Chemical Bonding
1. What This Chapter Covers
There are over 118 elements. The chapter opens by asking how they actually exist — as single atoms, or as groups of atoms? In class 9 you saw that oxygen, nitrogen and hydrogen exist as diatomic molecules. Something holds those atoms together.
The questions the chapter sets out to answer are worth listing, because the whole argument is built to meet them:
- Are there elements that exist as single atoms, and why do some exist as molecules instead?
- Why do some elements and compounds react vigorously while others are inert?
- Why is the formula of water H₂O and of common salt NaCl, and not HO₂ and NaCl₂?
- Why do some atoms combine while others do not?
The simplest demonstration comes first. When you shake salt from a shaker, it does not separate into sodium and chlorine. The sodium and chlorine are being held together — and the question is by what.
By the late nineteenth and early twentieth century scientists knew three kinds of force — gravitational, magnetic and electrostatic — and knew that electrons and protons exist. So electrostatic force was the natural candidate.
The picture that follows is a balance. When two atoms come close, the electrons of each feel the attraction of the other's nucleus. But the electrons repel each other, and the nuclei repel each other too. If attraction exceeds repulsion the atoms combine; if repulsion exceeds attraction they do not. The nucleus and inner-shell electrons are unaffected. Only the valence shell electrons are involved.
An early piece of evidence came from electrolysis: the metal part of a compound migrated to the negative electrode and the non-metal part to the positive. So metals were held responsible for positive particles and non-metals for negative ones, held together electrostatically. The book asks whether you agree — and then answers that while this explains NaCl and KCl, it cannot explain carbon compounds or the diatomic molecules of elements.
The Do you know? box here records Humphry Davy (1778-1819), professor of chemistry at the Royal Institution in London, who built a battery of over 250 metallic plates and in 1807 used it to extract the highly reactive metals potassium and sodium by electrolysis of fused salts.
The chapter is allotted 12 periods in October and runs from textbook page 150 to page 175.
2. Lewis Symbols and the Noble Gas Clue
The discovery of the noble gases, and of their configurations, is what made chemical bonding explainable. The noble gases of group 18 (VIIIA, the zero group) have almost negligible chemical activity. They are stable and do not form molecules with their own atoms or with others.
Table-1 shows why:
| Element | Z | K | L | M | N | Valence electrons |
|---|---|---|---|---|---|---|
| Helium (He) | 2 | 2 | 2 | |||
| Neon (Ne) | 10 | 2 | 8 | 8 | ||
| Argon (Ar) | 18 | 2 | 8 | 8 | 8 | |
| Krypton (Kr) | 36 | 2 | 8 | 18 | 8 | 8 |
Every noble gas has eight electrons in the outermost shell, except helium.
A Lewis symbol or electron dot structure shows this compactly. The symbol of the element stands for the nucleus and all the inner electrons, and the outer shell electrons are drawn as dots or crosses around it. The rule for placing them is to put up to two dots on each of the four sides of the symbol until all are used.
Argon, with eight valence electrons, gets two dots on each of its four sides. Sodium, with one, gets a single mark.
Activity 1 asks you to complete Table-2 for hydrogen, helium, beryllium, boron, carbon, nitrogen and oxygen, filling in the group number, the number of valence electrons and the Lewis structure.
Doing it reveals the shortcut the chapter wants you to notice: for groups 1-2 and 13-18 the group number gives the valence electrons directly. Group 1 has one outer electron, group 2 has two, group 13 has three, group 14 has four, and so on.
From this comes the central observation: elements that take part in chemical reactions end up with an octet, ns² np⁶, like a noble gas. The book is careful about its status — the octet rule is still a rule and not a law, because there are considerable exceptions to it.
3. The Electronic Theory of Valence
Many attempts were made to explain bonding in terms of electrons, but the satisfactory explanation came from Kossel and Lewis in 1916, independently of each other. Their basis was valence expressed in electrons, and their starting point was the chemical inactivity of the noble gases.
