Understand Electron Sharing Reactions — Understand Reactivity 3.3—Electron sharing reactions - IB Chemistry

SECTION 1

Historical Context & Motivation

For centuries, chemists wondered why some atoms bond together while others remain independent. Early models focused on the transfer of electrons between metals and nonmetals, but they could not explain how two nonmetal atoms — such as two hydrogen atoms — could form a stable molecule. The idea that atoms might share electrons, rather than simply giving or receiving them, was a breakthrough that transformed our understanding of chemical bonding.

1916 • Lewis's Shared Electron Pair

Gilbert N. Lewis proposed that atoms can achieve a stable octet by sharing pairs of electrons, introducing the concept of the covalent bond and the dot-structure notation that bears his name.

1927 • Heitler–London Theory

Walter Heitler and Fritz London applied quantum mechanics to the hydrogen molecule, showing mathematically how two atoms lower their energy by sharing electrons in overlapping orbitals.

1931 • Pauling's Hybridisation Model

Linus Pauling introduced hybridisation and the concept of electronegativity, explaining why covalent bonds have definite geometries and can have polar character.

1953 • Molecular Orbital Theory Matures

Robert Mulliken and Friedrich Hund refined molecular orbital theory, providing an alternative framework that explains bond order, magnetism, and delocalisation in electron-sharing reactions.

These advances raised a central question that IB Chemistry addresses in Reactivity 3.3: What determines whether electrons are shared equally, shared unequally, or transferred — and how does that sharing drive the mechanisms of chemical change?

SECTION 2

Core Principles of Electron Sharing

Electron sharing reactions occur when two atoms each contribute one or more electrons to form a covalent bond. Unlike ionic bonding, where electrons are transferred from one atom to another, covalent bonding involves a mutual attraction of both nuclei to a shared electron pair located between them. Understanding these reactions requires grasping several foundational ideas.

Covalent Bond Formation

A covalent bond forms when the bonding pair of electrons sits between two nuclei. Each nucleus is attracted to the shared pair, and this mutual attraction lowers the system's potential energy, creating a stable bond.

Bond Polarity & Electronegativity

When two atoms have different electronegativities, the shared electron pair is pulled closer to the more electronegative atom, creating a polar covalent bond with partial charges (δ⁺ and δ⁻).

Coordinate (Dative) Bonds

In a coordinate bond, both electrons in the shared pair come from the same atom (the donor). Once formed, a coordinate bond is indistinguishable from a regular covalent bond in strength.

Sigma (σ) and Pi (π) Bonds

A sigma bond results from head-on orbital overlap along the internuclear axis, while a pi bond results from sideways overlap above and below this axis. Double and triple bonds contain one σ bond plus one or two π bonds.

Bond Enthalpy & Reactivity

The energy required to break a covalent bond is called its bond enthalpy. In electron-sharing reactions, old bonds break (endothermic) and new bonds form (exothermic). The balance determines whether the overall reaction releases or absorbs energy.

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KEY TAKEAWAY

Think of a covalent bond like two people sharing a textbook at the same desk. If both people pull equally, the book stays centred — that's a nonpolar covalent bond. If one person is stronger and pulls the book closer, that's a polar covalent bond. If one person takes the book entirely, that's an ionic bond. The electronegativity difference between atoms decides which scenario occurs.

SECTION 3

Visualising Electron Sharing

How Atoms Share Electrons to Form Bonds

Top: Two separate hydrogen atoms each have one electron. When they approach each other, their electron clouds overlap and both electrons occupy the bonding region between the nuclei. Bottom: The potential energy curve shows that energy decreases as the atoms approach, reaching a minimum at the bond length (74 pm for H₂). The depth of the well equals the bond enthalpy (436 kJ mol⁻¹).

The upper portion of the diagram illustrates the fundamental process of covalent bond formation. Two isolated hydrogen atoms, each with a single electron, approach one another. As their 1s orbitals begin to overlap, both electrons are attracted to both nuclei simultaneously. This creates a region of high electron density between the two nuclei — the bonding region. The lower curve shows that the system's energy drops as the atoms move closer together, reaching a minimum at the equilibrium bond length. If the atoms are pushed even closer, strong nucleus-nucleus repulsion causes the energy to rise sharply. This balance between attraction and repulsion defines the strength and length of every covalent bond.

SECTION 4

Mechanisms of Electron Sharing Reactions

In Reactivity 3.3, the IB syllabus focuses on two key types of bond-breaking and bond-forming steps that involve shared electrons: homolytic fission and heterolytic fission. These terms describe how the shared electron pair is divided when a covalent bond breaks, and they determine what types of reactive intermediates are produced.

