Energy & Chemistry
• Burning peanuts
supply sufficient
energy to boil a cup
of water.
• Burning sugar
(sugar reacts with
KClO3, a strong
oxidizing agent)
Ch. 6 Energy and Chemistry
1. Vocabulary
2. Heat using specific heat and heat capacity
3. Enthalpy from single reaction
a) Single reagent
b) Limiting reagent
4. Enthalpy from multiple reactions
a) Hess’ Law
5. Enthalpy of reaction from ∆Hf°
Chapter 6: Energy and Chemical Change
• Energy is the ability to do work
(W = F×D or P×V) and/or supply heat
• Kinetic energy is the energy an object has
because of its motion
– For an object of mass m with velocity v
KineticEnergy KE  m v
1
2
PotentialEnergy  m gh
2
• The law of conservation of energy states
that energy cannot be created or destroyed
Kinetic Energy: The Energy Of
Motion
• KE=½mv2
• Energy can be transferred by moving
particles
• Collision of fast particles with slower
particles causes the slow particle to speed
up while the fast molecule slows
– this is why hot water cools in contact with
cool air
6.1 An object has energy if it is capable of doing work
5
Potential Energy
Breaking Bonds requires energy
Question:
A child on a swing has the highest potential
energy when:
a: the swing is at its highest point
b: the swing is at its lowest point
c: the swing is halfway between the
highest and lowest points
Law Of Conservation Of Energy
• Energy cannot be
created or destroyed
but can be transformed
from one form of
energy to another
• Also known as the first
law of thermodynamics
• How does water falling
over a waterfall
demonstrate this law?
6.1 An object has energy if it is capable of doing work
8
Internal Energy is Conserved
• 1st Law of Thermodynamics: For an isolated
system the internal energy (E) is constant:
Δ E = Ef - Ei = 0
Δ E = Eproduct - Ereactant = 0
• We can’t measure the internal energy of
anything, so we measure the changes in energy
• E is a state function
• E = work + heat
6.2 Internal energy is the total energy of an object’s molecules
9
Heat And Temperature Are Not The
Same
• The temperature of an object is
proportional to the average kinetic energy
of its particles—the higher the average
kinetic energy, the higher the temperature
• Heat is energy (also called thermal energy)
transferred between objects caused by
differences in their temperatures until they
reach thermal equilibrium
6.1 An object has energy if it is capable of doing work
10
Units
• The SI unit of energy is the joule (J)
• calorie (cal)
1 J  (2 kg)
1
2

1m 2
1s
 1 kg m s
2
-2
1 cal  4.184J (exactly)
– The dietary Calorie (note capital), Cal, is actually
1 kilocalorie
Energy Transfer
• When a cold and hot object come into
contact, they eventually reach thermal
equilibrium (the same temperature)
•
Movie1/06_heat transfer
•
energy transferred as heat comes from the object’s internal energy
Energy Distribution
Energy Transfer
a) Thermal equilibrium: the same average KE
for molecules in both objects
Question:
As the temperature of a substance increases,
what happens to the average kinetic energy
of the particles?
a: increases
b: stays the same
c: decreases
Question:
Which sample of a substance has the higher
most probable molecular speed?
a: the warmer sample
b: the cooler sample
c: both samples have the same most
probable speed
Definitions
• current condition: state
• Internal energy is a state function
• state functions: independent from the
mechanism or method by which a change
occurred
• object we are interested in: system
• Everything else: surroundings
• A boundary
• System and Surroundings together: universe
Open, closed, isolated
Question:
How is energy transferred as heat?
a: energy flows from cooler objects to
warmer ones.
b: energy flows from warmer objects to
cooler ones.
c: molecules from the warmer object
enter the cooler object.
Question:
What type of system exchanges energy but
not matter with the surroundings?
a: open
b: closed
c: isolated
A cast iron skillet is moved from a hot oven
to a sink full of water. Which of the
following is not true?
