Download g) Chemistry 30 - Mr. Jones LHS Science

CP Chemistry Unit 10 Thermodynamics-Test Plan
Essential Vocabulary:
 Work
 Heat
 System
 Surroundings
 Internal energy
 Specific heat
 Calorimeter
 Joule
 Celsius
 Enthalpy of reaction (ΔHrxn)
 Heat of formation (ΔHf)
 Hess’ Law
Essential Outcomes:
 Know the sign convention for heat and work entering or exiting the system.
 Be able to calculate the internal energy of a system (ΔE = q + w)
 Be able to relate the change in temperature of a substance to the heat
absorbed or released (q = mCΔT)
 Be able to calculate the specific heat capacity of a metal using calorimetry.
 Be able to measure the ΔHrxn using calorimetry
 Be able to convert joules to kilojoules
 Be able to calculate the enthalpy of reaction for different gram amounts of a
reactant or product.
 Be able to predict the change in enthalpy when a reaction is reversed or the
stoichiometry is changed.
 Be able to apply Hess’ Law.
 Be able draw a reaction that represents the heat of formation of a compound.
 Be able to use heats of formation to determine the ΔHrxn of a reaction.
(∑ΔHproducts- ∑ΔHreactants)
 Be able to identify the thermodynamics of a reaction from a reaction
coordinate (energy diagram)
 For common qualitative experiences (such as the melting of ice), be able to
identify the situation as exothermic or endothermic and be explain why.
Energy Unit Conversions
Practice Problems
Answer in the space provided. Show your work for full credit!
Remember:
1 cal = 4.184 J
Perform the following conversions:
1. 340. cal to joules
2. 75.6 joules to calories
3. 7200. kcal to cal
4. 12.9 kJ to J
5. 12755 J to kJ
6. Europeans food labels are often listed in kJ. How many calories were in this snack?
Heat, Work, and Internal Energy Practice
Word Bank
 Work Heat  System  Surroundings  Internal energy  Specific heat  Calorimeter  Joule  Enthalpy of reaction (ΔHrxn)
 Heat of formation (ΔHfus)
 Heat of dissolution (ΔHd)
 Heat of vaportization (ΔHv)
 Heat of formation (ΔHf)
 Heat of combustion(ΔHcomb)
 Hess’ Law
 Energy of activation
Endothermic and Exothermic Reactions
Answer the following questions in the space provided. Show your work on calculations!
1.
Identify the following as either endothermic or exothermic:
a. ice melting
b. a firecracker going off
c. natural gas is burned in a furnace
d. when solid KBr is dissolved in water, the solution gets colder.
e. frost forming on the windows of a car overnight
f. When concentrated sulfuric acid is added to water, the solution gets very hot
2.
The equation for the fermentation of glucose to alcohol and carbon dioxide is
C6H12O6  2C2H5OH + 2CO2
The enthalpy change for the reaction is -67 kJ.
a. Is this reaction endothermic or exothermic?
b How would the reaction feel to the touch?
b. Write the equation again below and include the energy term in the equation.
3.
For the reaction
S + O2  SO2
∆H = -296 kJ/mol
a. Is this reaction endo- or exothermic?
b. Rewrite the chemical equation with the energy term in the equation.
c. Draw a reaction coordinate diagram showing if the relative energy of the reactants and
products. . Label the ∆H on the diagram.
4.
Given the following reaction:
Fe3O4 + CO + 18 kJ  3FeO + CO2
a. Is this reaction endo- or exothermic?
b. How would the reaction feel to the touch?
c. What is the sign of the energy term when written in this form: ∆H =
d. Draw a reaction coordinate diagram showing the relative energy of the reactants and
products. Label the ∆H on the diagram
5.
For the following reaction
NO + O3  NO2 + O2 + 199 kJ
a. Is this reaction endo- or exothermic?
b. Write the energy term:
∆H =
c. How many moles of NO are reacted in the balanced equation?
d. Would you expect to produce more or less heat when 12 moles of NO are reacted?
e. How much heat would result from the 12 moles?
Enthalpy of Reaction (H) Practice
For each problem below, write the equation and show your work. Always use units.
1. The decomposition of limestone, CaCO3(s), into lime, CaO(s), and CO2(g) at constant pressure
requires the addition of 178.1 kJ of heat per mole of CaCO3.
a. Write a balanced thermochemical equation for the reaction showing energy as a reactant or
product.
b. Draw an enthalpy diagram for the reaction.
2. Consider the following reaction:
CH3OH (g) --> CO (g) + 2 H2 (g)
H = +90.7 kJ
a. Is the reaction exothermic or endothermic?
b. Calculate the amount of heat transferred when 45.0 g of CH3OH (g) are decomposed by this
reaction at constant pressure.
c.
If the enthalpy change is 16.5 kJ, how many grams of hydrogen gas are produced?
d. How many kilojoules of heat are released when 10.0 g of CO (g) reacts completely with H 2
(g) to form CH3OH (g) at constant pressure?
