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Chemistry 30 Unit 2: Energy

Name: __________________

Spring

2019

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Students will:

30–A1.1k use Q = mc∆t to the analysis of heat transfer

30–A1.2k explain, in a general way, how stored energy in the chemical bonds of hydrocarbons originated from the sun

30–A1.3k define enthalpy and molar enthalpy for chemical reactions

30–A1.4k write balanced equations for chemical reactions that include energy changes 30–A1.5k use and interpret ∆H notation to communicate and calculate energy changes in chemical reactions

30–A1.6k predict the enthalpy change for chemical equations using standard enthalpies of formation

30–A1.7k explain and use Hess’ law to calculate energy changes for a net reaction from a series of reactions

30–A1.8k use calorimetry data to determine the enthalpy changes in chemical reactions 30–A1.9k identify that liquid water and carbon dioxide gas are reactants in photosynthesis and products of cellular respiration and that gaseous water and carbon dioxide gas are the products of hydrocarbon combustion in an open system

30–A1.10k classify chemical reactions as endothermic or exothermic, including those for the processes of photosynthesis, cellular respiration and hydrocarbon combustion.

30–A2.1k define activation energy as the energy barrier that must be overcome for a chemical reaction to occur

30–A2.2k explain the energy changes that occur during chemical reactions, referring to bonds breaking and forming and changes in potential and kinetic energy

30–A2.3k analyze and label energy diagrams of a chemical reaction, including reactants, products, enthalpy change and activation energy

30–A2.4k explain that catalysts increase reaction rates by providing alternate pathways for changes, without affecting the net amount of energy involved; e.g., enzymes in living

systems.

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For this Course Energy Comes in Two Forms:

Potential energy (Ep) is stored energy (For much of our unit, it will be the stored energy in the bonds within and between molecules – chemical energy). During a phase change, Ep changes without a temperature change, and visually this is displayed as the flat lines on the heating and cooling curves below.

Kinetic energy (Ek) is the energy of motion of ions, atoms, molecules (Translational, rotational, vibrational motion). Ek is temperature dependent!

An Introduction to Terms:

Thermochemistry is the study of heat/energy changes associated with physical, chemical and nuclear processes.

HEAT is a form of energy that flows between two samples at different temperatures (Heat always moves from hot to cold)

TEMPERATURE is the average kinetic energy of the molecules. It does not reflect total energy, only the random movement of atoms and molecules.

THERMAL ENERGY is a measure of the total heat energy in a substance.

Heat and temperature are related however. If something absorbs heat energy, the temperature increases (except during phase changes).

SECOND LAW OF THERMODYNAMICS heat is spontaneously transferred from an object at a higher temperature to an object with lower temperature until they both reach the same temperature.

Most substances are solids or liquids at room temperature; few are gases.

Solids and liquids are called the condensed phase (the particles are close together).

The attractive forces between the molecules (intermolecular attractions) hold the molecules closely together.

Three Phases of Matter

Solid Phase Liquid Phase Gas Phase

In a solid, the intermolecular attractions are strong, resulting in a rigid structure:

There is only some back and forth movement (vibrational motion).

The particles have very strong attractions and are densely packed.

In liquids, there is enough kinetic energy to overcome some of the intermolecular attractions, resulting in a loosely bonded, flowing system:

The particles are freer to move in any direction (vibrational, rotational and translational motion).

There are still strong attractions between particles, and so particles are still densely packed.

Gas molecules have so much kinetic energy, that the intermolecular forces are overcome, and become negligible, resulting in a completely random system:

The molecules have vibrational(little), rotational and translational motion, but tend to have mostly translational motion.

Because there are essentially no intermolecular attractions, particles are very far apart.

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K inetic Energy, E

k

– Changes in Intermolecular Bonding

The heat capacity of a substance is a measure of the energy required to raise that amount of substance by one degree Celsius. The units for heat capacity are J/°C. Just like enthalpy, heat capacity will change depending on the amount of substance involved. For example, it will take more energy to heat up a bathtub of water to 55 °C than a cup of water to 55 °C. And just like enthalpy, we standardize our heat capacity measurements so they can be compared. Molar heat capacity (J/mol•°C) allows us to see which substances require more or less energy to heat up or cool down because they all involve the same amount. Unlike enthalpy, specific heat capacity (J/g•°C) is used, because, in industrial applications, we want to know how quickly something might heat up or cool down, and we can more easily measure the mass of a substance than moles.

SPECIFIC HEAT CAPACITY, c, The amount of heat energy needed to raise the temperature of 1 g of a substance 1°C (or 1K)

The specific heat is different for different substances and for different states. These values are found on page 3 of your Data Booklet.

Looking in the data booklet, we see that water can store more energy than Air (O2(g)) or Al(s) i.e. It takes longer for water to heat up, but it will also take longer for it to cool down.

This makes water a good coolant (stores heat well).

When 1g of water goes up 1°C it absorbs 4.19J of energy. Water has to lose more heat to lower its temperature by 1°C

When 1g of Al goes up 1°C, it only absorbs 0.903J of energy. When the temperature of Al drops by 1°C, only 0.903J of energy is given off. This explains why you can touch Al foil that has just come out of a hot oven.

As well, a drop of boiling water on your skin will not destroy skin tissue.

A CUP OF WATER will scald your hand.

The higher the specific heat capacity of a substance and the greater it’s mass, the more energy it can store.

We can calculate the heat transferred to or from a substance if we know the specific heat capacity, mass and temperature change of the substance.

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Remember 1mL of water is = 1g of water.

Examples of Problems You Will Be Expected To Do For Kinetic Energy Changes

E k = mc∆t

Where:

Ek = kinetic energy (J)

m = mass of substance that is heating or cooling (g) c = specific heat capacity of the substance (J/g·°C)

∆t = change in temperature (°C) where ∆t = tfinal - tinitial

Ek has a negative value when something loses heat (temperature decreases) EXOthermic Ek has a positive value when something gains heat (temperature increases) ENDOthermic Remember that Ek , m , c, t values are all for the same substance

Example #1: A propane stove heats up 150 g of water from 20.0

o

C to 53.2

o

C. How much energy is required to heat the water?

Example #2: A certain mass of tin is cooled from 1200C to 850C and releases 1.50 MJ of energy. What is the mass of tin used?

Example #3: If 20.0 kJ of energy is absorbed by a 500 g piece of aluminum at -30.0

o

C, what is the final temperature of the aluminum?

You may also see q = mc∆t, but we

think that using E

k

as a symbol

makes thermochemistry a little

easier to understand.

