Study Guide or: How I Learned to Stop Worrying and Love ... · Study Guide or: How I Learned to Stop Worrying and Love Thermochem . Math: First, determine what you’re being asked

Study Guide or: How I Learned to Stop Worrying and Love Thermochem

Math: First, determine what you’re being asked to find: ΔHrxn or something else?

1. Asked to find ΔHrxn? 2 Options!

Given? Asked to find? Use:

1. A lot of heats of formation (ΔHf) 2. A reaction without ΔHrxn

ΔHrxn ΔHrxn = ΣΔHf (products)− ΣΔHf (reactants)

1. Multiple reactions with ΔH 2. A goal reaction without ΔHrxn

ΔHrxn

Hess’s Law! Rearrange the equations to make the goal equation, then combine your new ΔH’s (remember, what you do to an equation you must do to ΔH!)

2. Asked to find something else? 3 Options!

Given? Asked to find? Use:

1. A reaction with ΔHrxn 2. Either:

a. g or mol of a substance b. energy change (J or kJ)

1. Either: a. g or mol of a

substance b. energy change (J or

kJ)

Stoich! Don’t forget to convert between moles of your substance and molesrxn

1. A phase change (vaporizing, condensing, freezing or melting)

2. Mass or moles of a substance

3. ΔHvap or ΔHfus

Energy change (heat absorbed or released) q = nΔH

1. A temperature change 2. Mass or moles of a substance

3. Heat capacity ( 𝐽𝐽𝑔𝑔 ℃

)

Energy change (heat absorbed or released)

q=mCΔT

Concepts:

Exothermic (−∆H): heat exits Endothermic (+∆H): heat enters

• System releases energy

• Heat is a product

• System absorbs energy • Heat is a reactant

• Magnitude of energy absorbed when bonds break is smaller than the energy released when bonds form

• Chemical potential energy converted to kinetic energy

• Potential energy of reactants is higher than the potential energy of products

• Magnitude of energy absorbed when bonds break is larger than the energy released when bonds form

• Kinetic energy converted to chemical potential

energy

• Potential energy of reactants is lower than the potential energy of products

Unit 12 Objectives: Thermochemistry

Heat Content Objectives:

• I can understand the law of conservation of energy and the processes of heat transfer. • I can perform calculations involving heat, mass, temperature change, and specific heat.

Criteria for Success: • I can explain the law of conservation of energy. • I can define heat, q, and explain how heat flows • I can solve for heat, mass, temperature change (or initial or final temp) or specific heat if given all but one of the

variables. Thermochemical Equations and Calculations Content Objective:

• I can use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic.

Criteria for Success: • I can write thermochemical equations. • I can use enthalpy change to determine whether a reaction is endothermic or exothermic. • I can perform stoichiometry calculations involving thermochemical equations. • I can use average bond dissociation energies to calculate the enthalpy change for a chemical reaction• I can use Hess’s Law to calculate the enthalpy (heat) of reaction. • I can use standard enthalpies (heats) of formation to calculate the enthalpy (heat) of reaction.

Heating and Cooling Curves Content Objectives:

• I can understand the law of conservation of energy and the processes of heat transfer. • I can perform calculations involving heat, mass, temperature change, and specific heat.

Criteria for Success: • I can label and explain parts of a heating or cooling curve. • I can define and calculate molar heat of fusion or molar heat of vaporization. • I can solve for heat, mass, temperature change (or initial or final temp) or specific heat if given all but one of the

variables.

Unit 12 Part 1: Temperature, Heat, & mCΔT

Energy: the capacity to do_____________

It comes in 2 main forms!

1) Kinetic:

energy due to object’s _____________

2) Potential:

________________ energy

a) Thermal energy (____________):

Determines how active the atoms are

a) ____________________ potential energy

From chemical reactions (bonds)

Stored within a substance

Examples:

b) ____________________ potential energy

From object’s height and mass

b) ___________________ energy:

Movement of electrons

c) _______________ energy (or EMR):

Type of electromagnetic energy

Law of Conservation of Energy – Energy CANNOT be _________________________ or ________________________ during chemical and physical processes. But, it can be transferred! Thermochemistry (T): a study of the __________________________ change that occur during chemical reactions.

• Temperature: a measure of the average __________________ energy of a substance

o Intensive property: does not depend on the ________________ of substance you have

• Heat (q): _____________ that transfers from one object to another because of a temperature difference between them.

o Extensive property: __________ depend on the amount of substance you have

o Heat always flows from a ___________ object to a ___________ object.

• The universe is divided into two halves:

o The ________________: the substance of interest

o The ___________________________: whatever is outside the system

Energy (heat) can be absorbed or released!

Exothermic (______): heat exits Endothermic (______): heat enters

Energy is ___________________ Energy is ________________

• System ___________ (lose) energy

• Surroundings absorb energy

• System ___________ (gain) energy

• Surroundings release energy

Chemical potential energy is converted to ______________ energy

________________ energy is converted to chemical

potential energy

Potential energy of reactants is ___________ than the potential energy of products.

Potential energy of reactants is _______________ than the potential energy of products.

Heat is a ________________ in a chemical reaction:

1. H2O(l) → H2O(s) + heat

2. CH4 + 2 O2 → CO2 + H2O + heat

Heat is a ________________ in a chemical reaction:

1. H2O(s) + heat → H2O(l)

2. 6 CO2 + 6 H2O + heat → C6H12O6 + 6 O2

State Changes Gas

Liquid

Solid

• Going from a ______________ to a __________________ energy state of matter is always ________thermic.

• Going from a ______________ to a __________________ energy state of matter is always ________thermic.

Specific Heat Capacity (C)

• The amount of heat (energy) required to raise temperature of 1 g of a substance by 1 K (1 oC)

• Units are J

g ℃ or

Jg K

• Metals have relatively _____ specific heats - relatively less energy is required to raise their temperatures.

• Water has a relatively ________ specific heat - requires much more energy to achieve a similar temperature change.

It does NOT matter if you are given °C or K, as long as initial and final temperatures are in the same unit.

• When adding a FINITE (________________) amount of energy:

o Matter with a low specific heat will change temperature __________

o Matter with a high specific heat will change temperature __________

• When adding an unlimited supply of heat (e.g. sitting in the sun)

o Matter with low specific heat will change temperature more _______________

o Matter with high specific heat will change temperature more _______________

Example 1: A candle of solid wax melts while the candle is burning. Write this process as a reaction with heat as a reactant or product. Is it endo- or exothermic?

