For the combustion reaction of methane, ΔH f is zero for

The content that follows is the substance of General Chemistry Lecture 23. In this lecture we further discuss Enthalpy and introduce its calculation using Heats of Formation and Hess's Law.

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More on Enthalpy
Enthalpy Values
Standard Enthalpy of Reaction

More on Enthalpy

As we defined it in the previous lecture, Enthalpy is a measure of the heat gained or lost by a system at constant pressure. It is also a state function, meaning its value is only concerned with the current status. The change in Enthalpy is therefore determined by the starting and ending amounts of heat and it does not care how the process was conducted in between these two points. We say that the value is "independent of path"

Further Properties of Enthalpy:

1.Enthalpy is an extensive property. The magnitude of ΔH is dependent upon the amounts of reactants consumed. Doubling the reactants, doubles the amount of enthalpy.

2.Reversing a chemical reaction results in the same magnitude of enthalpy but of the opposite sign.

3.The enthalpy change for a reaction depends upon the state of the reactants and products. The states (i.e. g, l, s or aq) must be specified.

Example Problem using an Enthalpy value:

CH4(g) + 2O2(g) -> CO2(g) + 2H2O(g) ΔH = -802 kJ

Given the above thermochemical equation for the combustion of methane, how much heat energy is released when 4.5 grams of methane is burned (in a constant pressure system)?

First determine the moles of methane: 4.5 g x 1 mole/16 g methane = 0.28125 mol CH4

Then multiply the amount of moles by the known per mole amount of Enthalpy shown: 0.28125 * -802 kJ = -225.56 kJ or -2.3e2 kJ

You may note that the units on the Enthalpy value are only shown as kJ and not kJ/mol in the reaction. This lack of the per mole unit is fairly common but the per mole unit is understood to be there even if it is not written, somewhat like a 1 coefficient in a balanced chemical reaction. Looking at the reaction above you can see that the enthalpy value for the reaction is for 1 mole of Methane or 2 moles of Oxygen or 1 mole of Carbon Dioxide etc.

Enthalpy Values

The Enthalpy values for many substances have already been determined experimentally and are readily available in tables of physical constants. The values are generally taken at what is called "standard state". This is the most stable state of a substance at 1 atm pressure and at a specified temperature, usually 25oC and 1M concentration for all substances in solution.

The thermodynamic standard state is defined so that different scientists can compare results

Standard Enthalpy of Reaction - ΔHoRxn - enthalpy change under standard conditions of 1 atm and 298.15 K.

Be careful not to confuse "Standard State" with STP as these are two different conditions.

There are many other types of Enthalpies as well:

Enthalpies of physical change

ΔHfus - enthalpy of fusion - amount of heat required to change a solid to a liquid
ΔHvap - enthalpy of vaporization - amount of heat required to change a liquid into a gas
ΔHsub - enthalpy of sublimation - heat required to change a solid into a gas

One Enthalpy of particular use is the Enthalpy of Formation. The Standard enthalpy of formation (ΔHoF) is the heat change that results when one mole of a compound is formed from its elements (in most stable form/natural) at a pressure of 1 atm. The standard enthalpy of formation of any element in its most stable form is zero.

Notice that the elements in their most stable or natural elemental form have a ΔHoF of zero while those forms that are not stable or require a process to form have a ΔHoF value.

Standard Enthalpy of Reaction

The standard enthalpy of reaction (ΔHoRxn) is the enthalpy of a reaction carried out at 1 atm. We have already learned one process by which we can calculate the Enthalpy of Reaction in Calorimetry. There are two other methods we will learn now:

1) Heat of Reaction from Standard Heats of Formation

2) Heat of Reaction from Hess' Law Calculation

Calculation of the Heat of Reaction from Standard Heats of Formation is based on the following equations:

The Heat of Reaction can be calculated from the Heats of Formation of each molecule in the reaction. The equation shown above shows that the ΔHoRxn value is calculated as the sum of the moles of the products times their ΔHoF values minus the sum of the the moles of the reactants times their ΔHoF values.