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Practice Questions:

1. What is the SI unit for energy?

2. Define:

a. Kinetic Energy (E

k

):

b. Potential Energy (E

p

):

3. What is the relationship between temperature and heat?

4. What is the Second Law of thermodynamics?

5. How is energy transferred from a hotter substance to a cooler substance according to the particle theory of matter?

6. A hot flame is use to raise the temperature of 500 g of water from 20.0°C to 40.0°C? How much energy is transferred from the flame to the water?

7. A 100g sample of copper is placed in cold water causing the copper to cool from 55.3°C to 12.8°C. How much energy is released by the copper?

8. If 250 g aluminium pot cools, after it was used to boil water on the stove, from 100°C down to 55.0°C. Find the amount of energy released to the surroundings.

9. It required 50.0 kJ of heat energy to heat 500 g of a solid from 5.0°C to 50°C. Calculate the specific heat

capacity of the solid.

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10. A red hot iron bar is lowered into a 500g sample of water which is at a temperature of 25.0°C. If the water sample absorbs 5.00 kJ of energy calculate the final temperature of the water.

11. If 200 g of water at 10.0°C absorbs 18.0 kJ of energy what is the temperature change of the water?

Potential Energy, E

p

– Changes in Intramolecular bonding

Calculating Potential Energy Changes (Enthalpy):

Enthalpy or Heat(∆H) is a measure of the total energy possessed by a substance. Enthalpy is a property of a system that reflects its capacity to exchange energy (usually heat – and this is why we use the symbol “H”, but sometimes other forms such as light or electricity) with the surroundings at constant pressure during a physical or chemical change.

Enthalpy can be different for a substance depending upon its state (gas > liquid > solid), and it will be different at different temperatures (the substance will possess greater Ek at greater temperature).

We CAN’T measure the enthalpy of a system.

We CAN measure the CHANGE in enthalpy (∆H) of a system. (∆H =Hfinal - Hinitial)

Usually we do this indirectly by looking at temperature changes as energy is given off or absorbed.

Communicating Enthalpy:

Exothermic Processes (∆H = negative) Endothermic Processes (∆H = positive)

Energy is lost by a system to its surroundings. Energy is gained by a system from its surroundings.

Energy is given off. Energy is absorbed

The products have less energy than reactants. The products have more energy than reactants.

Often Molar enthalpy, ∆Hº (kJ/mol) will have a subscript to help us understand the type of change occurring. You should be prepared to recognize the different possible ∆H subscripts such as:

f Molar Enthalpy of formation(Values found on page 4+5 of Data Book)

d Molar Enthalpy of decomposition (Found as the –(∆fHº ) )

c Molar Enthalpy of combustion

rxn Molar Enthalpy of reaction

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Enthalpy is most often measured in Joules or kiloJoules (J or kJ).

However, a substance will have more enthalpy if more of it takes part in a physical or chemical change (a

firecracker of TNT will have less enthalpy than a large stick of dynamite); and therefore we try to standardize our enthalpy measurements so they can be compared. Molar enthalpy (kJ/mol) allows us to see which substances release more of less energy during a change because they all involve the same amount.

One substance with a certain amount of energy is converted into another kind of substance with a different amount of energy.

Intramolecular Change Bonds between atoms (strong) Range from 100 kJ to <5000kJ

Ex. the burning of fuels

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + ENERGY (exothermic)

cHo = -802 kJ/mol or ∆H = -802 kJ

When calculating for Enthalpy we use:

E p = n Δ r H

Example 1: The molar enthalpy of combustion of methanol is -726 kJ/mol. What is the enthalpy change of combustion for methanol if 5.00g of methanol is burned?

Example 2: What mass of ethanol is formed if 800kJ of energy is released?

Example 3: What is the enthalpy change when 200 grams of sodium chloride is formed?

NOTE: They can cause a temperature change to occur but they do NOT undergo a temperature change!!!

Ep = potential Energy (kJ) n = moles (mol)

rHº = the molar enthalpy of a reaction (kJ/mol) The ° means it is standard conditions (25°C and 101.3 kPa)

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Practice Questions:

12. The molar enthalpy of combustion of butane, in an open air environment, is -2657.3 kJ/mol. How much energy would be transferred to the surroundings (i.e. what is the enthalpy change) if 20.0 mol of butane were burned?

13. The molar enthalpy of combustion of heptane ( C

7

H

16(l)

) is -4464.7 kJ/mol. What mass of heptane must be burned for an enthalpy change of -6000 kJ.

14. What is the molar enthalpy of reaction for 4.00 moles of substance XY which reacts with substance Z and absorbs 300 kJ of heat energy from the surrounding environment?

15. How much energy is required to decompose 100g of propane?

Chemical Reactions and Bond Energy

Breaking chemical bonds requires energy Endothermic and forming bonds releases energy Exothermic.

When more energy is absorbed to break bonds than is released when forming new bonds, the overall change is Endothermic.

When more energy is released forming new bonds than is required to break the existing bonds, the overall change is Exothermic

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Communicating Enthalpy Changes

There are 4 ways to communicate enthalpy change 1. Molar Enthalpy

2. Enthalpy Change

3. Thermochemical Equations 4. Potential Energy Diagrams

1. Molar Enthalpy (

r

H)

 The molar enthalpy of a substance is the enthalpy change that occurs for 1 mol of that substance.

 Molar enthalpies can be used to communicate exothermic or endothermic changes.

The symbol for molar enthalpy is ΔH (often with a subscript to indicate the type of change).

For example, the molar enthalpy for cell respiration is ΔH°respiration = - 2802.7 kJ/mol

2. Enthalpy Change (H)

• Enthalpy change depends on the number of moles of a reactant and product in a chemical reaction. Ex. the molar enthalpy of reaction on two moles of substance would be double that for one mole.

• Exothermic reactions have a negative H, and endothermic reactions have a positive H.

The formula H= nrH is used to determine enthalpy change.

• Example:

SO2 + ½ O2  SO3 H = -98.9 kJ 2 SO2 + O2  2 SO3 H = __________

3. Thermochemical Equations (energy is a term in the equation)

• If a reaction is exothermic (like the one above), it releases energy, so you can write the amount of energy released as part of your equation on the product side:

SO2(g) + ½ O2(g)  SO3(g) + 98.9 kJ

• Similarly, if a reaction is endothermic, it requires energy , so you can write the amount of energy it needs right into the equation on the reactant side:

H2O(l) + 285.8 kJ  H2(g) + ½ O2(g)

Thermochemical equations are chemical reactions with the energy in them. The units are kJ NOT kJ/mol.

• Energy will either be a product (if it's exothermic) or a reactant (if it's endothermic).

