thermochemistry: focus learning...
TRANSCRIPT
Thermochemistry: Focus Learning Targets
(1) Define temperature and heat and give the appropriate unit for each.
(2) Describe and graph the temperature changes for a heating or cooling curve, and label each part of the curve
with the appropriate phase(s). Determine the melting/freezing point and boiling/condensing point from a heating
or cooling curve. Describe what is happening to particles at each region of a heating/cooling curve including
where particles begin to move faster/slower and where bonds between particles are broken/formed.
(3) Define heat of vaporization and heat of fusion and give the appropriate unit for each.
(4) Perform calculations involving heat of vaporization and heat of fusion for phase changes of substances.
(5) Define specific heat capacity and give the appropriate units
(6) Perform calculations involving specific heat capacity for heating and cooling of substances.
(7) Define heat of reaction/enthalpy and give the appropriate units.
(8) Use data from a calorimetry experiment to determine the heat of a reaction in kJ/mol.
(9) Define exothermic and endothermic reactions and give examples of each.
(10) Determine if a reaction is exothermic or endothermic from the chemical equation or ∆H value. Given the
∆H value, add the heat term to the appropriate side of a chemical reaction.
(11) Draw Lewis structures for molecules to determine the number and type of bonds being broken and formed
in a chemical reaction. Calculate the heat of a reaction from tabulated bond energies.
(12) Calculate the heat of a reaction from tabulated heats of formation.
(13) Calculate the heat of a reaction using Hess’ Law.
Heating Curve for Water WS#1
Label the phases present (s, l , g) and the phase changes (s→l and l→g) on the heating curve for water. Label and
state the melting point and boiling point for water.
In the first slanted region of the curve, the temperature of the ice is _______________ and the molecules are
moving _______________. Once the temperature reaches ____ °C, the ice begins to _______________. When the ice
melts, heat energy is used to _______________ some of the hydrogen bonds between molecules. On the curve, the region
where the ice is melting is _______________ because the temperature _______________ increase. Once enough
hydrogen bonds have been broken, the water is liquid. In the second slanted region of the curve, the temperature of the
liquid water is _______________ and the molecules are moving _______________. Once the temperature reaches _____
°C, the water begins to _______________. When the water boils, heat energy is used to _______________ hydrogen
bonds between molecules. On the curve, the region where the water is boiling is _______________ because the
temperature _______________ increase. Once all the hydrogen bonds have been broken, the water is a gas. In the third
slanted region of the curve, the temperature of the steam is _______________ and the molecules are moving
_______________.
Cooling Curve for Water Label the phases present (s, l , g) and the phase changes (g→l and l →s) on the cooling curve for water Label and state
the freezing point and condensing point for water.
In the first slanted region of the curve, the temperature of the steam is _______________ and the molecules are moving
_______________. Once the temperature reaches ____ °C, the steam begins to _______________. When the steam
condenses, heat energy is released to _______________ some hydrogen bonds between molecules. On the curve, the
region where the steam is condensing is _______________ because the temperature _______________ decrease. Once
enough hydrogen bonds have been broken, the steam is liquid. In the second slanted region of the curve, the temperature
of the liquid water is _______________ and the molecules are moving _______________. Once the temperature reaches
_____ °C, the water begins to _______________. When the water freezes, heat energy is used to _______________
hydrogen bonds between molecules. On the curve, the region where the water is freezing is _______________ because
the temperature _______________ increase. Once all the hydrogen bonds have been formed, the water is solid ice. In the
third slanted region of the curve, the temperature of the ice is _______________ and the molecules are moving
_______________.
Heating, Cooling, and Phase Changes WS#2
Write the equation used to solve each problem and answer with the appropriate units. Show all of your work for each calculation.
Classify each change as endothermic or exothermic.