Example 2: Two aqueous substances in a glass beaker chemically react, and the temperature of the water in the beaker rises.

a. Have the surroundings gained or lost heat? How do you know?

b. Has the system gained or lost heat? How do you know?

c. Is the reaction endothermic or exothermic? _________________________

Why? Let’s try an example.

1. Calculate ΔT for a sample that started at 2oC and was heated up to 27oC:

2. Calculate ΔT for a sample that started at _______ K and was heated up to _______ K:

How to Calculate Amount of Heat Transferred: mCAT!

q = mC∆T q = heat transferred (J, joules)

m = mass of substance (g)

C = specific heat capacity (J/g°C or J/g K )

ΔT = Tfinal – Tinitial = change in temperature (°C or K)

• Energy can sometimes be measured as ________________ or ____________________.

• Use the following chart/reference table when performing conversions between different units of energy.

Example 1: The specific heat (in J/g oC) of Al(s) is 0.89, of Fe(s) is 0.45, of Hg(l) is 0.14, and of C(s) is 0.71. a) When the same amount of heat is applied to one gram of these substances, which one will reach the

highest temperature and why?

b) If each substance is heated until they are all the same temperature, which substance required the most heat energy? Explain.

Example 2: You are given a 45 cm3 sample of copper metal and a 30 cm3 sample of copper metal, both at 300 K. Which sample contains the most heat and why?

Example 1: The specific heat of aluminum is 0.89 J/g°C. How much energy is required to raise the temperature of a 15.0 gram aluminum can 18 K?

Practice Makes Perfect! 1. When barium hydroxide octahydrate, Ba(OH)2• 8H2O is mixed in a beaker with ammonium thiocyanate,

NH4SCN, a reaction occurs. The beaker becomes very cold. a. Have the surroundings gained or lost heat? How do you know?

b. Has the system gained or lost heat? How do you know?

c. Is the reaction endothermic or exothermic? _________________________

2. If 1.82 kJ of heat is required to raise the temperature of a sample of mercury 52°C , and the specific heat of mercury is 0.14 J/(g°C), what is the mass of the sample of mercury?

3. The temperature of a 95.4 g piece of copper increases from 25.0°C to 48.0°C when the copper absorbs 849 J of heat. What is the specific heat of copper?

Example 2: A 1.6 g piece of solid potassium is added to 150. mL of aqueous H2SO4, releasing 5.20 kJ of heat energy. Assume that the specific heat of the solution is 4.184 J/g°C, and the total mass of the solution with all reactants is 151.6 g. Calculate the temperature change of the solution. Example 3: If 0.596 kJ of heat are removed from 29.6 g of water at 22.9°C, what will be the final temperature of the water? The specific heat of water is 4.18 J/g K.

4. Identify each of the following phrases/pictures as describing an endothermic or exothermic process:

a. The reactants have more potential energy than the products.

b. A liquid evaporates.

c. q is negative.

d. An aqueous chemical reaction occurs, and the temperature of the water drops.

e.

f. The products have more potential energy than the reactants.

g. A chemical reaction occurs, and the container becomes hot to the touch.

h. q is positive.

i. A gas condenses.

j.

5. In an endothermic reaction, which of the following is true?

a. potential energy of reactants < potential energy of products

b. potential energy of reactants > potential energy of products

c. thermal energy of reactants < thermal energy of products

d. thermal energy of reactants > thermal energy of products

6. Two metals of equal mass with different heat capacities are subjected to the same amount of heat. Which undergoes the smallest change in temperature?

a. The metal with the lower heat capacity. c. Both undergo the same change in temperature.

b. The metal with the higher heat capacity. d. You need to know the initial temperature of the metals.

7. Which of the following changes is NOT exothermic?

a. Melting copper. c. Freezing water.

b. Combustion of butane. d. Condensing steam.

8. A student studies two solutions. Solution A has a volume of 10 mL and is at a temperature of 25.0oC. Solution B has a volume of 500 mL and is at a temperature of 24.0oC. Which of the following statements must be true about the two solutions?

a. Solution A has more thermal energy than solution B.

b. Solution A has a higher specific heat capacity than solution B.

c. If the solutions were mixed, heat would transfer from B to A.

d. Solution B has more thermal energy than solution A.

Unit 12 Part 2: Cheeto Lab (Energy of Foods)

Objective: To be able to calculate the number of calories in one Cheeto by using the formula for specific heat.

Materials:

Thermometer, Ring stand w/ ring clamp, 100 mL graduated cylinder, Food holder (paper clip), Lighter, Stir rod, Electronic balance, Soda can, Cheetos, Water

Note: ALWAYS leave the plastic boat on a balance. It helps protect the balance from stains and marks. DON’T set the can down anywhere - it is messy!

Procedure:

1) Obtain one Cheeto and place on the exposed end of the paper clip. Find and record the initial mass of the food sample + food holder (paper clip). Do NOT eat in the lab!

2) Measure about 70 mL of water with the graduated cylinder. Record the precise volume and mass of the water in the data table. Pour the water into the soda can.

3) Assemble calorimetry apparatus as shown: the can should be suspended about 2.5 cm above the food sample. Include a piece of aluminum foil underneath the food sample for easier clean-up.

4) Place the thermometer into the soda can. Record the initial temperature and place the information in the data table.

5) Contact your teacher to help set the Cheeto on fire. Be sure the burning Cheeto is directly beneath the center of the can. Make sure the ENTIRE Cheeto is burning before removing the lighter.

6) Record the highest temperature of the water as the final temperature in the data table.

7) Determine and record the final mass of the burned Cheeto + food holder (paper clip) in the data table.

8) Place burned food sample into the trash once it has cooled - keep the food holder (paper clip) and aluminum foil.

** ONLY throw away the Cheeto **

Data:

Flamin’ Hot Cheeto

Initial mass of food and holder

Final mass of food and holder

Volume of water

Mass of water

Initial temperature, T i

Final temperature, Tf

Calculations/Analysis:

1. Show the calculation for the change in temperature of the water, ΔT.

2. Calculate the heat absorbed by the water, q. For water, C is 4.18 J/g°C. Convert your final answer to kJ.

3. Show the calculation to determine the mass (in g) of your food sample that burned.

4. Show the calculation to calculate the energy content (in kJ/g) of your food sample. (Use answers #2 and #3).

5. Based on your answer to #4, calculate the number of kilojoules in a 28 g serving of Cheetos. USE DIMENSIONAL ANALYSIS TO RECEIVE CREDIT.

6. Based on your answer to #5, calculate the number of Calories in a 28 g serving of Cheetos. (There are 4.18 kJ in one Calorie.) USE DIMENSIONAL ANALYSIS TO RECEIVE CREDIT.