Watch what the main group elements actually do:
| Group | Behaviour | Example |
|---|---|---|
| IA (Li to Cs) | Lose one valence electron, form uni-positive ions with an octet | ₁₁Na 2, 8, 1 gives Na⁺ 2, 8 |
| IIA (Mg to Ba) | Lose two, form di-positive ions | ₁₂Mg 2, 8, 2 gives Mg²⁺ 2, 8 |
| IIIA | Lose three, form tri-positive ions | ₁₃Al 2, 8, 3 gives Al³⁺ 2, 8 |
| VIA | Gain two, form di-negative anions | ₈O 2, 6 gives O²⁻ 2, 8 |
| VIIA | Gain one, form uni-negative anions | ₉F 2, 7 gives F⁻ 2, 8 |
| VIIIA | Neither lose nor gain | ₁₀Ne 2, 8 stays as it is |
Is it accidental that groups IA to VIIA all end with eight outer electrons, exactly like the noble gases? The chapter answers flatly: no, it cannot simply be accidental.
The octet rule
The atoms of elements tend to undergo chemical changes that help to leave their atoms with eight outer-shell electrons.
Lewis pictured the atom as a positively charged kernel — the nucleus together with every electron except those in the outermost shell — surrounded by an outer shell that can hold a maximum of eight.
Helium is the exception that proves the point: its atom has only two electrons, but its only shell is completely filled, so it is just as stable.
Chemically active elements are exactly those whose atoms do not have an octet. Their reactivity is the tendency to get one, by bonding either with atoms of their own kind or with atoms of other elements.
That gives the definition the rest of the chapter uses:
A chemical bond is the force of attraction between any two atoms, or a group of atoms, that results in a stable entity.
There are many types. This chapter treats two: the ionic bond and the covalent bond.
4. The Ionic Bond
Kossel proposed the ionic or electrostatic bond on five facts:
- It forms between atoms of two dissimilar elements, by transfer of electrons from one to the other.
- The periodic table has highly reactive metals, the alkali metals of IA, on the left, and highly reactive non-metals, the halogens of VIIA, on the right.
- Noble gases except helium have eight valence electrons and are chemically inactive and stable.
- Metal atoms with one, two or three valence electrons lose them to reach the octet, forming cations.
- Non-metal atoms with five, six or seven valence electrons gain three, two or one respectively, forming anions.
Two Do you know? boxes give the arithmetic. The number of electrons lost from a metal atom is its valence, equal to its group number — Na and Mg have valence 1 and 2. The number gained by a non-metal is 8 minus its group number — chlorine gives 8 − 7 = 1.
| Cations | Anions |
|---|---|
| ₁₁Na 2, 8, 1 gives Na⁺ 2, 8 | ₁₅P 2, 8, 5 gives P³⁻ 2, 8, 8 |
| ₁₂Mg 2, 8, 2 gives Mg²⁺ 2, 8 | ₁₆S 2, 8, 6 gives S²⁻ 2, 8, 8 |
| ₁₃Al 2, 8, 3 gives Al³⁺ 2, 8 | ₁₇Cl 2, 8, 7 gives Cl⁻ 2, 8, 8 |
An ionic bond is the electrostatic attractive force that keeps a cation and an anion — formed from metal and non-metal atoms by transfer of electrons — together to form a new, electrically neutral compound.
Because the valence concept is expressed in electrons, the same bond is also called the electrovalent bond.
Four worked formations
Sodium chloride. Na(s) + ½Cl₂(g) -> NaCl(s). Sodium loses one electron, ₁₁Na(g) -> Na⁺(g) + e⁻, moving from 2, 8, 1 or [Ne] 3s¹ to 2, 8 or [Ne]. Chlorine is one electron short, so ₁₇Cl(g) + e⁻ -> Cl⁻(g), moving from [Ne] 3s² 3p⁵ to [Ne] 3s² 3p⁶, which is [Ar]. The two ions then attract: Na⁺(g) + Cl⁻(g) -> Na⁺Cl⁻(s).
Magnesium chloride. Mg(s) + Cl₂(g) -> MgCl₂(s). One Mg atom transfers two electrons, one each to two Cl atoms. Mg²⁺ reaches the neon configuration and each Cl⁻ the argon configuration.
Di sodium monoxide. 2Na⁺(g) + O²⁻(g) -> Na₂O(s). Two Na atoms transfer one electron each to a single oxygen atom. Here both ions end at the neon configuration.
