thermochemistry - penn state universitycourses.chem.psu.edu/chem110/summer/lectures/web...
TRANSCRIPT
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Mary J. Bojan Page 1 Chapter 5
Thermochemistry Energy
kinetic potential
First Law of Thermodynamics E = q + w
Enthalpy H Thermochemical Equations H for chemical reactions Calorimetry, Heat Capacity Hesss Law Heat of formation Hf Standard State H f Foods and Fuels Thermochemistry:
Study of energy changes in chemical processes
2H2 + O2 2H2O + energy
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Mary J. Bojan Page 2 Chapter 5
ENERGY Energy can be converted from one form to
another
Kinetic Energy energy of motion Mechanical moving mass (1/2mv2) Electrical moving charge Light photons Sound molecules moving
uniformly Heat molecules moving
randomly Potential Energy stored energy Mechanical mass in a place where
a force can act Chemical bonds Nuclear binding energy
Chemical energy the Potential Energy associated with bonding
HH + HH + O=O O
HH +
OH
H 3 bonds 4 bonds
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Mary J. Bojan Page 3 Chapter 5
Energy Units
SI Unit: ??? What is the SI Unit of energy? a) Cal b) kcal c) ergs d) joules e) BTUs
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Mary J. Bojan Page 4 Chapter 5
ENERGY but total energy remains constant.
First Law of thermodynamics: Law of conservation of energy Energy lost by the SYSTEM is gained by the SURROUNDINGS. The system: what you are interested in
(a chemical reaction) Surroundings: everything +w else
E +q E
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Mary J. Bojan Page 5 Chapter 5
Changes in Energy E = internal energy:
the capacity to do work or transfer heat. E =Efinal Einitial = q + w
w= work= action of force through a distance
(often PV) work done to the system (+)
q = heat (thermal energy) = energy transferred due to
a difference in temperature. heat added to the system (+)
energy
System Surroundings E = exothermic energy
Surroundings System E = + endothermic
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Mary J. Bojan Page 6 Chapter 5
Energy Changes - Heat and Work A B Dry ice (CO2) is heated to room temperature at constant P or at constant V What are the signs of q and w for each? Does Esystem increase or decrease? How do you know? Which system has higher E at the end? Where does the energy come from?
w = P!V
+q +q
!V = 0
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Mary J. Bojan Page 7 Chapter 5
90C H2O 100ml cool 25C A H2O B H2O heat Which one has more energy? 0C H2O
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Mary J. Bojan Page 8 Chapter 5
State function: a function whose value does not depend on the pathway used to get to the present state:
State Function: only depends on the current state (composition, T,P): does not depend on past history
Chemical Reaction
A
A A
B B
C
D
E
B
elevation
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Mary J. Bojan Page 9 Chapter 5
State Functions
State functions are written as uppercase letters (E, H, P, V, T, S) q and w are not state functions but E (= q + w) is a state function Changes in state functions are path-independent:
reactants
products
1
2 E
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Mary J. Bojan Page 10 Chapter 5
Chemical Reactions
When natural gas burns: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l)
Reaction produces heat and light
Is energy conserved? What are the signs of q and w for the overall reaction? Where does the energy come from?
Is energy stored or released when bonds are broken?
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Mary J. Bojan Page 11 Chapter 5
Relationship between Energy (E) and Enthalpy (H)
Energy transferred when V is constant
= E Usually run chemical reactions at constant
pressure (atmospheric).
E = q + w PV = work done on system at constant P
arises from expansion or contraction of the system: V = volume change
of system
E = qp PV P = constant qp = E + PV H
H = enthalpy
Energy transferred when P is constant
(=qp) = H
For many chemical processes, PV is small and
E H Like E, H is a state function
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Mary J. Bojan Page 12 Chapter 5
Thermochemical Equations
A balanced chemical equation that also includes the energy change
H2(g) + 1/2 O2(g) H2O(g) H = 241.8kJ H = enthalpy: heat given off or absorbed in the reaction
ENTHALPY 1) Enthalpy is an extensive property. 2) H for a reaction is equal in magnitude
and opposite in sign to H for the reverse reaction.
