Download Thermochemistry - Moorpark College

Thermochemistry Ch 6
Page |1
Unit 2:
Chapter 6: Thermochemistry
Homework: Read Chapter 6
Work out sample/practice exercises in the sections as you read,
Suggested Problems: 35, 37, 41, 45, 53, 59, 65, 71, 79, 81, 87, 91,103, 127, 135
Check the MasteringChemistry due dates
Chemical Thermodynamics:
Chemical thermodynamics is the study of the changes and the transfers of energy
that accompany chemical and physical processes.
It is VERY IMPORTANT to always identify the phases in chemical reactions. This unit
also will occasionally require balancing reactions with fractions.
Thermodynamics addresses 3 fundamental questions:
1) Will two or more substances spontaneously react when mixed under specified
conditions or will it be nonspontaneous (will not react)?
2) What energy changes and transfers are associated with the reaction?
3) To what extent does a reaction occur? (learn in Chem 1B)
Systems tend toward…
1) Minimizing potential energy
2) Maximizing entropy (degrees of freedom or ways of being).
Thermochemistry tells us nothing of the reaction rate. Reactions may be
spontaneously favored under specified conditions, but may not react at an
observable rate, or reactions may need a large initial amount of energy to initiate
and reach the activation energy.
When you see X where X can be anything remember that it is just the change from the
final minus the initial values. X = Xfinal – Xinitial
Thermochemistry Ch 6
Page |2
Thermo examples:
a) Methane will spontaneously react and combust in oxygen, but the reaction
does not initiate without a spark.
b) Carbon should be thermodynamically favorable to combust in oxygen to
create carbon dioxide, but fortunately diamond rings (carbon) will not
react at an observable rate.
c) Ice will spontaneously melt into water, but only if the temperature is
above freezing.
In this unit we will be studying physical and chemical processes accompanied by the
transfer of heat of work. Nuclear energy in which the atom changes is different.
Example 1:
Systems tend toward minimizing energy and maximizing entropy. Identify the
side that minimizes the energy. Identify the side that maximizes ways of being
(entropy, S). Assume the same temperature and atmospheric pressure for each.
a) Top of waterfall
or
Bottom of waterfall
b) 1 mole H2 molecules
or
2 mole of H atoms
c) H2O (g) at 100°C
or
H2O (l) at 100°C
Nature of Energy:
Energy is the capacity to do work or transfer heat. We will generally use joules.
SI units of energy: 1 Joule = 1 kg m2/s2
Energy Conversions
1 calorie (cal)
=
4.184 joules (J) (exact)
1 Calorie (Cal, food)
=
1000 cal = 1 kcal = 4184 J
1 kilowatt-hour (kWh)
=
3.60 x 106 J
Energy can be classified into two major categories, potential and kinetic.
Thermochemistry Ch 6
Page |3
Potential Energy is the energy by virtue of its position of composition.
a) For an object that has height, potential energy is the mass x gravity (9.8 m/s2)
x height; Ep = mgh.
b) Another form is electrostatic potential energy from interactions with charged
particles; Eel = kQ1Q2/d, where k = 8.99 x 109 J m/C2 and C is coulomb that
is a unit of charge, Q is charge, and d is distance. Like charges repel and
opposites attract.
c) Potential energy also comes from the arrangement of chemical compositions.
When bonds break and new ones form in a chemical reaction, the energy
change generally is due to the changes in potential energy/composition.
Kinetic Energy is the energy of motion. For a moving object kinetic energy is half
the mass x velocity squared; Ek = ½ mv2
Example 2: What is the potential energy of a 400 g ball on top of a building that
is 30.0 m tall?
Example 3: What is the kinetic energy of a 400 g ball moving at 30.0 m/s?
When studying Thermochemistry we need to understand the following terms…
System: the portion we study
Surroundings: everything else affected by the portion we study
Universe: the combined system plus surroundings. In the lab experiments our
universe will be a calorimeter and its contents.
