Chemical Bonding MCQs – 40 Practice Questions with Answers and Explanations

Master ionic, covalent and metallic bonding with 40 free practice MCQs, each with a worked explanation. The page starts with GCSE and IGCSE core ideas – electron transfer, giant structures and shapes of molecules – and builds up to A-Level, IB and AP stretch topics such as hybridisation, dative bonds and resonance. Every question is original and written for QuaidianPoint students.

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Section 1: Ionic, Covalent & Metallic Bonding

Q1. What happens when an ionic bond forms between sodium and chlorine?

(A) Each sodium atom transfers an electron to a chlorine atom, forming oppositely charged ions that attract each other
(B) Sodium and chlorine atoms share a pair of electrons equally
(C) Sodium atoms transfer protons to chlorine atoms
(D) The atoms are held together by a sea of delocalised electrons

Show Answer

Answer: A – Ionic bonding involves transfer of electrons from a metal to a non-metal. Sodium forms Na⁺ and chlorine forms Cl⁻, and the electrostatic attraction between the oppositely charged ions is the ionic bond.

Q2. When a magnesium atom forms a magnesium ion, what change takes place?

(A) It loses one electron and forms a +1 ion
(B) It loses two electrons and forms a 2+ ion
(C) It gains two electrons and forms a 2− ion
(D) It loses three electrons and forms a 3+ ion

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Answer: B – Magnesium is in Group 2, with two electrons in its outer shell. Losing both gives Mg²⁺ with a full outer shell.

Q3. Why does magnesium oxide have a very high melting point?

(A) It is made of molecules held together by weak forces
(B) It contains delocalised electrons that need a lot of energy to stop moving
(C) A large amount of energy is needed to break the strong electrostatic attractions between oppositely charged ions in the giant lattice
(D) Its ions are held together by hydrogen bonds

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Answer: C – In a giant ionic lattice every ion is strongly attracted to many oppositely charged neighbours. Breaking enough of these attractions to melt the solid takes a great deal of thermal energy.

Q4. Sodium chloride conducts electricity when molten but not when solid. Why?

(A) The solid contains no charged particles at all
(B) Electrons are free to move in the solid but not in the liquid
(C) Chloride ions lose their charge when the solid melts
(D) In the solid the ions are locked in fixed positions and cannot move; when molten the ions are free to move to the electrodes and carry the charge

Show Answer

Answer: D – Ionic compounds conduct only when their ions are mobile, that is, when molten or dissolved in water. In the solid lattice the ions cannot move, so no charge can be carried.

Q5. Methane, CH₄, has a very low boiling point. What is the best explanation?

(A) Only weak intermolecular forces between the molecules have to be overcome when it boils; the strong covalent bonds inside each molecule stay intact
(B) The covalent bonds between carbon and hydrogen are very weak
(C) Methane molecules carry full positive and negative charges
(D) Methane has no electrons, so there is nothing holding it together

Show Answer

Answer: A – Simple molecular substances have low boiling points because boiling only separates the molecules from each other. The covalent bonds within each molecule are not broken.

Q6. In diamond, each carbon atom forms covalent bonds with how many other carbon atoms?

(A) Two
(B) Four
(C) Three
(D) One

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Answer: B – Each carbon in diamond is bonded to four others in a giant tetrahedral network. This is why diamond is so hard and has a very high melting point.

Q7. Which statement best describes metallic bonding?

(A) Pairs of electrons shared between neighbouring atoms
(B) Electrons transferred from metal atoms to non-metal atoms
(C) A regular lattice of positive metal ions held together by attraction to a sea of delocalised electrons
(D) Weak attractions between separate molecules

Show Answer

Answer: C – In a metal, the outer electrons are delocalised and move freely through the lattice of positive ions. This attraction is the metallic bond, and the mobile electrons explain electrical conductivity.

Q8. Why can metals be bent and shaped (malleable) without breaking?

(A) Metallic bonds are weak, so the atoms fall apart easily
(B) The atoms in a metal are very small
(C) Positive metal ions repel each other, so layers push apart
(D) Layers of positive ions can slide over one another while the delocalised electrons continue to hold the structure together

Show Answer

Answer: D – Because the bonding electrons are delocalised rather than fixed between specific atoms, the ion layers can shift position without the bonding breaking.

