# Types of Chemical Reactions

> Synthesis, decomposition, single and double displacement, the activity series and predicting precipitates.

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

### Story time: The silver tree
In Mr Patel's class, Ruby dropped a coiled copper wire into a clear, colourless solution of silver nitrate and left it overnight.

The next morning the class crowded around the beaker. The copper wire was covered in sparkling, feathery silver crystals, like a tiny tree covered in frost. And the solution had turned pale **blue**.

"The copper pushed the silver out," explained Mr Patel. "Copper is more reactive than silver, so it takes silver's place in the compound. The silver comes out as metal, and copper ions go into the water and turn it blue."

That is a **single displacement** reaction. Chemists sort thousands of reactions into a handful of types like this one. Once you know the type, you can predict the products before you even start.

**In short:** Most reactions fit a few types: synthesis (A + B → AB), decomposition (AB → A + B), single displacement, double displacement and combustion. A single displacement only works if the element is higher in the activity series, and solubility rules predict when a double displacement forms a precipitate.

## Five main types

| Type | Pattern | Example |
|---|---|---|
| **Synthesis** (combination) | $\mathrm{A} + \mathrm{B} \rightarrow \mathrm{AB}$ | $\mathrm{2Mg} + \mathrm{O_2} \rightarrow \mathrm{2MgO}$ |
| **Decomposition** | $\mathrm{AB} \rightarrow \mathrm{A} + \mathrm{B}$ | $\mathrm{2H_2O} \rightarrow \mathrm{2H_2} + \mathrm{O_2}$ |
| **Single displacement** | $\mathrm{A} + \mathrm{BC} \rightarrow \mathrm{AC} + \mathrm{B}$ | $\mathrm{Zn} + \mathrm{CuSO_4} \rightarrow \mathrm{ZnSO_4} + \mathrm{Cu}$ |
| **Double displacement** | $\mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB}$ | $\mathrm{AgNO_3} + \mathrm{NaCl} \rightarrow \mathrm{AgCl} + \mathrm{NaNO_3}$ |
| **Combustion** | Fuel + $\mathrm{O_2}$ → oxides | $\mathrm{CH_4} + \mathrm{2O_2} \rightarrow \mathrm{CO_2} + \mathrm{2H_2O}$ |

Combustion has its own chapter next. This chapter looks at the other four.

### Exam tip: Spot the type by counting
- **Two or more in, one out**: synthesis.
- **One in, two or more out**: decomposition.
- **An element + a compound**: single displacement.
- **Two compounds swap partners**: double displacement.
- **Something + oxygen, with heat and flame**: combustion.

## Synthesis

Two or more substances join to make **one** product.

$$\mathrm{Fe(s)} + \mathrm{S(s)} \rightarrow \mathrm{FeS(s)}$$
$$\mathrm{2Na(s)} + \mathrm{Cl_2(g)} \rightarrow \mathrm{2NaCl(s)}$$

Metal + non-metal gives an ionic compound. Metal + oxygen gives a metal oxide.

## Decomposition

One compound **breaks down** into two or more simpler substances. It usually needs energy: heat (**thermal decomposition**), electricity (**electrolysis**) or light.

$$\mathrm{CaCO_3(s)} \rightarrow \mathrm{CaO(s)} + \mathrm{CO_2(g)}$$

Heating limestone (calcium carbonate) makes lime for cement. Passing electricity through water (electrolysis) splits it into hydrogen and oxygen.

Hydrogen peroxide slowly decomposes in sunlight, which is why it is sold in brown bottles:

$$\mathrm{2H_2O_2(l)} \rightarrow \mathrm{2H_2O(l)} + \mathrm{O_2(g)}$$

## Single displacement and the activity series

A **more reactive** element **displaces** (pushes out) a **less reactive** element from its compound.

The **activity series** ranks metals from most to least reactive:

$$\text{K} > \text{Na} > \text{Ca} > \text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Sn} > \text{Pb} > (\text{H}) > \text{Cu} > \text{Ag} > \text{Au}$$

### Key idea: The rule
A metal can only displace a metal (or hydrogen) **below** it in the activity series.
- Zinc in copper sulfate solution: zinc is higher, so it **reacts**. $\mathrm{Zn} + \mathrm{CuSO_4} \rightarrow \mathrm{ZnSO_4} + \mathrm{Cu}$
- Copper in zinc sulfate solution: copper is lower, so there is **no reaction**. Write $\mathrm{Cu} + \mathrm{ZnSO_4} \rightarrow$ no reaction.

Metals **above hydrogen** react with acids to release hydrogen gas:

$$\mathrm{Mg(s)} + \mathrm{2HCl(aq)} \rightarrow \mathrm{MgCl_2(aq)} + \mathrm{H_2(g)}$$

Copper, silver and gold are below hydrogen, so they do not react with dilute acids. That is why they are used for coins and jewellery.

**Halogens** follow the same idea, with their own order: $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$. Chlorine displaces bromine from potassium bromide, but iodine cannot displace chlorine.