Thermochemistry Ch 6
Page |4
Some Forms of Energy:
• Electrical
kinetic energy associated with the flow of electrical charge
• Heat or thermal energy
kinetic energy associated with molecular motion
• Light or radiant energy
kinetic energy associated with energy transitions in an atom
• Nuclear
potential energy in the nucleus of atoms
• Chemical
potential energy due to the structure of the atoms, the attachment between
atoms, the atoms’ positions relative to each other in the molecule, or the
molecules, relative positions in the structure
Laws of Thermodynamics:
Laws of thermodynamics are useful in predicting outcomes, but unlike a theory do
not explain the expected behavior.
Laws of Thermodynamics:
Zeroth Law: Temperature Concept: Temperature measures the intensity of
hotness or coldness of an object. When two objects are brought
together heat always flows spontaneously from a hotter object to a
colder object until thermal equilibrium is reached. Think of the heat
flow direction when placing your hand in ice water verses hot water.
First Law: Law of Conservation of Energy: Energy cannot be created or
destroyed. The energy of the universe is constant.
This law works for ordinary chemical and physical processes. It was
later combined with another (conservation of matter) to include
nuclear reactions after E =mc2. It was changed to the Law of
Conservation of Matter and Energy
Second Law: Entropy Principle: Every spontaneous change increases the entropy
of the universe. This does not mean local decreases in entropy cannot
take place. For the entropy of a system to decrease, the entropy of the
surroundings must increase more.
Third Law: Zero Entropy Established: The entropy of a pure, perfect, crystalline
substance (no disorder) is zero at the temperature of absolute zero, 0
Kelvin.
It is impossible to reach absolute zero. It has been tried in the lab and
gets close.
Thermochemistry Ch 6
Page |5
Example 4: What is the difference between temperature and heat?
Example 5: Review your temperature conversions between °F, °C, and Kelvin.
K = (°C) + 273.15 °F = 1.8(°C) + 32
°C = (K) - 273.15
°C = (°F-32)/1.8
a) 37.0°C = ? °F = ? K
b) -40.0 °F = ? °C = ? K
c) 800 K = ? °C = ? °F
More Definitions:
Internal Energy, E or U; the sum of all the kinetic and potential energies of all
components. This includes all motions of vibration, rotation, movement through
space of the object, its atoms, nuclei, and even including subatomic particles like
electrons. We cannot truly measure all of this, so we generally try to solve for the
difference or change in internal energy in our studies. E = Efinal – Einitial
When solving for E we work with a constant volume system. Units are generally
kJ/mol. When the energy is measured in J or kJ for a particular system, q v may be
used.
Enthalpy, H; the energy accounting for heat flow in constant pressure processes.
Units are generally kJ/mol. When the energy is measured in J or kJ for a particular
system, qp may be used.
Work, w; the energy used to cause an object to move. We will mostly deal with
pressure-volume work which involves the compression or expansion of gases.
Work = -P V ; if using atm and liter: convert to J with R/R (101.3 J/L atm)
Work = - nRT; use the R with joules: R = 8.314 J/mol K
For most reactions the energy from work is quite small.
Expansion work is (-); work is done by the system on the surroundings.
Compression work is (+); work is done on the system by the surroundings
Thermochemistry Ch 6
Page |6
Relating together; qv = qp + w , or it can be written E = H + w
Endothermic: a process that absorbs heat (+)
Exothermic: a process that releases heat (-)
State Function: a property that
depends only on the present
state and not the path the
system took to reach the state.
In Thermochemistry state
functions are generally
capitalized. (Notice work is a
small w and is not a state
function) Some examples of
state functions include: E,
H, V, P, T.
Example 6:
a) As 1.00 mol of H2O gas converts to H2O liquid, 40.7 kJ of energy are
transferred at 100°C temperature and 1.00 atm pressure. What is the sign
for this conversion? Is this internal energy ( E) or enthalpy ( H)? Is this
process endothermic or exothermic?
b) At 1.00 atm constant pressure, the volume of 1.00 mole of H2O gas
decreases from 30.6 liters down to 18 ml of H2O liquid, what is the work
energy in L atm and in J.
c) Identify or solve for work, enthalpy and internal energy for this example.
