9th Class Chemistry Chapter # 3 Exercise Solutions – Punjab Board
Chemical Bonding
Multiple Choice Questions (MCQs)
1. When molten copper and molten zinc are mixed together, they give rise to a new substance called brass. Predict what type of bond is formed between copper and zinc.
Options:
(a) Coordinate covalent bond
(b) Covalent bond
✅ (c) Metallic bond
(d) Ionic bond
✅ Correct Answer: (c) Metallic bond
Explanation:
Brass is formed by mixing two metals — copper and zinc. In metallic bonding, metal atoms share a ‘sea of delocalized electrons’, giving rise to properties like conductivity and luster. Hence, bonding in brass is metallic in nature.
2. Which element is capable of forming all the three types of bonds?
Options:
✅ (a) Carbon
(b) Oxygen
(c) Magnesium
(d) Silicon
✅ Correct Answer: (a) Carbon
Explanation:
Carbon is a highly versatile element. It can form:
Covalent bonds (e.g., CH₄)
Coordinate covalent bonds (e.g., CO)
Ionic bonds (e.g., CaC₂ where C acts as carbide ion)
Due to this tetravalency and bonding versatility, carbon can exhibit all three major types of bonds.(iii) Why is H₂O a liquid while H₂S is a gas?
Options:
(a) Because in water, the atomic size of oxygen is smaller than that of sulphur
✅ (b) Because water is a polar compound and there exists strong forces of attraction between its molecules
(c) Because H₂O molecule is lighter than H₂S
(d) Because water can easily freeze into ice✅ Correct Answer: (b)
Explanation:
Water (H₂O) forms strong hydrogen bonds due to its polarity, which leads to high intermolecular attraction. This makes water a liquid at room temperature, whereas H₂S, being less polar, is a gas.(iv) Which of the following bonds is expected to be the weakest?
Options:
(a) O−O
(b) Cl−Cl
(c) N−N
✅ (d) F−F✅ Correct Answer: (d)
Explanation:
The F–F bond is weak due to repulsion between lone pairs on fluorine atoms, despite its small size. Hence, it is the weakest single bond among the options.(v) Which form of carbon is used as a lubricant?
Options:
(a) Coal
(b) Diamond
✅ (c) Graphite
(d) Charcoal✅ Correct Answer: (c)
Explanation:
Graphite has layers of carbon atoms that can slide over one another due to weak forces, making it useful as a lubricant.(vi) Keeping in view the intermolecular forces of attraction, indicate which compound has the highest boiling point.
Options:
✅ (a) H₂O
(b) H₂S
(c) HF
(d) NH₃✅ Correct Answer: (a)
Explanation:
Water (H₂O) forms extensive hydrogen bonding, giving it the highest boiling point among the listed compounds.(vii) Which metal has the lowest melting point?
Options:
(a) Li
(b) Na
(c) K
✅ (d) Rb✅ Correct Answer: (d)
Explanation:
Rubidium (Rb) is the largest alkali metal listed and has the weakest metallic bonding, resulting in the lowest melting point among them.(viii) Which ionic compound has the highest melting point?
Options:
(a) NaCl
(b) KCl
✅ (c) LiCl
(d) RbCl✅ Correct Answer: (c)
Explanation:
Li⁺ and Cl⁻ have the smallest ionic sizes, resulting in stronger electrostatic forces and hence a higher melting point compared to others.(ix) Which compound contains both covalent and ionic bonds?
Options:
(a) MgCl₂
✅ (b) NH₄Cl
(c) CaO
(d) PCl₅✅ Correct Answer: (b)
Explanation:
NH₄Cl contains a covalent bond within the ammonium ion (NH₄⁺) and forms ionic bonds between NH₄⁺ and Cl⁻ — so it has both bond types.(x) Which among the following has a double covalent bond?
Options:
(a) Ethane
(b) Methane
✅ (c) Ethylene
(d) Acetylene✅ Correct Answer: (c)
Explanation:
Ethylene (C₂H₄) contains a double bond (C=C) between the two carbon atoms, making it the correct choice.
(Acetylene has a triple bond.)
Short Questions with Answers
i. What type of elements lose their outer electron easily and what type of elements gain electron easily?
Answer:
Metals lose their outer electrons easily to form positive ions (cations).
Non-metals gain electrons easily to form negative ions (anions).
ii. Why do lower molecular mass covalent compounds exist as gases or low boiling liquids?
Answer:
Lower molecular mass covalent compounds exist as gases or low-boiling liquids because they have weak intermolecular forces of attraction (like van der Waals forces), which require less energy to overcome.
iii. Give one example of an element which exists as a crystalline solid and it has covalent bonds between its atoms.
Answer:
Diamond is an example of a crystalline solid where each carbon atom is covalently bonded to four other carbon atoms, forming a rigid 3D structure.
iv. Which property of metals makes them malleable and ductile?