• ΔH in the equation is NOT the same as experimental ΔH Standard Enthalpies of Formation, ∆ f

 We can also find the enthalpy of any given reaction using the Molar Enthalpy of Formation (∆fH°) of each compound.

 The reactions for the formation of many compounds from their elements are recorded already in the data book pages 4 and 5.

These formation reactions are balanced, and record values based on one mole of the compound.

e.g. ½ H2(g) + ½ I2(s)  HI(g) ∆H = +26.5 kJ e.g. 6C(s) + 6H2(g) + 3O2(g)  C6H12O6(s) ∆H = -1273.1kJ The ∆H means the energy absorbed or released (on a per mole basis)

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It makes a difference what state the substances are in, so please record the state!!!

e.g. H2(g) + ½ O2(g)  H2O(l) ∆H =-285.8 kJ versus

H2(g) + ½ O2(g)  H2O(g) ∆H =-241.8 kJ

 Molar enthalpies of formation are given for compounds only, the molar enthalpy of formation for elements are assumed to be zero for this course.

Visually the standard enthalpies of formation look like this:

Important NOTES:

All ∆ fH° of elements are = 0 kJ/mol

 Most values are negative for molar enthalpy of formation.

 In reality, all exothermic reactions have activation energy. This energy requirement prevents a spontaneous reaction from occurring.

Example 1: Write a thermochemical equation for the formation of water vapor.

Progress of Reaction Ep

(kJ)

½ H2(g) + ½ I2(s)

HI(g)

∆H = +26.5 kJ

½ H2(g) + ½ I2(s)  HI(g) ∆ f H° = +26.5 kJ

0 265

Progress of Reaction Ep

(kJ) H2(g) + ½ O2(g)

H2O(l)

∆H =-285.8 kJ

H2(g) + ½ O2(g)  H2O(l) f H° =-285.8 kJ

0

285.8

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Example 2: Write a thermochemical equation for the decomposition of calcium sulfate.

Example 3: Rewrite the equations expressing the enthalpy in ∆H notation for one mole of the underlined substance.

a) 2NH

3(g)

+ 92.2 kJ  N

2(g)

+ 3H

2(g)

b) C

6

H

6(l)

+ 15/2 O

2(g)

 6CO

2(g)

+ 3H

2

O

(g)

+ 3135 kJ

Example 4: Use the reaction to answer the following questions: 4 Al(s) + 3O

2

(g)  2Al

2

O

3

(s) + 3200 kJ a) What is the molar enthalpy of formation of Al

2

O

3

(s)?

b) What is the molar enthalpy of reaction for aluminum?

Example 5: When sucrose is burned in our bodies with excess oxygen, carbon dioxide, water and 5640.3 kJ of energy are produced according to the following equation:

C

12

H

22

O

11

(aq) + 12O

2

(g)  12CO

2

(g) + 11H

2

O(l)

Rewrite the equation by (a) using ∆H notation, (b) placing the energy term within the equation, and (c)

drawing a potential energy diagram. (d) How much energy is released per mole of CO

2

(g)

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Example 6: The molar enthalpy of combustion of ethanol is -1600kJ/mol.

a. Write a thermochemical equation for the reaction.

b. Using the above info, calculate the molar enthalpy of combustion for oxygen.

c. How much energy is produced if 25.0g of oxygen is used?

d. How much energy is produced if 50.0 g of water is produced?

e. What mass of carbon dioxide is produced if 500kJ of energy is released?

Example 7: In making ammonia, what is the molar enthalpy of reaction for hydrogen, if the standard heat of formation for ammonia is -45.9kJ/mol?

Example 8: When 1.00 kg of pure carbon is burned in oxygen to produce carbon dioxide, 3.28 x 10

4

kJ of heat

energy is released. From this data, write a balanced thermochemical equation and sketch a E

p

graph.

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Practice Questions:

16. In each of the following examples the enthalpy change for the reaction is given. Find the molar enthalpy change of reaction for the underlined substance.

(a) 2Al

(s)

+ 3/2O

2(g)

 Al

2

O

3(s)

∆H = -1680kJ

r

H =

(b) 2H

2(g)

+ O

2(g)

 2H

2

O

(g)

∆H = -483.6 kJ

r

H =

(c) 2C

(s)

+ 2H

2

O

(g)

 2CO

(g)

+ 2H

2(g)

∆H = 262.6 kJ

r

H =

(d) 3C

(s)

+ 2Fe

2

O

3(s)

+ 466 kJ  4Fe

(s)

+ 3CO

2(g)

r

H =

(e) 4Mg

(s)

+ 2O

2(g)

 4MgO

(s)

+ 2406 kJ

r

H =

(f) 3X

(s)

+ 4Y

(g)

+ 1200 kJ  X

3

Y

4(s)

r

H =

17. In each of the following the molar enthalpy change is given for a specific substance. Find the enthalpy change of the reaction.

a) ∆

r

H (Fe

(s)

) = 120 kJ/mol

3C

(s)

+ 2Fe

2

O

3(s)

 4Fe

(s)

+ 3CO

2(g)

∆H = ?

b) ∆

r

H (O

2(g)

) = - 1125 kJ/mol

4Al

(s)

+ 3O

2(g)

 2Al

2

O

3(s)

∆H = ?

18. Write out thermochemical equations for the following reactions

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(a) The formation of chromium III oxide. Use ∆H notation.

(b) The decomposition of magnesium carbonate. Include energy in the equation.

(c) The combustion of copper II sulfide. The molar enthalpy of copper II sulfide is -800kJ/mol.

Write this in two ways, energy in the equation and using ∆H notation.

19. Energy, which originates from the sun, is stored in chemical bonds by plants using photosynthesis. The molar enthalpy of reaction for CO

2(g)

during photosynthesis is +467 kJ/mol

Sun’s Energy + 6CO

2(g)

+ 6H

2

O

(l)

 C

6

H

12

O

6(s)

+ 6O

2(g)

a. Describe what it means for CO

2(g)

to have a molar enthalpy of reaction of +467 kJ/mol?

b. What is the enthalpy change for the reaction (kJ)?

20. The energy stored in chemical bonds by plants during photosynthesis can then be accessed by animals, who consume the plants, in the process of cellular respiration. Cellular respiration is the reverse of photosynthesis. Using the information in the previous question write the thermochemical equation for cellular respiration. Is cellular respiration endothermic or exothermic?

21. The molar enthalpy of combustion for ethane (C

2

H

6(g)

), in an open air environment, is -1428.4kJ/mol.

a. Write the Thermochemical equation

b. What is the molar enthalpy of reaction for CO

2(g)

in this equation?

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c. What mass of CO

2(g)

will be produced if 30.0 MJ of heat is released from this reaction?

