Week 13: Lectures 37 – 39
Lecture 37: W 11/16
Lecture 38: F 11/18
Lecture 39: M 11/28
Reading:
BLB Ch 4.6; 10.5; 8.8; 5.3 – 5.7
Homework:
BLB 10: 57; 5: 4, 17, 29, 37, 39, 53, 55, 83, 85;
8: 65a, 67ac, 72ab, 74, 92a;
Supp Rxns: 12 – 18; 8: 12 – 14; 5: 1 – 7
Reminder:
Angel Quiz 12 on Thur 11/17
ALEKS Objective 13 due on Tues 11/29
Jensen Office Hour: 501 Chemistry Building
Tuesdays and Thursdays 10:30 – 11:30 am
Thanksgiving break: Nov 21 – 25
Final Exam: Monday Dec 12 2:30 – 4:20 pm
40 questions. A 100% Score earns 58 points out
of 200 course points (29% of final grade).
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Jensen
Chem 110 Chap 5
Page: 1
Solution Reactions
Basic skills:
• How to calculate formula weight (molar mass)
• How to do the following conversions:
gram ⇔ mole
concentrations ⇔ mole
Titration: find concentration of unknown solution
Can be used with:
√ Acid-Base reactions
(neutralization)
√ Precipitation
reactions
√ Redox reactions
Method:
!""#$%"&#$'()%*'+
,-&.&/#0&100#2#3!456
$'()%*'+#'7#)+8+'9+
:'+:-+%;!%*'+
,-&.&/#<<#2#5=(6
• React solution of unknown concentration with
solution of known concentration (standard
solution)
• At the end point or equivalence point (using an
indicator), the reaction is stoichiometrically
complete
• At the end point, you know the exact moles of
each reactant in solution from the balanced
equation.
Jensen
Chem 110 Chap 5
Page: 2
Example: How many mL of a 0.827 M KOH
solution is required for neutralization of a
35.00 mL sample of 0.737 M H2SO4 solution?
A.
B.
C.
D.
E.
35.0
1.12
25.8
62.4
39.3
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mL
mL
mL
mL
mL
Chem 110 Chap 5
Page: 3
Practice Example: One common component of
antacids is Mg(OH)2. If an upset stomach
contains 125 mL of 0.115 M HCl, how many
grams of Mg(OH)2 is required to completely
neutralize the acid?
A.
B.
C.
D.
E.
0.144
0.286
0.419
0.525
0.884
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g
g
g
g
g
Chem 110 Chap 5
Page: 4
Gas Phase Reactions
• When chemical reactions involve gases, the
balanced equation provides the number of moles
of reactants and products.
• The ideal gas equation provides the link between
number of moles and P, V, T of gases.
Example: Air Bag
How many liters of N2 at 735 mm Hg and
26°C are produced from 126 g NaN3 (sodium
azide)?
2 NaN3 (s) →
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2 Na(s) + 3 N2(g)
Chem 110 Chap 5
Page: 5
Practice Example: Humans consume glucose to
produce energy, and the products are CO2
and H2O.
C6H12O6(s) +
O2(g) →
CO2(g) + H2O(l)
[unbalanced]
What volume of CO2 is produced during the
consumption of 4.65 g glucose at body
temperature (37 ºC) and 1 atm?
A.
B.
C.
D.
E.
3.94 L
1.97 L
0.657 L
3.47 L
7.88 L
Jensen
Chem 110 Chap 5
Page: 6
Thermochemistry
KINETIC ENERGY
POTENTIAL ENERGY
Mechanical moving mass
1/2mv2
mass in a place
where force can act
Electrical
moving charge
electrostatic:Q1Q2/d
Light
photons
Chemical
Sound
bonds
molecules moving
uniformly
Nuclear
Heat
binding energy
molecules moving
randomly
Gravitational
mgh,
where g=9.8 m/s2
Chemical energy:
Potential Energy associated with Bonding
HH + HH + O=O →
H
3 bonds
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O
H+ H
O
H
4 bonds
Chem 110 Chap 5
Page: 7
Energy can be converted between various forms,
but total energy remains constant
Law of Conservation of Energy
First Law of Thermodynamics
• All energy lost by a system under observation
must be gained by the surroundings (and vice
versa)
system: what you are interested in
Surroundings: everything else
• During energy conversion, some heat is always
produced
Jensen
Chem 110 Chap 5
Page: 8
Changes in Energy
ΔE = Efinal – Einitial
• Both E and ΔE are state functions
State functions: only depend on the current
state (composition, T, P), does not depend on
path or history
"
#
!