Answer:
The presence of delocalized electrons in a metallic bond allows metal atoms to slide over each other without breaking the bond. This makes metals malleable (can be hammered) and ductile (can be drawn into wires).
v. Is coordinate covalent bond a strong bond?
Answer:
Yes, a coordinate covalent bond is a strong bond, similar in strength to a regular covalent bond. The only difference is that both electrons in the shared pair come from the same atom.
vi. Write down dot and cross formula of HNO₃.
Answer:
Here’s the description:
H is bonded to one O via a single bond (O–H).
N is bonded to two O atoms — one via double bond (O=), and one via a coordinate covalent bond (with lone pair donation from O⁻).
The structure is:
- O
||
H–O–N→O⁻ Dot & cross notation:
Use dots (•) for nitrogen electrons
Use crosses (×) for oxygen electrons
Hydrogen has one electron (× or •)
Coordinate bond is shown by an arrow (→) from oxygen to nitrogen.
Constructed Response Questions
i. Why is HF a liquid while HCl is a gas?
Answer:
HF (hydrogen fluoride) is a liquid at room temperature, while HCl (hydrogen chloride) is a gas, due to hydrogen bonding:
In HF, fluorine is highly electronegative, so it forms strong hydrogen bonds between HF molecules. These intermolecular forces hold the molecules close, making HF a liquid.
In HCl, chlorine is less electronegative than fluorine and does not form hydrogen bonds, so the only intermolecular forces are weak van der Waals forces, making it exist as a gas.
Conclusion: Strong hydrogen bonding in HF results in higher boiling point and liquid state.
ii. Why are covalent compounds generally not soluble in water?
Answer:
Most covalent compounds are non-polar or only slightly polar, while water is a polar solvent.
According to the rule “like dissolves like”, polar substances dissolve in polar solvents, and non-polar substances dissolve in non-polar solvents.
Covalent compounds do not form ions in water and cannot interact strongly with water molecules.
Result: Covalent compounds do not dissolve well in water (e.g., oil, wax, benzene).
iii. How do metals conduct heat?
Answer:
Metals conduct heat due to the presence of free electrons in their metallic structure:
In metallic bonding, valence electrons are delocalized and form a “sea of electrons”.
These electrons absorb heat energy, become excited, and transfer energy quickly through the metal.
The kinetic energy is rapidly passed from atom to atom, resulting in efficient thermal conductivity.
Conclusion: The mobility of delocalized electrons is responsible for metals being good conductors of heat.
iv. How many oxides does nitrogen form? Write down the formulae of oxides.
Answer:
Nitrogen forms five common oxides due to its ability to show variable oxidation states.
Oxides of nitrogen:
N₂O – Dinitrogen monoxide (nitrous oxide)
NO – Nitric oxide
N₂O₃ – Dinitrogen trioxide
NO₂ – Nitrogen dioxide
N₂O₅ – Dinitrogen pentoxide
Each oxide has different chemical and physical properties and is used in various industrial and environmental processes.
v. What will happen if NaBr is treated with AgNO₃ in water?
Answer:
When sodium bromide (NaBr) is mixed with silver nitrate (AgNO₃) in aqueous solution, a double displacement reaction (precipitation) occurs.
Chemical Equation:
NaBr (aq) + AgNO₃ (aq) → NaNO₃ (aq) + AgBr (s) ↓
Silver bromide (AgBr) is formed as a yellowish-white precipitate.
Sodium nitrate remains dissolved in solution.
Conclusion: This reaction confirms the presence of bromide ions in NaBr.
vi. Why does iodine exist as a solid while Cl₂ exists as a gas?
Answer:
The physical state of a substance at room temperature depends on the strength of intermolecular forces:
Iodine (I₂) is a larger molecule than chlorine (Cl₂), and its electron cloud is bigger.
Larger molecules have stronger London dispersion forces (van der Waals forces).
These strong forces hold iodine molecules tightly together, making it a solid.
In contrast, Cl₂ molecules have weaker intermolecular forces, so chlorine remains a gas at room temperature.
Conclusion: Stronger intermolecular forces in iodine cause it to exist as a solid.
Descriptive Questions
i. Explain the formation of an ionic bond and a covalent bond.
Ionic Bond:
An ionic bond is formed when one atom transfers its electrons to another atom.
Usually occurs between a metal and a non-metal.
The metal loses electrons to form a positive ion (cation), and the non-metal gains electrons to form a negative ion (anion).
These oppositely charged ions attract each other to form an ionic bond.
🔹 Example: Sodium (Na) + Chlorine (Cl) → NaCl
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
Na⁺ + Cl⁻ → NaCl
Covalent Bond:
A covalent bond is formed when two non-metal atoms share electrons to complete their outer shells.
