Alkenes: Structure and Bonding Revision Notes | A Level AQA Chemistry

1. What Makes Alkenes Different?

Alkanes are the workhorses of organic chemistry — stable, relatively unreactive, and saturated with hydrogen. Alkenes are something else entirely. One structural feature — a carbon-carbon double bond — transforms their reactivity and makes them central to industrial chemistry, polymer production, and biological processes.

The key word here is unsaturated. An alkene does not carry the maximum number of hydrogen atoms its carbon skeleton could hold. That missing hydrogen is the clue that a double bond is present — and that double bond is where all the interesting chemistry happens.

Compare ethane ($C_2H_6$) with ethene ($C_2H_4$). Both have two carbons, but ethene has two fewer hydrogen atoms. Those two missing hydrogens are the direct consequence of the double bond — one bond in the C=C pair takes the place of two C-H bonds. This is the defining structural signature of every alkene.

2. The Carbon-Carbon Double Bond: Sigma and Pi

The C=C double bond is not two single bonds side by side. It is made up of two fundamentally different types of bond — a sigma ($\sigma$) bond and a pi ($\pi$) bond — and understanding the difference between them is the foundation of all alkene chemistry at A Level.

Start with what you already know from alkanes: every C-C single bond is a sigma bond, formed by the direct, end-on overlap of atomic orbitals along the line connecting the two nuclei. The sigma bond is the stronger component of the C=C double bond and allows free rotation around the bond axis. In alkenes, the first bond between the two carbons is still a sigma bond — that part is unchanged.

Now for the second bond — the pi bond. Each carbon in the C=C double bond has one unhybridised p orbital, oriented perpendicular to the plane of the molecule. These two p orbitals, one on each carbon, overlap sideways — not end-on — to form the pi bond. The result is two regions of electron density: one above and one below the plane of the molecule.

This sideways overlap is less effective than the end-on overlap of a sigma bond, which is why the pi bond is weaker than the sigma bond. The C=C bond enthalpy (approximately 614 kJ mol$^{-1}$) is greater than a C-C single bond (approximately 347 kJ mol$^{-1}$), but the pi bond alone contributes only the difference — roughly 267 kJ mol$^{-1}$ — which is considerably less than the sigma component.

3. High Electron Density: Why the Double Bond Matters

The AQA specification describes the C=C double bond as a centre of high electron density. This phrase is not decorative — it is the mechanistic explanation for why alkenes react the way they do. The pi electrons are exposed, sitting above and below the molecular plane rather than being shielded between two nuclei as sigma electrons are.

Think of it this way: sigma bond electrons are tucked between the two carbon nuclei, relatively well shielded and harder for an incoming species to access. Pi electrons are out in the open — accessible, mobile, and electron-rich. This makes the C=C double bond an attractive target for electron-deficient species, which are called electrophiles.

This high electron density has two direct consequences that you need to be able to explain. First, the C=C double bond acts as a nucleophile — it donates its pi electrons to electrophiles in reactions. Second, the exposed pi electrons make alkenes susceptible to attack by reagents that would leave alkanes untouched. This is the chemical basis of electrophilic addition, which you will study in detail when you cover alkene reactions.

4. Structural Consequences: Planarity and Restricted Rotation

The pi bond does more than make alkenes reactive — it shapes them. For the sideways overlap of the two p orbitals to occur, those orbitals must remain parallel to each other. This means the two carbon atoms and all four atoms directly attached to them must lie in the same plane. Ethene is a flat molecule, and this planarity is a direct result of the pi bond.

Rotation around the C=C bond is restricted because rotating one carbon relative to the other would break the sideways overlap of the p orbitals, destroying the pi bond. This is fundamentally different from a C-C single bond, around which free rotation occurs at room temperature. The restricted rotation in alkenes is the structural basis for geometric (E/Z) isomerism, which you will encounter when studying isomerism in more depth.

5. Mechanism Overview: How the Pi Bond Reacts

Now that the bonding picture is clear, it is worth seeing how it connects to reactivity — because the AQA specification expects you to link structure to chemical behaviour. The pi bond, as a centre of high electron density, donates its electron pair to an incoming electrophile. This is the first step of electrophilic addition.

Notice that the sigma bond between the two carbons is never broken — it persists throughout the mechanism and is present in the final product. Only the pi bond breaks. This is consistent with the idea that the pi bond is the reactive, exposed component of the double bond, while the sigma bond provides the structural backbone.

6. Exam Technique and Mark-Winning Precision

AQA questions on alkene bonding tend to use command words such as 'describe', 'explain', and 'compare'. Each demands a different type of response, and the mark scheme rewards specific language. The following guidance will help you translate your understanding into marks.

Worked example: Answering an 'Explain' Question on Alkene Bonding

Explain why the C=C double bond in alkenes acts as a centre of high electron density. [3 marks]

1. The C=C double bond consists of one sigma bond (end-on orbital overlap) and one pi bond (sideways overlap of p orbitals). — [B1] Naming both bond types with their overlap descriptions earns the first mark. Do not write 'double bond' — be specific.
2. The pi bond electrons are located above and below the plane of the molecule, rather than between the two nuclei. — [B1] The spatial location of the pi electrons is the key point. 'Above and below the plane' is the AQA-standard phrasing.
3. Because the pi electrons are exposed and not shielded between nuclei, they are accessible to electrophiles, making the C=C a centre of high electron density. — [B1] Connecting the exposed position of pi electrons to the concept of high electron density and electrophilic accessibility completes the explanation.

Answer: The C=C double bond contains a sigma bond (end-on overlap) and a pi bond (sideways overlap of p orbitals). The pi electrons are located above and below the molecular plane, not shielded between the nuclei. This exposed electron density makes the C=C bond a centre of high electron density, accessible to electrophilic attack.

When a question asks you to compare alkenes and alkanes, structure your answer around specific, contrasting points. A vague statement such as 'alkenes are more reactive' will not score marks on its own — you need to link the structural difference (pi bond) to the chemical consequence (electrophilic attack on the exposed pi electrons).

The two specification requirements for this topic — that alkenes are unsaturated hydrocarbons, and that the double bond is a centre of high electron density — are both testable in isolation and in combination. Practise writing answers that use both ideas together, because the highest-mark questions will ask you to connect them to reactivity rather than recall definitions.