3) H for a reaction depends on the states
of reactants and products (gas, liquid, solid).
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Mary J. Bojan Page 13 Chapter 5
ENTHALPY OF REACTION
H = H(products) H(reactants) If H0 (+) endothermic
(heat absorbed) Ba(SCN)2(aq) + 2NH3(aq)+10H2O(l) H H = + Ba(OH)28H2O(s) + 2NH4SCN(s)
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Mary J. Bojan Page 14 Chapter 5
H2(g) + 1/2 O2(g) Hrxn = 241.8kJ H2O(g) H2O(l) 5.1 Is this reaction
1 exothermic 2 endothermic
5.2 How much heat is given off per mole of H2?
1 241.8 kJ 2 120.9 kJ 3 not enough information
5.3 How much heat is given off per mole of O2?
1 483.6 kJ 2 241.8 kJ 3 120.9 kJ 4 not enough information
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Mary J. Bojan Page 15 Chapter 5
H2(g) + 1/2 O2(g)
Hrxn = 241.8kJ H2O(g) H2O(l) 5.4 If I convert water to H2 + O2, the reaction
will be ___________
1 exothermic 2 endothermic
5.5 How much heat will be needed to convert
9g of water into H2 + O2? 1 483.6kJ 2 241.8 kJ 3 120.9 kJ
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Mary J. Bojan Page 16 Chapter 5
H2(g) + 1/2 O2(g) Hrxn = 241.8kJ H2O(g) H2O(l) 5.6 How much heat will be given off if liquid
water is formed instead of gaseous water. Hvap of water = 44kJ/mol
1 241.8kJ 2 +241.8kJ 3 285.8kJ 4 44kJ
5.7 How much heat will be given off if 10g of H2 is consumed?
1 24.18kJ 2 241.8kJ 3 1209 kJ 4 2418 kJ
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Mary J. Bojan Page 17 Chapter 5
CALORIMETRY Experimental measure of heat flow (used to determine Hrxn)
q = C m T q = heat flow C = specific heat (heat capacity per gram) m = mass T= Tfinal - Tinitial For H2O: C = 4.184 J/g C
Molar heat capacity = 75.2 J/mole C
Note: H2O is usually part of the surroundings
qsurr = Csurr m T
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Mary J. Bojan Page 18 Chapter 5
Mix 50ml of 1M NaOH+50ml of 1M HCl What is H rxn ? =
1. Write the balanced reaction. 2. Ti = Tf = T = Tf Ti = 3. Is qp (H rxn) positive or negative? 4. qp =H rxn= CmT
V=100ml, assume d = 1g/ml Then: m = (100ml)(1g/ml) = 100g
qp =(4.184J/goC)(100g) T
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Mary J. Bojan Page 19 Chapter 5
HESSS LAW
H for a sum of steps is the same as H for the overall process.