REFERENCE INFORMATION ΔH=mxCpx ΔT ΔH=mxHfus ΔH=mxHvap
(1) How much heat is released when 20.0 g of liquid water freezes? (Endothermic or exothermic)
_______________
(2) How much heat is required to boil 75.00 g of water? (Endothermic or exothermic)
_______________
(3) How much heat is required to increase the temperature of 200.0 g of aluminum by 10.0 ºC? (Endothermic or exothermic)
_______________
(4) How much heat is required to increase the temperature of 30.0 g of copper from 20.0 ºC to 60.0 ºC?
(Endothermic or exothermic)
_______________
(5) What mass of ethanol (C2H6O) can be boiled with 2.93x104 J of heat energy? How many moles is this?
(Endothermic or exothermic)
__________________
___________________
(6) What mass of iron can be heated from 100.0 ºC to 150.0 ºC with 3375 J of heat energy? How many moles is this? (Endothermic or
exothermic)
_______________
_______________
(7) What mass of gold can be can be heated from 500.0 ºC to 800.0 ºC with 7.80 kJ of heat energy? How many moles is
this?(Endothermic or exothermic)
_______________
_______________
(8) How much heat is released when 2.00 kg of aluminum solidifies? How many kJ is this?
(Endothermic or exothermic)
_______________
_______________
(9) What mass of acetic acid (HC2H3O2) can be condensed with 7.9x103 J of heat energy? How many moles is this? (Endothermic or
exothermic)
_______________
_______________
(10) How much heat is required to boil 260 g of liquid mercury? How many kJ is this? (Endothermic or exothermic)
_______________
(11) The temperature of a 20.0 g sample of iron is raised by adding 135 J of heat. If the initial temperature of the iron is 21.5 ºC, what
is the final temperature? (Endothermic or exothermic)
______________
(12) A 32.0 g sample of copper is cooled by releasing 156 J of heat. What is the temperature change of copper? If the initial
temperature of the copper is 45.2 ºC, what is the final temperature? (Endothermic or exothermic)
_______________ _______________
(13) If freezing 500 g of ammonia releases 1.66x105 J of heat energy, determine the heat of fusion for ammonia.
(Endothermic or exothermic)
_______________
(14) If boiling 20.0 g of methanol requires 21,600 J of heat energy, determine the heat of vaporization for methanol. (Endothermic or
exothermic)
_______________
(15) A 15.0 g sample of manganese metal is warmed by adding 118.8 J of heat. If the temperature increases from 23.2 ºC to 39.7 ºC,
what is the specific heat capacity of manganese? (Endothermic or exothermic)
______________
(16) A 25.0 g sample of tin metal is warmed by adding 201.5 J of heat. If the temperature increases from 24.8 ºC to 55.8 ºC, what is
the specific heat capacity of tin? (Endothermic or exothermic)
_______________
17)If you want to heat a metal plate to as high a temperature as possible when adding 100 J of heat energy, would you use a copper
plate(cp=0.385 J/g°C) , an iron plate (cp =0.449 J/g°C), an aluminum plate(cp =0.903J/g°C), or any of these plates because they would
all heat to the same temperature? Explain. Assume all plates have the same initial temperature and mass.
Answers:
(1) 6.68x103 J, exothermic
(2) 1.692x105 J, endothermic
(3) 1.84x103 J, endothermic
(4) 468 J, endothermic
(5) 50.0 g, 1.09 mol endothermic
(6) 150 g, 2.69 mol, endothermic
(7) 200 g, 1.02 mol, endothermic
(8) 7.94x105 J, 794 kJ, exothermic
(9) 20 g, 0.33 mol, exothermic
(10) 7.67x104 J, 76.7 kJ, endothermic
(11) 36.5 ºC, endothermic
(12) 32.7 ºC, exothermic
(13) 332 J/g, exothermic
(14) 1.08x103 J/g, endothermic
(15) 0.480 J/g ºC, endothermic
(16) 0.260 J/g ºC,
(17) Copper, It has the lowest Cp and will reach the highest
temperature. The other metals require more energy to raise
the temperature.