Homolytic Fission

In homolytic fission, the bonding pair splits evenly — one electron goes to each atom. This produces two free radicals, each carrying an unpaired electron. Free radicals are extremely reactive and drive chain reactions such as the combustion of hydrocarbons and the depletion of ozone. Homolytic fission is represented with single-headed curly arrows ("fishhook" arrows), each showing the movement of one electron.

HOMOLYTIC FISSION

A–B → A• + B•

The dot (•) represents an unpaired electron. Each fragment is a radical. Example: Cl₂ → 2 Cl• when UV light provides the activation energy.

Heterolytic Fission

In heterolytic fission, the bonding pair stays together — both electrons go to the more electronegative atom. This produces a positive ion (cation) and a negative ion (anion), or, in organic chemistry, a carbocation and a leaving group. Heterolytic fission is represented with double-headed curly arrows showing the movement of an electron pair.

HETEROLYTIC FISSION

A–B → A⁺ + B⁻ (if B is more electronegative)

Both electrons from the shared pair go to atom B. Example: In the reaction of a haloalkane with OH⁻, the C–Br bond breaks heterolytically: C–Br → C⁺ + Br⁻.

Curly Arrow Notation

The IB expects you to use curly arrows to show electron movement in reaction mechanisms. A full curly arrow (double-headed) shows the movement of an electron pair, and a half curly arrow (single-headed, or fishhook) shows the movement of a single electron. Arrows always start from the electron source (a lone pair or a bond) and point to the electron sink (an electrophilic centre or an atom).

BOND ENTHALPY AND REACTION ENTHALPY

ΔH = Σ(bond enthalpies broken) − Σ(bond enthalpies formed)

Breaking bonds requires energy (positive), forming bonds releases energy (negative). If ΔH < 0 the reaction is exothermic; if ΔH > 0 it is endothermic.

SECTION 5

Classifying Electron Sharing: Single, Double, Triple & Coordinate Bonds

Atoms can share one, two, or three pairs of electrons, forming single, double, or triple bonds respectively. Each additional shared pair increases the bond order, making the bond shorter and stronger. Additionally, when one atom supplies both electrons to the shared pair, a coordinate (dative) bond is formed. The following diagram and table compare these bond types.

This diagram compares four types of electron-sharing bonds. Single bonds share one pair; double bonds share two pairs; triple bonds share three pairs. In a coordinate bond, the donor atom provides both electrons. Notice the arrow pointing from donor to acceptor.

Bond Type	Shared Pairs	Bond Order	Relative Strength	Relative Length	Example
Single (σ)	1	1	Weakest	Longest	H–H (436 kJ mol⁻¹)
Double (σ + π)	2	2	Moderate	Shorter	O=O (498 kJ mol⁻¹)
Triple (σ + 2π)	3	3	Strongest	Shortest	N≡N (945 kJ mol⁻¹)
Coordinate (dative)	1 (from donor)	1	Same as regular σ	Same as regular σ	NH₄⁺ (N→H)

Comparison of covalent bond types

SECTION 6

Worked Example: Free Radical Substitution of Methane

One of the most important electron-sharing reactions in the IB syllabus is the free radical substitution of methane with chlorine. This reaction illustrates homolytic fission, radical intermediates, and the chain-reaction mechanism. Let's walk through it step by step.

Free Radical Substitution: CH₄ + Cl₂ → CH₃Cl + HCl

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Step 1 — Initiation

Ultraviolet light provides the energy needed to break the Cl–Cl bond homolytically. Each chlorine atom retains one electron, producing two chlorine radicals.

Cl₂ → 2 Cl• (homolytic fission, UV light required)
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Step 2 — Propagation (first step)

A chlorine radical attacks a methane molecule, abstracting a hydrogen atom. The Cl• radical takes one electron from the C–H bond, and the hydrogen brings its own electron. This forms HCl and leaves behind a methyl radical (CH₃•).

Cl• + CH₄ → HCl + CH₃•
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Step 3 — Propagation (second step)

The methyl radical now reacts with another Cl₂ molecule. The CH₃• takes one electron from the Cl–Cl bond, forming chloromethane and regenerating a chlorine radical. This new Cl• can repeat Step 2, creating a self-sustaining chain.

CH₃• + Cl₂ → CH₃Cl + Cl•
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Step 4 — Termination

The chain reaction ends when two radicals collide and combine their unpaired electrons to form a covalent bond. There are three possible termination steps depending on which radicals meet.