!
"
!
!
$
ΔE for path 1 = ΔE for path 2
System
energy
Surroundings
Surroundings
energy
System
ΔE is ___
ΔE is ___
• ΔE = q + w
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Chem 110 Chap 5
Page: 9
Energy & Enthalpy
When changes occur at constant pressure:
ΔE = qp + wexpansion
~ negligible
ΔH = qp
H is enthalpy
• ΔH is the change in enthalpy
Both H and ΔH are also state functions
• ΔH is the quantity of thermal energy (heat)
absorbed or released by a system at constant
pressure
Examples of enthalpy:
Energy transfers accompany physical changes;
ΔHfusion (heat of fusion)
ΔHvaporization (heat of vaporization)
Energy transfers accompany chemical changes
Enthalpy of reactions ΔHrxn
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Chem 110 Chap 5
Page: 10
Thermochemical Equation: A balanced chemical
equation that also includes the enthalpy change
ΔH = Hproducts – Hreactants = ΔHrxn
_____thermic reaction
ΔH < 0
_____thermic reaction
ΔH > 0
Characteristics of Enthalpy:
1) Enthalpy is an extensive property
2) ΔHrxn is equal in magnitude but opposite in sign
for ΔH of reverse reaction
3) ΔHrxn depends on states of reactants & products
(e.g., gas, liquid …)
ΔH° (delta H standard)
standard P (1 atm) & T (usually 25°C)
Note: NOT the same as STP for gases!!!
Jensen
Chem 110 Chap 5
Page: 11
Study the thermochemical equation of:
The Hydrogen Balloon
2H2(g) + O2(g) → 2H2O(g) ΔH = −483.6 kJ
2 mol H2(g) reacts with 1 mol O2(g), produces
2 mol of H2O(g) and gives off 483.6 kJ heat
A. Is this reaction exothermic or endothermic?
B. How much heat is given off per mole of O2?
C. How much heat is given off per mole of H2?
D. How much heat will be given off if 10.0 g of H2 is
consumed?
E. What is ΔH for 2H2O(g) →2H2(g) + O2(g)?
F. How much heat will be needed to convert 9.0 g of
water into hydrogen and oxygen?
Jensen
Chem 110 Chap 5
Page: 12
Practice Example: How much heat is released
when 25.0 g of sodium peroxide (Na2O2)
undergo this reaction?
2 Na2O2(s) + 2 H2O(l) → 4 NaOH (s) + O2(g)
∆H° = −126 kJ
A.
B.
C.
D.
E.
20.2 kJ
40.4 kJ
67.5 kJ
80.8 kJ
126 kJ
Jensen
Chem 110 Chap 5
Page: 13
Calorimetry: Experimental measure of heat
flow (used to determine ∆Hrxn)
• review molar heat capacity & specific heat
Note: H2O is usually part of the surroundings
qsurr = Csurr m ΔT
Energy Conversion:
qsystem = — qsurr
Measure ∆T for surroundings in a controlled
environment (calorimeter)
Jensen
Chem 110 Chap 5
Page: 14
Coffee Cup Calorimeter (Constant Pressure
calorimeter): Measure temperature change in
solutions to calculate enthalpy
change (heat lost or gained)
qrxn for reactions under
constant pressure
qrxn = – qsoln = – Csoln m ΔT
Bomb calorimeter (constant volume calorimetry):
Heat evolved during combustion
is absorbed by calorimeter
contents causing a rise in water
temperature
qrxn = – Ccal × ΔT
Ccal = heat capacity of bomb
calorimeter
NOTE: because bomb calorimetry
is at constant volume (not constant pressure), heat
transferred is ΔE, not ΔH
Jensen
Chem 110 Chap 5
Page: 15
Quantitative Calorimetry Example
When a student mixes 50 mL of 1.0 M HCl
and 50 mL of 1.0 M NaOH in a coffee-cup
calorimeter, the temperature of the
resultant solution increases from 21.0 °C to
27.5 °C. Calculate the enthalpy change for
the reaction in kJ per mol of HCl, assuming
that the total volume of the solution is 100
mL, its density is 1.0 g/mL, and its specific
heat is 4.18 J/(gK).
A.
B.
C.
D.
E.