Hess’s Law: When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps. Remember that Enthalpy is a state function. It doesn’t matter how you get there, only where you start and end.

Here is an example of how to complete a calculation of this type:

The third process by which you can calculate the Heat of a Reaction is by using a process called Hess' Law. Another way to state Hess' Law is The enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps.

The process needed to answer the question above is based on the principle that if you add two or more equations to get a new equation, you must add the ΔH’s to get the ΔH for the new equation.

A couple of rules before we start:

1) If you multiply an equation by a value to get the number of moles to match the reaction needed, you have to multiply the ΔH by the same value.

2) If you turn a reaction around to get a molecule on the correct side to match the reaction needed, you change the sign of the ΔH value for that reaction.

The "target" reaction is S (s) + 3/2 O2 (g) → SO3 (g) ΔH = ?

In the two reactions we are given, the Sulfur is located in Equation 1 and is in the right form so we just copy that equation in as is. The second equation contains the Sulfur Trioxide we need but in the wrong amount. Since there are two in the equation we will have to divide the entire equation by two to get it into the correct form remembering to cut the enthalpy value in half as well:

S (s) + O2 (g) → SO2 (g) H1 = -296.0 kJ

1/2(2 SO2 (g) + O2 (g) → 2SO3 (g) H2 = -198.2 kJ)

After division the two equations are ready to be added together. We can cancel anything that appears on both sides of the equation in equal amounts:

S (s) + O2 (g) → SO2 (g) H1 = -296.0 kJ

SO2 (g) +1/2 O2 (g) → SO3 (g) H2 = -99.1 kJ

S (s) + 3/2 O2 (g) → SO3 (g) ΔH = -395.1 kJ

Here is a website with lots of practice quizzes on all three types of enthalpy calculations: Here

In Conclusion, at this point we have learned 3 ways to calculate the Heat of a Reaction:

In about a month, once we have learned how to draw Lewis structures you will learn a 4th method using bond energy values.

Objects can possess energy as ________.

(a) endothermic energy (b) potential energy

A) a onlyB) b onlyC) c onlyD) a and c

E) b and c

The internal energy of a system is always increased by ________.A) adding heat to the systemB) having the system do work on the surroundingsC) withdrawing heat from the systemD) adding heat to the system and having the system do work on the surroundings

E) a volume decompression

The internal energy of a system ________.A) is the sum of the kinetic energy of all of its componentsB) is the sum of the rotational, vibrational, and translational energies of all of its componentsC) refers only to the energies of the nuclei of the atoms of the component moleculesD) is the sum of the potential and kinetic energies of the components

E) none of the above

Which one of the following conditions would always result in an increase in the internal energy of a system?A) The system loses heat and does work on the surroundings.B) The system gains heat and does work on the surroundings.C) The system loses heat and has work done on it by the surroundings.D) The system gains heat and has work done on it by the surroundings.

E) None of the above is correct.

When a system ________, ΔE is always negative.A) absorbs heat and does workB) gives off heat and does workC) absorbs heat and has work done on itD) gives off heat and has work done on it

E) None of the above is always negative.

Which one of the following is an endothermic process?A) ice meltingB) water freezingC) boiling soupD) Hydrochloric acid and barium hydroxide are mixed at 25 °C: the temperature increases.

E) Both A and C

Which one of the following is an exothermic process?A) ice meltingB) water evaporatingC) boiling soupD) condensation of water vapor

E) Ammonium thiocyanate and barium hydroxide are mixed at 25 °C: the temperature drops.

Of the following, which one is a state function?A) HB) qC) wD) heat

E) none of the above

Which of the following is a statement of the first law of thermodynamics?
A) Ek = 1/2mv2B) A negative ΔH corresponds to an exothermic process.

C) ΔE = Efinal - Einitial

D) Energy lost by the system must be gained by the surroundings.

