Single Replacement Reactions- Net Ionic Equations
What Are Single Replacement Reactions?
Single replacement reactions (also called single displacement reactions) occur when one element replaces another element in a compound. One reactant is an element; the other is a compound. The result is a new element and a new compound.
The general form is:
A + BC → AC + B
Element A replaces element B in the compound BC. Whether this happens depends on the reactivity of the elements involved.
The Three Types of Single Replacement Reactions
1. Metal Replaces Metal
When a more reactive metal displaces a less reactive metal from its compound.
Example: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Zinc is more reactive than copper, so zinc takes copper's place in the sulfate compound. Copper metal drops out as a solid.
2. Metal Replaces Hydrogen
Very reactive metals can displace hydrogen from acids or water.
With acid: Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)
With water (for very reactive metals like sodium): 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
The hydrogen gas bubbles out of solution.
3. Halogen Replaces Halogen
More reactive halogens displace less reactive halogens from their compounds.
Example: Cl₂(g) + 2NaBr(aq) → 2NaCl(aq) + Br₂(l)
Chlorine is more reactive than bromine, so chlorine takes bromine's place in the sodium bromide compound.
Activity Series: The Reactivity Hierarchy
You need a reactivity series (or activity series) to predict if a single replacement reaction will occur. The metal or halogen higher on the list displaces the one lower on the list.
Common Metal Activity Series (Most Reactive to Least Reactive)
- Lithium (Li)
- Potassium (K)
- Calcium (Ca)
- Sodium (Na)
- Magnesium (Mg)
- Aluminum (Al)
- Zinc (Zn)
- Iron (Fe)
- Nickel (Ni)
- Lead (Pb)
- Copper (Cu)
- Silver (Ag)
- Gold (Au)
Metals above hydrogen can displace it from acids. Metals below hydrogen cannot.
Halogen Activity Series
- Fluorine (F₂) — most reactive
- Chlorine (Cl₂)
- Bromine (Br₂)
- Iodine (I₂) — least reactive
If the replacing element is lower on the list than the element it's trying to replace, no reaction occurs. Don't force it.
Writing and Balancing Single Replacement Equations
Balancing these equations follows the same rules as any chemical equation. Count atoms on both sides. Adjust coefficients until they match.
Unbalanced: Al + Fe₂O₃ → Al₂O₃ + Fe
Count: Left side has 1 Al, 2 Fe, 3 O. Right side has 2 Al, 1 Fe, 3 O.
Balanced: 2Al + Fe₂O₃ → Al₂O₃ + 2Fe
Now both sides have 2 Al, 2 Fe, 3 O. Balanced.
What Are Net Ionic Equations?
Net ionic equations show only the species that actually participate in the reaction. They omit spectator ions — ions that are present but don't change during the reaction.
This matters because many reactions happen in solution. Ionic compounds dissolve in water, splitting into their component ions. Most of those ions just float around doing nothing. Net ionic equations filter out the noise.
Key Terms You Need
Complete ionic equation: Shows all dissolved ions explicitly. Every aqueous compound breaks apart.
Spectator ions: Ions that appear on both sides of the complete ionic equation. They don't participate.
Net ionic equation: The complete ionic equation minus the spectator ions. Shows only what actually reacts.
How to Write Net Ionic Equations
Here's the step-by-step process:
Step 1: Write the Balanced Molecular Equation
Start with the standard chemical equation.
Example: NaCl(aq) + AgNO₃(aq) → AgCl(s) + NaNO₃(aq)
Step 2: Write the Complete Ionic Equation
Break apart all aqueous compounds into their ions. Keep solids, liquids, and gases intact.
Na⁺(aq) + Cl⁻(aq) + Ag⁺(aq) + NO₃⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)
Step 3: Identify and Remove Spectator Ions
Spectator ions appear on both sides unchanged. Here, Na⁺ and NO₃⁻ appear on both sides. Remove them.
Step 4: Write the Net Ionic Equation
What's left is the actual reaction:
Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
Silver ions and chloride ions combine to form solid silver chloride. That's what actually happened.
Common Single Replacement Reactions and Their Net Ionic Forms
| Molecular Equation | Net Ionic Equation |
|---|---|
| Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) | Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) |
| 2Na(s) + 2HCl(aq) → 2NaCl(aq) + H₂(g) | 2Na(s) + 2H⁺(aq) → 2Na⁺(aq) + H₂(g) |
| Cl₂(g) + 2NaBr(aq) → 2NaCl(aq) + Br₂(l) | Cl₂(g) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(l) |
| Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) | Mg(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂(g) |
Getting Started: Practice Problems
Try writing the net ionic equation for this reaction:
Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
Break it down:
- Step 1: Identify aqueous compounds. All except PbI₂ are aqueous.
- Step 2: Write the complete ionic equation with all ions separated.
- Step 3: Find spectators — ions appearing on both sides.
- Step 4: Remove spectators and write the net ionic equation.
Answer: Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s)
The potassium and nitrate ions did nothing. Lead and iodide ions formed the precipitate.
Solubility Rules: Knowing What Forms a Solid
You need to know which compounds are soluble (dissolve in water) and which are insoluble (form solids) to write accurate net ionic equations.
Generally soluble:
- Compounds with Na⁺, K⁺, NH₄⁺
- Nitrates (NO₃⁻)
- Most chlorides, bromides, iodides
- Most sulfates (except BaSO₄, PbSO₄, CaSO₄)
Generally insoluble (form solids):
- Carbonates (CO₃²⁻) — except with Na⁺, K⁺, NH₄⁺
- Hydroxides (OH⁻) — except with Na⁺, K⁺, Ca²⁺, Ba²⁺
- Sulfides (S²⁻) — except with Na⁺, K⁺, NH₄⁺, Ca²⁺, Mg²⁺, Ba²⁺
- Phosphates (PO₄³⁻) — except with Na⁺, K⁺, NH₄⁺
- Silver salts (AgCl, AgBr, AgI)
- Lead salts (PbCl₂, PbBr₂, PbI₂, PbSO₄)
If a product is insoluble, it stays as a solid in the net ionic equation. If it's soluble, it breaks into ions.
Why Net Ionic Equations Matter
Net ionic equations reveal the actual chemical change. They strip away the ions that are just passengers.
In the reaction between silver nitrate and sodium chloride, you might think sodium nitrate formed. But the net ionic equation shows silver chloride is the actual product. Sodium and nitrate ions were just along for the ride.
This matters in:
- Precipitation reactions — identifying which ions form solids
- Acid-base reactions — identifying H⁺ and OH⁻ neutralization
- Laboratory work — understanding what's actually in your beaker
Quick Reference: Reaction Type Comparison
| Reaction Type | General Form | Example |
|---|---|---|
| Single Replacement (metal) | A + BC → AC + B | Zn + CuSO₄ → ZnSO₄ + Cu |
| Single Replacement (halogen) | A₂ + BC → BA + CA | Cl₂ + 2NaBr → 2NaCl + Br₂ |
| Double Replacement | AB + CD → AD + CB | AgNO₃ + NaCl → AgCl + NaNO₃ |
| Synthesis | A + B → AB | 2H₂ + O₂ → 2H₂O |
| Decomposition | AB → A + B | 2H₂O → 2H₂ + O₂ |
Single replacement reactions involve one element and one compound. Double replacement involves two compounds swapping parts. Don't mix them up.
The Bottom Line
Single replacement reactions happen when one element kicks another out of a compound. Check the activity series to see if the reaction is possible. Balance the equation. Then convert to net ionic form by removing spectator ions.
The net ionic equation shows the actual chemistry — what's changing and what's forming. Everything else is just background noise.