2.7 kJ/mol
−2.7 kJ/mol
54.4 kJ/mol
−54.4. kJ/mol
−108 kJ/mol
Jensen
Chem 110 Chap 5
Page: 16
Scratch paper:
Jensen
Chem 110 Chap 5
Page: 17
Hess’s Law
• Enthalpy is an extensive property
• ΔH for a reaction is equal in magnitude and
opposite in sign to ΔH for the reverse reaction
• ΔH for a reaction depends on the states of the
reactants and products (gas, liquid, solid)
Hessʼs Law: ΔH for an overall reaction is
equal to the sum of the individual steps
• Consequence of ΔH being a state function
C
ΔH2
ΔHrxn
ΔH1
A
Jensen
C
B
A
ΔHrxn
A
B
ΔH1
B
C
ΔH2
A+ B
B +C
ΔHrxn = ΔH1 + ΔH2
Chem 110 Chap 5
Page: 18
Example: Given the following information
A. H2(g) + F2(g) → 2HF(g)
ΔHA = –537kJ
B. 2H2(g) + O2(g) → 2 H2O(g) ΔHB = –572kJ
Determine ΔH for the reaction:
C. 2F2(g) + 2H2O(g)→ 4HF(g) + O2(g) ΔHC=?
Idea: find combinations of reactions
such that
n A + m B = C
then
n ΔHA + m ΔHB = ΔHC
Solve:
__ x A:
__ H2(g) + __ F2(g) → __ HF(g)
__ x B:
__ H2O(g) → __ H2(g) + __ O2(g)
C : 2 F2(g) + 2 H2O(g) → 4 HF(g) + O2(g)
__ x A: ΔH1 = __ x ΔHA = __________ kJ
__ x B: ΔH2 = __ x ΔHB = __________ kJ
ΔHC = ΔH1 + ΔH2 = _________ kJ
Jensen
Chem 110 Chap 5
Page: 19
Practice Example:
Given the following information:
2 SO2(g)+ O2(g) → 2 SO3(g)
ΔHA=–196kJ
2 S(s) + 3 O2(g) → 2 SO3(g)
ΔHB=–790kJ
What is ΔHrxn for the following reaction?
S(s) + O2(g) → SO2(g)
A. + 986 kJ
B. – 986 kJ
C. – 594 kJ
D. + 594 kJ
E. – 297 kJ
Jensen
Chem 110 Chap 5
Page: 20
Heat of formation
ΔHf (enthalpy of formation): heat given off (or
absorbed) when elements combine to form a
compound
combine
Elements
Compound
ΔHf
ΔHf° (standard enthalpy of formation): heat given
off (or absorbed) to form 1 mole compound when
all elements are in their standard states.
Definition of Standard State
1. P = 1 atm
2. T = 25°C (298K)
3. element is in its most stable state
(gas/liquid/solid)
For an element in its standard state: ΔHf°= 0
NOTE: ΔHf° values can be found in BLB Table 5.3 or
Appendix C
Two requirements for a reaction (at 25° C and 1
atm) to have ΔHrxn = ΔHf°:
1.
2.
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Chem 110 Chap 5
Page: 21
Standard States of the elements
(Phase under standard conditions
of 298K and 1 atm)
1. Metals: all are solid at 298K and 1 atm
except one (Hg)
2. Semi metals (metalloids): all are solids at
298K and 1 atm
3. Nonmetals at 298K and 1 atm
A: Noble gases (Group 8):
atomic gases
B: Diatomics: H2, N2, O2, Group 7
(F2, Cl2, Br2, I2)
H2, N2, O2, F2, Cl2, are gases
Br2 is a liquid
I2 is a solid
C: all other non-metals are solids:C
(graphite), S8, P, Se
Jensen
Chem 110 Chap 5
Page: 22
Examples of ΔHf°:
1. ΔHf° for methanol, CH3OH (l)
_ C(graphite) + _ H2(g) + _ O2(g)
1 atm
25°C
_ CH3OH(l)
• Balance the equation for 1 mole of product
• Reactants all in standard state
ΔHf° = –238.6 kJ/mol
Note: reactants might not have whole numbers for
coefficients & phase is important (l, s, g)
2. For which of the following reactions (at 25° C
and 1 atm) is ΔHrxn = ΔHf° ?
A. H2 (g) + F2 (g) → 2HF (g)
B. NO (g) + 1/2 O2 (g) → NO2 (g)
C. Pb (s) + Cl2 (g) →PbCl2 (s)
D. 2 Na (s) + N2 (g) + 3 O2 (g) → 2 NaNO3 (s)
E. S (s) + O3 (g) → SO3 (g)
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Chem 110 Chap 5
Page: 23
Standard Enthalpy of a Reaction: ΔHrxn°
Heat of reaction (ΔHrxn) when all reactants
and products are in the standard state.