Cl• + Cl• → Cl₂ | CH₃• + CH₃• → C₂H₆ | CH₃• + Cl• → CH₃Cl
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Step 5 — Overall Equation

The net reaction involves one C–H bond breaking and one C–Cl bond forming. Use bond enthalpies to estimate ΔH: Bonds broken = C–H (414 kJ mol⁻¹) + Cl–Cl (242 kJ mol⁻¹) = 656 kJ mol⁻¹. Bonds formed = C–Cl (324 kJ mol⁻¹) + H–Cl (431 kJ mol⁻¹) = 755 kJ mol⁻¹. ΔH ≈ 656 − 755 = −99 kJ mol⁻¹ (exothermic).

CH₄ + Cl₂ → CH₃Cl + HCl ΔH ≈ −99 kJ mol⁻¹

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IB Exam Tip

When drawing mechanisms for free radical substitution, always use single-headed (fishhook) curly arrows for homolytic steps. Label each step as initiation, propagation, or termination. The IB commonly asks for all three stages.

SECTION 7

Homolytic vs. Heterolytic Fission — Strengths & Limitations

Both homolytic and heterolytic fission are ways of breaking covalent bonds, but they produce different intermediates and occur under different conditions. Understanding when each type is favoured is essential for predicting reaction mechanisms in organic chemistry.

Feature	Homolytic Fission	Heterolytic Fission
Electron division	One electron to each fragment	Both electrons to one fragment
Products formed	Two free radicals (uncharged)	A cation and an anion
Arrow notation	Single-headed (fishhook) arrows	Double-headed curly arrows
Typical conditions	Gas phase, UV light, high temperature	Polar solvents, room temperature
Bond polarity	Favoured in nonpolar bonds	Favoured in polar bonds
Example reactions	Free radical substitution, combustion	Nucleophilic substitution (Sₙ1, Sₙ2)

Comparing the two modes of covalent bond cleavage

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KEY TAKEAWAY

Imagine snapping a wishbone at Thanksgiving. If it breaks perfectly in half — that's homolytic fission; each person gets an equal piece (one electron each). If one person keeps the entire wishbone and the other gets nothing — that's heterolytic fission; one fragment keeps both electrons. The stronger puller (more electronegative atom) wins the pair.

SECTION 8

Connecting to Advanced Bonding Theory

The electron-sharing model you have learned so far relies on Lewis structures and the valence bond approach. At higher levels, chemists use molecular orbital (MO) theory to explain phenomena that Lewis structures cannot capture, such as the paramagnetism of O₂ or the stability of delocalised π systems in benzene. The table below shows how the concepts you've studied connect to more advanced ideas.

Concept (This Lesson)	Advanced Extension
Shared electron pair between two atoms	Bonding molecular orbital (σ or π) formed from constructive overlap of atomic orbitals
Bond order (1, 2, or 3)	MO bond order = ½(bonding e⁻ − antibonding e⁻); can be fractional (e.g., 1.5 in O₂⁻)
Coordinate bond (lone pair donation)	Ligand-to-metal coordination in transition metal complexes (coordination chemistry)
Homolytic / heterolytic fission	Concerted mechanisms (e.g., pericyclic reactions) where bond breaking and forming occur simultaneously
Electronegativity and bond polarity	Dipole moments, partial atomic charges from computational chemistry

From IB-level electron sharing to university-level bonding theory

For now, focus on mastering Lewis structures, curly arrow notation, and the distinction between homolytic and heterolytic fission. These tools are sufficient to explain the mechanisms required at the IB level and provide a strong foundation for studying organic reaction mechanisms in greater depth at university.

SECTION 9

Practice Problems

Question 1: CONCEPTUAL

Explain the difference between homolytic and heterolytic fission. In your answer, describe the type of curly arrow used for each and the products formed.

Show answer

In homolytic fission, the bonding pair is split equally — one electron goes to each atom, producing two free radicals. This is shown with single-headed (fishhook) curly arrows, each representing the movement of one electron. In heterolytic fission, both electrons go to one atom, producing a cation (which lost the electrons) and an anion (which gained them). This is shown with a double-headed curly arrow representing the movement of the electron pair. Homolytic fission is favoured in nonpolar bonds and gaseous conditions, while heterolytic fission is favoured in polar bonds and polar solvents.

Question 2: BASIC CALCULATION

Using the bond enthalpies C–H = 414 kJ mol⁻¹, Cl–Cl = 242 kJ mol⁻¹, C–Cl = 324 kJ mol⁻¹, and H–Cl = 431 kJ mol⁻¹, calculate the enthalpy change for the reaction CH₄ + Cl₂ → CH₃Cl + HCl.