Specific Heat and Change in Energy Calculations worksheet #3
Hf(J/g) HV(J/g) Cp(J/g°C) Substance
205 4,726 0.387 copper
109 879 2.45 Ethyl alcohol
64.5 1,578 0.129 gold
24.7 858 0.128 lead
88 2,300 0.233 silver
334 2,260 2.06 Water(g)
334 2,260 4.18 Water(l)
334 2,260 2.02 Ice(s)
1. What is the specific heat of aluminum if the temperature of a 28.4 g sample of aluminum is increased by 8.1 oC
when 207 J of heat is added? 0.90J/g oC
2. How much heat must be added to a 8.21 g sample of gold to increase its temperature by 6.2 oC? The specific heat
of gold is 0.13 J/goC. 6.6 J
3. A 218 g sample of steam at 121oC is cooled to ice at –14oC. Find the change in heat content of the system.
–6.71 x 105 J
4. A 445 g sample of ice at –58oC is heated until its temperature reaches –29oC. Find the change in heat content of
the system. 2.68 x 104 J
5. A 152 g sample of ice at –37oC is heated until it turns into liquid water at 0oC. Find the change in heat content of
the system. 6.23 x 104 J
7. If 161 g of water at 85oC is cooled to ice at 0oC, find the change in heat content of the system. (–1.11 x 105 J)
8. Find the energy change of the system required to change 150 g of ice at –15 oC to ice at –63 oC
( –2.75 x 106 J)
9. A 2.50 g sample of zinc is heated, then placed in a calorimeter containing 65.0 g of water. Temperature of water
increases from 20.00 oC to 22.50 oC. The specific heat of zinc is 0.390 J/goC. What was the initial temperature of the zinc
metal sample? (final temperatures of zinc and water are the same) (720. oC)
10. A 28.4 g sample of aluminum is heated to 39.4 oC, then is placed in a calorimeter containing 50.0 g of water.
Temperature of water increases from 21.00 oC to 23.00 oC. What is the specific heat of aluminum? (0.898J/g oC)
11. If 20 g of silver at 350oC are mixed with 200 g of water at 30oC, find the final temperature of the system. ( 31.8oC )
12. 240 g of water (initially at 20oC) are mixed with an unknown mass of iron (initially at 500oC). When thermal
equilibrium is reached, the system has a temperature of 42oC. Find the mass of the iron. ( 107.3 g)
Standard Enthalpy of formation WS#4
Every substance is shown in its standard state. This is the physical state (solid, liquid, or gas) that a substance would be in under
standard conditions. Thus, the standard state for carbon is solid, for water is liquid and for hydrogen is gas because at 1.00 atm. and 25
°C, these substances are in the physical state specified.
Standard - this means a very specific temperature and pressure: one atmosphere and 25 °C (or 298 K). If a solution is being
discussed, then everything in solution will be at a 1.00-molar concentration. Formation - this word means a substance, written as the
product of a chemical equation, is formed DIRECTLY from the elements involved.
The heat of formation of an element is arbitrarily assigned a value of zero.
Use a standard enthalpies of formation table to determine the change in enthalpy for each of these reactions.
1. NaOH(s) + HCl(g) ----> NaCl(s) + H2O(g) -133.8 kJ
2. 2 CO(g) + O2(g) ---> 2 CO2(g) -566 kJ
3. CH4(g) + 2 O2(g) ---> CO2(g) + 2 H2O(l) -891.1 kJ
4. 2 H2S(g) + 3 O2(g) ---> 2 H2O(l) + 2 SO2(g) -1123.6 kJ
5. 2 NO(g) + O2(g) ---> 2 NO2(g) -113 kJ
Hess’s Law WS#5
Hess’s Law says that if a series of reactions are added together, the enthalpy change for the net reaction will be the
sum of the enthalpy changes for the individual steps. This law enables you to calculate enthalpy changes for an
enormous number of chemical reactions by imagining that each reaction occurs through a series of steps for which the
enthalpy changes are known.