A. The water heats
B. The skillet cools
C. The heat transfer for the skillet has a (-)
sign
D. The heat transfer for the skillet is the same
as the heat transfer for the water
E. None of these are untrue
6.3 Heat can be determined by measuring temperature changes
22
The symbol for heat is “q” – two ways to calculate
Question:
Which process is exothermic?
a: N2  2N
b: CO2  CO + O
c: BF3 + NH3  BF3NH3
Ch. 6 Energy and Chemistry
1. Vocabulary
2. Heat using specific heat and heat capacity
3. Enthalpy from single reaction
a) Single reagent
b) Limiting reagent
4. Enthalpy from multiple reactions
a) Hess’ Law
5. Enthalpy of reaction from ∆Hf°
Heat Capacity (C) relates the
heat (q) to an objects
temperature change
q  C (tfinal - tinitial )  C t
The heat capacity is the amount
of heat needed to raise the
object’s temperature by one
degree Celsius and has the
units J/°C
James Joule
1818-1889
q  C (tfinal - tinitial )  C t
• If my large ceramic coffee cup
has a heat capacity of 35 J/ºC,
how much heat does it take to
raise the temperature by 50ºC
James Joule
1818-1889
Specific Heat (s) is an intensive property, and
is unique for each substance
• q=m×Δt×s
Specific Heat
Subst ance J g C (25 C)
-1
-1
Copper
Gold
0.387
0.129
Silver
W at er
0.235
4.18
q  m s t
heat gain/lost = q = (mass)(SHeat)(T)
If 25.0 g of Al (SH=.902 J/gK) cools from
310 oC to 37 oC, how many joules of heat
energy are lost by the Al?
where T = Tfinal - Tinitial
Problem Solving
A sample of nickel weighing 425 grams was
initially at a temperature of 26.20C. It
required 975 joules of heat energy to increase
the temperature to 31.55 C. What is the
specific heat of nickel?
Problem Solving
• A 500.0 gram sample of water is initially at
25.0C. It absorbs 50.0 kJ of heat from its
surroundings, what is the final temperature
in C? specific heat of water = 4.184 J/gC
If my large ceramic cup (heat
capacity = 35 J/ºC) is at 25ºC
when 250 mL of 92ºC coffee is
added to it. What will be the final
temperature of my cup with coffee
assuming no other heat loss/gain?
James Joule
1818-1889
Problem Solving
Determine the final temperature when 255 g
of iron is heated to 100.0°C in boiling
water and it is then quickly placed into
100.0 g water at 23.4°C. SHiron=0.449 J/gK
q (water) = -q (iron)
Question:
How much energy is required to heat 50
grams of water from 24ºC to 34ºC?
a: 50 cal
b: 500 cal
c: 500 J
Question:
How many dietary calories equals 1000 kJ?
a: 240
b: 240,000
c: 4184
d: 4.184
Question:
Adding 5 calories of heat energy to 1 gram of copper raises
the temperature of the copper by 54.3 ºC. Adding 5
calories of heat energy to 1 gram of water raises the
temperature of the water by 5.0 ºC. Which statement must
be true?
a: Copper has a larger specific heat capacity than water.
b: Water has a larger specific heat capacity than copper.
c: There must be some mistake in these measurements.
d: Mass is being converted into energy in this process.
Question:
100 grams of water cool from 14 ºC to 13 ºC.
What is the heat change for the water?
a: 100 cal
b: -100 cal
c: 418 cal
d: - 418 cal
FIRST LAW OF
THERMODYNAMICS
heat energy transferred
E = q + w
work done
by the
system
energy
change
first law of thermodynamics, which says that
energy cannot be created or destroyed
In reactions
where gases
are produce
or consumed
qv and qp can
be very
different
w   PV whereP is theopposingpressure
The heat produced by a combustion reaction is
called the heat of combustion usually in a
bomb calorimeter. The reaction is run at
constant volume so that ΔE = qV
Heats of reactions in solution are usually run in
open containers at constant pressure, so that qP =
ΔE + PΔV = ΔH
A calorimeter has metal parts (heat capacity of 850.0
J/degree) and 1020 grams of oil (specific heat 2.248
J/gC, both at 24.50C.
Two metal slugs, one a 460.0 g piece of cobalt (specific
heat 25.12 J/mol-degree), one a 360.0 g piece of
cadmium (specific heat=25.34 J/mol-degree), were
removed from an oven maintained at 240.0C and
added to the calorimeter.
If no heat was lost to the surroundings, what would be
the final temperature of inside the calorimeter?