Q9. Alloys are usually harder than the pure metals they are made from. Why?

(A) The different-sized atoms distort the regular layers of ions, so the layers cannot slide over each other as easily
(B) Alloys contain covalent bonds, which are always stronger
(C) Alloy atoms have more protons than pure metal atoms
(D) Alloys contain no delocalised electrons

Show Answer

Answer: A – In a pure metal the regular layers slide easily. Adding atoms of a different size disrupts the layers and blocks the sliding, making the alloy harder and less malleable.

Q10. Which substance has a giant covalent structure?

(A) Carbon dioxide, CO₂
(B) Silicon dioxide, SiO₂
(C) Water, H₂O
(D) Methane, CH₄

Show Answer

Answer: B – Silicon dioxide is a giant covalent lattice in which every silicon is bonded to four oxygens and every oxygen to two silicons. CO₂, H₂O and CH₄ are simple molecules.

Section 2: Shapes of Molecules & Intermolecular Forces

Q11. What are the shape and bond angle in a methane molecule, CH₄?

(A) Trigonal planar, 120°
(B) Linear, 180°
(C) Tetrahedral, 109.5°
(D) Trigonal pyramidal, 107°

Show Answer

Answer: C – Carbon has four bonding pairs and no lone pairs in methane. They repel to the corners of a tetrahedron with H-C-H angles of 109.5°.

Q12. What are the shape and bond angle in an ammonia molecule, NH₃?

(A) Tetrahedral, 109.5°
(B) Trigonal planar, 120°
(C) Bent, 104.5°
(D) Trigonal pyramidal, 107°

Show Answer

Answer: D – Nitrogen has three bonding pairs and one lone pair. The lone pair pushes the N-H bonds closer together, giving a trigonal pyramidal shape with angles of about 107°.

Q13. What are the shape and bond angle in a water molecule, H₂O?

(A) Bent, 104.5°
(B) Bent, 109.5°
(C) Linear, 180°
(D) Trigonal pyramidal, 107°

Show Answer

Answer: A – Oxygen has two bonding pairs and two lone pairs. The two lone pairs repel strongly, squeezing the O-H bonds to an angle of 104.5° in a bent (non-linear) shape.

Q14. Why is the bond angle in ammonia (107°) smaller than the bond angle in methane (109.5°)?

(A) Nitrogen is more electronegative than carbon, so it pulls the bonds in
(B) The lone pair on nitrogen repels the bonding pairs more strongly than bonding pairs repel each other, pushing the N-H bonds closer together
(C) Hydrogen atoms in ammonia attract each other
(D) Ammonia has fewer atoms than methane, so there is more space

Show Answer

Answer: B – Lone pair-bonding pair repulsion is stronger than bonding pair-bonding pair repulsion. The lone pair in ammonia compresses the H-N-H angle from the tetrahedral 109.5° down to 107°.

Q15. Boron trifluoride, BF₃, has three bonding pairs and no lone pair on boron. What is its shape and bond angle?

(A) Trigonal pyramidal, 107°
(B) Tetrahedral, 109.5°
(C) Trigonal planar, 120°
(D) Linear, 180°

Show Answer

Answer: C – Three electron pairs around the central atom repel to the corners of a flat triangle, giving a trigonal planar molecule with F-B-F angles of 120°.

Q16. Carbon dioxide, CO₂, is a linear molecule with a bond angle of 180°. Why?

(A) Oxygen atoms attract each other across the molecule
(B) Carbon has two lone pairs that balance each other
(C) The molecule contains ionic bonds, which are always straight
(D) The two double bonds are regions of electron density that repel each other as far apart as possible, on opposite sides of the carbon atom

Show Answer

Answer: D – Each C=O double bond counts as one region of electron density. Two regions repel to opposite sides of the carbon, giving a linear molecule.

Q17. The boiling points of the halogens increase down the group: Cl₂ −34°C, Br₂ 59°C, I₂ 184°C. Why?