3. At one time, a common means of forming small quantities of oxygen gas in the laboratory was to
heat KClO3:
2KClO3 (s) --> 2 KCl (s) + 3 O2 (g)
H = -89.4 kJ
For this reaction, calculate H for the formation of
a. 1.24 mol of O2
b. 4.89 g of KCl
c.
1.76 g of KClO3 from KCl and O2
4. Consider the decomposition of liquid benzene, C6H6 (l), to gaseous acetylene, C2H2 (g):
1/3 C6H6 (l) --> C2H2 (g)
H = +210 kJ
a. What is the enthalpy change for the reverse reaction?
b. What is H for the decomposition of 1 mol of benzene to acetylene?
c.
Which is more likely to be thermochemically favored, the forward reaction or the reverse
reaction?
d. If C6H6 (g), were consumed instead of C6H6 (l), would you expect the magnitude of H to
increase, decrease, or stay the same? Explain.
Hess’ Law Practice
1. Calculate the heat of reaction for:
PbCl2 (s) + Cl2 (g)  PbCl4 (l) ∆ H = ?
by combining:
Pb (s) + Cl2 (g)  PbCl2 (s) ∆ H = - 359.4 kJ
Pb (s) + 2 Cl2 (g)  PbCl4 (l) ∆ H = - 329.3 kJ
2. From the following heats of reaction:
2 SO2(g) + O2(g)  2 SO3(g)
2 S(s) + 3 O2(g)  2 SO3(g)
Calculate the heat of reaction for:
S(s) + O2(g)  SO2(g)
H = – 196 kJ
H = – 790 kJ
H = ? kJ
3. The standard heats of formation of HCl (g) and HBr (g) are -92.0 kJ/mol and -36.4 kJ/mol
respectively. Using this information, calculate ΔH for the following reaction:
Cl2 (g) + 2 HBr (g) → 2 HCl (g) + Br2 (g)
4. Given the following equations:
4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(l) ΔH° = -1170 kJ
4 NH3(g) + 3 O2(g) → 2 N2(g) + 6 H2O(l) ΔH° = -1530 kJ
Using these two equations, determine the heat of formation, ΔHf° , for nitrogen monoxide,
NO.
5. From the following heats of reaction:
2 H2 (g) + O2 (g)  2 H2O (g)
3 O2 (g)  2 O3 (g)
Calculate the heat of the reaction for:
3 H2 (g) + O3 (g)  3 H2O (g)
H = - 483.6 kJ
H = +284.6 kJ
6. Given the following data:
N2 (g) + O2 (g)  2 NO (g)
H = + 180.7 kJ
2 NO (g) + O2 (g)  2 NO2 (g) H = - 113.1 kJ
2 N2O (g)  2 N2 (g) + O2 (g) H = - 162.3 kJ
Use Hess’s law to calculate Hrxn for the following reaction.
N2O (g) + NO2 (g)  3 NO (g)
Heat of Formation Practice
1. Calculate Ho (in kJ) for these reactions. In each case, state whether the reaction is
exothermic or endothermic, rewrite the equation as a thermochemical equation to include
the heat term.
For example:
2 NO(g) + O2(g) ---> 2 NO2(g)
Hf = [2(33.2kJ) – ((2(90.3kJ))+(0kJ)) = -114.2 kJ
The reaction is EXOTHERMIC, and the thermochemical reaction should
read:
2 NO(g) + O2(g) ---> 2 NO2(g) + 114.2 kJ
a) NaOH(s) + HCl(g) ----> NaCl(s) + H2O(g)
b) SO2(g) + ½ O2(g) ---> SO3(g)
c) CaO(s) + H2O(l) ---> Ca(OH)2(s)
d) N2(g) + 3 H2(g) ---> 2 NH3(g)
e) C6H6(l) + 1½ O2(g) ---> 6 C(s) + 3 H2O(l)
f) NH3(g) + HCl(g) ---> NH4Cl(s) Chemistry 30
2. Write heats of formation reactions for each of the following compounds in two ways, one
as a reaction with the appropriate Hf value, and secondly, as a thermochemical reaction.
You will need to refer to a Table of Thermochemical Data.
CuCl2 (g), N2H4 (l)
3. The standard heat of formation, ΔHof, for sulfur dioxide (SO2) is -297 kJ/mol. How many kJ
of energy are given off when 25.0 g of SO2 (g) is produced from its elements?
4. The heat of reaction for the combustion of 1 mol of ethyl alcohol is -9.50 × 102 kJ:
C2H5OH (l) + 3 O2 (g) → 2 CO2 (g) + 3 H2O (l) + 9.5 × 102 kJ
How much heat is produced when 11.5 g of alcohol is burned?
5. ΔH for the complete combustion of 1 mol of propane is -2.22 × 103 kJ:
C3H8 (g) + 5 O2 (g) → 3 CO2 (g) + 4 H2O (l)
Calculate the heat of reaction for the combustion of 33.0 g of propane.