22. Methane is the main component in natural gas used to heat our homes. The molar enthalpy of combustion for methane in an open air environment is -802.5 kJ/mol.

a. Write the Thermochemical equation

b. If heating the average Edmonton home requires 7.50 GJ of energy, each month, what is the monthly mass of oxygen removed from the environment?

23. Consider the reaction below for the next question:

3C

(s)

+ 2Fe

2

O

3(s)

+ 46.9 kJ  4Fe

(s)

+ 3CO

2(g)

a. How much heat is required to produce 100 g of Fe

(s)

?

b. What mass of iron III oxide could be reduced using 200 kJ of heat?

c. How much heat is required to reduce 50.0 g Fe

2

O

3(s)

?

24. 2C

2

H

2(g)

+ 5O

2(g)

 4CO

2(g)

+ 2H

2

O

(g)

+ 2600 kJ Referring to the above equation, answer the following:

a. What is the enthalpy change for the reaction?

b. What is the molar enthalpy change of the reaction for acetylene?

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c. What is the molar enthalpy change of the reaction for carbon dioxide?

d. If 520 kJ of heat is released, how many grams of acetylene were burned?

e. If 10.4 MJ of heat is released, how many moles of CO

2(g)

were produced?

f. What would be the enthalpy change for the reaction if one mole of acetylene is represented?

25. Given: 2A + 3B + 600 kJ  6C + 5D a. What is the enthalpy change for the reaction?

b. What is the molar enthalpy change of the reaction for B?

c. What is the molar enthalpy change for D?

d. What would be the enthalpy change for 10.0 mol of A?

e. How many moles of C would be produced if 1200 kJ of heat were used?

26. Given: 3X + 4Y  2Q + 5P ∆H = -520 kJ a. Is the above reaction endothermic or exothermic?

b. What is the molar enthalpy change of the reaction for Q?

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c. How much heat would be involved if 1.00 mole of Y reacted?

d. What is the molar enthalpy change for P?

e. How many moles of X would be required to produce 26.0 kJ of heat?

4. Potential Energy (PE) Diagrams

Steps for Drawing Potential Energy Diagrams:

1. Write a thermochemical equation

2. If it is an exothermic reaction, then the reactants are higher in potential energy than the products 3. If it is an endothermic reaction, the products are higher in potential energy than the reactants.

4. Label ΔH, the axes (include units for potential energy), and title the graph.

Example 1: The molar enthalpy of combustion for methanol is -725 kJ/mol. Draw the potential energy diagram.

Remember:

Kinetic Energy: is the energy of motion. It deals with temperature change E

k

= mcΔt

Potential energy: energy of chemical reactions (formation, combustion, etc.) - ΔH = nΔrH

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Example 2: Draw a P.E. diagram for the decomposition of potassium chlorate.

Practice Questions:

27. During a chemical reaction energy is absorbed in order to ____________ chemical bonds, while energy is released when bonds are _____________.

28. During an exothermic reaction the net energy change is ____________(+/-) and energy is released.

During an endothermic reaction the net energy change is__________(+/-) and energy is absorbed.

29. What is the difference between the enthalpy change during a chemical reaction and the molar enthalpy change? What are the units for each?

30. For each of the following reactions decide whether the reaction is endothermic or exothermic then sketch a potential energy diagram.

a. C

3

H

8

(g) + 5O

2

(g)  3CO

2

(g) + 4H

2

O(g) + 2043.9 kJ

b. 2NaCl(s) + 411.2kJ  2Na(s) + Cl

2

(g)

c. A + B  C + D + 200 kJ

d. 2C

6

H

14(l)

+ 19O

2(g)

 12 CO

2(g)

+ 14H

2

O(l) ΔH = -3854.9 kJ

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e. Draw a potential energy for the formation of barium sulfate.

f. Draw a potential energy diagram for the decomposition of sucrose.

Activation Energy and Catalysts

Collision Theory: reactant particles must collide for a reaction to occur, but only a certain portion of the total collisions will result in a chemical reaction.

A collision will only form products if it has:

1. The right orientation at the moment of impact

2. The minimal amount of collision energy needed (called activation energy)

All reactant collisions below activation energy will not result in reactions.

increase temperature -> more collisions -> more reactions

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For the following diagrams we are going to label Activation Energy (E

a

), ∆H, and E

a reverse

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Catalysts

can be used to speed up reactions. e.g. Decomposition of hydrogen peroxide

It takes three years for a bottle of hydrogen peroxide to decompose, but if manganese dioxide (MnO2(s)) is added, the decomposition takes only minutes.

H2O2(l)  H2O(l) + O2(g) ∆H = -145.4kJ (MnO2)

H2O2(l)  H2O(l) + O2(g) ∆H = -145.4 kJ

Catalysts simply reduce the activation energy required to move the reaction forward, and they get us to equilibrium (completion) more quickly, but do not release or absorb any more energy.

Visually an uncatalyzed reaction and a catalyzed reaction look like this:

In the

reverse direction

, this reaction would have a ∆H = +1273.1 kJ Ea = +1393 kJ Ea(catalyzed)= +1333 kJ Progress of Reaction

Ep (kJ)

6C(s) + 6H2(g)+3O2(g)

C6H12O6(s)

∆H =-1273.1kJ Uncatalyzed Reaction (No Enzyme)

Activation Energy, Ea

120kJ

Progress of Reaction Ep

(kJ)

6C(s) + 6H2(g)+3O2(g)

C6H12O6(s)

∆H =-1273.1kJ Catalyzed Reaction (Enzyme)

Activation Energy, Ea

60kJ

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Example 1:

a. What is the Ea (fwd) of the uncatalyzed reaction? Catalyzed reaction?

b. What is the Ea (rev) of the uncatalyzed reaction? Catalyzed reaction?

c. What is the ∆H of the forward reaction?

d. What is the ∆H of the reverse reaction?

Summary of Catalysts:

 speed up reactions

 decrease activation energy

 do not change ∆H

 are NOT used up in reactions

 Provides an alternate pathway

(requiring less energy) for reaction to occur

Common Catalysts to Know

 CFC – catalyze the conversion of O

3

(g) into O

2

(g)

 Sulfuric Acid - in the oil industry

 Nitric acid – in making fertilizer

 Catalytic converters in cars  Convert harmful NO

x

(g) and CO(g) into less harmful N

2

(g) and CO

2

(g)

 Enzymes – in biological processes

Catalyst in a chemical equation: Reactant(s) 

catalyst

Product(s)

Progress of Reaction Ep

(kJ)

6C(s) + 6H2(g)+3O2(g)

C6H12O6(s)

Catalyzed Reaction (Enzyme)

Ea = + 1393.1kJ 120

60 0

∆H = +1273.1kJ

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24

Practice Questions:

31. Describe two criteria that must be met in order for a reaction to take place according to collision theory.

32. How does an increase in temperature affect the rate of a chemical reaction? Explain your answer in terms of activation energy.