Thermochemistry Ch 6
Page |7
Example 7:
Predict the sign for work, determine if the system is expanding or compressing and
state if work is done on or by the system.
a)
2 NO (g) + O2 (g)

2 NO2 (g)
b)
PCl5 (g)

PCl3 (g) + Cl2 (g)
c)
2 C4H10 (g) + 13 O2 (g)

10 H2O (l) + 8 CO2 (g)
Example 8:
Which of the following are state functions and which are not state functions ?
a)
The temperature of an ice cube.
b)
The enthalpy of fusion for water.
c)
The time to travel 1-mile
Enthalpies of Reaction:
Thermochemical Equations with Enthalpy:
1) Always include the phases (s, l, g, aq) of the substances. Energy varies with
phase.
2) Assume temperature is constant for the reactants at the beginning and the
products at the end, even though the temperature may change during the
reaction. H is the energy required to return a system to the starting
temperature at the completion of the reaction.
3) Reactions may need to be balanced using fractions and not whole numbers.
4) Enthalpy ( H) is an extensive property.
When a reaction is doubled, the enthalpy doubles.
When a reaction is reversed the enthalpy changes its sign.
5) Change in enthalpy can be experimentally calculated.
6)
H°; the ° indicates standard conditions which is 1 atm and 25°C.
Thermochemistry Ch 6
Page |8
7) Scientists arbitrarily agreed to set a zero Hf° defined as the enthalpy of
formation of elements in their most common form under standard state
conditions. All other Hf° values are calculated from these.
8) Enthalpy of formation Hf°) is the formation of just 1 mole of a compound
from its elements in their most common form under standard conditions.
9)
Hf° values are found in the Appendix.
Example 9:
Given: 4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g);
H = -906 kJ/mol
a) Is the reaction given endothermic or exothermic?
b) How many grams of NH3 gas are required to produce 374 kJ of energy?
c) What is the enthalpy term for… 2 NO (g) + 3 H2O (g)2 NH3 (g) + 5/2 O2 (g)
Calorimetry:
Calorimetry reactions use the First Law of Thermodynamics to solve for energy
transfers.
Heat capacity (C) is the Joules needed to raise an object by 1°C or 1K
Specific heat capacity (C or s) is J/g°C
Molar heat capacity is J/mol °C
Qgain = - Qlost
Qcalorimeter = - Qreaction
Coffee- cup calorimetry: pressure is constant; Qp = mc T for calorimeter
Bomb calorimetry: volume is constant; Qv = C T for calorimeter
It is important to remember that the system you are studying is going to have equal
energy, but opposite sign when compared to what happens in the calorimeter.
WATCH THE SIGNS!
Thermochemistry Ch 6
Page |9
Example 10:
When 9.55 g sample of solid sodium hydroxide dissolves completely in 100.0 g of
water in a coffee-cup calorimeter, the temperature rises from 23.6˚C to 47.4˚C.
Calculate the Qcalorimeter and the molar enthalpy
in kJ/mol NaOH for the solution process.
Assume the specific heat of the solution is the
same as for pure water, 4.184J/mol K.
Example 11:
A cold pack used in athletic departments for injuries consists of an inner bag
containing solid ammonium nitrate inside another bag containing water. When the
cold pack is needed, the inner bag is broken, and the contents are mixed. If the bag
contains 250 g of water, how many grams of ammonium nitrate must the inner bag
contain to lower the temperature of the water from 25.0°C to 5.1°C?
The heat of solution, H°soln, of NH4NO3 is +25.69 kJ/mol.
The specific heat of the water/this solution is 4.184 J/g°C .
(Neglect the contribution of the ammonium nitrate to both the mass and specific heat of water, these
resulting errors cancel)
Thermochemistry Ch 6
P a g e | 10
Example 12:
A 0.900-g sample of glycerol trilaurate (a typical fat, C39H74O6) was burned in
excess oxygen in a bomb calorimeter. The temperature increased from 24.13°C to
26.93°C. The heat capacity of the bomb calorimeter is 12.0 kJ/°C.
a)
Balance the combustion reaction for 1 mole of the fat.