Show answer

Bonds broken: one C–H (414 kJ mol⁻¹) + one Cl–Cl (242 kJ mol⁻¹) = 656 kJ mol⁻¹. Bonds formed: one C–Cl (324 kJ mol⁻¹) + one H–Cl (431 kJ mol⁻¹) = 755 kJ mol⁻¹. Using ΔH = Σ(bonds broken) − Σ(bonds formed) = 656 − 755 = −99 kJ mol⁻¹. The negative sign indicates the reaction is exothermic, meaning it releases 99 kJ of energy per mole of CH₃Cl produced.

Question 3: INTERMEDIATE

Write the initiation, propagation, and termination steps for the free radical substitution of ethane (C₂H₆) with bromine (Br₂) in the presence of UV light. Identify the organic product.

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Initiation: Br₂ → 2 Br• (UV light causes homolytic fission of the Br–Br bond). Propagation step 1: Br• + C₂H₆ → HBr + C₂H₅• (a bromine radical abstracts a hydrogen from ethane, forming an ethyl radical). Propagation step 2: C₂H₅• + Br₂ → C₂H₅Br + Br• (the ethyl radical reacts with another Br₂ molecule, regenerating a bromine radical). Termination: any two radicals combine, e.g. 2 Br• → Br₂, or C₂H₅• + Br• → C₂H₅Br, or 2 C₂H₅• → C₄H₁₀. The main organic product is bromoethane (C₂H₅Br).

Question 4: APPLIED

Chlorofluorocarbons (CFCs) such as CCl₂F₂ damage the ozone layer via a free radical mechanism. The C–Cl bond in CCl₂F₂ has a bond enthalpy of about 318 kJ mol⁻¹, while the C–F bond is about 485 kJ mol⁻¹. Explain why UV radiation breaks the C–Cl bond but not the C–F bond, and write the initiation step for ozone depletion.

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UV radiation in the stratosphere carries enough energy to break the weaker C–Cl bond (318 kJ mol⁻¹) via homolytic fission, but not the stronger C–F bond (485 kJ mol⁻¹). The energy of UV photons in the stratosphere is sufficient to exceed the C–Cl bond enthalpy but falls short of the C–F bond enthalpy. The initiation step is: CCl₂F₂ → •CClF₂ + Cl• (UV). The chlorine radical produced then attacks ozone: Cl• + O₃ → ClO• + O₂. The ClO• radical is regenerated in a subsequent step, allowing one Cl• to destroy thousands of ozone molecules in a catalytic chain.

Question 5: CRITICAL THINKING

When methane reacts with Cl₂ via free radical substitution, the products include not only CH₃Cl but also CH₂Cl₂, CHCl₃, and CCl₄. Explain why multiple substitution products form and suggest one way to maximise the yield of CH₃Cl specifically. Use your understanding of the mechanism to justify your answer.

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Multiple substitution occurs because the product CH₃Cl still contains C–H bonds that can be attacked by chlorine radicals in subsequent propagation steps: Cl• + CH₃Cl → HCl + CH₂Cl•, followed by CH₂Cl• + Cl₂ → CH₂Cl₂ + Cl•, and so on up to CCl₄. Each successive substitution is another chain process. To maximise the yield of the monosubstituted product CH₃Cl, you should use a large excess of methane relative to chlorine. When CH₄ molecules greatly outnumber Cl₂ molecules, a chlorine radical is statistically much more likely to collide with unreacted CH₄ than with the product CH₃Cl. This shifts the product distribution towards CH₃Cl and minimises further substitution. This is a practical example of how controlling reagent ratios can steer a reaction mechanism toward a desired product.

SUMMARY

Summary — Electron Sharing Reactions

Summary

Covalent bonds form when atoms share one or more pairs of electrons to achieve a stable electron configuration. The degree of sharing depends on electronegativity: equal sharing produces nonpolar covalent bonds, unequal sharing produces polar covalent bonds, and complete transfer produces ionic bonds. Coordinate (dative) bonds occur when one atom donates both electrons. Sigma (σ) bonds arise from head-on overlap and pi (π) bonds from sideways overlap; double bonds contain one σ + one π, and triple bonds contain one σ + two π.

When covalent bonds break, homolytic fission splits the pair evenly to produce free radicals (shown with fishhook arrows), while heterolytic fission sends both electrons to one atom to produce ions (shown with double-headed curly arrows). Bond enthalpy values allow you to calculate overall reaction enthalpies using ΔH = Σ(bonds broken) − Σ(bonds formed). These principles underpin reaction mechanisms such as free radical substitution, which proceeds through initiation, propagation, and termination steps.