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Chapter 17
Reaction Energy and Reaction Kinetics
I. Thermochemistry
 study of the transfer of heat that accompany chemical reactions and physical changes
Temperature—a measure of the average kinetic energy of the particles in a sample of matter
Do objects at the same temperature contain the same amount of heat?
Enthalpy (H)
 the heat content of a system at constant pressure
 cannot be directly measured
Heat (q)
 energy transferred between samples with different temperatures
 ALWAYS from higher T to lower T
 Units:
o joules (J)
o calories (cal)
NOTE: 1 Calorie (food) = 1000 cal = 1 kcal = 4.18 kJ
II. Measuring Heat
We cannot directly measure the heat content (H) of a substance. We can measure the CHANGE in heat
content (H). When this happens, there is heat transfer between a system and its surroundings.
+q = Hsystem = -Hsurroundings
If two different objects with the same mass absorb the same quantity of heat, do they have to end up at the
same temperature?
Specific heat capacity (c)—the amount of energy required to raise the temperature of one gram of a
substance by 1 degree C (or K).
 the larger the s.h., the SLOWER it will heat up (and cool down)
 Units: J/g K
If we know the s.h. of a substance and the CHANGE in temperature (Tfinal – Tinitial) it experiences, we can
determine how much heat was transferred into (+) or out of (-) it.
q = m x c x T
Table 17-1 (p. 513) give s.h. values
III. Heat of Reaction
Just about all chemical reactions will show a change in enthalpy between the reactants and products
If reactants GAIN heat  endothermic reaction
 products have more enthalpy than the reactants
 H is (+)
If reactants LOSE heat  exothermic reaction
 products have less enthalpy than the reactants
 H is (-)
REMEMBER, we don’t know the enthalpy of the reactants or products, only the CHANGE.
H = Hproducts - Hreactants
The H can be put into the reaction (a reactant is H is (+) or a product if H is (-)).
EX.
2 H2(g) + O2(g)  2 H2O(g) + 483.6 kJ
This is known as a Thermochemical equation
NOTE: coefficients can ONLY stand for MOLES
IV. Heat of Formation
Molar heat of formation: H for the reaction when ONE MOLE of a substance is formed from its elements
in their standard state.
Standard state (25oC & 1atm)
 Hfo (values in Table A-14, page 902)
EX.
2 H2(g) + O2(g)  2 H2O(g)
Ho = -483.6 kJ
o
Hf (H2O) = -241.8 kJ
WHY?
NOTE: the Hfo of any element in its standard state is ZERO
H2(g)  H2(g) Ho = ?
V. Heat of Combustion
 Ho for a combustion reaction (ONE MOLE of the reactant)
 Hoc
 EX. What is Hoc for hydrogen?
H2(g) + 2 O2(g)  H2O(g)
o
H c = -241.8 kJ
VI. Hess’s Law and Heats of Reaction
Ho =  (Hfo products) -  (Hfo reactants)
NOTE: Hfo values are PER MOLE (multiply values by coefficients)
Ho is independent on the “path” the substance took to be formed.
EX.
A+CD
E+GX+A
2A + X  Y + D
Y+CG+E
Ho1
Ho2
Ho3
Ho4
Ho1 = Ho2 + Ho3 + Ho4
NOTE:
2A + 2C  2D
DA+C
Ho5 = 2 Ho1
Ho6 = -Ho1
VII. Entropy
Spontaneous reaction: reaction that occurs “as written” (under the given conditions with no “outside
assistance”)
System will naturally want to get to a LOWER state of energy
 exothermic reactions (reactants lose energy)
Why are some endothermic reactions spontaneous?
A second “driving force” behind spontaneous chemical reactions is entropy.
Entropy (S)
 a measure of the degree of randomness (disorder) of the particles in a system.
 CAN be measured directly
 unit: J/mol K
An INCREASE in the entropy of the system during a chemical reaction will promote spontaneity.
 if the products have a HIGHER entropy than the reactants, spontaneity is promoted
Low Entropy
High Entropy
solids
gases
less moles (particles)
more moles (particles)
low T
high T
S0 =  (S0 (products)) -  (S0 (reactants))
S0 – standard molar entropy (tabulated values)
Why are some reactions that produce a DECREASE in entropy still spontaneous?
VIII. Free Energy
The two driving forces behind a reaction are:
 enthalpy changes
 entropy changes
Favorable changes
(always spontaneous)
-H0; + S0
Unfavorable changes (never spontaneous)
+H0; - S0
Questionable:
 +H0; + S0 or -H0; - S0
 changing T can change reaction from spontaneous to nonspontaneous (or vice versa)
Free Energy (G)
G0 = H0 - TS0
T MUST be in KELVIN.
If G0 is NEGATIVE, the reaction is SPONTANEOUS.
NOTE: Tabulated values for S0 are usually in J/mol K and H0 in kJ/mol.
IX. Collision Theory
In order for reactions to occur between substances, their particles must collide. The collisions must also
result in INTERACTIONS (correct orientations and energies).
 how often these interactions occur will determine the reaction rate
Reaction Mechanisms: show the step-by-step sequence of reactions that the overall chemical reaction
follows.
EX.
I2(s) + H2(g)  2 HI(g)
Mechanism 1:
I2  2 I
2 I + H2  2 HI
Mechanism 2:
I2  2 I
I + H2  H2I
H2I + I  2 HI
NOTE: I and H2I are intermediates (stable substances that do NOT appear in the overall reaction)
Each step (elementary reactions) in a reaction mechanism must form an ACTIVATED COMPLEX
 transitional structure that results from an effective interaction (collision) and that persists while old
bonds are broken and new ones are formed.
Activated complexes can only from if the collision has a certain minimum energy
 ACTIVATION ENERGY
X. Energy Diagrams:
Endothermic Reactions:
 (+) activation energy
 (+) H
Exothermic Reactions:
 (+) activation energy
 (-) H
XI. Reaction Rates
Chemical kinetics: study of reaction rates and reaction mechanisms
Reaction rate: the change in the concentration of REACTANTS per unit of time.
Since the nature of reactant collisions determine how often reactions occur, changing the frequency and
energy of these collisions will change reaction rates.
5 Factors will influence the reaction rate:
1) Nature of the reactants—in general, the stronger the bonds in the reactants, the slower the reaction will
occur.
2) Surface area (solids)—since only the surface particles can react, increasing surface area (crushing) will
increase reaction rates.
3) Concentration—the higher the concentration of reactants, the higher the number of collisions produced.
4) Temperature—increasing T increases K.E. of particles  more collisions and they occur with more
energy.
5) Catalyst—increases the reaction rate by LOWERING the activation energy for the reaction (provides a
different reaction pathway)
XII. Rate Determining Steps
For multi-step reaction mechanism, the reaction cannot go any faster than the slowest step (elementary
reaction).
Thus,
Rate Determining Step of the reaction is the SLOWEST elementary reaction.
XIII. Rate Laws
AC+D
rate = - [A] / T
=  [C] / T =  [D] / T
rate = k [A]x
k is called the RATE CONSTANT.
x is the ORDER OF REACTION with respect to A.
For the general reaction:
aA+bBcC+dD
rate = k [A]x [B]y
In general, x and y will be 0, 1, or 2.
x and y MUST be determined experimentally.
NOTE: if the reaction has a single-step mechanism (or is an elementary reaction), x = a and y = b