Here are some rules for using Hess’s law in solving problems:
Make sure to rearrange the given equations so that reactants and products are on the appropriate sides of the
arrows.
If you reverse equations, you must also reverse the sign of ∆H.
If you multiply equations to obtain a correct coefficient, you must also multiply the ∆H by this coefficient.
Any terms that are common to both sides of the combined equation should be canceled, just like an algebra
problem.
Example Problem:
Here are the two reactions we need:
C (s, graphite) + O2(g) ---> CO2(g) ΔH° = -393.5 kJ
C (s, diamond) + O2(g) ---> CO2(g) ΔH° = -395.4 kJ
Determine the enthalpy change for the following reaction
C (s, graphite) ---> C (s, diamond)
ΔH° = ? kJ ΔH = +2 kJ
Given the following chemical equations, calculate the energy change for the reaction that produces SO3.
(A) S + O2 SO2 ∆H = - 297 kJ
(B) 2SO3 2SO2 + O2 ∆H = 198 kJ
Determine the enthalpy (∆H) for the following reaction:
2S + 3O2 2SO3 ∆H = -792 kJ
Practice Problems:
1. From the following heats of reaction:
2 SO2(g) + O2(g) 2 SO3(g) ∆H = – 196 kJ
2 S(s) + 3 O2(g) 2 SO3(g) ∆H = – 790 kJ
Calculate the heat of reaction for:
S(s) + O2(g) SO2(g) ∆H = ? kJ
2. From the following heats of reaction:
2 H2 (g) + O2 (g) 2 H2O (g) ∆H = - 483.6 kJ
3 O2 (g) 2 O3 (g) ∆H = +284.6 kJ
Calculate the heat of the reaction for:
3 H2 (g) + O3 (g) 3 H2O (g) ∆H = -867.7 kJ
3. Calculate the heat of reaction for the following reaction:
2NO2(g) N2O4(g) ∆H = -58.0 kJ
The following is known. N2(g) + 2O2(g) 2NO2(g) ∆H = 67.7 kJ N2(g) + 2O2(g) N2O4(g) ∆H = 9.7 kJ
4. Calculate the heat of reaction for the following reaction:
2P(s) + 5Cl2(g) 2PCl5(s) ΔH = -750. kJ
The following is known.
PCl5(s) PCl3(g) + Cl2(g) ΔH = 87.9 kJ
2P(s) + 3Cl2(g) 2PCl3 ΔH = -574 kJ
.
5. Find the ΔH for the reaction below:
2CO2(g) + H2O(g) → C 2H2(g) + 5/2O2(g) ΔH =235 kJ
Use the following information:
C2H2(g) + 2H2(g) → C2H6(g) ΔH =-94.5 kJ
H 2O(g) → H2(g) + 1/2O2 (g) ΔH =71.2 kJ
C2H6(g) + 7/2O2(g) → 2CO2(g) + 3H2O(g) ΔH =-283 kJ
6. What is the standard enthalpy change for the reaction below? 2N2(s) + 5O2(g) 2N2O5(g) The following is known: H2(g) + ½O2(g) H2O(l) ∆H = -285.8 kJ
N2O5(g) + H2O(l) 2HNO3(l) ∆H = -76.6 kJ
½N2(g) + 3/2 O2(g) + ½H2(g) HNO3(l) ∆H = -174.1 kJ ΔH = 28.4 kJ
Bond Energy Dissociation WS#6
Bond energy is defined as the amount of energy required to break a bond. These values are positive, indicating
that bond breaking is endothermic. Bond energies are reported in kilojoules per mole (kJ/mol). For example the
energy for breaking a hydrogen-hydrogen bond is 436 kJ/mol. In a chemical reaction several bonds are broken
and formed.
Energy is always required to break a bond and energy is released when a bond is made.
Carbon-Carbon Triple Bond: 839 kJ/mole
H-I Bond: 297 kJ/mole
I-I Bond: 151 kJ/mole
ΔH = Σ Ereactant bonds broken ---- Σ Eproduct bonds broken
Example: The heat of reaction can be calculated from a reaction, based on bond energies.
