Calcium carbonate (marble chips) reacts with dilute hydrochloric acid to produce carbon dioxide, water, and calcium chloride. Using equal acid volume and equal marble mass but different chip sizes, students measure carbon dioxide evolved (by mass loss) and compare initial rates to show that greater surface area gives a faster reaction.

1. Gather a conical flask, top-loading balance, dilute hydrochloric acid of fixed concentration, three marble chip size ranges (same total mass for each run), cotton wool, timer, and graph paper or spreadsheet.
2. Place a measured volume of the acid in the flask and set the flask on the balance; insert a loose cotton-wool plug to reduce spray while allowing gas to escape. Tare the balance to 0.00 g.
3. Start the timer, quickly add the preweighed marble chips (smallest size first), replace the cotton plug, and record mass every 5–10 s until the mass change per interval becomes very small.
4. Repeat the run with the same acid volume and concentration, the same marble mass, and the same temperature, but with medium-sized chips, then large chips.
5. Plot “mass of CO2 produced” versus time (mass loss from 0.00 g). Draw a best-fit curve for each chip size.
6. Note that with the same acid volume and concentration, all runs should approach the same total CO2 at long times, even though the initial rates differ.

The Effect of Surface Area on Reaction Rate - Rugby School Chemistry:

Effect of Surface Area on the Rate of Reaction | Chemistry Practicals - Science with Hazel:

📄 The effect of surface area on the rates of chemical reactions - Chemguide: Core Chemistry: [https://www.chemguide.uk/14to16/rates/surface.html]]

• Collect CO2 with a gas syringe or inverted buret and compare initial volumes per second instead of mass loss.
• Hold surface area constant and vary acid concentration to compare how concentration and surface area each affect initial rate.
• Use powdered CaCO3 versus large lumps to illustrate an extreme surface area change; shorten sampling intervals for powder.
• Investigate stirring or gentle swirling as a mass-transfer variable; keep all other conditions constant.

• Wear splash goggles, lab coat, and appropriate gloves; tie back hair and secure loose clothing.
• Hydrochloric acid is corrosive; avoid skin and eye contact and rinse spills with plenty of water.
• Do not seal the flask; CO2 must vent freely. Use only a loose cotton-wool plug to limit spray.
• Handle marble dust carefully; avoid inhalation if using finely divided CaCO3.
• Keep glassware stable on the balance; protect the balance with a tray in case of spills.
• Wash hands after the experiment and remove gloves before touching shared equipment.

• Why do smaller chips react faster than larger chips of the same total mass? (They provide more surface area, so more CaCO3 particles are exposed for collisions with H+ at any moment.)
• Why do all curves level off at roughly the same total CO2 for a fixed acid volume and concentration? (The acid is the limiting reagent; once H+ is consumed, the reaction stops regardless of chip size.)
• Why use the initial rate rather than the whole curve to compare conditions? (At t = 0 the concentrations and temperature are best defined; later, changes make comparisons less fair.)
• How does the cotton-wool plug improve data quality in the mass-loss method? (It reduces liquid spray loss while still letting CO2 escape, so mass loss mostly reflects gas release.)
• What controlled variables are essential here? (Acid concentration and volume, total CaCO3 mass, temperature, mixing, and sampling interval.)
• If you switched to gas collection, what units would you use for rate and how would you find the initial rate? (cm³ s⁻¹; draw a tangent at t = 0 to find the initial slope of volume vs time.)