C39H74O6 (s) +
O2 (g) 
CO2 (g) +
H2O (l)
Fat
b)
In a bomb calorimeter identify what is
constant (P or V) and identify if the
experiment generally measures heat in terms
of enthalpy or internal energy?
c)
Solve for the heat of reaction in the bomb
for the combustion of the fat in units of
kJ/mol. Label your answer as either E or
H as appropriate.
d)
Calculate the amount of dietary Calories (kcal) per gram of fat.
e)
Calculate the theoretical work in kJ/mol fat under standard conditions,
25°C for the reaction written in part (a).
f)
Rewrite or solve for internal energy, enthalpy, and work for the combustion
of 1 mole of this fat. For H and work, assume standard conditions.
Thermochemistry Ch 6
P a g e | 11
Example 13:
A 1.096 g sample of the sugar arabinose, C5H10O5, is burned in excess oxygen in
a bomb calorimeter, causing the temperature to rise from 20.00°C to 21.08°C. Heat
capacity of the bomb calorimeter and its contents is 15.8 kJ/°C.
a)
Calculate E in kJ/mol of sugar
b)
Write the balanced combustion reaction for 1 mole of arabinose.
c)
Assuming the reaction occurred under standard conditions of 25°C and 1 atm
pressure, solve for work (w), and H.
Example 14:
Enthalpy (Heat) of Reaction
50.0 ml of 0.500 M NaOH and 50.0 ml of 0.500 M HC2H3O2 both initially at
20.00 C are added together in a coffee-cup calorimeter and the final temperature is
found to be 23.37 C. Assume the volumes are additive, the density of the solutions
are 1.00 g/ml, the specific heat of the solution is 4.18 J/g C and it is an isolated
system in which no matter or energy is exchanged outside of the calorimeter.
a) Write the balanced reaction.
b) Solve for the Qcalorimer and the Qrxn.
c) Solve for the molar enthalpy of neutralization in kJ/mol reaction, Hrxn
Thermochemistry Ch 6
P a g e | 12
Example 15:
Before the days of refrigerators and icemakers, drinking water in hot countries was
often cooled by storage in clay pots. Evaporation of water through the pot cooled
the remaining water. How many grams of water must evaporate to cool 946 grams
(1 quart) of water from 33.5°C to 19.2°C? For water, specific heat is 4.184 J/g°C;
Hvap is 44.0 kJ/mol.
Enthalpies of Formation:
Enthalpy (or Heat) of Formation, H°f: The enthalpy change associated with the
formation of one mole of a substance from its constituent elements in their most
common form and state. H°f, the f represents formation and ° indicates the
standard state conditions of 25°C and 1 atm pressure.
Example 16:
Fix the following to make it an enthalpy of formation equation…
2K (l) + H2 (g) + C (diamond) + O3 (g)  KHCO3 (s)
Thermochemistry Ch 6
P a g e | 13
Example 17:
Write balanced equations that describe the formation of the following compounds
from their elements in their standard states. Use the appendix to obtain the values
of their standard enthalpy of formation:
a) Ca3(PO4)2 (s)
b) N2O5 (g)
c) NH4NO3 (s)
d) C3H8 (g)
Hess’s Law:
Hess’s Law: If a reaction is carried out in a series of steps, H for the overall
reaction will equal the sum of the enthalpy changes in the individual steps.
It is possible to calculate enthalpy of a reaction in both the lab setting, and by using
Appendix values. Enthalpy is an extensive property, so the amount is important. If
a series of reactions are used to create an overall reaction, the enthalpy values can
be added to solve for the overall enthalpy of a reaction.
Hrxn = Ha + Hb + Hc
Any reaction can be broken down to formation reactions from elements. When this
is done we have…
H°rxn = n H°f (products) - n H°f (reactants)
Thermochemistry Ch 6
P a g e | 14
Example 18: Use Hess’ Law to solve for the following…
a) Calculate Hrxn for… 3 H2 (g) + O3 (g)  3 H2O (g)
Given:
3 O2 (g)  2 O3 (g)
H = +284.6 kJ/mol
˚
Hf for H2O(g)
Hf = -241.8 kJ/mol
b) Calculate Hrxn for… N2O (g) + NO2 (g)  3 NO (g)
Given:
N2 (g) + O2 (g)  2 NO (g)
H = +180.7 kJ/mol
2 NO(g)+ O2(g) 2 NO2(g)
H = -113.1 kJ/mol
2 N2O(g) 2 N2(g) + O2(g)
H = -163.2 kJ/mol
c) Many cigarette lighters contain liquid butane, C4H10 (l). Using enthalpies of
formation in the appendix, calculate the quantity of heat produced when 1.0 g of
butane is completely combusted in air.