Aluminium chloride. Al³⁺(g) + 3Cl⁻(g) -> AlCl₃(s). Each aluminium atom loses three electrons and three chlorine atoms gain one each.
How the ions are arranged in the solid
Do pairs of Na⁺Cl⁻ exist as units in the crystal? No. Electrostatic forces are non-directional, so one Na⁺ cannot be attracted by one particular Cl⁻ and no other.
Instead, depending on the size and charge of the ion, a definite number of oppositely charged ions surround each one. In the sodium chloride crystal each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ by 6 Na⁺, in an orderly three-dimensional arrangement. NaCl has a face centred cubic lattice structure.
The coordination number of an ion is the number of oppositely charged ions surrounding it. In NaCl it is 6 for both.
What decides whether an atom forms a cation or an anion
Four properties from the previous chapter decide it: atomic size, ionisation potential, electron affinity and electronegativity.
| Forms cations | Forms anions |
|---|---|
| Low ionisation energy | High ionisation potential |
| Low electron affinity | High electron affinity |
| Large atomic size | Small atomic size |
| Low electronegativity | High electronegativity |
There is also a numerical criterion worth remembering: an ionic bond forms between atoms whose electronegativity difference is equal to or greater than 1.9.
5. The Covalent Bond
G.N. Lewis (1916) proposed that atoms of some elements can reach an octet without any transfer at all — by sharing valence electrons with one or more atoms. The shared electrons belong to both atoms.
A covalent bond is the chemical bond formed between two atoms by mutual sharing of a pair of valence shell electrons, so that both attain an octet or a duplet in their valence shell.
The name says the mechanism. The prefix co- marks things that are equal or joined, as in coexist and cooperate — here each atom contributes one electron, so the contribution of valence electrons is equal.
Take fluorine. Each atom contributes one electron, the pair formed is shared by both, and each atom in F₂ then has an octet.
| Molecule | What each atom needs | Electrons contributed | Result |
|---|---|---|---|
| F₂ | 1 each | 1 + 1 | Single bond |
| O₂ (₈O is 2, 6) | 2 each | 2 + 2 | Double bond, two shared pairs |
| N₂ (₇N is 2, 5) | 3 each | 3 + 3 | Triple bond, three shared pairs |
| CH₄ | C needs 4, each H needs 1 | C gives 4, each H gives 1 | Four C-H single bonds |
| NH₃ | N needs 3, each H needs 1 | N gives 3, each H gives 1 | Three N-H single bonds |
| H₂O | O needs 2, each H needs 1 | O gives 2, each H gives 1 | Two O-H single bonds |
The covalency of an element is the total number of covalent bonds that one of its atoms forms.
Bond length and bond energy
Bond length, or bond distance, is the equilibrium distance between the nuclei of the two atoms forming a covalent bond. It is given in nm or in angstroms. The Do you know? box fixes the units: 1 angstrom = 10⁻¹⁰ m = 0.1 nm = 100 pm, and 1 nm = 10⁻⁹ m.
Bond energy, or bond dissociation energy, is the energy needed to break a covalent bond between two atoms of a diatomic covalent compound in the gaseous state.
Table-3 gives both, and it is the table that destroys the simple sharing picture:
| Bond | Bond length (Å) | Bond dissociation energy (kJ mol⁻¹) |
|---|---|---|
| H-H | 0.74 | 436 |
| F-F | 1.44 | 159 |
| Cl-Cl | 1.95 | 243 |
| Br-Br | 2.28 | 193 |
| I-I | 2.68 | 151 |
| H-F | 0.918 | 570 |
| H-Cl | 1.27 | 432 |
| H-Br | 1.42 | 366 |
| H-I | 1.61 | 298 |
| H-O (of H₂O) | 0.96 | 460 |
| H-N (of NH₃) | 1.01 | 390 |
| H-C (of CH₄) | 1.10 | 410 |
6. Two Drawbacks of the Electronic Theory of Valence
First. If every covalent bond is simply the sharing of two identical electrons, then all covalent bonds should have the same length and the same energy whatever the atoms. Table-3 says otherwise — the values differ, and they differ systematically with the atoms involved.
