# Alcohols

> Primary, secondary and tertiary alcohols, their properties, combustion and measuring the heat of combustion, dehydration, substitution and oxidation, and making ethanol by fermentation or from ethene.

From HSC Chemistry on WhizzNotes, free study notes for NSW students: https://whizznotes.com/chemistry/alcohols/. Updated 3 November 2026.
Year 11 and Year 12 boxes go deeper than the Year 10 text around them.

### Story time: Fuel from sugar cane
In the 1970s, oil prices shot up and Brazil had little oil of its own. So the country turned to something it had plenty of: sugar cane. Yeast turns the sugar into **ethanol**, and Brazil built cars that could run on it.

Today most Brazilian cars can run on petrol, ethanol, or any mix of the two. Ethanol is an **alcohol**, a family of compounds with an −OH group. This chapter shows how alcohols are made, how they burn, and how they can be turned into other useful chemicals.

**In short:** Alcohols are organic compounds with an −OH group, like ethanol. Hydrogen bonding gives them high boiling points, primary alcohols oxidise to aldehydes and then carboxylic acids, secondary alcohols oxidise to ketones, and tertiary alcohols do not oxidise.

## The hydroxyl group

An **alcohol** has an **−OH (hydroxyl) group** bonded to a carbon. The general formula of the alkanols is $\mathrm{C_nH_{2n+1}OH}$.

Note: the −OH in an alcohol is **not** a hydroxide ion. Alcohols are neutral, not basic.

## Primary, secondary and tertiary

Look at the carbon that holds the −OH. Count how many **other carbons** are bonded to it.

| Type | Carbons bonded to the C−OH carbon | Example |
|---|---|---|
| Primary (1°) | 1 (or 0 for methanol) | butan-1-ol $\mathrm{CH_3CH_2CH_2CH_2OH}$ |
| Secondary (2°) | 2 | butan-2-ol $\mathrm{CH_3CH(OH)CH_2CH_3}$ |
| Tertiary (3°) | 3 | 2-methylpropan-2-ol $\mathrm{(CH_3)_3COH}$ |

## Physical properties

- Alcohols form **hydrogen bonds** (O−H). So they have **much higher boiling points** than alkanes of similar mass. Ethanol boils at 78 °C; propane (similar mass) at −42 °C.
- **Small alcohols mix completely with water**, because their −OH hydrogen bonds with water.
- As the carbon chain gets longer, the non-polar part grows, so **solubility falls** and boiling point rises.

## Combustion

Alcohols burn cleanly with a blue flame:

$$\mathrm{C_2H_5OH(l)} + \mathrm{3O_2(g)} \rightarrow \mathrm{2CO_2(g)} + \mathrm{3H_2O(l)}$$

The **molar heat of combustion** $\Delta H_c$ is the heat released when **one mole** burns completely in oxygen. It is always negative.

**Measuring it** with a spirit burner and a can of water:

1. Weigh the spirit burner.
2. Measure the water's mass and starting temperature.
3. Burn the alcohol under the can, stirring, until the temperature rises about 20 °C.
4. Reweigh the burner. The mass lost is the alcohol burnt.

$$q = mc\Delta T \qquad n = \frac{\text{mass burnt}}{MM} \qquad \Delta H_c = -\frac{q}{n}$$

### Worked example: Ethanol in the lab
1.00 g of ethanol (46.07 g/mol) heats 200 g of water by 22.5 °C.

$q = 200 \times 4.18 \times 22.5 = 18\,810$ J $= 18.8$ kJ

$n = 1.00 \div 46.07 = 0.0217$ mol

$\Delta H_c = -\dfrac{18.8}{0.0217} = -867$ kJ/mol

The accepted value is −1367 kJ/mol. Only $867 \div 1367 = 63\%$ of the heat reached the water.

Lab values are low because of **heat lost** to the air and the can, **incomplete combustion** (soot on the can), and alcohol **evaporating** from the wick. Using a draught shield, a lid, and a copper can close to the flame helps.

As alcohols get longer, each extra $\mathrm{CH_2}$ adds about **650 kJ/mol** to $\Delta H_c$.

## Reactions of alcohols

| Reaction | Reagent and conditions | Product |
|---|---|---|
| **Dehydration** (elimination) | Concentrated $\mathrm{H_2SO_4}$, heat | Alkene + water |
| **Substitution** | HCl, HBr (or $\mathrm{PCl_3}$) | Haloalkane + water |
| **Oxidation** | Acidified $\mathrm{K_2Cr_2O_7}$ or $\mathrm{KMnO_4}$, heat | Depends on type |
| **Esterification** | Carboxylic acid, conc. $\mathrm{H_2SO_4}$, reflux | Ester + water |

$$\mathrm{CH_3CH_2OH} \xrightarrow{\text{conc. } \mathrm{H_2SO_4}\text{, heat}} \mathrm{CH_2{=}CH_2} + \mathrm{H_2O}$$

Dehydration is the **reverse of hydration**.

## Oxidation

| Alcohol | Oxidised to | Colour of dichromate | Colour of permanganate |
|---|---|---|---|
| Primary | **Aldehyde**, then **carboxylic acid** | Orange → green | Purple → colourless |
| Secondary | **Ketone** | Orange → green | Purple → colourless |
| Tertiary | **No reaction** | Stays orange | Stays purple |

To stop a primary alcohol at the **aldehyde**, distil it off as it forms. To go all the way to the **carboxylic acid**, heat under **reflux**.

This gives a test: tertiary alcohols do **not** change the colour of acidified dichromate.