(A) The molecules become larger with more electrons, so the London (van der Waals) forces between molecules become stronger
(B) The covalent bonds inside the molecules become stronger down the group
(C) Hydrogen bonding appears between the molecules of the heavier halogens
(D) The molecules become ionic further down the group

Show Answer

Answer: A – Boiling a halogen only overcomes the forces between molecules, not the covalent bonds within them. Bigger molecules have more electrons, stronger temporary dipoles and so stronger London forces.

Q18. Water has a much higher boiling point than hydrogen sulfide, H₂S, even though H₂S has the larger relative molecular mass. Why?

(A) The covalent O-H bonds are much stronger than S-H bonds, and they must break on boiling
(B) Hydrogen bonds form between water molecules, but not between H₂S molecules, so extra energy is needed to separate water molecules
(C) Water molecules are heavier than H₂S molecules
(D) Sulfur is a metal, so H₂S is held by metallic bonds

Show Answer

Answer: B – Oxygen is small and highly electronegative, so water molecules hydrogen-bond to each other. Sulfur is too large and not electronegative enough, so H₂S has only weaker intermolecular forces and boils far lower.

Q19. Which substance has hydrogen bonding between its molecules?

(A) Methane, CH₄
(B) Hydrogen sulfide, H₂S
(C) Ammonia, NH₃
(D) Hydrogen chloride, HCl

Show Answer

Answer: C – Hydrogen bonding needs hydrogen attached directly to a small, highly electronegative atom (N, O or F) and a lone pair on such an atom. Ammonia qualifies; CH₄, H₂S and HCl do not.

Q20. Ethanol, C₂H₅OH, boils at 78°C. Its isomer dimethyl ether, CH₃OCH₃, has the same relative molecular mass but boils at −24°C. Why?

(A) Ethanol contains stronger covalent bonds than the ether
(B) The ether molecules form hydrogen bonds but ethanol molecules cannot
(C) Ethanol molecules are much larger than ether molecules
(D) Ethanol molecules hydrogen-bond to each other through their -OH groups; the ether has no O-H bond, so its molecules are held only by weaker forces

Show Answer

Answer: D – With the same M_r, London forces are similar in both isomers. The difference is hydrogen bonding, present in ethanol because of its O-H group and absent in dimethyl ether.

Section 3: Electronegativity, Polarity & Bond Properties

Q21. Which element has the highest electronegativity?

(A) Fluorine
(B) Chlorine
(C) Oxygen
(D) Nitrogen

Show Answer

Answer: A – Fluorine, at the top right of the Periodic Table, has the greatest power to attract the electrons in a covalent bond (Pauling value about 4.0).

Q22. Electronegativity increases across a period. Why?

(A) The shielding by inner electrons increases much faster than the nuclear charge
(B) The nuclear charge increases while shielding stays similar, so the attraction for bonding electrons becomes stronger
(C) Atoms become larger across a period, so their electrons spread out
(D) An extra electron shell is added to each atom

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Answer: B – Across a period, protons are added but electrons go into the same shell, so shielding barely changes. The increasing effective nuclear charge pulls bonding electrons in more strongly.

Q23. Chlorine is more electronegative than hydrogen. What does this mean for the bond in hydrogen chloride, HCl?

(A) The bond is fully ionic, with complete electron transfer
(B) The shared pair of electrons is shared exactly equally
(C) The shared pair is pulled closer to chlorine, giving chlorine a partial negative charge (δ−) and hydrogen a partial positive charge (δ+)
(D) Hydrogen gains an electron from chlorine

Show Answer

Answer: C – Unequal sharing in a covalent bond creates a polar bond with partial charges. It is not complete transfer, so HCl is a polar covalent molecule, not an ionic compound.

Q24. The C=O bonds in carbon dioxide are polar, yet the CO₂ molecule is non-polar overall. Why?

(A) The C=O bonds are not really polar
(B) Oxygen has no lone pairs in CO₂
(C) Carbon dioxide is an ionic compound
(D) The molecule is linear and symmetrical, so the two equal bond dipoles point in opposite directions and cancel out

Show Answer

Answer: D – Whether a molecule is polar depends on its shape as well as its bonds. In linear CO₂ the two dipoles are equal and opposite, so they cancel.