E) 1 cal = 4.184 J (exactly)

The internal energy can be increased by ________.

(a) transferring heat from the surroundings to the system (b) transferring heat from the system to the surroundings

A) a onlyB) b onlyC) c onlyD) a and c

E) b and c

A ________ ΔH corresponds to an ________ process.A) negative, endothermicB) negative, exothermicC) positive, exothermicD) zero, exothermic

E) zero, endothermic

A ________ ΔH corresponds to an ________ process.A) negative, endothermicB) positive, exothermicC) positive, endothermicD) zero, exothermic

E) zero, endothermic

ΔH for an endothermic process is ________ while ΔH for an exothermic process is ________.A) zero, positiveB) zero, negativeC) positive, zeroD) negative, positive

E) positive, negative

For a given process at constant pressure, w is positive. This means that the process involves ________.A) work being done by the system on the surroundingsB) work being done by the surroundings on the systemC) no work being doneD) an equal amount of work done on the system and by the system

E) work being done against a vacuum

Which one of the following statements is true?A) Enthalpy is an intensive property.B) The enthalpy change for a reaction is independent of the state of the reactants and products.C) Enthalpy is a state function.D) H is the value of q measured under conditions of constant volume.

E) The enthalpy change of a reaction is the reciprocal of the ΔH of the reverse reaction.

All of the following statements are true except ________.A) Internal energy is a state function.B) The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to the enthalpy change for the reverse reaction.C) The enthalpy change for a reaction depends on the state of the reactants and products.D) The enthalpy of a reaction is equal to the heat of the reaction.

E) Enthalpy is an intensive property.

A chemical reaction that absorbs heat from the surroundings is said to be ________ and has a ________ ΔH at constant pressure.A) endothermic, positiveB) endothermic, negativeC) exothermic, negativeD) exothermic, positive

E) exothermic, neutral

A chemical reaction that releases heat to the surroundings is said to be ________ and has a ________ ΔH at constant pressure.A) endothermic, positiveB) endothermic, negativeC) exothermic, negativeD) exothermic, positive

E) exothermic, neutral

The reaction

4Al (s) + 3O2 (g) → 2 Al2O3 (s) ΔH° = -3351 kJis ________, and therefore heat is ________ by the reaction.A) endothermic, releasedB) endothermic, absorbedC) exothermic, releasedD) exothermic, absorbed

E) thermoneutral, neither released nor absorbed

When ________ is constant, the enthalpy change of a process equal to the amount of heat transferred into or out of the system?A) temperatureB) volumeC) pressure and volumeD) temperature and volume

E) pressure

The units of heat capacity are ________.A) K/J or °C/JB) J/K or J/°CC) J/g-K or J/g-°CD) J/mol

E) g-K/J or g-°C/J

The units of specific heat are ________.A) K/J or °C/JB) J/K or J/°CC) J/g-K or J/g-°CD) J/mol

E) g-K/J or g-°C/J

The term Btu which stands for ________ is commonly used in engineering applications.A) Best thermal unitB) Bunsen thermal unitC) British thermal unitD) Bake thermal unit

E) Brush thermal unit

Which of the following is a statement of Hess's law?A) If a reaction is carried out in a series of steps, the ΔH for the reaction will equal the sum of the enthalpy changes for the individual steps.B) If a reaction is carried out in a series of steps, the ΔH for the reaction will equal the product of the enthalpy changes for the individual steps.C) The ΔH for a process in the forward direction is equal in magnitude and opposite in sign to the ΔH for the process in the reverse direction.D) The ΔH for a process in the forward direction is equal to the ΔH for the process in the reverse direction.

E) The ΔH of a reaction depends on the physical states of the reactants and products.