7. According to the package, how many Calories are in a 28 g serving of Cheetos? ________________

8. Calculate your percent error:

9. Identify one or more sources of error that might have led to your calculated percent error.

Unit 12 Part 3: Calorimetry

Calorimetry: experimental technique used to measure the change in _____________ of a chemical reaction or phase

change.

• Step 1: Put a chemical reaction or phase change in contact with a _______________ bath.

• Step 2: Measure the temperature change of the water bath, then calculate the energy gained or lost by water

• Energy gained by water = energy lost by object!

+𝑞𝑞𝐻𝐻2𝑂𝑂 = −𝑞𝑞𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜 or − 𝑞𝑞𝑟𝑟𝑟𝑟𝑟𝑟

+[mCΔT]𝐻𝐻2𝑂𝑂 = −[mCΔT]𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜𝑜 or − [mCΔT]𝑟𝑟𝑟𝑟𝑟𝑟

Remember the specific heat of water: ____________________

Warning: Experimental Error with Calorimetry!

• We make the assumption in the above equation that ALL energy lost by the object is gained only by the

__________, but that’s not true!

• The calorimeter (container, usually a cup) can also ____________ heat, or heat can be lost to the _____.

• Both of these errors would lead to a calculated heat (q) that was ____ than the actual heat exchange.

Types of Calorimeters:

1. _________________ cup calorimeter: coffee cups are commonly used as insulators in intro chemistry classes

to measure temperature changes without a substantial loss of energy to the surroundings (they can be VERY

effective!)

2. _____________ calorimeters are used at the professional level. Bomb calorimeters provide greater

insulation and reduce heat loss to the surroundings (thus minimizing error).

Practice Makes Perfect!

1. When 25.0 g of a metal at 90.°C is added to a quantity of water at 25°C, the temperature of the water rises to29.8°C. What mass of water was used in the calorimeter? Assume no heat was lost to the surroundings. Thespecific heat capacity of the metal is 0.70 J/g°C, and the specific heat capacity of water is 4.18 J/g°C.

2. A 50 g sample of a metal is heated to 100oC and then placed in a calorimeter containing 100.0 g of water (CH2O =4.18 J/g°C) at 20oC. The final temperature of the water is 24oC. Which metal was used?

a. Lead (c = 0.14 J/g°C) c. Iron (c = 0.45 J/g°C)

b. Copper (c = 0.20 J/g°C) d. Aluminum (c = 0.89 J/g°C)

Example 1: A 5.037 g piece of iron heated to 100.°C is placed in a coffee cup calorimeter that initially contains 27.3 g of water at 21.2°C. If the final temperature is 22.7°C, what is the specific heat capacity of the iron (J/g°C)? The specific heat capacity of water is 4.18 J/g K.

Example 2: A 376 g sample of gold at 400. K is placed in a coffee cup calorimeter containing 50.0 mL of water at 300. K. Determine the final temperature of the water (assuming that no heat is lost to the surroundings). Thespecific heat capacity of gold is 0.128 J/g°C, and the specific heat capacity of water is 4.18 J/g°C.

Unit 12 Part 4: Enthalpy (∆H)

Enthalpy Change (∆H): amount of energy ________________________or ________________________ as heat by a

system when the pressure is constant; measured in units of ___________ =

∆𝐻𝐻𝑟𝑟𝑟𝑟𝑟𝑟 =𝑞𝑞

𝑚𝑚𝑚𝑚𝑚𝑚𝑟𝑟𝑟𝑟𝑟𝑟Standard Enthalpy Change: (∆Ho): enthalpy changed measured at ______________________ conditions

• Thermochemistry standard conditions are _______ the same as gas laws STP

• Thermochemistry standard conditions are: _____________ and _____________

Enthalpy Changes of Different Types of Reactions

• You will encounter a variety of _______________________ following the ΔH, however, they are simply

indicating a ___________________ type of reaction or change of state.

Examples

∆𝐻𝐻𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐°R = Enthalpy of Combustion (Heat Energy Released during Combustion Reactions)

∆𝐻𝐻𝑟𝑟𝑛𝑛𝑛𝑛𝑛𝑛°R = Enthalpy of Neutralization (Heat Energy Released during Acid-Base Neutralization Reactions)

∆𝐻𝐻𝑠𝑠𝑐𝑐𝑠𝑠𝑟𝑟°R = Enthalpy of Solution (Heat Energy Released/Absorbed Dissolving a Solute in Water)

∆𝐻𝐻𝑣𝑣𝑣𝑣𝑣𝑣°R = Enthalpy of Vaporization (Heat Energy Absorbed to Convert from Liquid to Gas Phase)

∆𝐻𝐻𝑓𝑓𝑛𝑛𝑠𝑠°R = Enthalpy of Fusion (Heat Energy Absorbed to Convert from Solid to Liquid Phase)

∆𝐻𝐻𝑓𝑓° R = Enthalpy of Formation (Heat Energy Released during Formation of 1 Mole of a Substance)

Standard Enthalpy (Heat) of Formation (______): change in enthalpy that accompanies the formation of _______

mole of the compound in its standard state from its component _____________________ their standard states.

Note: you will see ___________________ coefficients to ensure only ____ mole of compound is formed.

The ∆𝑯𝑯𝒇𝒇 ° for _______________ (in their standard state) is always _____ kJ/molrxn!

Practice Makes Perfect!

1. Write the formation reaction of liquid NH3:

2. Write the formation reaction of gaseous P2O5:

Example 1: → CH4(g) ∆𝐻𝐻𝑓𝑓° = −74.9 kJ/molrxn

Example 2: → SO3(g) ∆𝐻𝐻𝑓𝑓° = −396 kJ/molrxn

Unit 12 Part 5: Enthalpy of Formation (∆𝑯𝑯𝒇𝒇° )

∆𝑯𝑯𝒇𝒇 ° for _______________ (in their standard state) is always _____ kJ/molrxn!

Fun Fact: The enthalpy change of a given reaction (∆𝐻𝐻𝑟𝑟𝑟𝑟𝑟𝑟 ° ) can be calculated by combining the ∆𝐻𝐻𝑓𝑓

° for thereactants and products in the reaction!

Δ𝐻𝐻𝑟𝑟𝑟𝑟𝑟𝑟° = Σ �Δ𝐻𝐻𝑓𝑓°(𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝𝑝)� − Σ [Δ𝐻𝐻𝑓𝑓°(𝑝𝑝𝑟𝑟𝑟𝑟𝑝𝑝𝑝𝑝𝑟𝑟𝑟𝑟𝑝𝑝𝑝𝑝)]

Example 1: Use the data regarding the standard enthalpies of formation to calculate ∆𝐻𝐻𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐° for the following

reaction:

2 C3H6(g) + 9 O2(g) → 6 CO2(g) + 6 H2O(l)

Substance ΔH°f C3H6(g) 20.9 kJ/molrxn CO2(g) −393.5 kJ/molrxn

H2O(l) −286 kJ/molrxn

Example 2: Use the information provided and the balanced equation to determine ∆𝐻𝐻𝑓𝑓° of carbon tetrachloride.