Question:
Which equals qP?
a: w
b: ΔE
c: ΔH
d: none of these
Hindenburg, 6 May 1937
http://www.flatrock.org.nz/topics/flying/historic_craft_flies_again.htm
http://video.google.com/videoplay?docid=-2380118142773657669&q=hindenburg
A sample of 50.00mL of 0.125M HCl at 22.36 ºC is added to
a 50.00mL of 0.125M Ca(OH)2 at 22.36 ºC. The
calorimeter constant was 72 J/g ºC. The temperature of
the solution (s=4.184 J/g ºC, d=1.00 g/mL) climbed to
23.30 ºC. Which of the following is not true?
A. qczl=67.9 J
B. qsolution= 393.3J
C. qrxn = 461.0 J
D. qrxn = -461.0 J
E. None of these
6.5 Heats of reaction are measured at constant volume or constant pressure
46
A sample of 50.00mL of 0.125M HCl at 22.36
ºC is added to a 50.00mL of 0.125M
Ca(OH)2 at 22.36 ºC. The calorimeter
constant was 72 J/g ºC. The temperature
of the solution (s=4.184 J/g ºC, d=1.00
g/mL) climbed to 23.30 ºC. Which of the
following is not true?
A. qczl=67.9 J
B. qsolution= 393.3J
C. qrxn = 461.0 J
D. qrxn = -461.0 J
E. None of these
6.5 Heats of reaction are measured at constant volume or constant pressure
47
Ch. 6 Energy and Chemistry
1. Vocabulary
2. Heat using specific heat and heat capacity
3. Enthalpy from single reaction
a) Single reagent
b) Limiting reagent
4. Enthalpy from multiple reactions
a) Hess’ Law
5. Enthalpy of reaction from ∆Hf°
Energy is in chemical bonds
Consider the decomposition of water:
H2O(g) + 243 kJ  H2(g) + 1/2 O2(g)
Endothermic reaction — heat is a “reactant”
HRXN = + 243 kJ
2H2(g) + O2(g)  2H2O(g) + 486 kJ
HRXN = - 243 kJ
An enthalpy change for standard conditions
is denoted H 
N 2 ( g )  3H2 ( g )  2NH3 ( g )
H   92.38 kJ
• The physical states are important
• The law of conservation of energy requires
2 NH3 ( g )  N 2 ( g )  3H 2 ( g )
H   92.38kJ
Given 2H2(g) + O2  2 H2O for which
HRXN = -483.6, how much heat is evolved
when one mole water is formed? How is much
heat is consumed when 3 moles of water are
decomposed?
How much heat is evolved when 2.500 moles of
water are formed by the same reaction?
Question:
Consider the following information:
N2 (g) + 3 H2 (g)  2 NH3 (g) ΔH º = - 92.38 kJ
What is ΔH º when 1.0 mole of liquid ammonia is
produced?
a: - 92.38 kJ
b: - 46.19 kJ
c: - 184.8 kJ
d: more information is needed
Question:
How much heat would be released if 1.0 mole
of compound X were produced?
2A + Z  2X
ΔH º = - 80 kJ
a: - 80 kJ
b: - 40 kJ
c: - 160 kJ
The thermite reaction produces a
tremendous amount of heat. If you
begin with 10.0 g of Al and excess
Fe2O3, how much heat in kilojoules
is evolved at constant pressure?
2 Al (s) + Fe2O3 (s)  Al2O3 (s) + 2 Fe (s)
Hrxn = -851.5 kJ
Ch. 7 Energy and Chemistry
1. Vocabulary
2. Heat using specific heat and heat capacity
3. Enthalpy from single reaction
a) Single reagent
b) Limiting reagent
4. Enthalpy from multiple reactions
a) Hess’ Law
5. Enthalpy of reaction from ∆Hf°
An enthalpy diagram is a graphical representation of
alternate paths between initial and final states
Remember to include the physical states of reactants and
products in thermochemical equations.
USING ENTHALPY
Making H2 from H2O involves two steps.
H2O(liq) + 44 kJ  H2O(g)
H2O(g) + 242 kJ  H2(g) + 1/2 O2(g)
-----------------------------------------------------------------------
H2O(liq) + 286 kJ  H2(g) + 1/2 O2(g)
HESS’S LAW—
If a rxn. is the sum of 2 or more others, the net
H is the sum of the H’s of the other rxns.