For which one of the following reactions is ΔH°rxn equal to the heat of formation of the product?A) N2 (g) + 3H2 (g) → 2NH3 (g)B) (1/2)N2 (g) + O2 (g) → NO2(g)C) 6C (s) + 6H (g) → C6H6 (l)D) P (g) + 4H (g) + Br (g) → PH4Br (l)

E) 12C (g) + 11H2 (g) + 11O (g) → C6H22O11 (g)

Of the following, ΔH°f is not zero for ________.A) O2 (g)B) C (graphite)C) N2 (g)D) F2 (s)

E) Cl2 (g)

Of the following, ΔH°f is not zero for ________.A) Sc (g)B) Si (s)C) P4 (s, white)D) Br2 (l)

E) Ca (s)

Consider the following two reactions:

A → 2B ΔH°rxn = 456.7 kJ/mol
A → C ΔH°rxn = -22.1 kJ/mol

Determine the enthalpy change for the process:

2B → CA) -478.8 kJ/molB) -434.6 kJ/molC) 434.6 kJ/molD) 478.8 kJ/mol

E) More information is needed to solve the problem.

In the reaction below, ΔH°f is zero for ________.
Ni (s) + 2CO (g) + 2PF3 (g) → Ni(CO)2(PF3)2 (l)

A) Ni (s)B) CO (g)C) PF3 (g)D) Ni(CO)2(PF3)2 (l)

E) both CO (g) and PF3 (g)

For the combustion reaction of methane, ΔH°f is zero for ________.
CH4 (g) + O2 (g) → 2H2O(g) + CO2 (g)

A) O2 (g)B) CH4 (g)C) CO2 (g)D) H2O (g)

E) Both O2 (g) and CH4 (g)

For the following reactions, the ΔH°rxn is NOT equal to ΔH°f for the product except for ________.A) CH4 (g) + 2Cl2 (g) → CH2Cl2 (l) + 2HCl (g)B) N2 (g) + O3 (g) → N2O3 (g)C) Xe (g) + 2F2 (g) → XeF4 (g)D) 2CO (g) + O2 (g) → 2CO2 (g)

E) C (diamond) + O2 (g) → CO2 (g)

For the following reactions, the ΔH°rxn is NOT equal to ΔH°f for the product except for ________.A) 2Ca (s) + O2 (g) → 2CaO (s)B) 3Mg (s) + N2 (g) → Mg3N2 (s)C) C2H2 (g) + H2 (g) → H4 (g)D) 2C (graphite) + O2 (g) → 2CO (g)

E) C (diamond) + O2 (g) → CO2 (g)

For the following reactions, the ΔH°rxn is NOT equal to ΔH°f for the product except for ________.A) N2 (g) + O2 (g) → 2NO (g)B) 2H2 (g) + O2 (g) → 2H2O (l)C) 2H2 (g) + O2 (g) → 2H2O (g)D) 2C (s, graphite) + 2H2(g) → C2H4 (g)

E) H2O (l) + 1/2 O2 (g) → H2O2 (l)

For the following reactions, the ΔH°rxn is NOT equal to ΔH°f for the product except for ________.A) H2O (l) + 1/2O2 (g) → H2O2 (l)B) N2 (g) + O2 (g) → 2NO (g)C) 2H2 (g) + O2 (g) → 2H2O (l)D) 2H2 (g) + O2 (g) → 2H2O (g)

E) none of the above

For the following reactions, the ΔH°rxn is NOT equal to ΔH°f for the product except for ________.A) H2 (g) + 1/2 O2 (g) → H2O (l)B) H2 (g) + O2 (g) → H2O2 (l)C) 2C (s, graphite) + 2H2 (g) → C2H4 (g)D) 1/2 N2 (g) + O2 (g) → NO2 (g)

E) all of the above

The term standard conditions with respect to enthalpy change means ________.A) 1 atm and 0 KB) 1 atm and 1 LC) 1 atm and 298 KD) 1 L and 0 K

E) 1 atm and 1 °C

The energy released by combustion of ________ of a substance is called the fuel value of the substance.A) 1 kJB) 1 kgC) 1 lbD) 1 J

E) 1 g

Fuel values of hydrocarbons increase as the ________ increases.A) C atomic ratioB) H/C atomic ratioC) H atomic ratioD) C/C atomic ratio

