Introduction
The Experiment 10: Double Displacement Reactions is a staple in high‑school and introductory college chemistry labs, designed to illustrate how ionic compounds exchange partners to form new products, often precipitates or gases. Also, understanding the outcomes of each reaction not only reinforces concepts of solubility rules and net‑ionic equations but also hones analytical skills essential for future laboratory work. This article provides complete answers for the ten double‑displacement reactions typically performed in this experiment, explains the underlying reasoning, and offers tips for interpreting results correctly Not complicated — just consistent..
Overview of Double Displacement Reactions
In a double displacement (or metathesis) reaction, two ionic compounds exchange their cations and anions:
[ \text{AB (aq)} + \text{CD (aq)} \rightarrow \text{AD (aq)} + \text{CB (aq)} ]
The reaction proceeds only if at least one of the following occurs:
- A precipitate forms (insoluble product).
- A gas evolves (e.g., CO₂, H₂S, NH₃).
- A weak electrolyte is produced (e.g., water in acid‑base neutralizations).
Applying solubility rules and acid‑base theory allows us to predict which of the four possible products will be observable.
General Procedure Recap
- Label all test tubes (A–J) and prepare the required solutions (usually 0.1 M).
- Mix equal volumes (typically 2 mL each) of the two reactants in each tube.
- Observe immediately for cloudiness, color change, gas bubbles, or temperature shift.
- Record the visible result and write the complete molecular, complete ionic, and net‑ionic equations.
- Confirm precipitates by filtration or gas evolution by odor/lead‑acetate test when appropriate.
Below are the answers for each of the ten reaction pairs, presented in the order commonly used in laboratory manuals.
Reaction 1: Sodium chloride + Silver nitrate
Observation: White, curdy precipitate forms instantly.
Molecular equation:
[ \text{NaCl (aq)} + \text{AgNO}{3}\text{ (aq)} \rightarrow \text{AgCl (s)} + \text{NaNO}{3}\text{ (aq)} ]
Complete ionic equation:
[ \text{Na}^{+}\text{ (aq)} + \text{Cl}^{-}\text{ (aq)} + \text{Ag}^{+}\text{ (aq)} + \text{NO}{3}^{-}\text{ (aq)} \rightarrow \text{AgCl (s)} + \text{Na}^{+}\text{ (aq)} + \text{NO}{3}^{-}\text{ (aq)} ]
Net‑ionic equation:
[ \boxed{\text{Ag}^{+}\text{ (aq)} + \text{Cl}^{-}\text{ (aq)} \rightarrow \text{AgCl (s)}} ]
Why it occurs: Silver chloride is insoluble in water, while sodium nitrate remains dissolved.
Reaction 2: Potassium iodide + Lead(II) nitrate
Observation: Bright yellow precipitate appears.
Molecular equation:
[ \text{2 KI (aq)} + \text{Pb(NO}{3}){2}\text{ (aq)} \rightarrow \text{PbI}{2}\text{ (s)} + 2 \text{KNO}{3}\text{ (aq)} ]
Complete ionic equation:
[ 2\text{K}^{+}\text{ (aq)} + 2\text{I}^{-}\text{ (aq)} + \text{Pb}^{2+}\text{ (aq)} + 2\text{NO}{3}^{-}\text{ (aq)} \rightarrow \text{PbI}{2}\text{ (s)} + 2\text{K}^{+}\text{ (aq)} + 2\text{NO}_{3}^{-}\text{ (aq)} ]
Net‑ionic equation:
[ \boxed{\text{Pb}^{2+}\text{ (aq)} + 2\text{I}^{-}\text{ (aq)} \rightarrow \text{PbI}_{2}\text{ (s)}} ]
Why it occurs: Lead(II) iodide is sparingly soluble, giving the characteristic yellow solid.
Reaction 3: Sodium carbonate + Calcium chloride
Observation: White precipitate forms; solution becomes milky.
Molecular equation:
[ \text{Na}{2}\text{CO}{3}\text{ (aq)} + \text{CaCl}{2}\text{ (aq)} \rightarrow \text{CaCO}{3}\text{ (s)} + 2 \text{NaCl (aq)} ]
Complete ionic equation:
[ 2\text{Na}^{+}\text{ (aq)} + \text{CO}{3}^{2-}\text{ (aq)} + \text{Ca}^{2+}\text{ (aq)} + 2\text{Cl}^{-}\text{ (aq)} \rightarrow \text{CaCO}{3}\text{ (s)} + 2\text{Na}^{+}\text{ (aq)} + 2\text{Cl}^{-}\text{ (aq)} ]
Net‑ionic equation:
[ \boxed{\text{Ca}^{2+}\text{ (aq)} + \text{CO}{3}^{2-}\text{ (aq)} \rightarrow \text{CaCO}{3}\text{ (s)}} ]
Why it occurs: Calcium carbonate is practically insoluble; sodium chloride stays in solution.