True because H is a state function. C H2 Hrxn B
H1 A A B H1 B C H2 A+B B + C H1 + H2= Hrxn
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Mary J. Bojan Page 20 Chapter 5
Example: Given the following information A H2(g) + F2(g) 2HF(g) HA = 537kJ B 2H2(g) + O2(g) 2 H2O(g) HB = 572kJ Determine H for the reaction: C 2F2(g) + 2H2O(g) 4HF(g) + O2(g) HC =? IDEA: find combinations of reactions
such that
n A + m B = C then
n HA + m HB = HC Here: 2xA 2H2(g) + 2F2(g) 4HF(g) H = 2(537kJ) 1xB 2 H2O(g) 2H2(g) + O2(g) H = (572kJ) _________________________________________________________________________________________________
2F2(g) + 2H2O(g) 4HF(g) + O2(g) HC = 2(537kJ) 1(572kJ) = 502kJ
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Mary J. Bojan Page 21 Chapter 5
Given the following information: H 2SO2(g)+ O2(g) 2SO3(g) 196kJ 2S(s) + 3 O2(g) 2SO3(g) 790kJ What is Hrxn for
S(s) + O2(g) SO2(g) 1. + 986 kJ 2. 986 kJ 3. 594 kJ 4. + 594 kJ 5. 297 kJ
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Mary J. Bojan Page 22 Chapter 5
Heat of Formation
Hf enthalpy of formation heat given off (or absorbed) when elements combine to give a compound
combine Elements Compounds Hf H f standard enthalpy of formation
heat given off (or absorbed) when all substances are in their Standard State
STANDARD STATE
P = 1 atm T = 25C (298K) most stable state (gas/liquid/solid)
For an element in its standard state: H f = 0 (by definition)
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Mary J. Bojan Page 23 Chapter 5
For which of the following reactions (at 25 C and 1 atm) is Hrxn = H f ? a) H2(g) + F2(g) 2HF(g) b) NO(g)+ 1/2 O2(g) NO2(g) c) 2C(graphite)+3H2(g)+1/2O2(g)C2H5OH(l) d) Pb(s) + Cl2(g) PbCl2(s) e) S(s) + O3(g) SO3(g) f) Br2(l) Br2(g)
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Mary J. Bojan Page 24 Chapter 5
Standard Enthalpy of a reaction: H rxn
Heat of reaction (Hrxn) when all reactants and products are in the standard state.
H rxn from H f H rxn = n H f (prod) m H f (react)
n, m are stoichiometric coefficients of products and reactants. (This is an application of Hesss Law.)
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Mary J. Bojan Page 25 Chapter 5
Using the information in Table 5.3 of your text book, what is H rxn for the reaction:
C2H5OH(l) + 3O2(g)2CO2(g) + 3H2O(l)
a) 115.8kJ b) 115.8kJ c) 1366.7kJ d) 13366.7kJ e) There is not enough information in
Table 5.3 to answer this. If a piece of fruit contains 16.0 g of fructose, how many food calories does it contribute? C6H12O6(s) +6O2 6CO2+ 6H2O Hrxn = 2803kJ 1 cal = 4.184J 1kcal = 1 Cal (= 1 food calorie)
a) 2803Cal b) 669.9Cal c) 10.72Cal d) 59.5Cal
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Mary J. Bojan Page 26 Chapter 5
Bond properties Review
COVALENT BOND LENGTHS and ENERGIES
Bond length: distance between nuclei bond Bond energy
kJ/mol Bond length pm
CC 348 154 C=C 614 134 CC 839 121 more electrons shared, shorter bond length
bond Bond energy
kJ/mol Bond length
pm CH 413 110 CCl 328 176 CBr 276 196 Shorter bond length, stronger the bond Table 8.4 and 8.5
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Mary J. Bojan Page 27 Chapter 5
BOND ENERGY bond (dissociation) energy: D
enthalpy of bond breaking reaction in the gas phase.
D > 0 (H > 0) for diatomics, D is H of one reaction: HH(g) 2H(g) DH-H= Hrxn = 436kJ/mol for polyatomics, D is an averaged quantity
HOH(g) HO(g) + H(g) +494kJ/mol HO(g) H(g) + O(g) +424kJ/mol
DOH = 463 kJ/mol *
* value obtained from averaging over many molecules
not exact for any one case (like H f)
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Mary J. Bojan Page 28 Chapter 5
Estimating Hrxn
Bond energies provide estimates of reaction enthalpies:
Hrxn nDbroken mDformed
n, m = # bonds Energy is given off () when bonds form. Helps understand origins of Hrxn Example:
2 H2O2 2 H2O + O2 Hrxn = ? Draw Lewis structures of reactants and products:
Reactants(broken) Products (formed)
bond # D bond # D