When chemical reactions or processes occur, there is always an exchange of energy. Some of these reactions or processes give off energy as heat; these are called exothermic (‘exo’ meaning outside, ‘thermic’ meaning heat). Other reactions and processes absorb energy, making the surroundings cooler; these are called endothermic (‘endo’ meaning inside).
First, we need to briefly discuss chemical reactions. A chemical reaction is when one or more chemical compounds are changed into one or more different compounds. In any chemical reaction, some bonds need to be broken, and others need to be formed—this is how the reaction produces new compounds. If we know how much energy is required to break the bonds in the reactants (the compounds present before the reaction takes place), and we know how much energy is released on formation of the bonds in the products (the compounds present after the reaction takes place), we can compare them to see how much energy will be produced or consumed by the reaction. Fortunately for us, there are tables we can use to figure out the energy of the reactants and products. These are called bond energy tables, similar to the one below (1).
The same is true for chemical processes. A chemical process is what happens when there is a change in the state of one or more chemical compounds (like changing from a liquid to a gas, or dissolving in water), but there is no formation of a new compound. If we know how much energy the compounds have before they undergo the process (such as melting, or dissolving in water), and how much energy they have after this process, we can discover if the process is endothermic or exothermic. For example, if we have an ice cube sitting at room temperature, we know the ice cube will begin to melt. The warmth of the room is melting the ice because the water molecules are absorbing the thermal energy from the air in the room, and this energy is making the molecules move faster and farther away from each other, bringing them from a solid state (ice) to a liquid state (water). Because this process absorbs energy, it is endothermic.
1. Energy is added to the solution to pull the ions away from each other: in order to pull the positively and negatively charged ions apart, energy must be added. This energy needed to pull the ions apart is called the Lattice Energy.
2. Energy is released into solution when the water molecules surround the ions: as the water molecules are attracted to and surround the ions, energy is released into the solution. This energy released as water molecules surround the ions is called the Hydration Energy.
Whether the dissolving of a salt is exothermic or endothermic depends on which is greater, the Lattice Energy, or the Hydration Energy. These are usually expressed in units describing the amount of energy released per set amount of salt, such as kilocalories per mole (kcal/mol) or kilojoules per mole (kJ/mol). We can usually look up the values of the Lattice and Hydration Energy values for a particular salt in tables, such as the one below (2).
779 kJ/mol – 774 kJ/mol = +5 kJ/mol
It takes just slightly more energy to separate the ions from one another than is released from the water molecules surrounding the ions. This means just slightly more energy must be put into the solution than is released back into the solution; therefore dissolving table salt in water is endothermic.
However, if we dissolve sodium hydroxide (NaOH) in water, it separates into Na+ and OH- ions. The Lattice Energy for this process is 737 kJ/mol, and the Hydration Energy is 779 kJ/mol. Subtracting as before, we get a change of -42 kJ/mol.
737 kJ/mol – 779 kJ/mol = -42 kJ/mol
More energy is released into the solution than is required to pull apart the ions; therefore dissolving sodium hydroxide in water is exothermic. If you dissolve sodium hydroxide in a small amount of water, be careful—the container may get hot enough to burn your hand!
Here’s what you’ll need:
1. Two small jars or drinking glasses
2. Two teaspoons
3. Two cups of distilled water
4. One half-cup of magnesium sulfate (MgSO4). You can purchase this online (click here for options).
5. One half-cup of ammonium chloride (NH4Cl). You can purchase this online (click here for options).
6. One thermometer that will measure temperatures from 70-150°F
7. Safety goggles, one pair for each person participating
8. Latex or nitrile gloves (you can get these in grocery or hardware stores)
Here’s what you need to do:
NOTE: Be very careful with the magnesium sulfate and ammonium chloride—they can cause irritation to the skin, lungs and eyes. Do not breathe them in or get them in your eyes! You should do this procedure in a well ventilated area, and wear the goggles and gloves to make sure your eyes and skin are protected.
1. Put on your goggles and gloves!
2. Pour one cup of distilled water into each of the small jars.
3. Measure and record the temperature of the water in each jar.
4. Pour one half cup of the magnesium sulfate into one of the jars. Stir carefully with a spoon for 20 seconds (don’t worry if not all the magnesium sulfate dissolves). Measure and record the temperature of the solution.
5. Pour one half cup of the ammonium chloride into one of the jars. Stir carefully with a spoon for 20 seconds (don’t worry if not all the magnesium sulfate dissolves). Measure and record the temperature of the solution.
Did the temperature of the water change each time? How much did it change? Did it get hotter or colder? Are these processes endothermic or exothermic? Did you observe anything else? BE SURE TO WRITE EVERYTHING DOWN IN YOUR JOURNAL!
(1) “Bond Enthalpy/Bond Energy”. Mr. Kent’s Chemistry Page. www.kentchemistry.com Accessed 8/28/14.
(2) “Chapter 13: Solutions”. intro.chem.okstate.edu. Accessed 8/29/14.