Second. The theory says nothing about shape. It cannot explain why the angle Cl-Be-Cl in BeCl₂ is 180°, F-B-F in BF₃ is 120°, H-C-H in CH₄ is 109° 28′, H-N-H in NH₃ is 107° 48′ and H-O-H in H₂O is 104° 31′.
Those two failures set up the rest of the chapter: VSEPR theory for the shapes, and valence bond theory with hybridisation for the strengths.
7. Valence Shell Electron Pair Repulsion Theory
VSEPRT was proposed by Sidgwick and Powell (1940) and improved by Gillespie and Nyholm (1957). It explains bond angles in molecules where three or more atoms are attached to a central atom by covalent bonds.
Its three governing points:
- Electrons in the valence shell — both those in covalent bonds and those in lone pairs — behave as charge clouds that repel one another and stay as far apart as possible. This is why molecules have specific shapes.
- Knowing the total number of electron pairs around the central atom, as bonds plus lone pairs, lets you predict their arrangement and hence the shape.
- Lone pairs occupy more space than bond pairs, because a lone pair is attracted to only one nucleus while a bond pair is shared between two. Lone pairs therefore distort the bond angles away from the regular shape.
Applying it gives the five standard cases:
| Bond pairs | Lone pairs | Arrangement | Shape | Angle | Example |
|---|---|---|---|---|---|
| 2 | 0 | Opposite sides | Linear | 180° | BeCl₂ |
| 3 | 0 | Three corners of a triangle | Trigonal planar | 120° | BF₃ |
| 4 | 0 | Four corners of a tetrahedron | Tetrahedral | 109° 28′ | CH₄ |
| 3 | 1 | Tetrahedral pairs, three bonds | Trigonal pyramidal | 107° 48′ | NH₃ |
| 2 | 2 | Tetrahedral pairs, two bonds | V shape, bent or angular | 104° 31′ | H₂O |
The ordering of the repulsions is what drives the last two rows. Lone pair-bond pair repulsion is greater than bond pair-bond pair repulsion, so ammonia's angle falls from the expected 109° 28′ to 107° 48′. Lone pair-lone pair repulsion is greater still, so water — with two lone pairs — falls further, to 104° 31′.
The book adds a note worth keeping. In BeCl₂ and BF₃ the central atoms do not have eight electrons in the valence shell; they have only four and six. These are electron deficient molecules.
A bond angle is the angle subtended at the central atom by two imaginary lines passing from the nuclei of the two atoms bonded to it, through the nucleus of the central atom.
VSEPRT has its own limit: it fails to explain the strengths of bonds, because it still rests on the Lewis picture and says nothing extra about the electronic nature of the covalent bond.
8. Valence Bond Theory
To describe covalent bonding properly a quantum mechanical model was needed. The book attributes the valence bond theory to Linus Pauling (1954). Its four points:
1. A bond is an overlap. A covalent bond forms when two atoms approach closely and one overlaps its valence orbital containing an unpaired electron with the valence orbital of the other containing an unpaired electron of opposite spin. The paired electrons in the overlapping orbitals are attracted to both nuclei, and that binds the atoms.
In H₂, the 1s orbital of one hydrogen atom overlaps the 1s orbital of the other.
2. More overlap, stronger bond. This also gives the bond a directional character whenever orbitals other than s are involved — which is exactly what the Lewis picture could not supply.
3. Each atom keeps its own atomic orbitals, but the electron pair in the overlapping region is shared.
4. Multiple bonds are not all alike. The first bond comes from overlap along the internuclear axis, giving a stronger sigma (σ) bond. Any further bonds come from sideways or lateral overlap, giving weaker pi (π) bonds. The σ bond is stronger because end-on overlap concentrates the shared pair between the two nuclei; lateral overlap of p orbitals is less extensive.