Reaction 4: Hydrochloric acid + Sodium bicarbonate
Observation: Vigorous bubbling (CO₂ gas) and fizzing.
Molecular equation:
[ \text{HCl (aq)} + \text{NaHCO}{3}\text{ (aq)} \rightarrow \text{NaCl (aq)} + \text{H}{2}\text{O (l)} + \text{CO}_{2}\text{ (g)} ]
Complete ionic equation:
[ \text{H}^{+}\text{ (aq)} + \text{Cl}^{-}\text{ (aq)} + \text{Na}^{+}\text{ (aq)} + \text{HCO}{3}^{-}\text{ (aq)} \rightarrow \text{Na}^{+}\text{ (aq)} + \text{Cl}^{-}\text{ (aq)} + \text{H}{2}\text{O (l)} + \text{CO}_{2}\text{ (g)} ]
Net‑ionic equation:
[ \boxed{\text{H}^{+}\text{ (aq)} + \text{HCO}{3}^{-}\text{ (aq)} \rightarrow \text{H}{2}\text{O (l)} + \text{CO}_{2}\text{ (g)}} ]
Why it occurs: The reaction is an acid‑base neutralization that produces carbonic acid, which instantly decomposes to water and carbon dioxide gas And that's really what it comes down to. Turns out it matters..
Reaction 5: Sodium sulfate + Barium nitrate
Observation: White, dense precipitate settles rapidly.
Molecular equation:
[ \text{Na}{2}\text{SO}{4}\text{ (aq)} + \text{Ba(NO}{3}){2}\text{ (aq)} \rightarrow \text{BaSO}{4}\text{ (s)} + 2 \text{NaNO}{3}\text{ (aq)} ]
Complete ionic equation:
[ 2\text{Na}^{+}\text{ (aq)} + \text{SO}{4}^{2-}\text{ (aq)} + \text{Ba}^{2+}\text{ (aq)} + 2\text{NO}{3}^{-}\text{ (aq)} \rightarrow \text{BaSO}{4}\text{ (s)} + 2\text{Na}^{+}\text{ (aq)} + 2\text{NO}{3}^{-}\text{ (aq)} ]
Net‑ionic equation:
[ \boxed{\text{Ba}^{2+}\text{ (aq)} + \text{SO}{4}^{2-}\text{ (aq)} \rightarrow \text{BaSO}{4}\text{ (s)}} ]
Why it occurs: Barium sulfate is extremely insoluble, making it a classic test for sulfate ions It's one of those things that adds up..
Reaction 6: Ammonium chloride + Sodium hydroxide
Observation: No visible precipitate; solution feels slightly basic.
Molecular equation:
[ \text{NH}{4}\text{Cl (aq)} + \text{NaOH (aq)} \rightarrow \text{NaCl (aq)} + \text{NH}{3}\text{ (g)} + \text{H}_{2}\text{O (l)} ]
Complete ionic equation:
[ \text{NH}{4}^{+}\text{ (aq)} + \text{Cl}^{-}\text{ (aq)} + \text{Na}^{+}\text{ (aq)} + \text{OH}^{-}\text{ (aq)} \rightarrow \text{Na}^{+}\text{ (aq)} + \text{Cl}^{-}\text{ (aq)} + \text{NH}{3}\text{ (g)} + \text{H}_{2}\text{O (l)} ]
Net‑ionic equation:
[ \boxed{\text{NH}{4}^{+}\text{ (aq)} + \text{OH}^{-}\text{ (aq)} \rightarrow \text{NH}{3}\text{ (g)} + \text{H}_{2}\text{O (l)}} ]
Why it occurs: Hydroxide ions deprotonate ammonium, releasing ammonia gas; the gas may be faintly detected by its characteristic odor or by turning moist red litmus paper blue.
Reaction 7: Potassium chromate + Barium chloride
Observation: Yellow precipitate forms instantly Small thing, real impact..