33. Describe the energy conversions during a chemical reaction. Use the terms kinetic energy, potential energy, bonds breaking and bonds forming in your answer.

34. Describe what is meant by activated complex.

35. Consider the following reaction:

A

2(g)

+ B

2(g)

 2AB

(g)

The energy of activation (E

a

) in the forward direction is 143 kJ. The energy of activation (Ea) for the reverse direction is 75 kJ.

a. Draw the potential energy diagram for the forward direction. Include ΔH.

b. Is this reaction endothermic or exothermic?

36. Explain why a catalyst will never affect the enthalpy change of a chemical reaction. Use a potential energy diagram in your explanation.

37. How do catalysts speed up a reaction?

38. Are catalysts consumed in a reaction? (Yes or No)

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25 39. The broken line on the energy diagram indicates that:

A. the temperature of the reaction has increased B. the pressure of the system has decreased C. more reactants have been added

D. a catalyst has been added

40. One of the byproducts of the cracking process used at Novacor is ethyne (C

2

H

2(g)

). In presence of a palladium catalyst, the ethyne forms ethene and ethane. This equation is represented by the unbalanced equation

C

2

H

2(g)

+ H

2(g)

 C

2

H

4(g)

+ C

2

H

6(g)

+ energy

The energy diagram that represents both the catalyzed (---) and uncatalyzed reactions ( ) is

41. Given the following potential energy diagram:

Progress of Reaction Ep

(kJ)

Reactants

Products 0

100 180

-400

Give values for:

a. Activation energy=

b. ∆H=

c. Catalyzed Activation Energy=

d. Ea in reverse reaction=

e. ∆H for reverse reaction=

(26)

26 42. Draw and label a potential energy diagram for the formation of ethyne, C

2

H

2

(g), from its elements

and include activation energy and a catalyst path.

43. Draw and label a potential energy diagram for the formation of methane, CH

4

(g), from its elements and include activation energy and a catalyst path.

Use the following information to answer the next question:

How do they get the caramel into the Caramilk bar? Solid sucrose C

12

H

22

O

11(s)

will turn into liquid sugars C

6

H

12

O

6(l)

and C

6

H

10

O

5(l),

but at an extremely slow rate. If a small amount of invertase catalyst is added, the rate is greatly increased.

44. With the invertase (catalyst) present, the ∆H is ___i___ and the forward activation energy is __ii___.

i ii

A. smaller same

B. same smaller

C. smaller smaller

D. same same

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27 45.

Consider the diagram.

a. The potential energy change for this reaction is

A. +170 kJ B. -90 kJ C. -80 kJ D. –170 kJ

b. The energy of activation for this reaction is

c. Draw a line for the catalyzed reaction.

Assorted Practice Questions (Test Your Knowledge So Far)

46. A student designed a calorimetry experiment to determine the mass of water that could be heated by 500kJ of energy released by burning butane gas, and recorded the following results:

Final temperature of water 52.1°C Initial temperature of water 20.6°C Energy released 500 kJ

The mass of water involved in this calorimetry experiment is _______________kg.

47.

Which of the following equations represents the burning of copper?

A. 2Cu

(s)

+ O

2(g)

+ 157.3kJ  2CuO

(s)

B. 2Cu

(s)

+ O

2(g)

 2CuO

(s)

+ 157.3kJ C. Cu

(s)

+ O

12 2(g)

 CuO

(s)

 H=  157.3kJ D. 2Cu

(s)

+ O

2(g)

 2CuO

(s)

 H=  314.6kJ

Burning copper in a camp fire produces a blue-green flame which many people enjoy. This

is not necessarily recommended if you are in an enclosed environment.

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28 Use the graph below for the next TWO questions.

48. The reaction represented by the graph above is _____i______ and the energy term would be included as a ____ii____ in the balanced chemical equation.

The statement above is completed by the information in row

Row i ii

A. exothermic product

B. exothermic reactant

C. endothermic product

D. endothermic reactant

49. Activation energy for the forward, catalyzed reaction is _____i_______ and ∆H is ____ii_____.

The statement above is completed by the information in row

i ii

A. II minus I IV minus I

B. II minus I III minus I

C. III minus IV IV minus I

D. III minus IV III minus I

Use this equation to answer the next 2 questions:

4(g) 2(g) 2(g) 2 ( )

CH + 2O  CO + 2H O g + 805kJ

50. The energy released when 1.95g of methane is burned is _______________ kJ.

51. The mass of water vapor required in order to produce 100kJ is ________________ g.

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29

Calorimetry in Chemical Reactions

(Measuring Energy Changes in the Laboratory is called Calorimetry) When we try to measure energy changes we make two assumptions:

1. Energy is conserved (not lost to the environment)

2. Heat flows from hot to cold until thermal equilibrium is reached.

These assumptions are derived from the first two laws of thermodynamics.

Heat lost = Heat gained

When you place ice in a glass of liquid, energy flows into the ice because it is colder. The energy breaks up the bonds holding the ice in a crystal structure.

We can measure the amount of heat entering the ice indirectly. All we need to do is find the heat lost by the liquid, and assume:

Heat lost by liquid = Heat gained by ice

Water is cheap, and has a high heat capacity, so it is a great liquid to use in calorimetry. Its high heat capacity allows us to keep water as a liquid, and then calculating energy gained or lost by water is only an Ek

= mc∆t calculation.

In the lab, to help these assumptions we use a calorimeter. For phase changes we often use a Styrofoam calorimeter like the one shown below:

1. Styrofoam calorimeter

(AKA - Thermos).

The calorimeter is insulated, and minimizes heat movement with the outer environment.

As well, because it is assumed to be perfectly insulated, the calorimeter does not need to be included in the calculations. Good for aqueous solutions, and reactions.

Ep phase change or reaction = Ek calorimeter

n∆rHº = mc∆T

To calculate the heat of fusion or reactions using a calorimeter, the following data are required:

1) Mass or moles of substance 2) Mass of water in inner container 3) Change in temperature

4) Specific heat of water 5) (If calorimeter is metal)

a) Mass of inner can of calorimeter

b) Specific heat of can substance (usually Al or Cu)

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30 2. Metal Can calorimeter

Usually these calorimeters are made of copper or aluminum (good conductors).