Rules for Manipulating Thermochemical Equations:
1) When an equation is reversed the sign of the
enthalpy change must also be reversed.
2) Formulas canceled from both sides of an
equation must be for substances in identical
physical states.
3) If all the coefficients of an equation are
multiplied or divided by the same factor, the
value of the enthalpy change must likewise be
multiplied or divided by that factor.
Using Standard Enthalpy Values
H2O(g) + C(graphite)  H2(g) + CO(g)
From reference books we find
• H2(g) + 1/2 O2(g)  H2O(g)
Hf of H2O vapor = - 242 kJ/mol
• C(s) + 1/2 O2(g)  CO(g)
H f of CO = - 111 kJ/mol
What is the energy of the following process:
6A + 9B + 3D + 1 F→2 G
Given that:
• C → A + 2B ∆H=20.2 kJ/mol
• 2C + D →E + B
∆H=30.1 kJ/mol
• 3E + F →2G ∆H=-80.1 kJ/mol
A. 70.6 kJ
B. -29.8 kJ
C. -111.0 kJ
D. None of these
6.7 Thermochemical equations can be combined because enthalpy is a state function
66
Problem solving:
Given C(s) + O2(g) --> CO2(g) ΔHrxn(-393.5 kJ)
and CO(g) + 1/2 O2(g) --> CO2(g) ΔHrxn(-283 kJ)
calculate enthalpy change for reaction:
C(s) + 1/2 O2(g) --> CO(g)
Question:
Consider the following information:
A + F  2G
ΔH º = +100 kJ
AG+E
ΔH º = +200 kJ
What is ΔH º for 2E + F  A?
a: +300 kJ
b: +100 kJ
c: -100 kJ
d: -300 kJ
e: +500 kJ
Problem solving:
Calculate the enthalpy of combustion of
ethylene C2H4 to CO2 and H2O
2C(s) + 2H2(g)  C2H4 ΔHrxn(52.3 kJ)
C(s) + O2(g)  CO2(g) ΔHrxn(-393.5 kJ)
H2(g) + 1/2O2(g)  H2O(l) ΔHrxn(-285.8 kJ)
Problem Solving
• Two oxides of copper can be made from
copper by the following reactions:
2 Cu(s) + O2(g)  2 CuO(s) H= -310 kJ
2 Cu(s) + ½O2(g)  Cu2O(s) H= -169 kJ
Construct an enthalpy diagram and find H
for the following
• Cu2O(s) + ½O2  2CuO(s)
Ch. 6 Energy and Chemistry
1. Vocabulary
2. Heat using specific heat and heat capacity
3. Enthalpy from single reaction
a) Single reagent
b) Limiting reagent
4. Enthalpy from multiple reactions
a) Hess’ Law
5. Enthalpy of reaction from ∆Hf°
o
H
f,
standard molar enthalpy
of formation
H2(g) + 1/2 O2(g)  H2O(g)
Hof = -241.8 kJ/mol
By definition,
Hof = 0 for elements in their
standard states.
Question:
For which equation is ΔH º equal to
ΔH º F for SO2 (g)?
a: S8 (s) + 8 O2 (g)  8SO2 (g)
b: S8 (s) + 8 O2 (g)  SO2 (g)
c: 1/8 S8 (s) + O2 (g)  SO2 (g)
standard states have been defined for reporting
heats of reactions
1 bar pressure for all gases and
1 M concentration for aqueous solutions
25 °C (298 K) is often specified as well
Measurements made under standard state conditions
have the ° mark: eg ΔH°
The standard heat of reaction is the value of the
enthalpy change occurring under standard
conditions involving the actual number of moles
specified the the equation coefficients
Horxn = Hof (products) - Hof (reactants)
In general, when ALL enthalpies of
formation are known,
Calculate ΔH
of reaction?