Obtain ΔHrxn° from ΔHf° (on data sheet):
ΔHrxn°= Σn ΔHf°(products)
− Σm ΔHf°(reactants)
n, m are stoichiometric coefficients of each
individual product and reactant, respectively
This is an application of Hessʼs Law
energy
elements
− Σm ΔHf°(reactants)
Σn ΔHf°(products)
reactants
ΔHrxn
products
ΔHrxn° = Σn ΔHf°(products) − Σm ΔHf°(reactants)
Jensen
Chem 110 Chap 5
Page: 24
Example: Sugar is broken down in the body to
produce carbon dioxide and water. Using the
information in the table, determine how much
energy is produced by the controlled combustion
of one mol of sugar in the body.
C12H22O11(s) + 12 O2(g) → 11 H2O(g) + 12 CO2(g)
∆Hf° (kJ/mol)
A.
B.
C.
D.
E.
− 1585.7
− 2856.3
− 5160.8
− 5644.8
− 9602.8
H2O(g)
-241.8
CO2(g) C12H22O11(s)
-393.5
-2221.0
kJ
kJ
kJ
kJ
kJ
NOTE: ΔHf° values can be found in BLB Table 5.3 or
Appendix C
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Chem 110 Chap 5
Page: 25
Practice Example: The thermite reaction below is
used for welding. What is the ΔHrxn° for the
reaction involving 1 mole of Al?
2 Al(s) + Fe2O3(s) → Al2O3(s) + 2 Fe (s)
ΔHf° of Al2O3(s) = – 1669.8 kJ/mol
ΔHf° of Fe2O3(s) = – 822.2 kJ/mol
A.
B.
C.
D.
E.
+ 847.6 kJ
– 847.6 kJ
+ 1895 kJ
– 423.8 kJ
– 2492 kJ
Jensen
Chem 110 Chap 5
Page: 26
Where does the Energy Come From?
• The overall reaction has two steps: breaking the
original bonds, and forming new ones.
Bond breaking is always ______thermic
Bond formation is always ______thermic
√ Bond Enthalpies (bond energy) can be used to
estimate ΔHrxn.
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Chem 110 Chap 5
Page: 27
Covalent Bond Length & Energy
Bond length: distance between nuclei
bond
C–C
C=C
C≡C
Bond energy
(kJ/mol)
348
614
839
Bond length
(pm)
154
134
121
More electrons shared, shorter bond length
Shorter bond length, stronger the bond
Bond (dissociation) energy (D): enthalpy of bond
breaking reaction in the gas phase.
√D>0
(ΔH > 0)
√ For diatomics, D is ΔH of one reaction:
H–H(g) → 2H(g)
DH-H= ΔHrxn = 436kJ/mol
√ For polyatomics, D is an averaged quantity
H–O–H(g) → H–O(g) + H(g)
ΔHrxn = 494kJ/mol
H–O(g)
ΔHrxn = 424kJ/mol
→ H(g) + O(g)
DH–O = 463 kJ/mol * (value obtained from averaging
over many molecules)
√ Exact values are known for some bonds, but the
average values are not exact for any one case
(unlike ΔHf°)
Jensen
Chem 110 Chap 5
Page: 28
Estimating ΔHrxn using Bond Energies
ΔHrxn ≅ ΣnDbroken − ΣmDformed; (n, m = # bonds)
Energy is given off (−) when bonds form.
Example: Based on the data below, what is
the estimated ΔHrxn for the combustion of 1
mo of propene (CH2=CH-CH3)?
Bond
C-C
C=O
O=O
A.
B.
C.
D.
E.
D (kJ/mol)
348
799
495
Bond
C=C
C–H
O–H
D (kJ/mol)
614
413
463
− 3809 kJ
+ 1905 kJ
− 1905 kJ
+ 3809 kJ
− 2438 kJ
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Chem 110 Chap 5
Page: 29
Scratch paper:
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Chem 110 Chap 5
Page: 30
Practice Example: The reaction below is used
to produce methanol:
CO(g) + 2H2(g) → CH3OH
∆Hrxn = –128kJ
Use the data below to calculate the
approximate C≡O bond enthalpy.
Bond
C-O
C–H
C≡O
A.
B.
C.
D.
E.
D (kJ/mol)
358
413
??
Bond
C=O
O–H
H–H
D (kJ/mol)
799
463
436
417 kJ/mol
574 kJ/mol
687 kJ/mol
896 kJ/mol
1060 kJ/mol
Jensen
Chem 110 Chap 5
Page: 31
Scratch paper:
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Chem 110 Chap 5
Page: 32