These calorimeters are great for finding heat of a flame, or combustion reactions. You need to include mc∆T of the can as well as the water. The temperature change is the same as the water, because metal is such a good conductor that the heat transfers from the water to the can.

Ep combustion reaction = Ek calorimeter

n∆xH = mc∆Twater + mc∆Tmetal can

3. Bomb Calorimeter (optional)

Reaction occurs in oxygen gas, and the reaction centre is surrounded by liquid water. The calorimeter is sealed and can withstand great pressures resulting from gases produced in exothermic reactions. Used for explosive reactions all heat needs to be transferred to the water, and so final products are CO2(g) and H2O(l).

n∆xH = C∆T

C is the heat capacity of the entire bomb calorimeter.

On diploma examinations, the general principle that will be followed is that if combustion reactions are performed empirically in a bomb calorimeter, liquid water will be the product, and if the combustion occurs in an ambient (open) environment and a theoretical heat of combustion is to be determined, the product will be water vapour.

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31 Why calorimetry works:

Law of conservation of energy – Energy is neither created nor destroyed, only transformed from one form to another. The energy released from a chemical reaction in a calorimeter changes the temperature of the calorimeter (water) surrounding it. Heat lost by a system = heat gained by the surroundings (calorimeter)

Second law of thermodynamics- thermal energy is spontaneously transferred from an object at a higher temperature to an object at a lower temperature until the two objects reach the same temperature (equilibrium)

Therefore enthalpy change is:

heat loss = heat gain - Ep = E

k

loss of Potential Energy = gain of kinetic energy

- n Δ r H = mct

OR

heat gain = heat loss Ep = - E

k

gain of Potential Energy = loss of kinetic energy

n Δ r H = - mct

You must calculate heat on both sides in the same units (we usually use kJ).

Combustion Calorimetry Problems

Example 1: 5.00 grams of butane is burned with excess oxygen from a butane-powered camp stove to

heat 10.0 L of water from 19.5°C to 30.0°C. In these conditions, what is the molar enthalpy change for

the combustion for butane?

(32)

32 Example 2: In a styrene calorimeter, 1.00 g of sulfur, S

8

(s), was burned in oxygen, O

2

(g), to produce sulfur dioxide, SO

2

(g). Using the data given, find the molar heat of reaction for sulfur and add to the balanced equation.

Data Given:

mass of water in calorimeter 100.0 g Initial temperature of water 21.5 °C Final temperature of water 43.6 °C S

8

(s) + 8O

2

(g)  8SO

2

(g) ∆H = ?

Solution: Energy Gained by the _______:

Example 3: Predict the final temperature of an 850 g copper bracelet that is initially at 21.0

o

C and is

heated by burning 3.50 g propane in an open system (∆

c

H of propane is -2043.9 kJ/mol).

Practice Questions:

52. Briefly describe what a calorimeter is.

53. Define the first and second law of thermodynamics and explain how they form the basis for calorimetry.

54. List three assumptions made when using a simple calorimeter

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33 55. Calculate the quantity (in moles) of methane gas that must be burned to increase the

temperature of water in a 120 L hot water tank from 5.00°C to 54.5°C. (The molar enthalpy change of combustion of methane is – 802.7 kJ/mol)

56. A 2.50 g sample of sucrose (C

12

H

22

O

11(s)

) was burned in excess oxygen in a bomb calorimeter which contained 2.19 kg of water. The temperature of the water increased from 20.50°C to 25.01°C.

Determine the molar enthalpy change of combustion for sucrose.

57. Calculate the final temperature of a 500g iron ring (c = 0.449 J/g°C) that is initially at 25.0°C and is heated by combusting 4.95 g of ethanol ( Δ

c

H = -1234.8 kJ/mol), in an open system.

58. How much propane (in grams) would have to be burned in an open system to raise the

temperature of 300 mL of water from 20.00°C to its boiling point (100°C)? The molar enthalpy of

combustion of propane is -2043.9 kJ/mol.

(34)

34 59. A 12.7 g sample of sulfur (S

8(s)

) is placed in a bomb which is then filled with oxygen under

pressure. The bomb is placed in the calorimeter which is filled with 2.20 kg of water at 21.08°C. The reaction mixture is ignited and the temperature of the water rises to a high of 33.88°C. From this data, calculate the molar enthalpy change of combustion of sulfur.

60. In a calorimetry experiment the burning of 5.08 g of benzene (C

6

H

6(l)

) released enough heat to raise the temperature of 5.00 kg of water from 10.0°C to 19.6°C. Calculate the molar enthalpy change of combustion of benzene.

61. A chemist wants to determine empirically the enthalpy change for the following reaction as written:

Mg

(s)

+ 2HCl

(aq)

 MgCl

2(aq)

+ H

2(g)

The chemist uses a simple calorimeter to react 0.50g of magnesium ribbon with 100mL of 1.00 mol/L HCl

(aq)

. The initial temperature of the HCl

(aq)

is 20.40°C. After the reaction the highest recorded temperature is 40.70°C.

a. Calculate the molar enthalpy change of Mg

(s)

b. State all of the assumptions that you made in order to determine the enthalpy change in

A.

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35 62. Sodium reacts violently to form sodium hydroxide when placed in water. Use the following data:

Mass of sodium 0.37g

Mass of water in the calorimeter 175g Initial temperature of the water 19.30°C Final temperature of the water 25.70°C a. Calculate the molar enthalpy of reaction for sodium.

b. Write the thermochemical equation for the reaction.

c. Does all the water react in the calorimeter? Explain your answer.

d. Make a small sketch of the potential energy diagram for this reaction. Remember to

include the enthalpy change.

(36)

36 If two things change temperature, you need to use two mct expressions.

Example 1: A 1.90 g sample of sodium metal was reacted with chlorine in a copper calorimeter. Calculate the molar enthalpy of reaction for sodium given the following information:

Mass of copper container 948 g

Mass of water 265 g

Specific heat capacity of copper

Initial temperature of water and calorimeter

19.0°C

Final temperature of water and calorimeter

23.0°C

Example 2: In an aluminum calorimeter, 20.0 g of nitrogen, N2(g) was burned in oxygen, O2(g) to produce nitrogen monoxide, NO(g). From the following data, find the molar enthalpy of reaction and add to the balanced equation.

Data:

Mass of N2 burned 20.0 g

Mass of inner container, Al 70.37 g

Specific Heat of Al 0.897 J/g°C

Volume of water in inner container 500.0 mL

Initial temperature 72.60 °C

Final temperature 12.80 °C

N2(g) + O2(g)  2NO(g) ∆H = ??

Solution:

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37

Practice Questions :

If the calorimeter is a metal, it will absorb heat too – what does that mean about the math you will do?

