Chemical Bonding Class 10 ICSE — Chemistry Chapter 2

Chemical bonding class 10 ICSE explains why atoms combine and how they reach a stable electron arrangement. This chapter from the Class 10 Chemistry book covers the three bond types: electrovalent, covalent and coordinate. It also compares their properties and shows how electricity affects them. Section by section, you will learn the definitions, conditions and electron dot structures that board questions repeatedly test.

Download Chemical Bonding Class 10 ICSE Chemistry Chapter 2 PDF

Download the Chemical Bonding Class 10 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 17 pages, 13 sections, 36 questions.

What is inside Chapter 2

17 pages · 13 sections · 1 worked examples · 36 exercise questions

• Intext Questions: 12 questions
• Intext Questions: 15 questions
• EXERCISE: 9 questions

2.1 INTRODUCTION

Atoms, like everything else, want stability. For an atom, stability means the electron arrangement of an inert gas. Helium holds two electrons, called a duplet. All other inert gases hold eight, called an octet. Table 2.1 lists the configurations of helium through radon.

Elements with a complete outermost shell are the least reactive. To become stable, atoms combine chemically. They redistribute their valence electrons to reach the nearest noble gas configuration. This redistribution creates a bond between atoms.

A chemical bond is the force of attraction between any two atoms in a molecule. It keeps the molecule stable. Atoms reach stability in three ways: transfer of electrons, sharing of electrons, and donation of a lone pair. These produce electrovalent, covalent and coordinate bonds respectively.

2.2 ELECTROVALENT (OR IONIC) BOND

Metals with 1, 2 or 3 valence electrons can lose them to non-metals with 7, 6 or 5 valence electrons. The metal atom becomes a cation, a positively charged ion. The non-metal atom becomes an anion, a negatively charged ion. An ion is a charged particle formed by the gain or loss of electrons.

The electrostatic attraction between the cation and anion is the electrovalent (or ionic) bond. Electrovalency is the number of electrons an atom loses or gains to form this bond. Three factors favour ionic bonding: low ionisation potential, high electron affinity, and a large electronegativity difference.

Metals of groups 1, 2 and 13 combine with non-metals of groups 15, 16 and 17. Caesium fluoride is the most ionic compound because fluorine is the most electronegative element. Ionic compounds stay stable because the attraction between opposite charges is stronger than the repulsion between like charges.

Take potassium and bromine as the pattern. Potassium (atomic number 19) has configuration 2, 8, 8, 1. It loses one electron to reach argon’s octet and becomes \mathrm{K^+}. Bromine (atomic number 35) has 2, 8, 18, 7. It gains one electron and becomes \mathrm{Br^-}. The bond between them gives potassium bromide, KBr.

\mathrm{K} \rightarrow \mathrm{K^+} + \mathrm{e^-} \quad \text{(oxidation)}

\mathrm{Br} + \mathrm{e^-} \rightarrow \mathrm{Br^-} \quad \text{(reduction)}

2.2.1 Structures of some electrovalent compounds

Sodium (2, 8, 1) loses one electron to reach neon’s octet, forming \mathrm{Na^+}. Chlorine (2, 8, 7) gains one electron to reach argon’s octet, forming \mathrm{Cl^-}. The two ions attract and form sodium chloride. The chapter compares the sodium atom with the sodium ion. The atom is silvery white and reactive, while the ion is colourless and inactive. Chlorine shows the same pattern.

The atom is a poisonous yellowish-green gas, but the chloride ion is colourless and harmless.

The electron dot symbol shows only the outermost electrons as dots around the element symbol. Unpaired electrons appear as single dots. Paired electrons appear as dot pairs. This is also called the Lewis symbol.

Magnesium (2, 8, 2) must lose two electrons. Each chlorine atom can accept only one. So one magnesium atom needs two chlorine atoms, giving the 1 : 2 ratio in \mathrm{MgCl_2}. Calcium (2, 8, 8, 2) also loses two electrons, but oxygen needs two. One calcium and one oxygen therefore give CaO, not \mathrm{Ca_2O_2}.

Electron transfer is a redox process. The electropositive atom loses electrons and is oxidised. The electronegative atom gains electrons and is reduced. An oxidising agent accepts electrons; a reducing agent donates them.

2.3 COVALENT (MOLECULAR) BOND

A covalent bond forms when two atoms share one or more pairs of electrons. Non-metals usually have 5, 6 or 7 valence electrons. They do not lose electrons easily, so transfer is not possible. Instead, each atom contributes an equal number of electrons and shares the pair. The compound formed is a covalent compound.

Covalent bonds come in three types. A single bond shares one pair, written as \mathrm{H-H}. A double bond shares two pairs, written as \mathrm{O=O}. A triple bond shares three pairs, written as \mathrm{N \equiv N}. Ethyne shows a triple bond with two single bonds: \mathrm{H-C \equiv C-H}.

Covalency is the number of electrons an atom uses in shared pairs. Hydrogen has covalency 1, oxygen 2, nitrogen 3 and carbon 4.