Worked molecules
| Molecule | Configuration used | Overlap |
|---|---|---|
| H₂ | 1s¹ each | 1s-1s, one σ bond |
| Cl₂ | ₁₇Cl 1s² 2s² 2p⁶ 3s² 3pₓ² 3p_y² 3p_z¹ | 3p_z of one with 3p_z of the other |
| HCl | H 1s¹, Cl 3p¹ | 1s of H with the 3p of Cl |
| N₂ | ₇N 1s² 2s² 2pₓ¹ 2p_y¹ 2p_z¹ | σ pₓ-pₓ along the axis, plus π p_y-p_y and π p_z-p_z laterally — a triple bond |
| O₂ | ₈O 1s² 2s² 2pₓ² 2p_y¹ 2p_z¹ | σ p_y-p_y along the axis, plus π p_z-p_z laterally — a double bond |
9. Hybridisation
Valence bond theory as stated still leaves a problem, and beryllium chloride shows it.
₄Be is 1s² 2s² and has no unpaired electrons, so it should form no covalent bonds at all. Yet it forms two. To allow that, an excited state is proposed in which one electron shifts from 2s to 2pₓ, giving 1s² 2s¹ 2pₓ¹.
But now a new discrepancy appears. One Be-Cl bond would be a σ 2s-3p overlap and the other a σ 2p-3p overlap, so the two bonds should have different strengths. In fact both are identical and the angle Cl-Be-Cl is exactly 180°.
Linus Pauling (1931) proposed hybridisation to resolve exactly this kind of discrepancy:
Hybridisation is the intermixing of atomic orbitals of almost equal energy present in the outer shell of an atom, and their reshuffling or redistribution into the same number of orbitals but with equal energy and shape.
| Molecule | Orbitals intermixed | Hybrid | Number formed | Angle | Shape |
|---|---|---|---|---|---|
| BeCl₂ | 2s and 2pₓ | sp | 2 | 180° | Linear |
| BF₃ | 2s, 2pₓ, 2p_y | sp² | 3 | 120° | Trigonal planar |
| NH₃ | 2s, 2pₓ, 2p_y, 2p_z | sp³ | 4 | 107° 48′ | Trigonal pyramidal |
| H₂O | 2s, 2pₓ, 2p_y, 2p_z | sp³ | 4 | 104° 31′ | V shape |
Beryllium chloride. The two sp orbitals separate by 180°, each taking one electron by Hund's rule. Each chlorine overlaps its 3p_z orbital with one sp orbital, giving two identical σ sp-p bonds of the same strength.
Boron trifluoride. ₅B is 1s² 2s² 2pₓ¹ and has one unpaired electron, so it should give only B-F. In practice we get BF₃. Boron is first excited to 1s² 2s¹ 2pₓ¹ 2p_y¹, then the three orbitals intermix into three sp² hybrids at 120° to one another, each taking one electron. Three fluorine atoms (₉F 1s² 2s² 2pₓ² 2p_y² 2p_z¹) overlap their 2p_z orbitals to give three σ sp²-p bonds.
Ammonia. If three hydrogens simply overlapped the three p orbitals of nitrogen the angle would be 90°, but it is 107° 48′. So nitrogen undergoes sp³ hybridisation: 2s and the three 2p orbitals redistribute into four identical sp³ orbitals, one of which gets a pair and three of which get one electron each. Three σ s-sp³ bonds form. sp³ alone would give 109° 28′; the lone pair-bond pair repulsion pulls it down to 107° 48′.
Water. The same argument. Two σ s-p bonds would give 90°, but the observed angle is 104° 31′. Oxygen undergoes sp³ hybridisation; of the four hybrids, two hold lone pairs and two hold single electrons. The two σ sp³-s bonds are then squeezed by both lone pair-lone pair and lone pair-bond pair repulsion, from 109° 28′ down to 104° 31′.
The book notes that CH₄, C₂H₄ and C₂H₂ are left for the Carbon and its Compounds chapter later in the same class.
10. Properties of Ionic and Covalent Compounds
Table-4 sets one of each type side by side, with a polar covalent compound in between:
| # | Property | NaCl (ionic) | HCl (polar covalent) | C₂H₆ (covalent) |
|---|---|---|---|---|
| 1 | Formula mass | 58.5 | 36.5 | 30.0 |
| 2 | Physical appearance | White crystalline solid | Colourless gas | Colourless gas |
| 3 | Type of bond | Ionic | Polar covalent | Covalent |
| 4 | Melting point | 801 °C | −115 °C | −183 °C |
| 5 | Boiling point | 1413 °C | −84.9 °C | −88.63 °C |
| 6 | Solubility | Soluble in polar solvents like water, insoluble in non-polar | Soluble in polar solvents like water, and to some extent in non-polar | Soluble in non-polar solvents, insoluble in polar solvents like water |
| 7 | Chemical activity | Highly reactive in polar solvents, reactions instantaneous | Moderately reactive | Slow or very slow at room temperature |
Why the columns differ
Ionic compounds have strong electrostatic attractions between oppositely charged ions, so they are solids with high melting and boiling points. Being highly polar, they dissolve in polar solvents. Their reactions in solution are merely a rearrangement of ions, so they are instantaneous or very fast.
