My PBL Project

Ammonium nitrate (NH4NO3), is classified as a salt. Chemically speaking, there are thousands of salts in addition to sodium chloride, common table salt. Salts contain ions, particles with electrical charges. Because the ions with positive charge are strongly attracted to those with negative charge, they form a solid crystal. Below is a diagram of a salt crystal dissolving in water. To the eye, it looks like a simple process. In fact, there are two distinct steps, and each involves energy changes. In the first step, the solid crystal separates into ions. Breaking the ionic bonds requires a lot of energy, which means that heat must be absorbed from the surroundings. In the second step the water molecules, which are attracted to the charged ions, attach themselves to the ions. This step releases energy, which means that heat flows to the surroundings. The steps can be written like this:

Step 1: NH4NO3 + Heat ⇒ NH4 + + NO3 -

Step 2: NH4 + + NO3 - + H2O ⇒ NH4 + (H2O)x + NO3 - (H2O)x + Heat

Several water molecules may bond to each ion, as indicated by (H2O)x. In the first step, heat is absorbed; in the second step, heat is released. Overall, because more heat is involved in the first step than in the second step, heat is absorbed from the surroundings (6 kilocalories per mole of ammonium nitrate). This leaves the surroundings with less thermal energy—colder.

Portable hot and cold packs depend on reactions that are spontaneous. Because the packs must be quick and easy to use, they require reactions that begin as soon as the reactants are placed together and that will continue on their own. Most spontaneous chemical reactions are exothermic—they give off heat. This is because chemical bonds have a tendency to shed their stored energy and release it as heat. The people who designed hot packs found that this natural flow of energy suited their needs perfectly. They selected the appropriate reactants, put them in the same package, but kept them separated. When warmth is needed, the reactants are simply mixed, and heat is produced automatically. The tendency of stored bond energy to emerge as heat would seem to rule out cold packs. Because endothermic reactions absorb heat, the bonds end up with more stored energy than they started with—which is against the natural flow of things. Yet, this occasionally happens. When ammonium nitrate dissolves in water it gets very cold— spontaneously. Why does this occur? Scientists explain it with a concept called entropy. Entropy is the degree of disorder in a system. Chemical changes tend to go from orderly arrangements of molecules and ions to disorderly arrangements. Nature tends to increase the amount of messiness, or disorder, or entropy. The natural tendency to increase entropy sometimes opposes the tendency to release heat. When the increase in entropy is great enough, it can drive the heat flow “backward.” The drive for high entropy overpowers the drive to release heat. Endothermic reactions happen spontaneously only when the reactions permits a large increase in entropy. In the case of the instant cold pack, the starting material were highly ordered: The water was pure and sealed in

its own container, and the ammonium ions and nitrate ions were arranged in an orderly pattern in solid crystals. The substances were sorted and organized—everything in its place. When the inner plastic bag was broken and the water dissolved the ammonium nitrate, the orderly arrangement of the ions was disrupted. The ions were dispersed randomly throughout the water, and the once-pure water became “contaminated.” Disorder reigned. The system went from very ordered to very disordered, and the reaction was partly driven by this increase in entropy.

Concepts

• • Enthalpy change

  • Heat of solution

  • Calorimetry

  • Dependent and Independent variables

Background

The energy or enthalpy change associated with the process of a solute dissolving in a solvent is called the heat of solution. In case of an iconic compound dissolving in water. The overall energy change is the net result of two processes the energy required to break the attractive forces (iconic bonds) between the ions in the crystal lattice, and the energy released when the dissociate (free) ions from ion-dipole attractive forces with the water molecules. The process of a solute dissolving in wter may either release heat into aqueous solution or absorb heat from the solution.

Experiment Overview

The purpose of this inquiry-based experiment is to design and carry out a procedure to determine the enthalpy change that occurs when a "cold pack solid" dissolves in water.

Materials

• Beaker 400 ML

• Cold pack solid 10 g

• Distilled or deionized water

• Graduated cylinder 100 ML

• Insulated foam (styrofoam) cups 6 oz 2

• Balance

• Thermometer or temperature sensor

• Spatula

• Stirring rod

• Weighing dishes

Safety Precautions

The cold pack solid is slightly toxic by ingestion and is a body tissue irritant. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Questions

1)What information is needed to calculate the heat energy change for a reaction?

If we know the enthalpy changes of a series of reactions that add up to give an overall reaction, we add these enthalpy changes to determine the enthalpy change of the overall rection.

• 1. Using the enthalpy change for the reaction of Fe with Cl₂ to give FeCl₂ and the enthalpy change for the reaction of FeCl₂ with Cl₂ to give FeCl₃, we can determine the enthalpy change for the reaction of Fe with Cl₂ to give FeCl₃.