Covalent compounds can be polar or non-polar. When identical atoms share equally, the molecule is non-polar, like \mathrm{H_2} or \mathrm{Cl_2}. When the shared pair sits closer to one atom, fractional charges develop. In hydrogen chloride, chlorine pulls the pair closer, giving \mathrm{Cl^{\delta-}} and \mathrm{H^{\delta+}}. This polar molecule is called a dipole molecule.

The larger the electronegativity difference, the more polar the bond. A very large difference makes the bond ionic.

2.3.1 Conditions for the formation of a covalent bond

Five conditions favour a covalent bond:

• Both atoms have four or more valence electrons (exceptions: H, Be, B, Al).
• Both atoms have high electronegativity.
• Both atoms have high electron affinity.
• Both atoms have high ionisation energy.
• The electronegativity difference is zero or negligible.

The chlorides table shows the trend. Group I gives NaCl, an ionic solid. Group III gives \mathrm{AlCl_3}, partially ionic and partially covalent. Group IV gives \mathrm{SiCl_4}, a covalent liquid. Moving across a period, the electrovalent character of chlorides and oxides decreases.

The oxides follow the same pattern. \mathrm{Na_2O} and MgO are ionic solids, while \mathrm{SO_2} and \mathrm{Cl_2O_7} are covalent. The group position alone often predicts the bond type. In the Class 10 ICSE syllabus, this trend is a quick way to answer bonding questions.

2.4 PROPERTIES AND COMPARISON OF ELECTROVALENT AND COVALENT COMPOUNDS

This section compares electrovalent and covalent compounds across seven properties. Each property has a reason rooted in the type of particles and the forces between them. The table below maps every property to its reason.

In this ICSE chapter, the comparison table is a favourite exam target. Board questions repeatedly ask why ionic compounds conduct when molten, why covalent compounds have low melting points, and why NaCl dissolves in water while \mathrm{CCl_4} does not. Each answer traces to the presence or absence of mobile ions.

2.3.2 Some covalent molecules and their structures

Seven molecules get full treatment here. Hydrogen, chlorine and nitrogen are homonuclear, meaning both atoms are the same element. Hydrogen needs one electron for a duplet, so two hydrogen atoms share one pair, giving \mathrm{H-H}. Chlorine (2, 8, 7) needs one electron for argon’s octet, giving \mathrm{Cl-Cl}. Nitrogen (2, 5) needs three electrons, so two nitrogen atoms share three pairs, giving \mathrm{N \equiv N}.

Water and ammonia are polar. Two hydrogen atoms share pairs with oxygen (2, 6), which needs two electrons. Each hydrogen reaches helium’s duplet and oxygen reaches neon’s octet. In ammonia, nitrogen (2, 5) shares three pairs with three hydrogen atoms.

Carbon tetrachloride and methane are non-polar. Carbon (2, 4) needs four electrons. In \mathrm{CCl_4} it shares four pairs with four chlorine atoms. In \mathrm{CH_4} it shares four pairs with four hydrogen atoms.

Mnemonic: remember the polar pair with “wet”. Water and ammonia are the only two polar molecules here. Hydrogen, chlorine and nitrogen are non-polar because both atoms are identical. Carbon tetrachloride and methane stay non-polar because their structures are symmetrical.

Common mistakes to avoid:

• Electron dot diagrams: draw dots only for the outermost shell, never the inner shells.
• Polarity: water is polar even though it has only single bonds.
• Electrovalency: count the electrons actually lost or gained, not the total valence electrons.

2.5 EFFECT OF ELECTRICITY ON ELECTROVALENT AND COVALENT COMPOUNDS

The chapter describes a simple experiment. A circuit holds a 6-volt battery, an ammeter, a bulb and platinum electrodes in series. Each solution is tested one by one.

Electrovalent solutions make the bulb glow. These include magnesium chloride, sodium chloride, sodium hydroxide and copper sulphate. The current passes because the solutions contain free mobile ions.

Covalent solutions keep the bulb off. These include distilled water, sugar solution, alcohol, chloroform, benzene and petrol. They contain only molecules and no ions.

The conclusion is clear. Ionic compounds conduct electricity in molten or aqueous states because their ions are mobile. Covalent compounds do not conduct because they have no free ions.

2.6 COORDINATE BOND

A coordinate bond forms when one atom supplies the entire shared pair of electrons. The other atom only accepts the pair. This bond is also called a dative bond.

It is known as a co-ionic bond because it carries properties of both covalent and ionic bonds. The two standard examples are the ammonium ion \mathrm{NH_4^+} and the hydronium ion \mathrm{H_3O^+}.

2.6.1 Conditions for the formation of coordinate bond

Two conditions must hold. First, one atom must carry at least one lone pair of electrons. Ammonia and water both qualify. A lone pair is a pair of electrons not shared with any other atom.

Second, the other atom must be short of at least one pair of electrons. The hydrogen ion \mathrm{H^+} has no electrons at all, so it qualifies perfectly.

The atom that provides the pair is the donor. The atom or ion that accepts it is the acceptor. A coordinate bond is shown with an arrow → pointing from donor to acceptor.

Common mistake: a coordinate bond is not a separate physical bond. Once formed, all four N–H bonds in the ammonium ion are identical.