Let’s examine the electrolysis/Decomposition of water.
The overall heat of reaction can be calculated as follows:
Write the Equation H2O → H2 + O2
Balance the Equation 2 H2O → 2 H2 + O2
Draw the Structures
Calculate bond energies
Total Required
Thus, this is the reaction absorbs 466 kJ/mole more energy than released and the thermochemical
equation is: 466 kJ/mole + 2 H2O → 2 H2 + O2
The heat of reaction is +466kJ/mole
1. Which process releases energy: breaking a bond or forming a bond?
2. Which process requires energy: breaking a bond or forming a bond?
3. If the energy used to break bonds is greater than the energy released in the formation of new
bonds, is the reaction endothermic or exothermic?
4. If the energy used to break bonds is less than the energy released in the formation of new
bonds, is the reaction endothermic or exothermic?
5. Draw the electron dor structure of CH4. How much energy does it take to break apart the molecule CH4
into it individual atoms?
6. Draw the electron dot structure of CO2. How much energy is released when the molecule CO2 is formed
7. Determine the approximate ΔH Hrxn using the bond energies above, for each of the following reactions
a. H2 + O2 --> 2 H2O
b. CH4 + Cl2 --> CH3Cl + HCl
c. CH4 +2 O2 --> CO2 + 2 H2O
d. H2 + I2 2HI
e. 2H2 + O2 2H2O
f. CH4 + 2O2 CO2 + 2H2O
g. C3H8 + O2
Stoichiometry using Enthalpy WS#7
1. ___CH4 (g) + ___O2 (g) ___H2O (g) + ___CO2 (g) + 212.8 kcal
Balance the above equation showing the combustion of methane (CH4) and answer these questions:
a. Is this reaction exothermic or endothermic?
b. How many moles of methane (CH4) are needed to make 5.98 x 105 kcal of heat?
c. If 445.2 kcal are released in this reaction, how many moles of oxygen were used? grams of oxygen?
d. If a reaction produced 3.88 x 104 g of CO2, how much heat did it produce?
Every endothermic reaction is an exothermic reaction in reverse and every exothermic reaction is an endothermic
reaction in reverse! 2. ___H2O (g) + ___CO2 (g) + 212.8 kcal ___CH4 (g) + ___O2 (g)
a. How much heat is needed to produce 10.5 moles of methane gas?
b. How much heat is needed to produce 4.5 x 103 grams of methane gas?
3. ___Fe (s) + ___O2 (g) ___Fe2O3 (s) + 1652 kJ
a. Is the reaction an endothermic or exothermic reaction?
c. Rewrite the reaction as the opposite reaction for what you answered in (a.) and put the heat in the appropriate
place in the reaction.
c. How much heat is released when 4.50 moles of iron is reacted with excess oxygen?
d. How much heat is released when 8.00 moles of iron (III) oxide is formed?
e. How much heat is needed to produce 5.00 moles of solid iron?
f. How much heat is released when 10.0 g of Fe and 2.00 g of O2 are reached?
Writing Thermochemical equation and calculating heat of the reaction.
• A Thermochemical Equation is a balanced stoichiometric chemical equation that includes the enthalpy
change, ΔH. In a thermochemical equation it is important to note phase labels because the enthalpy
change, ΔH, depends on the phase of the substances.
Write the thermochemical equation for each one of the reactions. Include ΔH in the reaction.
1. How much heat will be released when 6.44 g of Sulfur reacts with excess O2 according to the following equation?
2S + 3O2 → 2SO3 ΔH° = -791.4kJ -79.4kJ
2. How much heat will be released when 4.72 g of Carbon reacts with excess O2 according to the following equation?
C + O2 → CO2 ΔH° = -393.5kJ -155kJ
3. How much heat will be absorbed when 38.2 g of Bromine reacts with excess H2 according to the following equation?
H2 + Br2 → 2HBr ΔH° = 72.80kJ 17.4kJ
4. How much heat will be released when 1.48 g of Chlorine reacts with excess phosphorus according to the following equation?