Using Standard Enthalpy Values
Calculate the heat of combustion of methanol,
CH3OH(g) + 3/2 O2(g)  CO2(g) + 2 H2O(g)
Horxn =  Hof (prod) -  Hof (react)
TABLE 7.2
Substance
Ag(s)
Standard Enthalpies of Formation of Typical Substances
(kJ mol−1)
0
Substance
(kJ mol−1)
Substance
(kJ mol−1)
CaO(s)
−635.5
KCl(s)
−435.89
−1433.7
AgBr(s)
−100.4
Ca(OH)2(s)
−986.59
K2SO4(s)
AgCl(s)
−127.0
CaSO4(s)
−1432.7
N2(g)
0
CaSO4·½H2
O
−1575.2
NH3(g)
−46.19
−1669.8
CaSO4·2H2O(
s)
−2021.1
NH4Cl(s)
−315.4
Al(s)
Al2O3(s)
C(s, graphite)
0
0
Cl2(g)
0
NO(g)
90.37
CO(g)
−110.5
Fe(s)
0
NO2(g)
33.8
CO2(g)
−393.5
Fe2O3(s)
−822.2
N2O(g)
81.57
CH4(g)
−74.848
H2O(g)
−241.8
N2O4(g)
9.67
CH3Cl(g)
−82.0
H2O(l)
−285.9
N2O5(g)
11
CH3I(g)
14.2
H2(g)
0
Na(s)
0
CH3OH(l)
−238.6
H2O2(l)
−187.6
NaHCO3(s)
−947.7
CO(NH2)2(s) urea)
−333.19
HBr(g)
−36
Na2CO3(s)
−1131
CO(NH2)2(aq)
−391.2
HCl(g)
−92.30
NaCl(s)
−411.0
C2H2(g)
226.75
HI(g)
26.6
NaOH(s)
−426.8
C2H4(g)
52.284
HNO3(l)
−173.2
Na2SO4(s)
−1384.5
C2H6(g)
−84.667
H2SO4(l)
−811.32
O2(g)
0
C2H5OH(l)
−277.63
HC2H3O2(l)
−487.0
Pb(s)
0
Ca(s)
CaBr2(s)
0
Hg(l)
0
−682.8
Hg(g)
60.84
PbO(s)
S(s)
−219.2
0
Using Standard Enthalpy Values
CH3OH(g) + 3/2 O2(g)  CO2(g) + 2 H2O(g)
Horxn =  Hof (prod) -  Hof (react)
Horxn = 1 mol Hof (CO2) + 2 mol Hof (H2O)
- {3/2 mol Hof (O2) + 1 mol Hof (CH3OH)}
= (-393.5 kJ) + 2 (-241.8 kJ)
- {0 + (-201.5 kJ)}
Horxn = -675.6 kJ per mol of methanol
Problem Solving
Calculate H for the following reactions:
2NO(g) + O2(g)  2NO2(g)
Hf (kJ/mol)
NO (g) = 90.4
NO2 (g) = 33.85
N2O (g) = 81.56
Problem Solving:
Calculate the enthalpy of combustion of
acetylene C2H2 to produce carbon dioxide
and water, given the molar enthalpy of
formations for the following compounds:
C2H2 = +226.7 kJ/mol
CO2 = -393.5 kJ/mol
H2O = -285.8 kJ/mol
Question:
For which reaction is ΔH º equal to
ΔH ºCombustion for C2H2?
a: C2H2 (g) + 5/2 O2 (g)  2 CO2 (g) + H2O (g)
b: C2H2 (g) + 5/2 O2 (g)  2 CO2 (g) + H2O (l)
c: 2 C2H2 (g) + 5 O2 (g)  4 CO2 (g) + 2 H2O (g)
d: 2 C2H2 (g) + 5 O2 (g)  4 CO2 (g) + 2 H2O (l)
Question:
Which has ΔH º F = 0?
a: O (g)
b: H2 (l)
c: Hg (l)
d: F2 (s)
Question:
When a reaction is reversed and multiplied by
2, what happens to its ΔH º?
a: sign changed; multiplied by 1/2
b: sign changed; multiplied by 2
c: squared; then the sign changed
d: sign kept; multiplied by ½
Question:
What does the superscript º mean in notation
such as ΔH º?
a: measurements at absolute zero
b: degrees Celsius
c: measurements at standard conditions
d: measurements taken in dietary
calories
Problem Solving
• Calculate the
standard molar
enthalpy of
combustion for
ethanol C2H5OH
using the standard
molar enthalpies
of formation
Question:
How much heat would be released if 1.0 mole
of compound X were produced?