2.6.2 Formation of Ammonium ion (NH_4^+)

Nitrogen has five valence electrons and hydrogen has one. In ammonia, nitrogen shares one pair with each of three hydrogen atoms. Three N–H bonds form, and nitrogen keeps one lone pair.

A hydrogen ion \mathrm{H^+} has no electrons in its outermost shell. The lone pair on nitrogen is shared with this ion. The hydrogen ion gains helium’s duplet configuration.

The arrow → marks the coordinate bond. The dash — marks covalent bonds. All four N–H bonds become identical, and the ion carries a single positive charge.

Ammonium chloride shows all three bond types. The \mathrm{NH_4^+} cation holds covalent and coordinate bonds. The \mathrm{Cl^-} anion forms an ionic bond with it.

2.6.3 Formation of H₃O⁺ ion and lone pair effect of oxygen atom

Water has one oxygen atom and two hydrogen atoms. Oxygen (2, 6) shares one pair with each hydrogen. Two lone pairs remain unshared on oxygen.

Oxygen has electronegativity 3.5, while hydrogen has 2.1. The large difference polarises the O–H bonds. Oxygen gains a slight negative charge \delta^- and hydrogen a slight positive charge \delta^+. Water is therefore a polar molecule.

When an acid dissolves in water, dipole interaction releases a hydrogen ion. That ion adds to one lone pair of oxygen.

\mathrm{H^+} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^+}

The hydronium ion is a hydrated proton. This reaction is the foundation of the acids, bases and salts chapter that follows in this Chemistry book.

2.6.4 Formation of hydroxyl ion (OH^-)

The hydroxyl ion, or hydroxide ion, is written \mathrm{OH^-}. It forms when one hydrogen ion is removed from a water molecule.

\mathrm{H_2O} \rightarrow \mathrm{H^+} + \mathrm{OH^-}

When \mathrm{H^+} leaves, the shared pair of electrons stays with oxygen. Oxygen is more electronegative, so it keeps the pair. The ion therefore carries a negative charge.

Water self-ionises to a small extent. One water molecule releases \mathrm{H^+} and leaves \mathrm{OH^-}. The released ion transfers to another water molecule, forming \mathrm{H_3O^+}.

\mathrm{H_2O} \rightleftharpoons \mathrm{H^+} + \mathrm{OH^-}

\mathrm{H^+} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^+}

Board questions often ask for electron dot diagrams of hydronium, ammonium and hydroxyl ions. Practise drawing all three.

What to practise from Chemical Bonding Class 10 ICSE

This chapter carries three practice blocks. Working through all of them covers every definition and structure in the chapter.

• Intext Questions: 12 questions on noble gas configuration, bond types, electrovalent conditions and redox reactions.
• Intext Questions: 15 questions on covalent conditions, dipole molecules, property comparisons and electron dot diagrams.
• EXERCISE: 9 questions including past board questions from 2007 to 2014.

The board questions cluster around four themes. Coordinate bond definitions and conditions appear almost every year. Electron dot diagrams of \mathrm{NH_4^+}, \mathrm{H_3O^+} and \mathrm{OH^-} are regular asks. Property comparisons between electrovalent and covalent compounds recur constantly. Redox identification, naming oxidising and reducing agents, appears in multiple years.

For the full set of chapters, browse the Class 10 Chemistry books hub. The official CISCE site carries the syllabus and examination updates.

Chemical Bonding Class 10 ICSE frequently asked questions

What is chemical bonding class 10 ICSE about?

It explains why atoms combine to reach a stable inert gas configuration. The chapter covers the three bond types — electrovalent, covalent and coordinate — and compares their properties. It also shows how electricity affects these compounds.

How do I know if a bond is electrovalent or covalent?

Check the electronegativity difference. A large difference between a metal and a non-metal gives an electrovalent bond. A zero or negligible difference between non-metals gives a covalent bond. Conditions like ionisation potential and electron affinity also matter.

Why does ammonium chloride have all three types of bonds?

Ammonium chloride contains \mathrm{NH_4^+} and \mathrm{Cl^-}. The ammonium ion holds three covalent N–H bonds and one coordinate N–H bond. The electrostatic attraction between \mathrm{NH_4^+} and \mathrm{Cl^-} is the ionic bond.

What is the difference between a lone pair and a shared pair?

A shared pair is a pair of electrons used by two atoms in a covalent bond. A lone pair is a pair on one atom that is not shared with any other atom. In ammonia, nitrogen has three shared pairs and one lone pair.

Do polar covalent compounds conduct electricity in water?

Yes, some do. Polar covalent compounds like hydrogen chloride ionise in water, forming ions that carry current. Non-polar covalent compounds have no free ions and do not conduct.

Reference: ICSE Class 10 Chemistry textbook material, chapter 2: Chemical Bonding — CISCE-aligned.

More from this book

• Chapter 5: Mole Concept and Stoichiometry
• Chapter 8: Study of Compounds Hydrogen Chloride
• Chapter 9: Study of Compounds Ammonia
• Chapter 10: Study of Compounds Nitric Acid
• Chapter 11: Sulphuric Acid