## Double displacement

Two ionic compounds in solution **swap partners**. The positive ion of one pairs with the negative ion of the other.

$$\mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB}$$

A double displacement only happens if one of the products is:

1. A **precipitate** (an insoluble solid that appears in the solution), or
2. **Water** (a neutralisation), or
3. A **gas**.

### Worked example: Making a precipitate
Mix lead(II) nitrate and potassium iodide solutions. The ions swap:

$$\mathrm{Pb(NO_3)_2(aq)} + \mathrm{2KI(aq)} \rightarrow \mathrm{PbI_2(s)} + \mathrm{2KNO_3(aq)}$$

Lead iodide is insoluble, so a bright yellow solid appears. Potassium nitrate stays dissolved.

To predict whether a product is a precipitate, use **solubility rules**:

| Always soluble | Mostly insoluble |
|---|---|
| All sodium, potassium and ammonium compounds | Most carbonates, phosphates and hydroxides (except with Na, K, $\mathrm{NH_4}$) |
| All nitrates | Silver chloride, silver bromide, silver iodide |
| Most chlorides, bromides and iodides (except with Ag and Pb) | Lead chloride, lead iodide |
| Most sulfates (except barium, lead and calcium sulfate) | Barium sulfate |

**Neutralisation** is a double displacement that makes water:

$$\mathrm{HCl(aq)} + \mathrm{NaOH(aq)} \rightarrow \mathrm{NaCl(aq)} + \mathrm{H_2O(l)}$$

### Watch out: Single or double?
In a **single** displacement, one reactant is a lone **element** (like Zn or $\mathrm{Cl_2}$). In a **double** displacement, **both** reactants are **compounds**.

**Mind map**

```text
Types of reactions
  Synthesis
    A + B → AB
  Decomposition
    AB → A + B
    Heat, electricity, light
  Single displacement
    A + BC → AC + B
    Activity series
    Metal must be higher
  Double displacement
    AB + CD → AD + CB
    Needs precipitate, water or gas
    Solubility rules
  Combustion
    Fuel + oxygen
```

### Year 11: Ionic equations and spectator ions
In a double displacement, many ions do nothing. Write all dissolved ionic compounds as separate ions, then cancel the ions that appear unchanged on both sides (**spectator ions**).

For lead iodide:

$$\mathrm{Pb^{2+}} + \mathrm{2NO_3^-} + \mathrm{2K^+} + \mathrm{2I^-} \rightarrow \mathrm{PbI_2(s)} + \mathrm{2K^+} + \mathrm{2NO_3^-}$$

$\mathrm{K^+}$ and $\mathrm{NO_3^-}$ cancel, leaving the **net ionic equation**:

$$\mathrm{Pb^{2+}(aq)} + \mathrm{2I^-(aq)} \rightarrow \mathrm{PbI_2(s)}$$

Any soluble lead compound mixed with any soluble iodide gives this same reaction.

### Year 12: Displacement as redox
A single displacement is a **redox** reaction: one substance loses electrons (**oxidation**) and another gains them (**reduction**). "OIL RIG": Oxidation Is Loss, Reduction Is Gain.

In $\mathrm{Zn(s)} + \mathrm{Cu^{2+}(aq)} \rightarrow \mathrm{Zn^{2+}(aq)} + \mathrm{Cu(s)}$:

- Oxidation half: $\mathrm{Zn} \rightarrow \mathrm{Zn^{2+}} + \mathrm{2e^-}$
- Reduction half: $\mathrm{Cu^{2+}} + \mathrm{2e^-} \rightarrow \mathrm{Cu}$

The activity series is really a list of how easily metals lose electrons. Putting this reaction in two separate beakers, joined by a wire, makes a **galvanic cell**: a battery.

## Quick summary

- **Synthesis**: A + B → AB. **Decomposition**: AB → A + B.
- **Single displacement**: an element replaces another in a compound. It only works if the element is **higher in the activity series**.
- **Double displacement**: two compounds swap ions. It needs a precipitate, water or a gas to form.
- Use **solubility rules** to predict precipitates. Sodium, potassium, ammonium and nitrate compounds always dissolve.
- **Combustion**: fuel + oxygen gives oxides, heat and light.

## Quiz

Try these questions from the chapter quiz.

1. What type of reaction is $\mathrm{2H_2} + \mathrm{O_2} \rightarrow \mathrm{2H_2O}$?
   - A. Synthesis
   - B. Decomposition
   - C. Single displacement
   - D. Double displacement
   - Answer: A. Synthesis. Two substances join to make one product.

2. What type of reaction is $\mathrm{2KClO_3} \rightarrow \mathrm{2KCl} + \mathrm{3O_2}$?
   - A. Synthesis
   - B. Decomposition
   - C. Single displacement
   - D. Combustion
   - Answer: B. Decomposition. One compound breaks into two simpler substances.

3. $\mathrm{Fe} + \mathrm{CuSO_4} \rightarrow \mathrm{FeSO_4} + \mathrm{Cu}$ is a:
   - A. Synthesis reaction
   - B. Decomposition reaction
   - C. Single displacement reaction
   - D. Double displacement reaction
   - Answer: C. Single displacement reaction. Iron, an element, replaces copper in the compound.

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