Chapter 20 Section 20.1

Chemical Reactions and Energy; Basic
Thermodynamics
Section 20.1
Objectives ~
• Compare and contrast exothermic and endothermic chemical
reactions
• Analyze the energy diagrams for typical chemical reactions
• Illustrate the meaning of entropy and trace its role in various
processes
Energy that is transferred from an object at high temperature to
an object at lower temperature.
Measured in Joules (J) 1 kilojoule (kJ) = 1000J
MOLE = A unit of measure in a chemical reaction that contains the
same number of particles
Energy changes that occur in a chemical reaction are a result of
the bonds that are broken (reactants) and the bonds that are
formed (products)
• If more energy is released in forming new bonds than is
required to break bonds, the reaction is exothermic
• If reactants go from higher energy to products at lower energy
= exothermic
• In an endothermic reaction, the energy required to break the
bonds of the reactants is greater than the energy released by
forming the bonds of the products.
• If reactants go from lower energy to products at higher energy
= endothermic
• Symbolizes the difference in the energy between the products
and reactants
• At constant pressure, the enthalpy of a reaction or system is the
same as the heat that is gained or lost
• Exothermic = -∆H
• Endothermic = +∆H
ΔH reaction = ΔH products - ΔHreactants
• Is…
2H2(g) + O2(g) → 2H2O(l)
∆H = -572 kJ
• In the synthesis of water, 2 moles of hydrogen gas react with 1
mole of oxygen gas to produce 2 moles of liquid water. The
reaction is exothermic and produces 572 kJ of energy for every
2 moles water produced.
• What would the enthalpy of the decomposition be?
• What is the enthalpy value for 1 mole of water produced?
What does decomposition look like?
2H2O(l)  2H2(g) + O2(g)
∆H = +572 kJ
If only 1 mole of water produced?
572kJ = 2 x 286 kJ
286kJ
• the amount of energy that particles in a reaction must have
when they collide in order for the reaction to occur. Figures 20.4
and 20.5
• Recall that a catalyst will speed up the rate of a reaction
without being consumed
• this occurs because the presence of a catalyst lowers the
activation energy of a reaction.
• ∆S a measure of disorder
• The natural tendency of entropy is to increase. MORE disorder
is more favorable; Entropy is affected by temperature
• + ∆S means increasing disorder and is favorable
• -∆S means increasing order and is not favorable
• Once activated, a reaction that uses some of its released
energy to continue is considered spontaneous
• When one considers ∆S and ∆H together, we know whether a
reaction will be spontaneous or not
• Spontaneity depends on the balance between energy and
entropy factors.
• An exothermic reaction with increasing disorder is always
spontaneous, BUT an endothermic reaction with increasing order
is never spontaneous. (table 20.1 page 714)
• Are reactions that occur spontaneously at room temperature
generally exo or endothermic?
• Exothermic - Remember the tendency is to change from a state of high
energy to a state of low energy.
1) Water in an ice-cube tray freezing?
2) You pick up scattered trash along a highway and pack it in a
bag.
3) Your campfire burns, leaving gray ashes.
4) A cube of sugar dissolves in a cup of tea.
1)
2)
3)
4)
Decrease
Decrease
Increase
Increase
• Which has greater entropy, a tablespoon of dry salt or a
tablespoon of salt dissolved in water? Explain.
• The salt dissolved in water has greater entropy because it has a
less orderly structure.
• Assuming that temperature does not change, would there be an
increase or decrease in entropy in the following reaction:
• SO2(g) + H2O(l)
H2SO4(aq)