## Making ethanol

| | Fermentation | Hydration of ethene |
|---|---|---|
| Raw material | Sugars from plants (renewable) | Ethene from crude oil (non-renewable) |
| Conditions | Yeast, warm (about 35 °C), no oxygen, water | Steam, acid catalyst, 300 °C, high pressure |
| Speed | Slow (batch) | Fast (continuous) |
| Purity | About 15% ethanol; must be distilled | Very pure |

$$\mathrm{C_6H_{12}O_6(aq)} \xrightarrow{\text{yeast}} \mathrm{2C_2H_5OH(aq)} + \mathrm{2CO_2(g)}$$

Fermentation needs **no oxygen**, otherwise yeast makes $\mathrm{CO_2}$ and water instead. The mass of the flask drops as $\mathrm{CO_2}$ escapes, which lets you track the reaction.

## Biofuels

Ethanol made from plants is a **biofuel**.

- **For**: renewable; the $\mathrm{CO_2}$ released was recently taken in by the growing plants; burns more cleanly; can be blended with petrol (E10).
- **Against**: less energy per gram than petrol (about 30 kJ/g compared with 47 kJ/g); uses farmland and water that could grow food; fertilisers, harvesting and distillation use fossil fuel energy; can absorb water and corrode engines.

**Mind map**

```text
Alcohols
  Structure
    −OH hydroxyl
    1°, 2°, 3°
  Properties
    Hydrogen bonding
    High boiling points
    Small ones soluble
  Combustion
    ΔHc per mole
    q = mcΔT
    Heat loss, low results
  Reactions
    Dehydration: alkene
    Substitution: haloalkane
    Oxidation
    Esterification
  Oxidation
    1°: aldehyde, acid
    2°: ketone
    3°: none
  Making ethanol
    Fermentation
    Hydration of ethene
    Biofuel pros and cons
```

### Year 11: Energy per gram
Fuels are often compared by **energy per gram**:

$$\text{kJ/g} = \frac{|\Delta H_c| \text{ in kJ/mol}}{\text{molar mass in g/mol}}$$

Butan-1-ol: $\Delta H_c = -2676$ kJ/mol, molar mass 74.12 g/mol, so $2676 \div 74.12 = 36.1$ kJ/g.

To find how much fuel is needed: 200 g of water heated by 30.0 °C needs $q = 200 \times 4.18 \times 30.0 = 25\,080$ J. With butan-1-ol, $n = 25.08 \div 2676 = 0.00937$ mol, which is $0.00937 \times 74.12 = 0.695$ g (assuming no heat is lost).

### Year 12: Fermentation calculations
**Theoretical yield.** 180 g of glucose (180.16 g/mol) is about 1.00 mol. It makes 2 mol of ethanol: $2 \times 46.07 = 92.1$ g, and 2 mol of $\mathrm{CO_2}$: $2 \times 24.79 = 49.6$ L at 25 °C and 100 kPa.

**Volume percentage.** If the 92.1 g of ethanol (density 0.79 g/mL) is in 750 mL of solution, its volume is $92.1 \div 0.79 = 117$ mL, so the solution is $117 \div 750 \times 100 = 15.6\%$ v/v.

**Distinguishing alcohols with the Lucas test.** Zinc chloride in concentrated HCl turns tertiary alcohols **cloudy immediately**, secondary alcohols cloudy in about 5 minutes, and primary alcohols stay clear at room temperature. The cloudiness is the insoluble haloalkane formed by substitution.

**From mass lost.** The flask gets lighter by the mass of $\mathrm{CO_2}$ that escapes. 2.20 g is $2.20 \div 44.01 = 0.0500$ mol, which is $0.0500 \times 24.79 = 1.24$ L.

**Temperature rise from a fuel.** 1.26 g of methanol (32.04 g/mol) is 0.0393 mol. It releases $0.0393 \times 726 = 28.6$ kJ, so 250 g of water warms by $\dfrac{28\,600}{250 \times 4.18} = 27.3$ °C.

**Products from energy released.** Hexan-1-ol releases 3984 kJ per mole. 2350 kJ means $2350 \div 3984 = 0.590$ mol burnt, making $6 \times 0.590 = 3.54$ mol of $\mathrm{CO_2}$, or 156 g.

**Percentage of heat reaching the water** $= \dfrac{\text{heat the water gained}}{\text{moles burnt} \times |\Delta H_c|} \times 100$.

## Quick summary

- Alcohols have an −OH group. Primary, secondary or tertiary depends on carbons attached to the C−OH carbon.
- Hydrogen bonding gives high boiling points and water solubility for small alcohols.
- $\Delta H_c = -q/n$; lab values are low because of heat loss and incomplete combustion.
- Dehydration (conc. $\mathrm{H_2SO_4}$) gives alkenes. HX gives haloalkanes.
- Oxidation: 1° → aldehyde → acid; 2° → ketone; 3° no reaction (dichromate stays orange).
- Ethanol is made by fermentation (renewable, slow) or hydration of ethene (fast, pure, from oil).

## Quiz

Try these questions from the chapter quiz.

1. The functional group in an alcohol is:
   - A. −OH (hydroxyl)
   - B. −COOH
   - C. −CHO
   - D. $\mathrm{OH^-}$ ion
   - Answer: A. −OH (hydroxyl). Alcohols have a hydroxyl group bonded to carbon.

2. Butan-2-ol is a:
   - A. Primary alcohol
   - B. Secondary alcohol
   - C. Tertiary alcohol
   - D. Carboxylic acid
   - Answer: B. Secondary alcohol. The C−OH carbon is bonded to two other carbons.

3. 2-methylpropan-2-ol is a:
   - A. Primary alcohol
   - B. Secondary alcohol
   - C. Tertiary alcohol
   - D. Ketone
   - Answer: C. Tertiary alcohol. The C−OH carbon is bonded to three other carbons.

The whole chapter quiz is on the page, for signed-in students: https://whizznotes.com/chemistry/alcohols/#quiz