Q25. Unlike carbon dioxide, water is a polar molecule. Why?

(A) Water is bent, so its bond dipoles point in similar directions and do not cancel
(B) Oxygen forms ions when bonded to hydrogen
(C) Hydrogen bonds hold each individual water molecule together
(D) The O-H bonds in water are ionic

Show Answer

Answer: A – Water’s bent shape means the two O-H dipoles add together instead of cancelling, leaving the oxygen end δ− and the hydrogen end δ+. (Hydrogen bonds act between molecules, not within one molecule.)

Q26. Tetrachloromethane, CCl₄, contains four polar C-Cl bonds, but the molecule is non-polar. Why?

(A) The C-Cl bonds share their electrons equally
(B) The molecule is tetrahedral and symmetrical, so the four bond dipoles cancel out
(C) Chlorine is less electronegative than carbon
(D) The molecule is flat, so the dipoles balance

Show Answer

Answer: B – In a perfectly symmetrical tetrahedral molecule the four equal dipoles point to the corners and cancel, so CCl₄ has no overall dipole.

Q27. Which statement about bond length and bond strength is correct?

(A) A triple bond is longer than a single bond between the same two atoms
(B) Single bonds are the strongest type of covalent bond
(C) A double bond is shorter and stronger than a single bond between the same two atoms
(D) Bond length has no connection with how many electron pairs are shared

Show Answer

Answer: C – More shared electron pairs pull the nuclei closer together and hold them more strongly, so bond enthalpy increases and bond length decreases in the order single, double, triple.

Q28. Which compound has the most ionic character in its bonding?

(A) CH₄ (electronegativities: C 2.5, H 2.1)
(B) HCl (H 2.1, Cl 3.0)
(C) NaCl (Na 0.9, Cl 3.0)
(D) CsF (Cs 0.7, F 4.0)

Show Answer

Answer: D – The larger the difference in electronegativity, the more ionic the bond. CsF has the biggest difference (4.0 − 0.7 = 3.3), so its bonding is the most ionic.

Q29. Graphite conducts electricity, but diamond does not. Why?

(A) In graphite each carbon bonds to three others, leaving one electron per atom delocalised along the layers; in diamond all four outer electrons are used in bonds
(B) Graphite contains metal impurities that carry the current
(C) The layers in graphite are charged ions
(D) Diamond has no electrons

Show Answer

Answer: A – Graphite’s spare electrons are delocalised within each layer and can move and carry charge. In diamond every outer electron is locked into a covalent bond, so nothing can move.

Q30. Silicon dioxide, SiO₂, melts at about 1710°C. Why is its melting point so high?

(A) It has unusually strong intermolecular forces between its molecules
(B) It is a giant covalent structure, so a huge number of strong covalent bonds must be broken to melt it
(C) It is an ionic lattice with very small ions
(D) Its molecules are joined by hydrogen bonds

Show Answer

Answer: B – SiO₂ has no separate molecules. Melting it means breaking strong Si-O covalent bonds throughout the giant lattice, which needs a very large amount of energy.

Section 4: Advanced Bonding (A-Level / IB / AP Stretch)

Q31. What is the hybridisation of the carbon atom in methane, CH₄?

(A) sp
(B) sp²
(C) sp³
(D) No hybridisation takes place

Show Answer

Answer: C – Carbon in methane forms four identical sigma bonds pointing tetrahedrally, which come from mixing one s and three p orbitals to give four sp³ hybrid orbitals.

Q32. How many pi (π) bonds are there in one molecule of ethene, C₂H₄?

(A) None
(B) Two
(C) Five
(D) One

Show Answer

Answer: D – The C=C double bond consists of one sigma bond and one pi bond. All four C-H bonds are sigma bonds, so ethene has exactly one pi bond.

Q33. How many sigma (σ) and pi (π) bonds are there in one molecule of ethyne, HC≡CH?