2P + 5Cl2 → 2PCl5 ΔH° = -886kJ -3.75kJ
5. How much heat will be released when 4.77 g of ethanol (C2H5OH) reacts with excess O2 according to the following
equation?
C2H5OH + 3O2 → 2CO2 + 3H2O ΔH° = -1366.7kJ -142kJ
6. One way to cool down your cup of coffee is to plunge an ice-cold piece of aluminum into it. Suppose you store a 20.0 g piece
of aluminum (Cp of Al = .902J/g•°C) in the refrigerator at 4.40 °C and then drop it into your coffee. The coffee temperature
drops from 90.0 °C to 55.0 °C. How many joules of heat energy did the aluminum block absorb?
912J
Thermochemistry Review WS#8
1. Which of these changes are endothermic? Which are exothermic?
a) Vaporization of ice
b) CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) + 882 kJ
c) Sublimation of iodine
d) NH4Cl (s) NH4Cl (aq) ΔH° = + 14.7 kJ/mol
e) Citric acid reacting with baking soda causes a temperature decrease
2. Use the standard enthalpies of formation table to determine the ΔHrxn for each of these reactions. (You must know the enthalpy
of formation for all elements = 0.)
a) NaOH(s) + HCl(g) ----> NaCl(s) + H2O(g) - 133.8 kJ
b) 2 H2S(g) + 3 O2(g) ---> 2 H2O(l) + 2 SO2(g) - 1123.6 kJ
Determine which method applies and calculate the enthalpy for each of the problems below.
2. For the reaction
2 C4H10(g) + 13 O2(g) 8 CO2 (g) + 10 H2O(l) the ΔH°comb is – 2877.4 kJ.
a. What is the ΔHcomb° for the combustion of 1 mole of C4H10? - 1438.7 kJ
b. What is the ΔHcomb° for the production of 5 moles of H2O? - 1438.7 kJ
c. What is the ΔHrxn° for the production of 1 mole of C4H10? + 1438.7 kJ
3. If 540g of water condenses on a car during a cool night, calculate the amount of energy released to the air during this
condensation. Given: The molar heat of vaporization of water is 40.79kJ/mol
4. Calculate the mass of a sample of lead (CPb = 0.160 J/g.℃) when it loses 200. J cooling from 75.0℃ to 42.0℃. 37.9 g
5. Calculate the energy required to take 450.0 g of water from 27.5°C to 102.0°C. (CH2O = 4.184 J/g°C, Csteam = 2.006 J/g,
6. Hvap = 2260 J/g) 1.16 x 106 J
7. A 28.4 g sample of aluminum is heated to 39.4 oC, and placed in a calorimeter containing 50.0 g of water. The temperature of
water increases from 21.00 oC to 23.00 oC. What is the specific heat capacity, C, of aluminum? (for H2O(l), C = 4.184 J/g°C)
0.898 J/g.℃
8. A 248-g piece of copper initially at 314 °C is dropped into 390 mL of water initially at 22.6 °C. Assuming that all heat
transfer occurs between the copper and the water, calculate the final temperature. 38.8 °C
9. Calculate the heat of formation of ethane, C2H6 (g), if its heat of combustion, ΔH°comb, is - 3120. kJ as described by the equation
is
2 C2H6 (g) + 7 O2 (g) 6 H2O (l) + 4 CO2 (g). 84.4 kJ/mol
10. Calculate the enthalpy of the following chemical reaction:
CS2(ℓ) + 3O2(g) ---> CO2(g) + 2SO2(g) -973.7 kJ
Given:
C(s) + O2(g) ---> CO2(g) ΔH = -339.8 kJ/mol
S(s) + O2(g) ---> SO2(g) ΔH = -273.0 kJ/mol
C(s) + 2S(s) ---> CS2(ℓ) ΔH = +87.9 kJ/mol
11. Using bond enthalpies, calculate the reaction enthalpy (ΔH) for:
CH4 (g) + Cl2 (g) CH3Cl (g) + HCl (g) -101 kJ
12.