E) C/H atomic ratio

________ yields the highest fuel value.A) hydrogenB) charcoalC) bituminous coalD) natural gas

E) wood

All of the following are considered fossil fuels except ________.A) hydrogenB) anthracite coalC) crude oilD) natural gas

E) petroleum

The most abundant fossil fuel is ________.A) natural gasB) petroleumC) coalD) uranium

E) hydrogen

Calculate the kinetic energy in J of an electron moving at 6.00 × 106 m/s. The mass of an electron is 9.11 × 10-28 g.
A) 4.98 × 10-48 J
B) 3.28 × 10-14 J
C) 1.64 × 10-17 J
D) 2.49 × 10-48 J
E) 6.56 × 10-14 J

Calculate the kinetic energy in joules of an automobile weighing 2135 lb and traveling at 55 mph.
A) 1.2 × 104 J
B) 2.9 × 105 J
C) 5.9 × 105 J
D) 3.2 × 106 J
E) 3.2 × 10-6 J

Calculate the kinetic energy in joules of an automobile weighing 4345 lb and traveling at 75 mph.
A) 5.5 × 105 J
B) 5.5 × 10-5 J
C) 1.1 × 106 J
D) 2.2 × 106 J
E) 2.2 × 10-6 J

The kinetic energy of a 7.3 kg steel ball traveling at 18.0 m/s is ________ J.
A) 1.2 × 103 B) 66

C) 2.4 × 103

D) 1.3 × 102
E) 7.3

The kinetic energy of a 10.3 g golf ball traveling at 48.0 m/s is ________ J.
A) 1.20 × 103 B) 66C) 11.9

D) 1.3 × 102

E) 23.7

Calculate the kinetic energy in joules of a 150 lb jogger (68.1 kg) traveling at 12.0 mile/hr (5.36 m/s).
A) 1.96 × 103 JB) 365 JC) 978 JD) 183 J

E) 68.1 J

What is the kinetic energy of a 55.2 g object moving at 135 m/s.A) 503 J

B) 5.03 × 105 J

C) 1.01 × 103 JD) 3.73 J

E) 3.73 × 103 J

The kinetic energy of a 23.2-g object moving at a speed of 81.9 km/hr is ________ J.A) 1900B) 77.8C) 145

D) 1.43 × 10-3

E) 6.00

What is the kinetic energy of a 145 g baseball traveling at 89.9 mi/hr? A) 1.17 × 105 JB) 2.91 JC) 234 JD) 5.83

E) 117 J

A 100-watt electric incandescent light bulb consumes ________ J of energy in 24 hours. [1 Watt (W) = 1 J/sec]
A) 2.40 × 103
B) 8.64 × 103 C) 4.17

D) 2.10 × 103

E) 8.64 × 106

The ΔE of a system that releases 12.4 J of heat and does 4.2 J of work on the surroundings is ________ J.A) 16.6B) 12.4C) 4.2D) -16.6

E) -8.2

The ΔE of a system that absorbs 12.4 J of heat and does 4.2 J of work on the surroundings is ________ J.A) 16.6B) 12.4C) 4.2D) -16.6

E) 8.2

The change in the internal energy of a system that absorbs 2,500 J of heat and that does 7,655 J of work on the surroundings is ________ J.A) 10,155B) 5,155C) -5,155D) -10,155

E) 1.91 × 107

The change in the internal energy of a system that releases 2,500 J of heat and that does 7,655 J of work on the surroundings is ________ J.A) -10,155B) -5,155

C) -1.91 × 107

D) 10,155

E) 5,155

Hydrogen gas and bromine gas react to form hydrogen bromide gas. How much heat (kJ) is released when 155 grams of HBr is formed in this reaction? ΔH° = -72 kJ. A) 137B) 69C) -69D) -137

E) 1.12 × 105

The value of ΔH° for the reaction below is -126 kJ. ________ kj are released when 2.00 mol of NaOH is formed in the reaction?