63. In an aluminum calorimeter nitrogen was reacted with oxygen under high pressure to produce nitrogen monoxide gas.

a. From the following data, find the molar enthalpy change of reaction for nitrogen under these conditions.

b. Write the thermochemical equation for the reaction.

64. In calorimeter made from tin, 25.0 g of iodine was reacted with excess hydrogen to produce hydrogen iodide gas.

a. From the data below, calculate the molar enthalpy change of the reaction for iodine.

Mass of I

2(s)

25.0 g Mass of tin

calorimeter 70.0 g

Volume of water 100 mL Initial

temperature of calorimeter and water

6.00

o

C

Final temperature of calorimeter and water

23.8

o

C

b. Write the thermochemical equation.

Mass of N

2(g)

reacted 20.0 g Mass of aluminum

calorimeter 70.37 g

Volume of water in

calorimeter 500 mL

Initial temperature of calorimeter and water

72.6

o

C Final temperature of

calorimeter and water

12.80

o

C

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38 KEY

1. Kinetic energy 2. Potential energy 3. Water

4. Copper 5. Propane

65. Determine the enthalpy of combustion of an unknown fuel if a 2.75 g sample increased the temperature of 500 mL (assume 500 g) of hot chocolate (c = 3.75 J/g°C) in a 150 g glass mug (c = 0.84 J/g°C) from 10.00°C to 45.00°C.

a. Express the value for the enthalpy change (in KJ) of combustion for the unknown fuel.

b. What is the enthalpy change per gram for the unknown fuel?

66. A calorimeter is constructed from copper and contains 500 mL of water initially at 21.0°C. If 0.500 g of propane are burned in the calorimeter raising the temperature to 30.0°C, what is the mass of the copper calorimeter? (The molar enthalpy change of combustion of propane is -2043.9 kJ/mol).

67. Use the key to complete the following statements:

Answer: _____ _____ _____ _____

When solving the above problem…

 Water and ________ will require a temperature change calculation (place this number in the first blank below)

 Which substance will require a molar mass calculation? _________

(place this number in the second blank below)

 The propane is involved in what type of energy change? ______ (place this number in the third blank below)

 The copper is involved in what type of energy change? _______ (place

this number in the fourth blank below)

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39 More Energy questions involving non-standard units (kJ/°C, kJ/min, kJ/g)

68. An exercise researcher found that during the course of a three hour walk-run race, the average energy used for the participants was 50.8 kJ/min. How much energy would this person use in the 3 hour walk?

69. The heat capacity of a calorimeter and its contents is 3.60 kJ/°C. If substance X is burned in the calorimeter and the temperature increases by 15.0°C, how much energy is released?

70. Food energy is sometimes measured in units of kJ/g. A lab technician places food samples into a

calorimeter to determine the total energy available. Fats have approximately 37kJ/g available and proteins have approximately 17kJ/g. If 2.5g of protein is burned in the calorimeter, how much energy is available?

71. Calibrating a Calorimeter

Before enthalpies of reaction can be determined, a calorimeter must be calibrated using a primary standard of precisely known molar enthalpy. Complete the Analysis of the investigation report.

Problem

What is the specific heat capacity of a newly assembled calorimeter?

Experimental Design

A calorimeter is assembled and several samples of the primary standard, benzoic acid, are burned using a constant pressure of excess oxygen. The evidence that is collected determines the specific heat capacity of the calorimeter for future experiments.

Evidence

In The CRC Handbook of Chemistry and Physics, the molar enthalpy of combustion for benzoic acid is reported as Δ

c

H = -3231 kJ/mol

Calorimetric evidence for the burning of benzoic acid Mass of calorimeter (g) 9.885 kg

Mass of C

6

H

5

COOH

(s)

(g) 1.024

Initial temperature ºC 24.96

Final temperature ºC 27.99

Analysis (answer the problem)

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40 Calorimetry : Molar Enthalpy (Heat) of solution

Example 1: 4.24g of LiCl is dissolved in 100 mL of water at an initial temperature of 16.3C. The final temperature of the solution is 25.1C. Calculate the molar enthalpy of solution for LiCl.

Thermochemical equation:

72. When 100 g of potassium nitrate are dissolved in 400 mL of water, the water temperature goes down by 20.5

o

C. Calculate the molar heat of solution of potassium nitrate.

73. When 25.0 g of lithium iodide are dissolved in 500 mL of water at 20.0

o

C, the temperature rises to 26.0

o

C. From this information, calculate the molar enthalpy of solution of lithium iodide.

74. (a) If 20.0 g of potassium hydroxide dissolving in 300 mL of water changes the temperature from 30.0

o

C to 46.0

o

C, calculate the molar enthalpy of solution for potassium hydroxide.

(b) Write a thermochemical equation for the above change.

75. If 100 g of potassium hydroxide (  H

solution

  62.0 kJ/mol ) was dissolved in 1.00 L of water at 20.0

o

C, what

would be the final temperature of the solution?

(41)

41 Neutralization Calorimetry (acid and base in solution reacting)

-nΔ

n

H = mcΔt + mcΔt

NOTES:

• Most solutions in this course are (aq) thus water is the largest component. Assume a c value of 4.19 J/g°C

• If a volume of a solution is given, again because it’s mostly water you can assume that 1mL = 1 g (or 1L=1kg)

• You can calculate moles using volume and concentration

• ΔnH is molar enthalpy of neutralization

Example 1: Acid+Base – heat of neutralization

A student mixed 100.0 mL of 1.50 mol/L sulfuric acid with 200.0 mL of 1.50 mol/L sodium hydroxide solution. Both solutions were at 19.7

C initially, and the highest temperature reached was 34.1

C.

a) Calculate the molar heat of neutralization for the sulfuric acid solution.

b) Calculate the molar heat of neutralization for the sodium hydroxide solution

c) Write a thermochemical equation for the reaction.

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42 Practice Problems Involving Chemical Reactions in Solution

In these problems we assume the volume of calorimeter water is the two solutions added together! We also assume the specific heat of the solution is the same as the specific heat of water.

76. Lab Exercise

Problem: What is the molar enthalpy change of reaction of barium hydroxide in the following experiment?

Experimental Design: Barium hydroxide and hydrochloric acid are mixed and the maximum temperature increase is recorded.

Evidence:

Barium hydroxide solution Hydrochloric acid

Volume: 50.0 mL Volume: 200.0 mL

Concentration: 2.00 mol/L Concentration: 1.00 mol/L Initial Temperature: 21.0°C Initial Temperature: 21.0°C Final Temperature of the mixture: 32.5°C

a. Using data from above, calculate the molar enthalpy of neutralization for hydrochloric acid.

b. Write the thermochemical equation for the reaction.

c. Use your thermochemical equation to determine the molar enthalpy of neutralization for barium hydroxide.