CH4(g) + 4 Cl2(g) → CCl4(g) + 4 HCl(g) ΔH°rxn = −389 kJ/molrxn

Substance ΔH°f CH4(g) −75 kJ/molrxn

HCl(g) −92 kJ/molrxn

Practice Makes Perfect! For each reaction: (1) use the heats of formation in the chart below to determine ∆Hrxn

(2) predict whether the reaction is endothermic or exothermic & is energy absorbed or released

Name Formula ∆𝐻𝐻𝑓𝑓° Name Formula ∆𝐻𝐻𝑓𝑓°

hydrochloric acid HCl(aq) −167.16 kJ/molrxn zinc chloride ZnCl2(aq) −415.1 kJ/molrxn

carbon dioxide CO2(g) −393.51 kJ/molrxn calcium chloride CaCl2(aq) −877.3 kJ/molrxn

calcium carbonate CaCO3(s) −1207 kJ/molrxn liquid water H2O(l) −286 kJ/molrxn

Reaction #1: ZnCl2(aq) + H2(g) → Zn(s) + 2 HCl(aq)

Reaction #2: CaCO3(s) + 2 HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)

2 SO2(𝑔𝑔) + O2(𝑔𝑔) → 2 SO3(𝑔𝑔) ΔH𝑟𝑟𝑟𝑟𝑟𝑟𝑐𝑐 = −197.8 kJ/mol𝑟𝑟𝑟𝑟𝑟𝑟

3. The value of ΔHof at 298 K for SO3(g) is −395.7 kJ/mol. What is the value of the heat of formation of SO2(g)?

a. −594 kJ/mol b. −297 kJ/mol c. +297 kJ/mol d. +594 kJ/mol

4. If the standard enthalpies of formation of HBr(g) and Br2(g) are −36 kJ mol−1 and +31 kJ mol−1 (at 298 K)

respectively, what is ΔH°rxn for the following reaction?

H2(𝑔𝑔) + Br2(𝑔𝑔) → 2 HBr(𝑔𝑔)

a. −103 kJ/mol b. −67 kJ/mol c. +67 kJ/mol d. +103 kJ/mol

Unit 12 Part 6: Heating and Cooling Curves

Phase Changes

Endothermic changes = ________, require an addition of heat/energy to overcome IMFs and change state

Exothermic changes = ________, require heat/energy is released as IMFs form and the state changes

state change endothermic (+∆H) exothermic (−∆H)

solid ↔ liquid

liquid ↔ gas

solid ↔ gas

Heating Curve of Water

Energy Changes

1. _____________ (slanted) sections (1, 3, and 5): Temperature change

a. Temp (thus _____) changing

b. Potential energy (PE) constant

c. Constant ________ of matter

d. Use mCΔT

2. Flat (plateau) sections (2 and 4): 3 Ps: Plateau, Phase change, and Potential Energy Change

a. ___ change in temperature = KE constant

b. Potential energy increases (in the forward direction)

c. Two states of matter

d. Heat is used to overcome IMFs (attraction between particles), not change temperature!

e. Use q = n∆H (fusion or vaporization)

1

2

3

4

5

How to Calculate Energy Needed to Change Temperature or Phase

1. Temperature changes: 𝑞𝑞 = 𝑚𝑚 𝐶𝐶 ∆𝑇𝑇 This starts with grams and gives heat in Joules

2. Phase changes: 𝑞𝑞 = 𝑛𝑛 ∆𝐻𝐻 These start with moles and give heat in kilojoules

• Melting 𝑞𝑞 = 𝑛𝑛 ∆𝐻𝐻𝑓𝑓𝑓𝑓𝑓𝑓 Freezing 𝑞𝑞 = −𝑛𝑛 ∆𝐻𝐻𝑓𝑓𝑓𝑓𝑓𝑓 • Vaporizing 𝑞𝑞 = 𝑛𝑛 ∆𝐻𝐻𝑣𝑣𝑣𝑣𝑣𝑣 Condensing 𝑞𝑞 = −𝑛𝑛 ∆𝐻𝐻𝑣𝑣𝑣𝑣𝑣𝑣

How To Tell Whether You’re Doing A One-Step Or Multi-Step Calculation:

1. Is only one thing changing? (i.e. only temperature or only the state of matter?) One-step.

2. Are both temperature and state changing? Multi-step = multiple equations.

• For multi-step calculations, solve each equation separately and then add the results together, butmake sure you are adding the same units of energy! (J + J OR kJ +kJ)

Examples: Determine the amount of heat gained (+) or lost (−) during each of the following changes.

1. Melting 55.8 g of titanium at its melting point, 1677°C. (ΔHfus = 18.8 kJ/mol, ΔHvap = 425 kJ/mol, CTi = J/goC)

2. Converting 45.0 g of water at 20.0°C to steam at 115°C. (ΔHfus = 6.02 kJ/mol, ΔHvap = 40.7 kJ/mol, Csolid water =2.06 J/goC, Cliquid water = 4.18 J/goC, Cgaseous water = 2.02 J/goC)

n = moles of substance

ΔHfus = enthalpy of fusion

ΔHvap = enthalpy of vaporization

Practice Makes Perfect!: Calculate the heat changes associated with the following transitions. 1. 14.2 g of water at 100°C undergoes condensation. (ΔHfus = 6.02 kJ/mol, ΔHvap = 40.7 kJ/mol, Csolid water = 2.06

J/goC, Cliquid water = 4.18 J/goC, Cgaseous water = 2.02 J/goC)

2. 5.00 g of steam at 155°C is converted to 100% liquid water at 100.0°C. (ΔHfus = 6.02 kJ/mol, ΔHvap = 40.7 kJ/mol,Csolid water = 2.06 J/goC, Cliquid water = 4.18 J/goC, Cgaseous water = 2.02 J/goC

3. 6.9 g of solid aluminum is heated from 32oC to 320oC. The melting point of aluminum is 660oC. (ΔHfus = 10.8kJ/mol, ΔHvap = 284 kJ/mol, CAl = 0.903 J/goC)