Molecular equation:
[ \text{K}{2}\text{CrO}{4}\text{ (aq)} + \text{BaCl}{2}\text{ (aq)} \rightarrow \text{BaCrO}{4}\text{ (s)} + 2 \text{KCl (aq)} ]
Complete ionic equation:
[ 2\text{K}^{+}\text{ (aq)} + \text{CrO}{4}^{2-}\text{ (aq)} + \text{Ba}^{2+}\text{ (aq)} + 2\text{Cl}^{-}\text{ (aq)} \rightarrow \text{BaCrO}{4}\text{ (s)} + 2\text{K}^{+}\text{ (aq)} + 2\text{Cl}^{-}\text{ (aq)} ]
Net‑ionic equation:
[ \boxed{\text{Ba}^{2+}\text{ (aq)} + \text{CrO}{4}^{2-}\text{ (aq)} \rightarrow \text{BaCrO}{4}\text{ (s)}} ]
Why it occurs: Barium chromate is sparingly soluble, giving a bright yellow solid useful for identifying chromate ions.
Reaction 8: Hydrochloric acid + Sodium sulfide
Observation: H₂S gas with a rotten‑egg smell is produced; solution becomes acidic.
Molecular equation:
[ 2 \text{HCl (aq)} + \text{Na}{2}\text{S (aq)} \rightarrow 2 \text{NaCl (aq)} + \text{H}{2}\text{S (g)} ]
Complete ionic equation:
[ 2\text{H}^{+}\text{ (aq)} + 2\text{Cl}^{-}\text{ (aq)} + 2\text{Na}^{+}\text{ (aq)} + \text{S}^{2-}\text{ (aq)} \rightarrow 2\text{Na}^{+}\text{ (aq)} + 2\text{Cl}^{-}\text{ (aq)} + \text{H}_{2}\text{S (g)} ]
Net‑ionic equation:
[ \boxed{2\text{H}^{+}\text{ (aq)} + \text{S}^{2-}\text{ (aq)} \rightarrow \text{H}_{2}\text{S (g)}} ]
Why it occurs: Sulfide ions react with protons to form hydrogen sulfide, a toxic gas detectable by its odor or by turning lead(II) acetate paper black.
Reaction 9: Magnesium sulfate + Sodium hydroxide
Observation: No precipitate; solution remains clear Small thing, real impact..
Molecular equation:
[ \text{MgSO}{4}\text{ (aq)} + 2 \text{NaOH (aq)} \rightarrow \text{Mg(OH)}{2}\text{ (s)} + \text{Na}{2}\text{SO}{4}\text{ (aq)} ]
Note: In many textbooks the reaction is listed as “no observable change” because magnesium hydroxide precipitates only slightly at room temperature and may appear as a faint milky suspension. If the solution is warmed, a more noticeable white precipitate forms And that's really what it comes down to..
Complete ionic equation:
[ \text{Mg}^{2+}\text{ (aq)} + \text{SO}{4}^{2-}\text{ (aq)} + 2\text{Na}^{+}\text{ (aq)} + 2\text{OH}^{-}\text{ (aq)} \rightarrow \text{Mg(OH)}{2}\text{ (s)} + 2\text{Na}^{+}\text{ (aq)} + \text{SO}_{4}^{2-}\text{ (aq)} ]
Net‑ionic equation:
[ \boxed{\text{Mg}^{2+}\text{ (aq)} + 2\text{OH}^{-}\text{ (aq)} \rightarrow \text{Mg(OH)}_{2}\text{ (s)}} ]
Why it may appear “no reaction”: The precipitate is colloidal and can redissolve in excess NaOH, giving the impression of a clear solution Practical, not theoretical..
Reaction 10: Potassium chromate + Barium nitrate (acidic medium)
This variation explores the effect of pH on chromate equilibrium.
Observation: In acidic solution, the yellow precipitate disappears, turning the mixture orange‑red due to formation of dichromate Worth keeping that in mind..
Molecular equation (acidic shift):
[ 2 \text{K}{2}\text{CrO}{4}\text{ (aq)} + 2 \text{HNO}{3}\text{ (aq)} \rightarrow \text{K}{2}\text{Cr}{2}\text{O}{7}\text{ (aq)} + 2 \text{KNO}{3}\text{ (aq)} + \text{H}{2}\text{O (l)} ]
Net‑ionic equation (chromate ↔ dichromate):
[ 2\text{CrO}{4}^{2-}\text{ (aq)} + 2\text{H}^{+}\text{ (aq)} \rightarrow \text{Cr}{2}\text{O}{7}^{2-}\text{ (aq)} + \text{H}{2}\text{O (l)} ]
Why it occurs: Adding acid converts the yellow chromate ion to the orange‑red dichromate ion, demonstrating Le Chatelier’s principle. No solid precipitate forms unless the solution is later neutralized, allowing BaCrO₄ to reappear.
