EXPERIMENT : HEAT OF REACTION
Endothermic and Exothermic Reactions
Many chemical reactions release energy in the form of heat, light, or sound. These are exothermic reactions. Exothermic reactions may occur spontaneously and result in higher randomness or entropy (ΔS > 0) of the system. They are denoted by a negative heat flow (heat is lost to the surroundings) and decrease in enthalpy (ΔH < 0). In the lab, exothermic reactions produce heat or may even be explosive.
There are other chemical reactions that must absorb energy in order to proceed. These are endothermic reactions. Endothermic reactions cannot occur spontaneously. Work must be done in order to get these reactions to occur. When endothermic reactions absorb energy, a temperature drop is measured during the reaction. Endothermic reactions are characterized by positive heat flow (into the reaction) and an increase in enthalpy (+ΔH).
Examples of Endothermic and Exothermic Processes
Photosynthesis is an example of an endothermic chemical reaction. In this process, plants use the energy from the sun to convert carbon dioxide and water into glucose and oxygen. This reaction requires 15MJ of energy (sunlight) for every kilogram of glucose that is produced: sunlight + 6CO
2
(g) + H
2
O(l) = C
6
H
12
O
6
(aq) + 6O
2
(g)
An example of an exothermic reaction is the mixture of sodium and chlorine to yield table salt. This reaction produces 411 kJ of energy for each mole of salt that is produced:
Na(s) + 0.5Cl
2
(s) = NaCl(s)
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Exothermic Processes
 freezing water

 solidifying solid salts condensing water vapor
 making a hydrate from an anhydrous salt
 forming an anion from an atom in the gas phase
 Annihilation of matter
E=mc2
 splitting of an atom
Exothermic Reactions
 Combustion of hydrogen
 dissolving lithium chloride in
 water
 Burning of propane dehydration of sugar with
 sulfuring acid
 thermite decomposition of hydrogen peroxide
 decomposition od ammonium
 dichromate halogenation of acetylene
Endothermic Processes







 phase
 melting ice cubes melting solid salts evaporating liquid water making an anhydrous salt from a hydrate forming a cation from an atom in the gas splitting a gas molecule separating ion pairs cooking an egg baking bread
Endothermic Reactions
 Reaction of barium hydroxide octahydrate
 crystals with dry ammonium chloride
 dissolving ammonium chloride in water reaction of thionyl chloride (SOCl2) with
 cobalt(II) sulfate heptahydrate
 mixing water and ammonium nitrate mixing water with potassium chloride
 reacting ethanoic acid with sodium carbonate
 photosynthesis (chlorophyll is used to react carbon dioxide plus water plus energy to make glucose and oxygen)
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EXPERIMENTAL
In this experiment, you will examine the heat of two reactions. Hydration of the concentrated sulphuric acid and dissolving of ammonium chloride in water.
1.
Hydration of the concentrated sulphuric acid
Adding of the concentrated sulphuric acid to water is an exothermic reaction.
H
2
SO
4

H
2
O and H
2
SO
4

2H
2
O molecules occur in the result of the hydration reaction of the sulphuric acid. The formation of the bonds between water and acid is the reason of releasing of heat.
Pour 40 mL tap water using a graduated cylinder to the beaker. Measure the temperature of water with a thermometer. Add 4 mL concentrated sulphuric acid to water.
Mix and measure of the temperature of the maximum temperature reached.
2.
Dissolving of ammonium chloride in water
There are two major competitor heat effect when a salt dissolves. These are lattice energy and hydration energy. Lattice energy is defined as energy needed for breaking of the ionic bonds of the crystal lattice. The energy is released while the ions leave the crystal lattice and hydration occurs. This energy is hydration energy. If the lattice energy is bigger than the hydration energy, the reaction will be endothermic. . If the hydration energy is bigger than the lattice energy, the reaction will be exothermic.
Pour 20 mL tap water using a graduated cylinder to the beaker. Measure the temperature of water with a thermometer. Weigh out about 3 g of amonium chloride and add to the beaker. Mix and measure of the temperature of the minimum temperature reached.
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