4. Melting 27.3 g of Al at its melting point of 660°C. (ΔHfus = 10.8 kJ/mol, ΔHvap = 284 kJ/mol, CAl = 0.903 J/goC)

5. 220.0 g of ice at −35.0°C is converted to liquid water at 50.0°C. (ΔHfus = 6.02 kJ/mol, ΔHvap = 40.7 kJ/mol, Csolid

water = 2.06 J/goC, Cliquid water = 4.18 J/goC, Cgaseous water = 2.02 J/goC)

Directions: Answer the following questions using this heating curve for substance X:

1. In what part of the curve would substance X have a definite shape and definite volume?_________________

2. In what part of the curve would substance X have a definite volume but no definite shape? _______________

3. What part of the curve represents a mixed solid/liquid phase of substance X? _________________

6. What part of the curve represents a mixed liquid/vapor phase of substance X? _________________

7. What is the melting temperature of substance X? _________________

8. What is the boiling temperature of substance X? _________________

9. In what part(s) of the curve would increasing kinetic energy be displayed? _________________

10. In what part(s) of the curve would increasing potential energy be displayed? _________________

11. In what part of the curve would the molecules of substance X be farthest apart? _________________

12. In what part of the curve would the molecules of X have the lowest kinetic energy? _________________

13. In what part of the curve would the molecules of X have the greatest kinetic energy? _________________

Unit 12 Part 7: Thermochemical Equations

Thermochemical equation

• Chemical equation that includes the __________ change (the energy value)

Endothermic Exothermic

Energy is _____________________ Energy is _____________________

+q/molrxn = +ΔHrxn −q/molrxn = −ΔHrxn

_____________________bonds Form _____________________

Energy appears in _____________________ Energy appears in _____________________

The energy added (for endo AND exo) will always be ________________________!

Examples:

Equation with Separate ∆Horxn Thermochemical Equation Endo- or

exothermic?

CH4 + 2 O2 → CO2 + 2 H2O + 890 kJ

H2O(l) → H2O(g) ∆Hvap = 44 kJ/molrxn

2 H2O + 571.6 kJ → 2 H2 + O2

Energy Stoichiometry! • Enthalpy is commonly measured in kJ/molrxn, but what is a mole of reaction?

1 molrxn = 1 mole of reaction = stoichiometric # of reactants/ products

• For the combustion of ethane: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2Oo When ____ mole of reaction has occurred,

____ mol of C2H6 reacted ____ mol of O2 reacted 3120 kJ energy _________________

____ mol of CO2 were produced ____ mol of H2O were produced

• Luckily for us, the enthalpy of a reaction, when measured in kJ/molrxn, can act as a _______________________factor between the amount of chemicals which react and the energy that is absorbed or released by thereaction!

Example 1: Given the following reaction, 2 Fe + 3 CO2 → 3 CO + Fe2O3 (∆H = +25 kJ/molrxn) what energy change occurs when 6.00 moles of carbon dioxide react?

Example 2: Given the following reaction, N2 + 3 H2 → 2 NH3 (∆H = −324 kJ/molrxn) what mass of hydrogen must have reacted if 525 kJ of heat energy were released?

Example 3: Given the reaction 3 NaOH +H3PO4 3 H2O + Na3PO4 + 156 kJ: a) How much heat is absorbed or released if 200. mL of 3.0 M of NaOH reacts with 300. mL of H3PO4?

(Assume NaOH is limiting.)

b) When this reaction occurs in a calorimeter, what will the temperature change be? (Assume the total massof the calorimeter is 500. g and the specific heat of solution is 4.18 J/g°C).

Practice Makes Perfect!

Equation with Separate ∆Horxn Thermochemical Equation Endo- or

exothermic?

2 C2H6 + 7 O2 → 4 CO2 + 6 H2O ∆Horxn = −3120

kJ/molrxn

NH3 + HCl → 176 kJ + NH4Cl

4 NO + 6 H2O → 4 NH3 + 5 O2 ∆Horxn = 1170

kJ/molrxn

1. Solid sodium hydrogen carbonate decomposes into solid sodium carbonate, liquid water, and carbon dioxidegas. (ΔHrxn = +85 kJ/molrxn)

a. Write the balanced thermochemical equation for this reaction.

b. Is the reaction endothermic or exothermic? _______________

c. What is the energy change that occurs when 2.25 mol of NaHCO3(s) decomposes?

2. The heat of combustion of propane (C3H8) is −2220 kJ/molrxn.

a. Write the balanced thermochemical equation for the combustion of propane.

b. Is the reaction endothermic or exothermic? _______________

c. What mass of propane must be burned to release 5,550 kJ of heat?

3. A 1.25 g piece of solid magnesium is added to 25.0 mL of HBr(aq), releasing 22,300 J of heat energy. Assume thatthe specific heat of the solution is 4.184 J/g°C, and the total mass of the solution with all reactants is 26.25 g. Ifthe reactants were originally at 25.0oC, calculate the final temperature of the solution.

4. A 523 g cube of brass increases from 25.0°C to 266.5°C when the metal absorbs 48 kJ of heat. What is the specificheat of brass?

5. The heat of neutralization of phosphoric acid (H3PO4) with NaOH is −152.2 kJ/molrxn. How many grams of sodiumhydroxide would need to react with excess phosphoric acid to produce 867.5 kJ of heat energy?

6. Given the reaction Ba(OH)2 · 8H2O + NH4Cl + 164 kJ/molrxn BaCl2 + 2 NH3 + 10 H2O:a) How much heat is absorbed or released if 147 g of ammonium chloride react completely with barium oxide

octahydrate?

b) When this reaction occurs in a calorimeter, what will the temperature change be? (Assume the total mass ofthe calorimeter is 850. g and the specific heat of solution is 4.18 J/g°C).

Unit 12 Part 8: Bond Energy and Enthalpy

Covalent Bond Energies and Enthalpy

• The enthalpy change for a reaction, ΔHrxn, can also be understood in terms of bonds ___________________

(endothermic) and bond ___________________ (exothermic) during a chemical reaction.

• The total enthalpy change can be negative (exothermic) or positive (endothermic) depending on the relative

magnitude of two (breaking and making) processes.

Endothermic: +∆Ho (system gains energy) Exothermic: −∆Ho (system loses energy)

Breaking a chemical bond is alwaysendothermic

Hint: “end-ing a bond is endothermic”

Greater magnitude of energy absorbed whenbonds break

More energy needed to break the bonds inthe reactants than is released when newbonds form in the products

Forming a chemical bond is alwaysexothermic (more stable)

Greater magnitude of energy released whenbonds form

Less energy is needed to break the bonds inthe reactants than is released when newbonds form in the products

Practice Makes Perfect!