77. In order to neutralize 50.0 mL of 0.400 mol/L KOH(aq) in a simple calorimeter, 15.9mL of a 0.630 mol/L H

2

SO

4

(aq) is added. The initial temperature of the acid and the base before they are combined is 25.0°C.

Once the reaction has come to completion the highest temperature reached is 28°C.

a. Calculate the molar enthalpy of neutralization for the potassium hydroxide.

b. Write the thermochemical equation

c. Use the equation to calculate the molar enthalpy of neutralization for the sulfuric acid.

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43

Life and Thermochemical Changes

Cellular Respiration is very similar to the combustion reaction of glucose. It is an exothermic reaction that produces energy from the reaction of glucose with oxygen.

• Reaction: _________________________________________________________

Photosynthesis is an endothermic reaction to produce the sugar needed to burn in cellular respiration. The energy for the reaction comes from sunlight.

• Reaction: _________________________________________________________

** The original source of energy in fossil fuels is the sun!

Pick the sun!**

78. When methane gas is burned in a fireplace, the reaction that occurs is ____i____. The original source of energy stored in the methane gas is ____ii____.

The statements above are completed by the information in row

Row i ii

a. endothermic fossil fuel

b. endothermic the sun

c. exothermic fossil fuel

d. exothermic the sun

Hess’s Law (Additive Principle of Reaction Heats)

Hess’s Law lets us calculate the energy released or required by a reaction, simply by adding up the energy of formation for each reactant and compound.

More Formally, Hess’s Law states:

If a chemical reaction can be expressed as the algebraic sum of two or more reactions, then its enthalpy of reaction (∆H rxn)is also the algebraic sum of the separate enthalpy of reactions (∆Hf).

If you can add the equations, you can add the ∆H’s We need to follow some basic rules to make Hess’s law work:

1. Write a balanced equation for the reaction for which ∆H is needed (this is the original equation).

2. Using the heats of formation table, write out the balanced equation to form one mole of each product.

3. Using the heats of formation table, write out the balanced equation to decompose one mole of each reactant (change the sign of the ∆H!).

4. Multiply any of the equations so they match up with the moles of that compound in the original equation (remember to multiply the ∆H as well!).

5. Ignore elements on either side (they will sort themselves out by the end).

6. Add the equations, and they should add up to the original equation.

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44

Or simply…Stay, Flip, Multiple or Omit given reactions to produce the net reaction.

Example 1: Find the enthalpy of combustion for carbon monoxide, CO(g), burned in excess oxygen, O2(g) to produce carbon dioxide CO2(g).

CO(g) + ½ O2(g)  CO2(g) ∆H = ??

Given Equations:

C(s) + O2(g)  CO2(g) ∆H = -393.5 kJ C(s) + ½ O2(g)  CO(g) ∆H = -110.5 kJ

Example 2: Calculate H for the following:

CO

( )g

H

2( )g

O

2( )g

CO

2( )g

H O

2 ( )g , using the following equations:

( ) 2( ) ( )

2 C

s

O

g

 2 CO

g

   H 221.0 kJ

( )s 2( )g 2( )g

393.5 COCO    H kJ

2( ) 2( ) 2 ( )

2 H

g

O

g

 2 H O

g

   H 483.6 kJ

Example 3: Calculate the molar heat of formation for ammonia, NH3(g), using the reactions given.

Given:

2NH3(g) + 3/2 O2(g)  N2(g) + 3H2O(g) ∆H = -633.2kJ H2(g) + 1/2 O2(g)  H2O(g) ∆H = -241.8 kJ

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45

Practice Questions:

79. The standard molar enthalpy change of combustion of liquid ethyl alcohol is -1370 kJ/mol and that of acetic acid is -875.7 kJ/mol. The equations are:

C

2

H

5

OH

(l)

+ 3 O

2(g)

 2 CO

2(g)

+ 3 H

2

O

(l)

∆H = -1370.1 kJ HC

2

H

3

O

2(l)

+ 2 O

2(g)

 2 CO

2(g)

+ 2 H

2

O

(l)

∆H = -875.7 kJ What is the heat of reaction for the oxidation of ethyl alcohol to acetic acid?

C

2

H

5

OH

(l)

+ O

2(g)

 HC

2

H

3

O

2(l)

+ H

2

O

(l)

∆H = ?

80. Given: 1/2 N

2(g)

+ 1/2 O

2(g)

 NO

(g)

∆H = 90.5 kJ 1/2 N

2(g)

+ O

2(g)

 NO

2(g)

∆H = 33.9 kJ Calculate ∆H for:

NO

(g)

+ 1/2 O

2(g)

 NO

2(g)

∆H = ?

81. Use the following information to determine the molar enthalpy of combustion for ammonia.

NH

3(g)

 1/2 N

2(g)

+ 3/2 H

2(g)

∆H = 46.1 kJ 1/2 N

2(g)

+ O

2(g)

 NO

2(g)

∆H = 33.9 kJ H

2(g)

+ 1/2 O

2(g)

 H

2

O

(g)

∆H = -242.0 kJ

82. Find the heat of formation of one mole of SO

3

gas from the following data:

S

8(s)

+ 8O

2(g)

 8 SO

2(g)

+ 2378.4 kJ

SO

2(g)

+ 1/2 O

2(g)

 SO

3(g)

+ 99.0 kJ

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46 83. Given:

2 Y

(s)

+ 3/2 O

2(g)

 Y

2

O

3(s)

+ 2095 kJ 2 M

(s)

+ 3/2 O

2(g)

 M

2

O

3(s)

+ 419 kJ

The heat of reaction for:

2 Y

(s)

+ M

2

O

3(s)

 2 M

(s)

+ Y

2

O

3(s)

is:

84. Given:

N

2(g)

+ 5/2 O

2(g)

+ 15 kJ  N

2

O

5(g)

H

2(g)

+ 1/2 O

2(g)

 H

2

O

(l)

+ 286 kJ

1/2 H

2(g)

+ 1/2 N

2(g)

+ 3/2 O

2(g)

 HNO

3(l)

+ 173 kJ The heat of the reaction for:

N

2

O

5(g)

+ H

2

O

(l

 2 HNO

3(l)

is:

85. Given:

H

2

O

(l)

 H

2(g)

+ 1/2 O

2(g)

∆H = 286 kJ H

2

O

(g)

 H

2(g)

+ 1/2 O

2(g)

∆H = 242 kJ Find:

H

2

O

(l)

 H

2

O

(g)

∆H = ?

References

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