Download Endothermic and Exothermic Reactions

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 + 6CO2(g) + H2O(l) = C6H12O6(aq) + 6O2(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) + 0.5Cl2(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





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
phase




splitting a gas molecule
separating ion pairs
cooking an egg
baking bread
Endothermic Reactions


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

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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.
H2SO4H2O and H2SO42H2O 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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