(A) 3 sigma and 2 pi
(B) 2 sigma and 3 pi
(C) 5 sigma and 0 pi
(D) 3 sigma and 1 pi

Show Answer

Answer: A – Ethyne has three sigma bonds (two C-H and one C-C). The triple bond is made of that one C-C sigma bond plus two pi bonds formed by sideways overlap of p orbitals.

Q34. An ammonium ion, NH₄⁺, forms when ammonia reacts with a hydrogen ion. Which statement describes the new N-H bond?

(A) The hydrogen ion donates a pair of electrons to nitrogen
(B) Both electrons in the bond come from the lone pair on nitrogen; it is a dative (coordinate) bond
(C) Nitrogen and hydrogen each supply one electron, as in an ordinary covalent bond
(D) An electron is transferred from nitrogen to hydrogen, forming ions

Show Answer

Answer: B – H⁺ has no electrons, so both electrons in the fourth N-H bond must come from nitrogen’s lone pair. Once formed, this dative bond is identical to the other three N-H bonds.

Q35. Which species contains a dative (coordinate) covalent bond?

(A) Na⁺
(B) Mg²⁺
(C) H₃O⁺
(D) Cl⁻

Show Answer

Answer: C – In the hydroxonium ion, H₃O⁺, a lone pair on the oxygen of a water molecule is donated to an H⁺ ion, forming a coordinate bond. The other species are simple ions with no such bond.

Q36. All six carbon-carbon bonds in benzene are the same length, intermediate between a single bond and a double bond. Why?

(A) The bonds quickly swap between single and double positions
(B) Benzene contains three single bonds and three double bonds of equal length by coincidence
(C) Each carbon is joined by one and a half bonds because the atoms share protons
(D) The pi electrons are delocalised in a ring above and below the plane of the molecule, spreading the bonding evenly over all six C-C bonds

Show Answer

Answer: D – Benzene does not contain alternating single and double bonds. Its six pi electrons are delocalised around the ring, so every C-C bond is equivalent with a bond order of about 1.5.

Q37. In the carbonate ion, CO₃²⁻, all three carbon-oxygen bonds are equal in length. What is the average bond order of each C-O bond?

(A) About 1.33, because the pi bonding is delocalised (resonance) over all three C-O bonds
(B) Exactly 1, because all the bonds are single
(C) Exactly 2, because all the bonds are double
(D) 3, because there are three oxygen atoms

Show Answer

Answer: A – One C=O pi bond is shared out over three positions by resonance, so each C-O bond has bond order 1 + 1/3, about 1.33, and all three bonds are the same length.

Q38. Which molecule contains a central atom with an expanded octet (more than eight electrons in its outer shell)?

(A) BF₃
(B) SF₆
(C) NH₃
(D) CH₄

Show Answer

Answer: B – Sulfur in SF₆ forms six bonds, so 12 electrons surround it – an expanded octet, possible for elements in Period 3 and below. Boron in BF₃ has only six electrons (an incomplete octet).

Q39. What is the formal charge on the nitrogen atom in the ammonium ion, NH₄⁺? (Formal charge = valence electrons − non-bonding electrons − half the bonding electrons; nitrogen has 5 valence electrons.)

(A) 0
(B) −1
(C) +1
(D) +5

Show Answer

Answer: C – Nitrogen in NH₄⁺ has no lone pair and four bonds (8 bonding electrons). Formal charge = 5 − 0 − (8 / 2) = 5 − 4 = +1, which matches the charge on the whole ion.

Q40. What is the hybridisation of the carbon atom in carbon dioxide, CO₂?

(A) sp³
(B) sp²
(C) Carbon does not use hybrid orbitals in CO₂
(D) sp

Show Answer

Answer: D – Carbon in CO₂ has two regions of electron density (two double bonds) arranged linearly at 180°, which corresponds to sp hybridisation; the two unhybridised p orbitals form the two pi bonds.

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Chemical Bonding MCQs – 40 Practice Questions with Answers and Explanations

Section 1: Ionic, Covalent & Metallic Bonding

Section 2: Shapes of Molecules & Intermolecular Forces

Section 3: Electronegativity, Polarity & Bond Properties

Section 4: Advanced Bonding (A-Level / IB / AP Stretch)

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