Calorimetry Practice
q=m c ∆T
1. A 25.0 g sample of an alloy was heated to 100.0 oC and dropped into a beaker
containing 90 grams of water at 25.32 oC. The temperature of the water rose to a final
value of 27.18 oC. Neglecting heat losses to the room and the heat capacity of the
beaker itself, what is the specific heat of the alloy?
2. The specific of solid iron is 0.477 J/goC. If 230 g of Fe at 95oC were placed in 1 L of water
at 15oC, what would Tf be?
3. 101 g of an unknown metal at 50oC is placed in 40 g of water at 25.0oC. The
temperature of the system at equilibrium is 29.7 oC. What is the unknown metal? Your
choices, and their Specific Heat values (in J/goC) are:
a. Ba (0.284)
b. Cr (0.464)
c. Cu (0.385)
d. Au (0.130)
e. Sn (0.213)
f. Rb (0.360)
g. Ti (0.594)
Heat of Reaction Lab: Hess’s Law
In this experiment, you will determine and compare the enthalpy changes in three
thermodynamic reactions. Start by writing balanced chemical equations for each of the
reactions.
Reaction 1: Solid sodium hydroxide dissolves in water to form an aqueous solution of
ions.
Reaction 2: An aqueous solution of sodium hydroxide reacts with an aqueous solution
of hydrogen chloride to form water and an aqueous solution of sodium chloride.
Reaction 3: Solid ammonium nitrate dissolves in water to form an aqueous solution of
ions.
.
Procedure
Reaction 1
1. Place 10 mL of water into a test tube cup. Record the temperature of water as
T1.
2. QUICKLY measure approximately 0.20 g of NaOH and add it to the water stirring
gently until the pellets have dissolved.
3. Record the highest temperature reached as T2.
4. Flush the solution down the sink and rinse the cup and thermometer.
Reaction 2
1. Measure 5 mL of 1.0 M HCl into a test tube.
2. Obtain an equal volume of 1.0 M sodium hydroxide in another test tube.
3. Record the temperature either solution.
4. Pour the NaOH solution into the HCl. Stir and record the highest temperature
reached as T2.
5. Discard the solution and clean all materials by rinsing with water.
Reaction 3
1. Place 10 mL of water into a test tube cup. Record the temperature of water as
T1.
2. QUICKLY measure approximately 0.20 g of ammonium nitrate and add it to the
water stirring gently until the pellets have dissolved.
3. Record the highest temperature reached as T2.
4. Flush the solution down the sink and rinse the cup and thermometer.
Data Table
Reaction
Total
volume of
water used.
Mass of
water used.
Tinitial
Tfinal
∆T
(final –
initial)
(should be the
same!)
q (heat absorbed or released by
water.) What sign should you use?
What units!
1
2
3
Calculations Table
Reaction
∆Hrxn
Mass (g)
1
Mass of NaOH
2
Mass of NaOH
3
Mass of NH4NO3
Moles (of each
∆Hrxn (in
from previous
column)
kJ/mol)
Was the reaction
endo or exothermic?
1. Draw an energy coordinate diagram for reaction #1.
2. Re-write the reaction for reaction #3 but include the energy in the equation on the
correct side.
The Specific Heat of a Metal Experiment
Purpose:
To Determine the Specific Heat of a Metal
Equipment Needed:
Hot plate, two beakers, Styrofoam cup, string, 100-ml graduated cylinder, electronic balance,
thermometer, water, and metal object(s)
Procedure:
1. Put on your goggles!
2. Measure and record the mass of the metal object in the data table. Attach a piece of string
about 15 cm long to the metal object.
3. Fill a 250-ml beaker about half-full with water. Place the metal object into the beaker of
water with some string protruding over the beaker's lip.
4. Place the beaker on the hot plate and allow the water to come to a rolling boil. Allow the
metal object to remain in the boiling water for at least five minutes.
5. Record the temperature of the boiling water. This will be the initial temperature of the metal
object.
6. Use a graduated cylinder to pour 100 ml of tap water into the Styrofoam cup. Since the
density of water is 1.0 g/ml, record this mass in the data table. For support, place the cup inside a
second 250-ml beaker.
7. Measure and record the initial temperature of the tap water in the Styrofoam cup.
8. Use the protruding string to remove the metal object from the boiling water and quickly lower
it into the water in the Styrofoam cup. When the temperature reaches a constant, record it. This
will be the final temperature of the metal object and the water.
9. Repeat the above procedure for a second trial.
Data Table:
System
Surroundings
Data Table
System
Surroundings
Calculations:
a) Determine the specific heat of your metal by using:
Heat lost by the metal object = Heat gained by the water
mmetal cmetal ∆T = mwater cwater ∆T
Trial 1 Calculations
Trial 2 Calculations
b) Determine the Percentage error by using:
Percent. Error = (Experimental value - Theoretical value) X 100 %
(Theoretical value)