1. If, during a reaction, more energy is released in bond formation than is absorbed in bond breaking, the reaction__________.

a. must be exothermic b. must be endothermic

2. When bonds are broken, ____________.a. energy is always required b. energy is always released

3. If the following reaction is exothermic, which is greater, the energy required to break the existing bonds, or theenergy released when new bonds are formed?

2Na(s) + Cl2(g) 2NaCl(s)

a. energy released when new bonds are formedb. energy required to break the bonds

Examples 1. For each example below, identify if the process is endothermic or exothermic, and explain why.

Process Endo or Exothermic? Why?

H(g) + H(g) H2(g)

F(g) + e− F−(g)

F−(g) F(g) + e−

N2(g) N2(l)

Cl2(g) Cl(g) + Cl(g)

2. Complete the chart below.

Equation with Separate ∆Horxn ∆Horxn within the Equation Endo- or exothermic?

Which is greater in

energy: bonds broken or

bonds formed?

K + M → N + 45 kJ

D → E + F ∆Horxn = 127kJ/molrxn

A + B → C + D ∆Horxn = −35 kJ/molrxn

4. Which of the following is a graph that describes the pathway of a reaction that is endothermic and has a highactivation energy?

5. When solid ammonium chloride, NH4Cl(s), is added to water at 25oC, it dissolves and the temperature of thesolution decreases. Which of the following is true for the type of energy conversion that happens during thedissolving process?

Type Energy Conversion

a. Endothermic thermal energy → chemical potential energy b. Exothermic thermal energy → chemical potential energy c. Endothermic chemical potential energy → thermal energy d. Exothermic chemical potential energy → thermal energy

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) ΔHocomb = −890 kJ/molrxn

6. How much heat is absorbed or released when 2.0 mol of CH4(g) reacts with 2.0 mol of O2(g)?

a. 890 kJ of heat is released. c. 1780 kJ of heat is released.

b. 890 kJ of heat is absorbed. c. 1780 kJ of heat is absorbed.

7. The value of ΔHofus for bromine is 10.6 kJ mol−1. To which process does this value refer?

a. Br(s) → Br(l) c. Br2(s) → Br2(l)

b. Br(l) → Br(s) c. Br2(l)→ Br2(s)

8. In an endothermic reaction, which of the following is true?

a. energy content of reactants < energy content of products < activation energy

b. energy content of reactants < activation energy < energy content of products

c. energy content of products < energy content of reactants < activation energy

d. energy content of products < activation energy < energy content of reactants

9. Breaking bonds , while forming bonds .

a. releases energy; requires energy

b. releases energy; releases energy

c. requires energy; releases energy

d. requires energy; requires energy

10. Which one of the following statements best describes the enthalpy change of a reaction?

a. The energy released when chemical bonds are formed during a chemical reaction

b. The energy consumed when chemical bonds are broken during a chemical reaction

c. The difference between the energy released by bond formation and the energy consumed by bond cleavage

during a chemical reaction

d. The increase in disorder of the system as a reaction proceeds

11. Label each of the following processes as exothermic or endothermic. Label each as having a positive or negativevalue of ΔH.

Process Endo or Exothermic? Positive or Negative value of ΔH?

NH3(g) → NH3(l)

Br(g) + e− Br−(g)

I−(g) I(g) + e−

N2(l) N2(g)

O(g) + O(g) O2(g)

Br2(l) → Br2(s)

CaS(s) Ca(s) + S(s)

Unit 12 Part 9: Energy Warm-Up

Glucose (C6H12O6) can be formed in the cells of green plants through an endothermic reaction of carbon dioxide and water. This process, known as photosynthesis, occurs using energy provided by the sun.

2,800 kJ + 6 CO2 + 6 H2O C6H12O6 + 6 O2

a) In the process of photosynthesis, which requires a greater magnitude of energy: breaking bonds or formingbonds? How can you tell?

b) What is the enthalpy value for photosynthesis, ∆Hrxn, in kJ/mol?

c) How much heat would be absorbed or released if 17.2 g of carbon dioxide reacted with excess water?Assume 100% of the carbon dioxide reacts.

d) If this reaction occurs in a calorimeter, what will the temperature change be? (Assume the total mass of thecalorimeter, including the reacting chemicals, is 245 g and the specific heat of solution is 4.18 J/goC).

e) The initial temperature of the solution in the calorimeter is 28.3oC. What would the final temperature beafter the reaction?

Unit 12 Part 10: Hess’s Law Hess’s Law

• Sometimes it is impossible or impractical to measure the ΔH of a reaction by using a calorimeter. In these

situations, ΔH can be calculated using Hess’s Law!

• Hess’s Law: Combining two or more reactions to achieve a goal reaction

o When adding given reactions, they combine to produce the _______ reaction.

o ΔHnew rxn = ΔHrxn 1 + ΔHrxn 2 + ΔHrxn 3 + …

But… life isn’t always that straightforward. Sometimes you have to manipulate your given reactions they sum to your goal reaction. If so, you also need to manipulate ΔHrxn using the following rules:

Rule 1: If you reverse the reactions, then change the sign of ΔH. For example,

Rule 2: If you multiply the reaction by a coefficient, then multiply ΔH by same coefficient. For example,

Rule 3: Rule 1 and 2 can be combined! For example, if the first reaction is tripled and reversed,

Example 1: X + B → D ΔHrxn = ?

A + B → C + 3 D ΔHrxn = 270 kJ/molrxn 2 D + C + X → A ΔHrxn = −630 kJ/molrxn

Example 2: Find the enthalpy change for the formation of PbCl4 by the reaction of lead(II) chloride with chlorine.

PbCl2(s) + Cl2(g) → PbCl4(l) ∆H = ?

Use the following thermochemical equations:

Pb(s) + 2 Cl2(g) → PbCl4(l) ∆H = –329.2 kJ/molrxn

Pb(s) + Cl2(g) → PbCl2(s) ∆H = –359.4 kJ/molrxn

Strategy

• Find things in your goal equation that appear in only one of the available reactions and make them match byflipping equations or multiplying/dividing coefficients. Arrange equations to cancel out things that do ______appear in the goal.

Whatever you do to the equation, you must do to _______!

• When combining reactions:

o Reagents on the same side: add together

o Reagents on the opposite side: subtract (from the side with the greatest amount)

Watch out for ____________: you can only add and subtract substances in the same state of matter!

