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Bond Enthalpies
How does a chemical reaction have energy?
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Bond Energy Energy required to make/break a chemical bond
Endothermic reactions
Products have more energy than reactants
More energy to BREAK bonds
Exothermic reactions
Reactants have more energy than products
More energy to FORM bonds
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Bond Enthalpy Focuses on the energy/heat between products and
reactants as it relates to chemical bonding
Amount of energy absorbed to break a chemical bond---amount of energy released to form a bond.
Multiple chemical bonds take more energy to break and release more energy at formation
Amount of energy absorbed = amount of energy released
to break chemical bond to form a chemical bond
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Calculating ΔHrxn. by bond enthalpies (4th method)
Least accurate method
ΔH = ΣBE (bonds broken) - ΣBE (bonds formed)
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Example 1: Using average bond enthalpy data,
calcaulate ΔH for the following reaction.
CH4 + 2O2 CO2 + 2H2O ΔH = ?
Bond Average Bond Enthalpy
C-H 413 kJ/mol
O=O 495 kJ/mol
C-O 358 kJ/mol
C=O 799 kJ/mol
O-H 467 kJ/mol
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Entropy
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Spontaneous vs. Nonspontaneous
1) Spontaneous Process
Occurs WITHOUT help outside of the system, natural
Many are exothermic—favors energy release to create an energy reduction after a chemical reaction
Ex. Rusting iron with O2 and H2O, cold coffee in a mug
Some are endothermic
Ex. Evaporation of water/boiling, NaCl dissolving in water
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Spontaneous vs. Nonspontaneous
2) Nonspontaneous Process
REQUIRES help outside system to perform chemical reaction, gets aid from environment
Ex. Water cannot freeze at standard conditions (25°C, 1atm), cannot boil at 25°C
**Chemical processes that are spontaneous have a nonspontaneous process in reverse **
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Entropy (S) Measure of a system’s disorder
Disorder is more favorable than order
ΔS = S(products) - S(reactants)
ΔS is (+) with increased disorder
State function
Only dependent on initial and final states of a reaction
Ex. Evaporation, dissolving, dirty house
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Thermodynamic Laws
1st Law of Thermodynamics
Energy cannot be created or destroyed
2nd Law of Thermodynamics
The entropy of the universe is always increasing.
Naturally favors a disordered state
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When does a system become MORE disordered from a
chemical reaction? (ΔS > 0)1) Melting
2) Vaporization
3) More particles present in the products than the reactants
4C3H5N3O9 (l) 6N2 (g) + 12CO2 (g) + 10H2O (g) + O2 (g)
4) Solution formation with liquids and solids
5) Addition of heat
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When does a system become LESS disordered from a
chemical reaction? (ΔS < 0)
1) Solution formation with liquids and gases
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3rd Law of Thermodynamics
The entropy (ΔS) of a perfect crystal is 0 at a temperature of absolute zero (0°K).
No particle motion at all in crystal structure
All motion stops
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How do we determine if a chemical reaction is
spontaneous?1) Change in entropy (ΔS)
2) Gibbs Free Energy (ΔG)
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Change in entropy (ΔS)
For a chemical reaction to be spontaneous (ΔST > 0), there MUST be an increase in system’s entropy (Δssys> 0) and the reaction MUST be exothermic (Δssurr > 0).
Exothermic reactions are favored, NOT endothermic reactions.
Exothermic (ΔH < 0, ΔS > 0)
Endothermic (ΔH > 0, ΔS < 0)
ΔST = Δssys + Δssurr
If ΔST > 0, then the chemical reaction is spontaneous
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Example 1:
Will entropy increase or decrease for the following?
a) N2 (g) + 3H2 (g) 2NH3 (g)
b) 2KClO3 (s) 2KCl (s) + 3O2 (g)
c) CO(g) + H2O(g) CO2 (g) + H2 (g)
d) C12H22O11 (s) C12H22O11
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How do we calculate the entropy change (ΔS) in a chemical reaction?
Same method as using the enthalpies of formation to calculate ΔH and use the same table.
aA + bB cC + dD
ΔS° =[c (ΔS°C) + d(ΔS°D)] - [a (ΔS°A) + b (ΔS°B)]
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Example 2: Calculate ΔS° for the following reaction at
25°C….
4HCl(g) + O2 (g) 2Cl2 (g) + 2H2O (g)
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Homework pp. 382-383 #69, 71-73
pp. 742-743 #19, 27