A Visual Introduction to Ionic and Net Ionic Equations
A Carolina Essentials™ Activity
Total Time: 75-90 mins
Prep: 30 mins | Activity: 45-60 mins
Physical Science, Chemistry
- Total Time: 75-90 minutes [ Prep: 30 mins | Activity: 45-60 mins ]
- Subject: Physical Science, Chemistry
- Grade: High School
Overview
This activity explores the phenomenon of chemical precipitation and asks students to construct an atomic level model of precipitation using ionic and net ionic equations. Initially, students run a series of reactions. Some reactions produce precipitates and some don’t. Using only the reactions that produce precipitates, students then write ionic equations, cross out spectator ions, and conclude with the net ionic equation. The activity may be used as an extension to the reference kit, Mystery Chemical Reactions or as a stand-alone visual introduction to precipitates, solubility, ionic and net ionic equations.
Essential Question
How is the phenomenon of precipitate formation explained on the atomic level?
Activity Objectives
- Identify reactions that produce precipitates.
- Model the formation of a precipitate using ionic and net ionic equations.
Next Generation Science Standards* (NGSS)
HS-PS1-7. Use mathematical representations to support the claim that atoms, and therefore mass, are conserved during a chemical reaction.
Analyzing and Interpreting Data
- Use mathematical representations of phenomena to support claims.
PS1.B Chemical Reactions
- The fact that atoms are conserved, together with knowledge of the chemical properties of the elements involved, can be used to describe and predict chemical reactions.
Patterns
- The total amount of energy and matter in closed systems is conserved.
- Macroscopic patterns are related to the nature of microscopic and atomic-level structure.
Materials
- 8 Dropper bottles per group
- 120 mL of 0.1 M sodium phosphate Na3PO4
- 120 mL of 0.1 M copper(II) sulfate CuSO4
- 120 mL of 0.1 M potassium iodide KI
- 120 mL of 0.1 M lead(II) nitrate Pb(NO3)2
- 120 mL of 0.1 M sodium carbonate Na2CO3
- 120 mL of 0.1 M silver nitrate AgNO3
- 120 mL of 0.1 M calcium chloride CaCl2
- 120 mL of 0.1 M sodium hydroxide NaOH
- 1 Reaction Mat per group (Reaction Mat template below) or 1 Spot plate/ microplate per group
- Absorbent paper towels
Safety & Disposal
Use safety goggles, gloves, and apron. Wash hands with soap and water when finished. Use this kit only in accordance with established laboratory safety practices, including appropriate personal protective equipment (PPE) such as gloves, chemical splash goggles, and lab coats or aprons. Ensure that students understand and adhere to these practices. Students should not eat, drink, or chew gum in the lab and should wash their hands before or after entering and exiting the lab. Because they might irritate or burn the skin, avoid contact with the dilute solutions in this lab. If contact occurs, flush the affected area with water.
Solutions in dropper bottles may be stored for additional classes or additional activities. Know and follow all federal, state, and local regulations as well as school district guidelines for the disposal of laboratory wastes.
Procedures
- Prepare dropper bottles of 0.1 M test chemicals. (See the Solution Preparation Resource if making solutions from stock chemicals) Label every bottle with the formula, concentration, and date prepared. Dropper bottles can be prepared the day before and stored for up to a year.
- If using the reaction mat, make the needed number of copies of the template on copier transparency film.
- At the central materials station, get 1 dropper bottle of each of the 8 chemicals listed in the materials list.
- If using a Reaction Mat transparency, place the Reaction Mat transparency on top of the Known Reaction Grid, so that the boxes line up. If using spot plates, position the spot plates to align with an ordered list of reactants.
- Place one drop of sodium phosphate in the reaction square or well in the top, left corner of the grid or spot plate. Then, add one drop of copper(II) sulfate to that reaction square or well. Do not let the dropper bottle tip touch the drop of the chemical you have already placed in the block or in the well.
- Record observations.
- Add the remaining chemicals, one at a time, to the columns and rows of reaction squares, as the chemicals are listed.
- Record observations after each reaction.
- If chemicals are placed in the wrong reaction square or well, use a rolled-up piece of absorbent paper towel to remove the chemical.
- Chemicals may be made early and stored in capped dropper bottles.
- Place a lab set of dropper bottles in a small basket for easy pick-up by students.
- Make sure students are keeping the drops within the reaction area on the mat or in the well on the spot plate so no secondary reactions occur.
- Remind students to look for cloudiness or particles. For this experiment color change or bubbles are not indicative of a precipitate.
- *Even though students are using small amounts of chemicals they should still wear goggles, gloves, and aprons.
Data and Observations
*If the silver nitrate solution is old and the silver appears tarnished, silver precipitate colors may vary slightly.
Analysis & Discussion
Identify all reactions that produced a precipitate and write a balanced chemical reaction to model the bonds being broken and reformed at the atomic level and conservation of matter.
See below. Note the (ppt) designation may be written as (s)
What evidence do you have that all the reactants are soluble?
All the reactant solutions were transparent and showed no signs of the reactant falling out of solution as a solid.
Using the chart of Solubility Rules above, identify the product that is the precipitate and place a (ppt) or (s) to the right of the product formula. Place an (aq) to the right of all chemicals that are soluble.
Convert the balanced chemical equation to an ionic equation to model the process of dissociation. Split apart all soluble chemicals into a cation and anion. Show the charge on the ion. If needed, change the coefficient to reflect the total number of ions in solution.
For example:
2 Ca(NO3)2 (aq)→ 2Ca2+ (aq) + 4NO3-(aq)
With a single line, mark out the spectator ions with the coefficients. Write the net ionic equation modeling the formation of the precipitate on the atomic level. Make certain it is balanced to illustrate conservation of matter. See above for the strikethroughs.
Use an ionic and net ionic equation to explain why equations are not written for reactions in this activity that do not produce precipitates
For reactions in this activity that do not produce a precipitate, all ions remain in the aqueous state. No molecular compounds are formed either. Consequently, all ions get canceled and there is no net ionic equation that can be written.
An example equation follows:
NaOH(aq) + KI(aq) NaI(aq) + KOH(aq)
Na1+(aq) + OH1-(aq) + K1+(aq) + I1-(aq) Na1+(aq) + I1-(aq) + K1+(aq) + OH1-(aq)
No net ionic equation possible. All ions are spectator ions and are crossed out.
- Safety Apron Required
- Safety Goggles Required
- Safety Gloves Required
*Next Generation Science Standards® is a registered trademark of Achieve. Neither Achieve nor the lead states and partners that developed the Next Generation Science Standards were involved in the production of, and do not endorse, these products.
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