Let’s Practice! 1. Given the following information, find the heat of formation for methane: C(s) + 2 H2(g) → CH4(g)

Example3: Find the enthalpy change for the following reaction: 2 N2(g) + 5 O2(g) → 2 N2O5(g)

Use the following thermochemical equations:

N2O5(g) + H2O(l) → HNO3(l)

H2(g) + ½ O2(g) → H2O(l)

½ N2(g) + 32

O2(g) + ½ H2(g) → HNO3(l)

∆H = –76.7 kJ/molrxn

∆H = –286 kJ/molrxn

∆H = –171 kJ/molrxn

2. Find the enthalpy change for the formation of pentane, C5H12, by the reaction of carbon with hydrogen.5 C(s) + 6 H2(g) → C5H12(g)

Use the following thermochemical equations:

C(s) + O2(g) → CO2(g) ∆H = −393.5 kJ/molrxn

H2(g) + ½ O2(g) → H2O(l) ∆H = −285.8 kJ/molrxn

C5H12(g) + 8 O2(g) → 5 CO2(g) + 6 H2O(l) ∆H = −3535.6 kJ/molrxn

3. Calculate the heat of formation for sulfur dioxide, SO2, from its elements sulfur and oxygen. Use thebalanced chemical equation and the following information.

S(s) + 32

O2(g) → SO3(g) ∆H = −395.2kJ/molrxn

2 SO2(g) + O2(g) → 2 SO3(g) ∆H = −198.2kJ/molrxn

4. Based on the information given above, what is ∆Ho for the following reaction?

CH4(𝑔𝑔) + 2 O2(𝑔𝑔) → CO2(𝑔𝑔) + 2 H2O(𝑙𝑙)

a. x + y + z c. y + z – 2x

b. x + y – z d. y + 2z – x

Unit 12 ExamFree Response Review #1

Pre-AP Chemistry Directions: The suggested time is about 15 minutes for answering the constructed response section of the chemistry test. The parts within a question may not have equal weight. For calculations, show all your work in the spaces provided after each part. Pay particular attention to the proper use of units. Be sure your final answer is rounded to the correct number of significant figures. Make sure your work is legible. Illegible work will receive a grade of zero.

Question 1 [4 POINTS] The following is a heat curve for chloroform, a solvent for fats, oils, and waxes. It is an organic compound with formula CHCl3 and a molar mass of 119.368 g/mol. The heating curve for a 1.50 mol sample of chloroform is shown below.

(a) Label each section of the graph with statesand phase change (if applicable): [1.5 POINTS]

A:

B:

C:

D:

E:

(b) Determine the boiling point of chloroform. [0.5 POINTS]

(c) Calculate the amount of heat needed to completely vaporize a 1.50 mol sample of liquid chloroformoriginally at −60.0°C. The specific heat capacity of the substance in the liquid phase is 0.965J/(g°C), and the heat of vaporization of the substance is 32.5 kJ/mol. [2 POINTS]

___ - ___ - ___ = ___ /10 INCORRECT

SIG FIGS INCORRECT

UNITS

Question 2 [6 POINTS]

Heats of Formation

Iron (III) oxide Fe2O3 −824.2 kJ/mol

Carbon monoxide CO −110.5 kJ/mol

Carbon dioxide CO2 −393.5 kJ/mol

Iron (III) oxide reacts with carbon monoxide to create solid iron and carbon dioxide gas as shown below:

Fe2O3 (s) + 3 CO(g) 2 Fe(s) + 3 CO2(g)

(a) Use the heats of formation in the chart above to determine the heat of the reaction, ΔHrxn.[2 POINTS]

(b) Is the reaction endo or exothermic? Justify your answer. [1 POINT]

(c) Write the complete thermochemical equation for the reaction. [1 POINT]

(d) How much energy is absorbed/released if 2.00 moles of carbon monoxide reacted with iron (III) oxide?[2 POINTS]

Unit 12 ExamFree Response Review #2

Pre-AP Chemistry Directions: The suggested time is about 15 minutes for answering the constructed response section of the chemistry test. The parts within a question may not have equal weight. For calculations, show all your work in the spaces provided after each part. Pay particular attention to the proper use of units. Be sure your final answer is rounded to the correct number of significant figures. Make sure your work is legible. Illegible work will receive a grade of zero.

Question 1 [3 POINTS] A 15 g iron nail is heated to 95.0°C and placed into a beaker of water.

(a) Calculate the final temperature of the water if the amount of heat absorbed by the water is 0.251kJ. The specific heat capacity of iron = 0.449 J/g° C, and the specific heat capacity of water =4.18 J/g °C. [3 POINTS]

Question 2 [7 POINTS] Hydrochloric acid and barium hydroxide undergo a neutralization reaction, with ΔHrxn = −118 kJ/molrxn.

(a) Write the balanced thermochemical equation for this reaction. [2 POINTS]

(b) Sketch a potential energy diagram for this reaction on the axis below. [1 POINT]

___ - ___ - ___ = ___ /10 INCORRECT

SIG FIGS INCORRECT

UNITS

(c) Is the magnitude of the energy change greater for the bonds being broken in the reactants or thebonds being formed in the products? Justify your answer. [1 POINTS]

(d) When the two reactants are mixed, does the temperature of the water they’re dissolved in increase,decrease, or stay the same? Justify your answer. [1 POINT]

(e) Given the heats of formation provided in the table, calculate ΔHf of barium hydroxide. [2 POINTS]

Heats of Formation

Hydrochloric acid −167.2 kJ/mol

Water (l) −286 kJ/mol

barium chloride −855 kJ/mol

Unit 12: Energy

Multiple Choice Practice Directions: Each of the questions or incomplete statements below is followed by four suggested answers or

completions. Select the answer that is best in each case and then fill in the corresponding circle on the answer sheet.

Note: For all questions, assume that the temperature is 298K, the pressure is 1.00 atm, and solutions are aqueous unless otherwise specified.

3

4

1. Which of the following terms describes a reaction where more energy is released by bond formation than

is consumed by bond breaking?

a. Exothermic

b. Endothermic

How much energy must be added to 36 g of ice at 0oC in order to raise its temperature to 25oC? (The molecular mass of water is 18.0 g/mol, and ΔHfus(H2O) = 6.02 kJ/mol.)

7

5

9

8

6

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) ΔHocomb = −890 kJ/molrxn

Which of the following correctly describes the reaction above?

A. 890 kJ of energy released when bonds form C. 890 kJ of energy released when bonds break

B. 890 kJ of energy absorbed when bonds form D. 890 kJ of energy absorbed when bonds break

10

11

12

A. 1 B. 2 C. 3 D. 3

A. aluminum B. iron C. mercury D. silver

13

14

15

EZPZ Thermo Review 1. One form of energy can be converted to another. Here are some examples:

a) A flaming hot Cheeto is burned in chemistry lab.

As __________________________energy is converted to _________________________energy.

b) A space heater is plugged into an electrical outlet and used to warm up a room.