Each atom contributes one or more electrons.
No ions are formed; instead, a shared pair of electrons forms the bond.
🔹 Example: Hydrogen molecule (H₂)
Each H atom shares one electron to form a single bond → H–H
ii. How do ions arrange themselves to form NaCl crystal?
Answer:
In a solid NaCl crystal, sodium ions (Na⁺) and chloride ions (Cl⁻) are arranged in a regular 3D lattice structure known as crystal lattice.
Each Na⁺ ion is surrounded by six Cl⁻ ions, and each Cl⁻ ion is surrounded by six Na⁺ ions.
This regular, repeating pattern forms a strong electrostatic force of attraction between oppositely charged ions.
This gives NaCl its high melting point, hardness, and brittle nature.
iii. Explain the properties of metals keeping in view the nature of metallic bond.
Answer:
In metals, atoms are held together by metallic bonding, where positive metal ions are surrounded by a “sea of delocalized electrons”.
Properties due to metallic bonding:
Electrical conductivity:
Free-moving electrons allow metals to conduct electricity.Thermal conductivity:
Electrons transfer heat energy efficiently.Malleability and ductility:
Layers of atoms can slide over each other without breaking the bond.Luster:
Free electrons reflect light, giving shiny appearance.Strength and toughness:
Strong attraction between positive ions and electrons gives metals their hardness.
iv. Compare the properties of ionic and covalent compounds.
| Property | Ionic Compounds | Covalent Compounds |
|---|---|---|
| Bond type | Transfer of electrons (ionic bond) | Sharing of electrons (covalent bond) |
| Components | Metal + Non-metal | Non-metal + Non-metal |
| Physical state | Usually solids | Gases, liquids, or soft solids |
| Melting/boiling point | High | Low to moderate |
| Solubility in water | Mostly soluble | Mostly insoluble |
| Electrical conductivity | Conducts in molten/aqueous state | Usually poor conductor |
| Example | NaCl, MgO | H₂O, CO₂, CH₄ |
v. How will you explain the electrical conductivity of graphite crystals?
Answer:
Graphite is an allotrope of carbon in which each carbon atom is covalently bonded to three other carbon atoms in hexagonal layers.
The fourth electron of each carbon atom is free (delocalized) and moves between layers.
These free electrons allow electric current to pass through graphite, making it a good conductor of electricity.
This property makes graphite useful in electrodes, batteries, and conductive materials.
vi. Why are metals usually hard and heavy?
Answer:
Hardness and heaviness of metals are due to:
Strong metallic bonds:
The tight packing of metal atoms and the strong attraction between positive metal ions and delocalized electrons gives metals their hard and rigid structure.High density:
Metal atoms are closely packed, and many metals have high atomic masses, resulting in greater density (heaviness).
Conclusion: The compact crystal structure and strong bonding make most metals hard and heavy.
Investigative Questions
i. The formula of AlCl₃ in vapour phase is Al₂Cl₆ which means it exists as a dimer. Explain the bonding between its two molecules.
Answer:
In the vapour phase, aluminum chloride (AlCl₃) does not remain as a single molecule. Instead, it dimerizes to form Al₂Cl₆ due to electron deficiency in aluminum atoms.
🔹 Why does dimerization occur?
In AlCl₃, aluminum has only 6 electrons in its valence shell after bonding with three Cl atoms — it needs 2 more electrons to complete its octet.
To achieve stability, two AlCl₃ molecules combine and share two chloride ions through coordinate covalent bonds.
🔹 Structure of Al₂Cl₆:
Two Al atoms are bridged by two Cl atoms.
Each bridging Cl donates a lone pair to Al using a coordinate bond (Al ← Cl).
This results in both Al atoms completing their octet. Cl Cl
\ /
Cl–Al←Cl→Al–Cl
/ \
Cl Cl
Conclusion:
Al₂Cl₆ is stabilized in the vapour phase through two coordinate covalent bonds involving bridging chloride ions, allowing both aluminum atoms to complete their octets.
ii. Explain the structure of sand (SiO₂).
Answer:
Sand is mainly composed of silicon dioxide (SiO₂), which has a giant covalent (network) structure — not simple discrete molecules.
🔹 Structure:
Each silicon (Si) atom is covalently bonded to four oxygen (O) atoms in a tetrahedral geometry.
Each oxygen atom is shared between two silicon atoms.
This arrangement forms a continuous 3D network of strong Si–O bonds throughout the crystal.
🔹 Characteristics of SiO₂:
Hard and rigid structure due to strong covalent bonding.
Very high melting and boiling points.
Insoluble in water and most solvents.
Does not conduct electricity, as there are no free ions or electrons.
Conclusion:
Sand (SiO₂) has a giant covalent lattice, with each silicon atom bonded to four oxygen atoms, forming a strong, stable, and hard structure typical of network solids.