Consider the heat capacities of two cast iron frying pans. The heat capacity of the small cast iron frying pan is found by observing that
it takes 18,150 J of energy to raise the temperature of the pan by 50.0 °C
The heat capacity of the larger cast iron frying pan, while made of the same substance, requires 90,700 J of energy to raise its
temperature by 50.0 °C
a. Why is the heat capacity of the large pan five times greater than that of the small pan?
b. Find the specific heat capacity of the smaller cast iron frying pan that has a mass of 808 g.( Note the heat capacity is
363J/°C with a 50.0 °C temperature change
c. Find the specific heat capacity of the larger cast iron frying pan that has a mass of 4040 g.( Note the heat capacity is
1814J/°C with a 50.0 °C temperature change
d. Is specific heat capacity intensive or extensive?
Problem of the unit:
Thermochemistry of Ethanol Honors Chemistry
Properties of Ethanol:
Formula: C2H5OH
Structure:
Molar mass: 46.068 g/mol
State Specific Heat
Capacity (J/g°C)
solid 2.42
liquid 2.50
gas 1.80
Melting point: –114 °C
Heat of Fusion: 109 J/g
Boiling Point: 78 °C
Heat of Vapourization: 586 J/g
Heat of Formation: –235.2
kJ/mol 1) Label the heating curve to show the phase(s) present in each region. Label the melting point and boiling point.
Between zero minutes and three minutes, the temperature of the solid ethanol is _______________ which causes the molecules to move
_______________. Once the temperature reaches ____ °C, the ethanol begins to _______________. When the ethanol melts, heat
energy is used to _______________ some of the hydrogen bonds between molecules. The ethanol is melting between ____ minutes
and ____ minutes. Once enough hydrogen bonds have been broken, the ethanol is a liquid. Between seven minutes and nine minutes,
the temperature of the liquid ethanol is _______________ which causes the molecules to move _______________. Once the
temperature reaches _____ °C, the ethanol begins to _______________. When the ethanol boils, heat energy is used to
_______________ hydrogen bonds between molecules. The ethanol is boiling between ____ minutes and ____ minutes. Once all the
hydrogen bonds have been broken, the ethanol is a gas. Between fourteen minutes and sixteen minutes, the temperature of the gaseous
ethanol is _______________ which causes the molecules to move _______________.
(2) Determine the total amount of energy (in kJ) required to change 50 g of solid ethanol at –180 °C to ethanol vapour at 150 °C. Use
the data given above for the five thermochemical calculations needed to complete this problem.
3) Balance the equation for the combustion of ethanol:
____ C2H5OH (l) + ____ O2 (g) → ____ CO2 (g) + ____ H2O (g)
(4) Calculate the heat of reaction from the heats of formation of each chemical.
(5) (a) Complete the following using the calculated value of the heat of reaction:
∆H =
2 5 2 2 2
_________
__ , __ , __ , __
kJ
mol C H OH mol O mol CO mol H O
(b) Calculate the amount of heat released for reactions involving each of the following quantities:
(i) 0.200 mol C2H5OH
(ii) 92.0 g C2H5OH
(iii) 1.20 mol O2
(iv) 8.80 g CO2
(6) Calculate the heat of the reaction from bond energies. Is this value similar to the ∆H value calculated in #4?
(7) Calculate the heat of reaction using Hess’ Law.
Steps: C (s) + O2 (g) → CO2 (g) ∆H = –393.5 kJ/mol
H2 (g) + ½ O2 (g) → H2O (g) ∆H = –242.0 kJ/mol
2C (s) + 3H2 (g) + ½ O2 (g) → C2H5OH (l) ∆H = –235.2 kJ/mol
How does this value compare to the value calculated in #4?