Bond Dissociation Energies:
This is similar to Hess’ Law. Breaking bonds (reactants) take energy and are
endothermic (+), while forming bonds (products) release energy and are
exothermic(-). Hrxn = BEreactants- BEproducts. Bond dissociation energies assume
all species are in the gas phase. Averages are taken, so the enthalpy of the reaction
is only approximated.
Example 19: Use information on page 387(Table 9.3) to find Hrxn for the combustion of
methane (all reactants and products as gases). Solve again using Hf values.
Thermochemistry Ch 6
P a g e | 15
Foods and Fuels:
Foods:
Most of the energy we need for life comes from carbohydrates and fats in food.
Carbohydrates (starches) decompose in our intestines into glucose (C6H12O6, blood
sugar). Sugars and carbohydrates have an average fuel value of 17 kJ/g (4 kcal/g)
C6H12O6 (s) + 6 O2 (g)  6 CO2 (g) + 6 H2O (l)
H° = -2803kJ
Fats also react with oxygen to give energy. This energy can be stored for later use.
Fats average fuel value is 38 kJ/g (9 kcal/g).
Fuels:
Coal, petroleum and natural gas (fossil fuels) are the primary sources of energy. In
units of energy per gram, the fuel value generally increases for hydrocarbons as the
quantity of H is added.
In the United States, each person uses over 105 kWh of energy per year
C(s) + O2(g)
→ CO2(g)
H°rxn = −393.5 kJ
CH4(g) +2 O2(g)
→ CO2(g) + 2 H2O(g)
H°rxn = −802.3 kJ
C8H18(g)+12.5 O2(g) → 8 CO2(g)+9 H2O(g) H°rxn = −5074.1 kJ
• Most comes from the combustion of
fossil fuels, combustible materials that
originate from ancient life.
• Fossil fuels cannot be replenished
• At current rates of consumption, oil
and natural gas supplies will be
depleted in 50–100 years.
1 Quad = quadrillion British thermal units
= 1.06 x 1018 J.
The US consumes close to 100 quads of
energy annually.
Thermochemistry Ch 6
P a g e | 16
Example 20: (Chapter 6: # 93 on page 272). Citizens of the world burn the fossil
fuel equivalent of 7 x 1012 kg of petroleum each year.
a) If all 7 x 1012 kg was in the form of octane, calculate the kg of CO2 produced
annually. C8H18 (l) + O2 (g)  CO2 (g) + H2O (l) (unbalanced)
b) If the atmosphere currently contains 3 x 1015 kg of CO2, how long will it take
to double the atmospheric carbon dioxide, assuming no photosynthesis or
other reactions remove it?
c) Would we have a chance to double the atmospheric carbon dioxide before the
fossil fuels are depleted?
Global Warming:
• CO2 is a greenhouse gas. It allows light from the sun to reach the earth, but does
not allow the heat (infrared light) reflected off the Earth to escape into outer space.
It acts like a blanket
• CO2 levels in the atmosphere have been steadily increasing
• Current observations suggest that the average global air temperature has risen 0.6
°C in the past 100 yrs.
• Atmospheric models suggest that the warming effect could worsen if CO2 levels are
not curbed
• Some models predict that the result will be more severe storms, more floods and
droughts, shifts in agricultural zones, rising sea levels, and changes in habitats
Thermochemistry Ch 6
P a g e | 17
Renewable Energy:
• Our greatest unlimited supply of energy is the sun. New technologies are being
developed to capture the energy of sunlight: parabolic troughs, solar power towers,
and dish engines concentrate the sun’s light to generate electricity; solar energy
may also be used to decompose water into H2(g) and O2(g); the H2 can then be used
by fuel cells to generate electricity.
H2(g) + ½ O2(g) → H2O(l)
• Hydroelectric power
• Wind power
H°rxn = −285.8 kJ