Covalent compounds have only weak forces between molecules, so they are gases or liquids at room temperature with low melting and boiling points. They dissolve in non-polar solvents. Their reactions require bonds to break and new bonds to form, so they are moderate or very slow.
The organising principle behind the solubility row is "like dissolves in like" — a solute dissolves in the solvent whose molecules carry the same type of chemical bond.
The polar covalent bond in between
If a covalent bond joins atoms of two different elements, the shared pair shifts towards the more electronegative atom. That atom then carries a partial negative charge and the other a partial positive charge, while the molecule as a whole stays neutral. Such a molecule is polar, and the bond is a polar covalent bond — partly ionic and partly covalent:
H + Cl -> Hᵟ⁺ - Clᵟ⁻
This is why HCl sits between NaCl and C₂H₆ on every row of the table.
Key words from the chapter
Electrons, noble gases, Lewis dot structures, octet rule, chemical bond, ionic bond, covalent bond, cation, anion, electrostatic force, electrovalent, polar solvent, non-polar solvent, ionic compounds, covalent compounds, electropositive character, electronegative character, polar bonds, bonded pair, lone pairs, bond length, bond energy, shape of the molecule, linear, tetrahedral.
11. Summary
The position of an element in the periodic table predicts the bonding its atoms will do, because bonding is a valence-shell affair. Noble gases are stable because they have an octet — helium a filled duplet — and every other main-group element reacts in order to get one. That is the octet rule, and Lewis's kernel-and-dots notation is how it is written down.
Kossel and Lewis, independently in 1916, gave the electronic theory of valence. Kossel's ionic bond is transfer: metals with one, two or three valence electrons lose them to form cations, non-metals with five, six or seven gain electrons to form anions, and the electrostatic attraction between the two is the bond.
Because that attraction is non-directional, ionic solids are three-dimensional lattices — in NaCl each ion has a coordination number of 6 in a face centred cubic lattice. An ionic bond needs an electronegativity difference of 1.9 or more.
Lewis's covalent bond is sharing. Each atom contributes an electron to a shared pair, and one, two or three shared pairs give single, double and triple bonds — F₂, O₂ and N₂ respectively.
The electronic theory then fails twice. Bond lengths and bond energies are not the same for every covalent bond, as Table-3 shows, and the theory cannot explain molecular shapes.
VSEPRT, from Sidgwick and Powell (1940) and Gillespie and Nyholm (1957), supplies the shapes by treating bond pairs and lone pairs as charge clouds that stay as far apart as possible. Since lone pairs take more space, the angle falls from 109° 28′ in CH₄ to 107° 48′ in NH₃ to 104° 31′ in H₂O as lone pairs are added.
Valence bond theory, from Pauling, supplies the strengths by describing a bond as the overlap of orbitals carrying unpaired electrons of opposite spin. Head-on overlap gives the stronger sigma bond and lateral overlap the weaker pi bond, so N₂ has one σ and two π bonds.
Hybridisation, proposed by Pauling in 1931, explains why bonds that should differ are identical: orbitals of nearly equal energy intermix and redistribute into equivalent hybrids — sp at 180° for BeCl₂, sp² at 120° for BF₃, and sp³ for NH₃ and H₂O.
Finally, the properties follow from the bonding. Strong non-directional ionic attractions give high-melting crystalline solids that dissolve in polar solvents and react instantly; weak intermolecular forces in covalent compounds give gases and liquids that dissolve in non-polar solvents and react slowly. Polar covalent compounds like HCl, where the shared pair sits nearer the more electronegative atom, fall between the two on every property.
