Free Chemical Bonding MCQs with Answers

31. The bond in a hydrogen chloride molecule is

• A. purely ionic
• B. polar covalent, with the shared pair drawn towards chlorine
• C. non polar covalent
• D. metallic

Explanation: Chlorine is more electronegative than hydrogen, so the shared electrons sit closer to it, giving chlorine a partial negative charge and hydrogen a partial positive one. The electronegativity difference is not large enough for complete transfer, so the bond is polar rather than ionic. This polarity is why hydrogen chloride dissolves readily in water and ionises there.

32. The dipole moment of a molecule of carbon tetrachloride is zero because

• A. the carbon to chlorine bonds are non polar
• B. the four polar bonds are arranged symmetrically so their dipoles cancel
• C. chlorine and carbon have the same electronegativity
• D. the molecule is planar

Explanation: Each bond carries a substantial dipole, but the tetrahedral symmetry means the four vectors sum to zero, leaving the molecule non polar overall. Replacing one chlorine with hydrogen destroys the symmetry and gives chloroform a measurable dipole moment. Shape therefore matters as much as bond polarity in deciding whether a molecule is polar.

33. Which molecule contains a bond with the greatest ionic character?

• A. HF
• B. HCl
• C. HBr
• D. HI

Explanation: Fluorine is the most electronegative element, so the difference from hydrogen is largest in hydrogen fluoride and the shared pair is most unevenly distributed. Ionic character therefore falls down the group as the halogen becomes less electronegative. This trend also explains why hydrogen fluoride is the only one of the four that hydrogen bonds strongly.

34. The hybridisation of the carbon atoms in benzene is

• A. sp3
• B. sp2
• C. sp
• D. a mixture of sp2 and sp3

Explanation: Each carbon forms three sigma bonds, two to neighbouring carbons and one to a hydrogen, using sp2 hybrids at 120 degrees, which makes the ring flat and regular. The remaining p orbital on each carbon overlaps sideways all around the ring to form the delocalised pi system. That delocalisation gives benzene its exceptional stability.

35. Metallic bonding is best described as

• A. the sharing of electrons between two adjacent atoms
• B. the attraction between a lattice of positive ions and a sea of delocalised electrons
• C. the transfer of electrons from one metal atom to another
• D. hydrogen bonding between metal atoms

Explanation: The valence electrons are free to move throughout the whole structure, which accounts for electrical and thermal conductivity, lustre and the ability of the metal to be hammered into shape without shattering. Because the bonding is non directional, layers of ions can slide without breaking the structure. Strength increases with the number of delocalised electrons per atom.

36. The number of lone pairs on the central atom in a water molecule is

• A. zero
• B. one
• C. two
• D. three

Explanation: Oxygen has six valence electrons, two of which are used in bonding to the hydrogens, leaving two lone pairs. These compress the bond angle from the tetrahedral 109.5 degrees to about 104.5 and make the molecule bent and polar. They are also what allow water to accept two hydrogen bonds.

37. Which species is linear?

• A. H2O
• B. NH3
• C. BeCl2
• D. CH4

Explanation: Beryllium in beryllium chloride has only two bond pairs and no lone pairs, so they point in opposite directions at 180 degrees. Water is bent, ammonia pyramidal and methane tetrahedral, each shaped by the number of lone pairs present. Counting electron regions and lone pairs settles every VSEPR question of this kind.

38. The strength of a covalent bond generally increases as

• A. the bond length increases
• B. the bond length decreases and the bond order rises
• C. the atoms become larger
• D. the electronegativity difference falls to zero

Explanation: Shorter bonds hold the nuclei closer to the shared electrons, so more energy is needed to separate them, and each extra shared pair shortens the bond further. This is why the carbon to carbon triple bond is both the shortest and the strongest of the three. Larger atoms form longer and therefore weaker bonds, which is why bond strength falls down a group.

39. In the ammonium ion, all four nitrogen to hydrogen bonds are identical even though one was formed as a dative bond because

• A. the dative bond breaks immediately
• B. once formed, the origin of the electron pair no longer affects the bond
• C. nitrogen changes its hybridisation
• D. one hydrogen leaves the ion

Explanation: A bond is defined by the shared pair between the two nuclei, not by which atom supplied the electrons, so experimentally the four bonds have the same length and strength. The nitrogen is sp3 hybridised and the ion is tetrahedral. This is a favourite examination point because the formation and the final structure are so easily confused.

40. The forces holding molecules of iodine together in the solid state are

• A. covalent bonds
• B. ionic bonds
• C. London dispersion forces
• D. hydrogen bonds

Explanation: Iodine molecules are non polar, so the only attraction available is the temporary induced dipole force, which nonetheless grows large enough in such a heavy molecule to make iodine a solid at room temperature. The strong covalent bond inside each molecule is not broken when the solid sublimes. This distinction between intramolecular and intermolecular forces is the point of the question.