Scientific Explanation Behind the Patterns
Solubility Rules in Action
The experiment repeatedly confirms classic solubility guidelines:
| Ion Pair | Typical Solubility | Resulting Observation |
|---|---|---|
| Ag⁺ + Cl⁻ | Insoluble (AgCl) | White precipitate |
| Pb²⁺ + I⁻ | Insoluble (PbI₂) | Yellow precipitate |
| Ca²⁺ + CO₃²⁻ | Insoluble (CaCO₃) | Milky white solid |
| Ba²⁺ + SO₄²⁻ | Insoluble (BaSO₄) | Dense white solid |
| Ba²⁺ + CrO₄²⁻ | Insoluble (BaCrO₄) | Yellow solid |
| NH₄⁺ + OH⁻ | Forms gas (NH₃) | Faint odor, basic shift |
Understanding these patterns enables quick prediction of reaction outcomes without performing the experiment.
Acid‑Base Neutralization & Gas Evolution
Reactions 4, 6, and 8 illustrate proton transfer leading to gaseous products (CO₂, NH₃, H₂S). The net‑ionic equations strip away spectator ions, emphasizing the acid‑base nature of the process:
- H⁺ + HCO₃⁻ → H₂O + CO₂ (carbonic acid decomposition).
- NH₄⁺ + OH⁻ → NH₃ + H₂O (weak base formation).
- 2H⁺ + S²⁻ → H₂S (formation of a volatile sulfide).
These are textbook examples of neutralization where the weak conjugate base or acid drives gas release And it works..
Influence of pH on Chromate/Dichromate Equilibrium
Reaction 10 showcases a dynamic equilibrium:
[ \text{CrO}{4}^{2-} + \text{H}^{+} \leftrightarrow \text{HCrO}{4}^{-} \leftrightarrow \text{Cr}{2}\text{O}{7}^{2-} + \text{H}_{2}\text{O} ]
In neutral or basic media, the yellow chromate dominates; under acidic conditions, the equilibrium shifts to the orange‑red dichromate. This visual cue reinforces the concept that pH can change both solubility and color of transition‑metal oxyanions.
Frequently Asked Questions (FAQ)
Q1: How can I be sure a precipitate is truly insoluble and not just fine particles?
A: Filter a small sample of the mixture through a pre‑weighed filter paper, dry the residue, and weigh it. A consistent mass after washing indicates an insoluble product. Alternatively, add excess of a soluble ion that would dissolve a partially soluble salt; if the solid persists, it is truly insoluble And that's really what it comes down to..
Q2: Why does Reaction 9 sometimes appear to have “no change”?
A: Magnesium hydroxide has a low Ksp (≈5.6 × 10⁻¹²), precipitating only when the product of [Mg²⁺][OH⁻]² exceeds this value. At modest concentrations, the precipitate is colloidal and may stay suspended, giving a clear appearance. Heating or adding more OH⁻ pushes the equilibrium toward solid formation Most people skip this — try not to..
Q3: Are there safety concerns with the gases produced?
A: Yes. CO₂ is non‑toxic at lab scale but can displace oxygen in confined spaces. NH₃ is irritating to eyes and respiratory tract; work under a fume hood. H₂S is toxic and flammable; always use a hood, wear a mask, and have a H₂S detector if possible Simple as that..
Q4: Can I use any concentration of reactants?
A: The classic experiment uses 0.1 M solutions to balance visibility of precipitates with manageable gas evolution. Higher concentrations increase reaction rates and precipitate mass but may cause vigorous bubbling or splattering.
Q5: How do I write the net‑ionic equation correctly?
A: 1) Write the balanced molecular equation. 2) Split all soluble strong electrolytes into their constituent ions. 3) Cancel identical ions appearing on both sides (spectator ions). 4) The remaining species constitute the net‑ionic equation But it adds up..
Tips for Successful Lab Execution
- Label meticulously – Mislabeling leads to confusion when interpreting results.
- Use clean glassware – Residual ions from previous experiments can create false precipitates.
- Observe quickly – Some precipitates form within seconds; delayed observation may miss transient gases.
- Record temperature – Exothermic or endothermic reactions (e.g., neutralizations) can affect solubility.
- Perform a control – Mix each reactant with distilled water to confirm that any observed change is due to the partner ion, not an impurity.
Conclusion
The ten double‑displacement reactions in Experiment 10 provide a comprehensive showcase of solubility rules, acid‑base neutralizations, gas evolution, and equilibrium shifts. Because of that, by mastering the molecular, complete ionic, and net‑ionic equations, students gain a deeper appreciation for how ions interact in aqueous media. Because of that, the systematic approach outlined above—observing, recording, and rationalizing each outcome—builds a solid foundation for more advanced topics such as precipitation analysis, qualitative ion testing, and coordination chemistry. Armed with these answers and the underlying scientific principles, learners can approach future laboratory work with confidence and analytical precision.