2A + Z  2X
ΔH º = - 80 kJ
a: - 80 kJ
b: - 40 kJ
c: - 160 kJ
Problem Solving
Use the following thermodynamic equations to find
the enthalpy change for the following reaction:
M2O3(s) + 3CO(g)  3CO2 + 2M(s)
M(s) + CO2(g)  CO(g) + MO(s) H = -140.0 kJ
3MO(s) + CO2(g)  CO(g) + M3O4(s) H = +1.20 kJ
2M3O4(s) + CO2(g)  CO(g) + 3M2O3(s) H =-380.0 kJ
M2O3(s) + CO2(g)  CO(g) + 2MO2(s)
H = -121.8 kJ
2 MO2(s) + CO2(g)  CO(g) + M2O5(s) H = +344.5 kJ
Problem Solving
White phosphorus, P4, ignites in air to
produce heat, light and P4O10. If 3.56 g
of P4 is burned, 37.4 kJ of heat is evolved
at constant pressure. What is the molar
enthalpy of reaction for the burning of
white phosphorus?
Question:
When two atoms come together and form a
bond, what happens to the energy?
a: it disappears into space.
b: it is stored in the form of potential
energy.
Question:
An aqueous chemical reaction in a coffee cup
calorimeter causes the water temperature in
the cup to increase. Is the reaction
endothermic or exothermic?
a: endothermic
b: exothermic
c: neither
d: it’s impossible to tell
Question:
Which is not a state function?
a: internal energy
b: enthalpy
c: heat
d: temperature
Problem Solving:
A 0.3000 g sample of tin ore was dissolved in acid
solution converting all the tin to tin(II). In a
titration, 8.08 mL of 0.0500 M KMnO4 was
required to oxidize the tin(II) to tin(IV). What
was the percentage tin in the original sample?
The reaction also produced MnO2
2

4
3Sn  2MnO  8H  3Sn  2MnO2  4H2O
4
Problem Solving
When 250.00 mL of 0.200M of sodium sulfite is
reacted with 3.50 g of potassium dichromate
(K2Cr2O7 FW=294.18) in an acidic solution, the
products of the reaction include sulfate ion and
chromium(III) ion. What is the concentration of
chromium(III) ion if the total volume is now
251.25 mL?
How would the problem be different with 5.50 g of
potassium dichromate?
Problem Solving
A solution of sodium thiosulfate (Na2S2O3) reacts
with chlorine gas in an acidic solution to give
sulfate ion and chloride ion. How many mL of
0.100 M sodium thiosulfate would be needed to
react with 4.25 g of chlorine gas?
Problem Solving
All the iron in a 2.000 g sample of iron ore is
dissolved in an acidic solution and converted to
iron(II) ion. When titrated with 0.100 M
potassium permanganate the iron was oxidized to
iron(III) and the permanganate to manganese
dioxide. The titration required 27.45 ml of
permanganate solution to reach the endpoint.
Problem Solving
A student, using the same calorimeter, then reacted
100.0 mL of 1.00 M NaOH with 100.0 mL of 1.00
M HCl. The Tinitial of the base was 24.70 °C, and of
the acid 24.74 °C, while the Tfinal was 30.53 °C.
Calculate the molar enthalpy of neutralization for this
reaction. Density of solution 1.02 g/mL, specific
heat of solution 4.016 J/gK
Problem Solving
How much heat is required to raise the
temperature of 300. g of copper from
21°C to 75°C? SH(copper)=0.385 J/gK
If the same amount of heat is added to
each of the following metals, which
will experience the greatest
temperature change?
Specific Heat in J/gC: Cu(0.385)
Mg(1.02) Hg(0.138) Pb(0.129)
Bonus: Problem Solving
How much heat energy is needed to raise
the temperature of 500 gallon of water
in a whirlpool bath by 5C. Assume
the specific heat of the water is 1.00
cal/gC
The specific heat of lead is 0.129 J/gC.
If 74.0 J of heat is added to a 125 piece
of lead at 23 C, what is the final
temperature of the lead?
Bonus: Problem Solving
What is the specific heat of ethanol in J/gC if
560.0 J of heat is required to raise the
temperature from 22.0C to 39.0C for a
50.0 g sample?
A lake that is 2.0 miles square and has an
average depth of 10. ft contains 3.2 x1010
liters of water. How many joules of energy
must be transferred from the lake when its
temperature decreases by 10.°C. Assume
lake density 1.0 g/mL SHH20 = 4.184 J/gK