Welcome to the World of Alkanes!

In this chapter, we are diving into the simplest family of organic compounds: Alkanes. Think of alkanes as the "starting point" of organic chemistry. They are the main components of the fuels we use every day, like the petrol in cars and the gas in our stoves. By the end of these notes, you’ll understand how they are built, why they behave the way they do, and how they react with other chemicals.

Don’t worry if organic chemistry seems like a lot of symbols at first! We will break it down piece by piece.

1. What exactly are Alkanes?

Alkanes are saturated hydrocarbons. Let's break that name down:

  • Hydrocarbon: A compound containing only carbon and hydrogen atoms.
  • Saturated: This means all the carbon-carbon bonds are single bonds. The carbon atoms are "full up" and cannot bond to any more atoms.

The General Formula

All alkanes belong to a homologous series (a family with the same functional group where each member differs by a \(CH_2\) group). Because of this, they all follow the same "recipe" or general formula:

\(C_nH_{2n+2}\)

Example: If an alkane has 3 carbons (\(n = 3\)), it must have \((2 \times 3) + 2 = 8\) hydrogens. So, the formula is \(C_3H_8\) (Propane).

Naming and Formulas

You need to know the names of the first ten alkanes. Here is a quick memory trick for the first four:

Monkeys Eat Peeled Bananas = Methane (\(CH_4\)), Ethane (\(C_2H_6\)), Propane (\(C_3H_8\)), Butane (\(C_4H_{10}\)).

From 5 onwards, the names sound like shapes: Pentane (5), Hexane (6), Heptane (7), Octane (8), etc.

Quick Review: Alkanes are saturated (single bonds only) hydrocarbons (C and H only) with the general formula \(C_nH_{2n+2}\).

2. Bonding and Shape

How do these atoms actually stick together? Alkanes use covalent bonds, specifically a type called a \(\sigma\)-bond (pronounced "sigma bond").

What is a \(\sigma\)-bond?

A \(\sigma\)-bond is formed by the direct overlap of orbitals between two atoms. Imagine two people shaking hands—their hands meet directly in the middle. This bond is very strong and allows the atoms to rotate freely. This is why alkane chains can twist and bend into different shapes!

The Tetrahedral Shape

In an alkane, every carbon atom is surrounded by 4 electron pairs (4 bonds). Because these electron pairs are all negatively charged, they push each other away as far as possible. This is called Electron Pair Repulsion Theory.

  • The bonds spread out to the corners of a tetrahedron.
  • The bond angle around each carbon is exactly \(109.5^\circ\).

Analogy: Imagine four balloons tied together at the ends; they will naturally push each other into this 3D tetrahedral shape.

Key Takeaway: Carbon atoms in alkanes have a tetrahedral shape with \(109.5^\circ\) angles because the electron pairs in the \(\sigma\)-bonds repel each other equally.

3. Physical Properties: Boiling Points

Why is Methane a gas at room temperature, but Octane (in petrol) is a liquid? It all comes down to intermolecular forces.

Chain Length

As the carbon chain gets longer, the boiling point increases. This is because:

  1. The molecules have a larger surface area.
  2. There are more surface contacts between molecules.
  3. This creates stronger induced dipole-dipole interactions (also known as London forces).
  4. More energy is needed to overcome these stronger forces.

Branching

If an alkane is branched (like a tree with side groups), its boiling point is lower than a straight-chain alkane with the same number of carbons. This is because:

  • The molecules cannot pack together as closely.
  • There is less surface area for contact.
  • The London forces are weaker, so they are easier to break.

Analogy: Think of straight chains like uncooked spaghetti (they stack perfectly with lots of contact) and branched chains like ping-pong balls (they only touch at tiny points).

Quick Review: Longer chains = higher boiling point. More branching = lower boiling point. It's all about the strength of the London forces.

4. Chemical Reactivity and Combustion

Alkanes are generally unreactive. You could leave an alkane in a beaker with a strong acid or alkali, and usually... nothing would happen!

Why are they so "boring" (unreactive)?

  • High Bond Enthalpy: The \(C-C\) and \(C-H\) \(\sigma\)-bonds are very strong and require a lot of energy to break.
  • Low Polarity: Carbon and Hydrogen have very similar electronegativities, so the bonds are non-polar. There are no "partial charges" to attract other reagents.

Combustion (Burning)

The one thing alkanes do very well is burn! This reaction is exothermic (releases heat).

  1. Complete Combustion: In plenty of oxygen, alkanes burn to produce carbon dioxide (\(CO_2\)) and water (\(H_2O\)).
    Example: \(CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O\)
  2. Incomplete Combustion: If oxygen is limited, we get carbon monoxide (\(CO\)) or even just soot (C).
    Example: \(CH_4 + 1.5O_2 \rightarrow CO + 2H_2O\)

Did you know? Carbon monoxide (\(CO\)) is a "silent killer" because it is colourless, odourless, and toxic. It binds to the haemoglobin in your blood more strongly than oxygen does!

5. Radical Substitution

While alkanes don't like to react, they can be forced to react with halogens (like Chlorine or Bromine) if you hit them with Ultraviolet (UV) radiation. This reaction is called radical substitution.

The Mechanism: Three Steps

A mechanism shows the step-by-step "story" of a reaction. For radical substitution, there are three acts:

Step 1: Initiation

The UV light provides enough energy to break the halogen bond (\(Cl-Cl\)). This is homolytic fission—each atom takes one electron from the shared pair, creating two radicals.

\(Cl_2 \xrightarrow{UV} 2Cl \bullet\)

(Note: A radical is a species with an unpaired electron, represented by a dot \(\bullet\). They are extremely reactive!)

Step 2: Propagation

This is a chain reaction. Radicals are like "hot potatoes"—they keep reacting to pass the unpaired electron along.

1. \(CH_4 + Cl \bullet \rightarrow \bullet CH_3 + HCl\)
2. \(\bullet CH_3 + Cl_2 \rightarrow CH_3Cl + Cl \bullet\)

Notice how we start with a \(Cl \bullet\) radical and end with a new \(Cl \bullet\) radical? The cycle can continue!

Step 3: Termination

The reaction ends when two radicals collide and pair up their electrons. This removes the radicals from the system.

\(Cl \bullet + Cl \bullet \rightarrow Cl_2\)
\(\bullet CH_3 + \bullet CH_3 \rightarrow C_2H_6\)
\(\bullet CH_3 + Cl \bullet \rightarrow CH_3Cl\)

Limitations of this Reaction

This reaction is a bit messy for chemists because:

  • Further Substitution: The product (\(CH_3Cl\)) can be hit by more radicals to become \(CH_2Cl_2\), \(CHCl_3\), or \(CCl_4\). It's hard to get just one product.
  • Substitutions at different positions: If the carbon chain is long (like propane), the halogen could end up on the end carbon or the middle carbon, creating a mixture of isomers.

Key Takeaway: Radical substitution requires UV light and happens in three stages: initiation, propagation, and termination. It often results in a mixture of products.

Final Summary Checklist

Before you move on, make sure you can:

  • Explain that alkanes are saturated hydrocarbons with \(\sigma\)-bonds.
  • State the tetrahedral shape and \(109.5^\circ\) angle.
  • Describe how chain length and branching affect boiling points (London forces).
  • Explain why alkanes are unreactive (bond enthalpy/polarity).
  • Write equations for complete and incomplete combustion.
  • Outline the radical substitution mechanism (Initiation, Propagation, Termination).

Great job! You've mastered the basics of Alkanes. Keep practicing those radical mechanisms—they get easier the more you write them out!