As __________________________energy is converted to _________________________energy.

c) The process of photosynthesis is used to make a tasty carrot.

As __________________________energy is converted to _________________________energy.

2. Complete the chart below:

∆H (+ or - ?)

Heat (Absorbed or

released?)

Type of Reaction Endothermic or Exothermic?

energy+2H2S+ SO2 3S + 2H2O

N2(g)+ 3H2 (g) 2NH3 + energy

Ice freezing

Dry ice sublimation

3. Use Hess’s Law to determine ΔH for the following target reaction.a) 2 NOCl (g) → N2 (g) + O2 (g) + Cl2 (g) ΔH = ?

½ N2 (g) + ½ O2 (g) → NO (g) ΔH = 90.3 kJ NO (g) + ½ Cl2 (g) → NOCl (g) ΔH = –38.6 kJ

b) N2(g) + 2O2(g) → 2NO2(g) ΔH1 = ? N2(g) + O2(g) → 2NO(g) ΔH = 180 kJ 2NO(g) + O2(g) → 2NO2(g) ΔH = – 112 kJ

c) 2 N2 + 6 H2O 3 O2 + 4 NH3 ΔH = ?

4. Answer the questions below

1) Endothermic or exothermic?

2) Energy absorbed or released?

1) Endothermic or exothermic?

2) Energy absorbed or released?

5. Enthalpy of Reaction Practice

For each reaction: (1) Use the heats of formation in the chart below to determine the heat of each reaction, and (2) write the thermochemical equation.

Name Formula ΔHof name formula ΔHo

f

Calcium oxide CaO -634.9 kJ/mole Carbon monoxide CO -110.5 kJ/mole

Hydrochloric acid HCl -167.2 kJ/mole Iron (II) oxide FeO -272.0 kJ/mole

Iron (III) oxide Fe2O3 -824.2 kJ/mole Hydrobromic acid HBr -36.4 kJ/mole

Carbon dioxide CO2 -393.5 kJ/mole Ethane C2H6 -83.8 kJ/mole

Calcium carbonate CaCO3 -1207.6 kJ/mole Water (liquid) H2O -285.8 kJ/mole

Reaction #1: Calcium carbonate decomposes to calcium oxide and carbon dioxide.

Reaction #2: Hydrobromic acid reacts with chlorine.

Reaction #3: Fe2O3 (s) + 3CO(g) 2Fe(s) + 3CO2(g)

Reaction #4: 2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(l)

6. Solve these problems:a) A 1.5 g iron nail is heated to 95.0 °C and placed into a beaker of water. Calculate the heat gained by the water if

the final equilibrium temperature is 57.8 ° C. The specific heat capacity of iron = 0.449 J/ g° C, and the specificheat capacity of water = 4.18 J/ g ° C.

b) The main engines of the space shuttle burn hydrogen to produce water. How much heat (in kJ) is associatedwith this process if 1.32 × 105 kg of liquid H2 is burned?

2 H2 (l ) + O2 (l ) → 2 H2O (l ) ΔHºrxn = – 571.6 kJ

c) A 32.5 g cube of aluminum initially at 45.8 °C is submerged into 105.3 g of water at 15.4 °C. What is the finaltemperature of both substances at thermal equilibrium? The specific heat of aluminum is 0.903 J/ g° C.

d) A block of copper of unknown mass has an initial temperature of 65.4 °C. The copper is immersed in a beakercontaining 95.7 g of water at 22.7 °C. When the two substances reach thermal equilibrium, the finaltemperature is 24.2 °C. What is the mass of the copper block? (Use the specific heat chart to find Cs of waterand copper.)

e) An LP gas tank in a home barbeque contains 13.2 kg of propane, C3H8. Calculate the heat (in kJ) associated withthe complete combustion of all of the propane in the tank.

f) What mass of butane in grams is necessary to produce 1.5 × 103 kJ of heat? What mass of CO2 is produced?

g) If 5750 J of energy is added to 455g piece of glass, what was the temperature change of the glass? The specificheat of glass is 0.50 J/ g° C.

7. Determine the amount of heat gained (+) or lost (−) during each of the following changes.

a) Melting 55.8 g of Ti at 1677°C

b) Converting 45.0 g of water at 20.0°C to steam at 115°C

c) Condensing 14.2 g of water at 100°C

d) Heating 6.9 g of solid aluminum from 32°C to 32°C

e) Melting 27.3 g of Al at 660°C

f) 220.0 g of ice at −35.0°C is converted to liquid water at 50.0°C.

g) 5.00 g of steam at 155°C is converted to liquid water at 100.0°C

8. Given the reaction Ba(OH)2 · 8H2O + NH4Cl + 164 kJ/molrxn BaCl2 + 2 NH3 + 10 H2O:a) How much heat is absorbed or released if 223 g of ammonium chloride react completely with barium oxide

octahydrate?

b) When this reaction occurs in a calorimeter, what will the temperature change be? (Assume the total mass of thecalorimeter is 420. g and the specific heat of solution is 4.18 J/g°C).

9. Multiple Choice Questions1) Identical amounts of heat are applied to 50 g blocks of lead, silver, and

copper, all at an initial temperature of 25 °C. Which block will have thelargest increase in temperature?

a) Lead

b) Silver

c) Copper

d) None; all will be at the same temperature.

2) Which of the following statements about enthalpy is FALSE?

a) The value of ΔH for a chemical reaction is the amount of heat absorbed orevolved in the reaction under conditions of constant pressure.

b) An endothermic reaction has a positive ΔH and absorbs heat from thesurroundings. An endothermic reaction feels cold to the touch.

c) An exothermic reaction has a negative ΔH and gives off heat to thesurroundings. An exothermic reaction feels warm to the touch.

d) Statements (a), (b), and (c) are all false statements.

e) Statements (a), (b), and (c) are all true statements.

3) For the reaction below, the enthalpy change is +624.7 kJ. How would you classify this reaction?SiO2(g) + 3C(s) --> SiC(s) + 2CO(g)

a) endothermic reaction, heat is lost from the systemb) endothermic reaction, heat is gained by the systemc) exothermic reaction, heat is lost from the systemd) exothermic reaction, heat is gained by the system

4) Which one of the following would have an enthalpy of formation value (ΔHf) of zero?a) O2(g) b) H2O(l) c) NO(g) d) H2O(g)

5) Which of the following statements does NOT describe an exothermic reaction?a) 2 C2H6 + 7 O2 --> 4 CO2 + 6 H2O + heat c) Energy is a reactantb) Releases heat into the surrounding. d) ΔH = -267.9 kJ