count bond pairs and lone pairs first; the hybridisation and shape follow directly.
How many Chemistry questions are there in EAPCET?: 40 of the 160 questions in the Engineering paper are Chemistry, spread across 24 chapters.
What mistakes do students make in Chemical Bonding and Molecular Structure?: Counting all valence electron pairs for hybridisation including π bonds — count only sigma bonds plus lone pairs on the central atom. Calling a molecule non-polar just because its bonds are polar or vice versa — symmetric shapes like CO₂, BF₃ and CCl₄ have zero dipole moment, while NH₃ and H₂O do not. Using the O₂ orbital order for N₂ or B₂ — for molecules up to N₂, π2p lies below σ2p, which is why B₂ is paramagnetic. Naming the electron-pair geometry instead of the molecular shape — NH₃ is sp³ but its shape is trigonal pyramidal, and H₂O is bent.