2Na2O2 (s) + 2H2O (l) → 4NaOH (s) + O2 (g)

A) 252B) 63C) 3.9D) 7.8

E) -126

The value of ΔH° for the reaction below is -790 kJ. The enthalpy change accompanying the reaction of 0.95 g of S is ________ kJ.
2S (s) + 3O2 (g) → 2SO3 (g)

A) 23B) -23C) -12D) 12

E) -790

The value of ΔH° for the reaction below is -6535 kJ. ________ kJ of heat are released in the combustion of 16.0 g of (l)?
2C6H6 (l) + 15O2 (g) → 12CO2 (g) + 6H2O (l)

A) 1.34 × 103
B) 5.23 × 104 C) 669

D) 2.68 × 103

E) -6535

Carbon monoxide and oxygen gas react to form carbon dioxide. How much heat is released when 89.5 grams of O2 (g) reacts with excess CO? ΔH° = -482 kJ.
A) 1.35 × 103 kJ
B) 2.70 × 103 kJ
C) 1.35 × 10-3 kJD) 674 kJ

E) 4.31 × 104 kJ

The value of ΔH° for the reaction below is -336 kJ. Calculate the heat (kJ) released to the surroundings when 23.0 g of HCl is formed.

CH4 (g) + 3Cl2 (g) → CHCl3 (l) + 3HCl (g)

A) 177 kJ
B) 2.57 × 103 kJC) 70.7 kJD) 211 kJ

E) -336 kJ

How much heat is released when 29.5 grams of Cl2 (g) reacts with excess hydrogen?

H2 (g) + Cl2 (g) → 2HCl (g) ΔH° = -186 kJ.

A) 186 kJB) 310 kJC) -77.4 kJD) -186 kJ

E) 77.4 kJ

The enthalpy change for the following reaction is -483.6 kJ:

2H2 (g) + O2 (g) → 2H2O (g)

Therefore, the enthalpy change for the following reaction is ________ kJ. 4H2 (g) + 2O2 (g) → 4H2O (g)A) -483.6B) -967.2

C) 2.34 × 105

D) 483.6

E) 967.2

The value of ΔH° for the reaction below is +128.1 kJ:

CH3OH (l) → CO (g) + 2H2 (g)How many kJ of heat are consumed when 15.5 g of C OH (l) decomposes as shown in the equation?A) 0.48 kJB) 62.0 kJ

C) 1.3 × 102 kJ

D) 32 kJ

E) 8.3 kJ

The value of ΔH° for the reaction below is +128.1 kJ:

CH3OH (l) → CO (g) + 2H2 (g)How much heat is consumed when 87.1 g of hydrogen gas is formed?

A) 2.76 × 103 kJ

B) 5.52 × 103 kJ
C) -5.52 × 103 kJ
D) -2.76 × 103 kJ
E) -128.1 kJ

The value of ΔH° for the reaction below is +128.1 kJ:

CH3OH (l) → CO (g) + 2H2 (g)How many kJ of heat are consumed when 5.10 g of CO (g) is formed as shown in the equation?A) 0.182 kJB) 162 kJC) 8.31 kJD) 23.3 kJ

E) 62.0 kJ

CH3OH (l) decomposes into carbon monoxide and hydrogen gas in the presence of heat. How much heat is consumed when 5.75 g of CO (g) is formed? ΔH° = +128.1 kJ.A) 26.3 kJB) 23.3 kJC) 62.0 kJD) 162 kJ

E) 8.3 kJ

The value of ΔH° for the reaction below is -1107 kJ:

2Ba (s) + O2 (g) → 2BaO (s)How many kJ of heat are released when 5.75 g of Ba (s) reacts completely with oxygen to form A) 96.3 kJB) 26.3 kJC) 46.4 kJD) 23.2 kJ

E) 193 kJ

The value of ΔH° for the reaction below is -1107 kJ:

2Ba (s) + O2 (g) → 2BaO (s)How many kJ of heat are released when 5.75 g of BaO (s) is produced?A) 56.9 kJB) 23.2 kJC) 20.8 kJD) 193 kJ

E) 96.3 kJ

How many kJ of heat are released when 15.75 g of Ba (s) reacts completely with oxygen gas to form BaO (s)? ΔH° = -1107 kJ.A) 63.5 kJB) 20.8 kJC) 114 kJD) 70.3 kJ

E) 35.1 kJ

The molar heat capacity of an unknown substance is 92.1 J/mol-K. If the unknown has a molar mass of 118 g/mol, what is the specific heat (J/g-K) of this substance?A) 1.28B) -92.1

C) 1.09 × 104

D) 0.781

E) 92.1

The specific heat capacity of lead is 0.13 J/g-K. How much heat (in J) is required to raise the temperature of of lead from 22 °C to 37 °C?A) 2.0 JB) -0.13 J

C) 5.8 × 10-4 J

D) 29 J

E) 0.13 J

The temperature of a 15-g sample of lead metal increases from 22 °C to 37 °C upon the addition of 29.0 J of heat. The specific heat capacity of the lead is ________ J/g-K.A) 7.8B) 1.9C) 29D) 0.13

E) -29

What is the molar heat capacity (in J/mol-K) of liquid bromine? The specific heat of liquid bromine is 0.226 J/g-K.A) 36.1 J/mol-KB) 707 J/mol-KC) 18.1 J/mol-KD) 9.05 J/mol-K

E) 0.226 J/mol-K

The specific heat of liquid bromine is 0.226 J/g-K. How much heat (J) is required to raise the temperature of 10.0 mL of bromine from 25.00 °C to 27.30 °C? The density of liquid bromine: 3.12 g/mL.A) 5.20 JB) 16.2 JC) 300 JD) 32.4 J

E) 10.4 J

ΔH for the reaction

IF5 (g) → IF3 (g) + F2 (g)

is ________ kJ, given the data below.
IF (g) + F2 (g) → IF3 (g) ΔH = -390 kJ

IF (g) + 2F2 (g) → IF5 (g) ΔH = -745 kJ

A) +355B) -1135C) +1135D) +35

E) -35

Given the following reactions
Fe2O3 (s) + 3CO (s) → 2Fe (s) + 3CO2 (g) ΔH = -28.0 kJ

3Fe (s) + 4CO2(s) → 4CO (g) + Fe3O4(s) ΔH = +12.5 kJ

the enthalpy of the reaction of Fe2O3 with CO
3Fe2O3 (s) + CO (g) → CO2 (g) + 2Fe3O4 (s)

is ________ kJ.A) -59.0B) 40.5C) -15.5D) -109

E) +109

Given the following reactions
N2 (g) + 2O2 (g) → 2NO2 (g) ΔH = 66.4 kJ

2NO (g) + O2 (g) → 2NO2 (g) ΔH = -114.2 kJthe enthalpy of the reaction of the nitrogen to produce nitric oxide

N2 (g) + O2 (g) → 2NO (g)

is ________ kJ.A) 180.6B) -47.8C) 47.8D) 90.3

E) -180.6

Given the following reactions
2NO → N2 + O2 ΔH = -180 kJ

2NO + O2 → 2NO2 ΔH = -112 kJ

the enthalpy of the reaction of nitrogen with oxygen to produce nitrogen dioxide

N2 + 2O2 → 2NO2

is ________ kJ.A) 68B) -68C) -292D) 292

E) -146

Given the following reactions
2S (s) + 3O2 (g) → 2SO3 (g) ΔH = -790 kJ

S (s) + O2 (g) → SO2(g) ΔH = -297 kJ

the enthalpy of the reaction in which sulfur dioxide is oxidized to sulfur trioxide

2SO2 (g) + O2 (g) → 2SO3 (g)

is ________ kJ.A) 196B) -196C) 1087D) -1384

E) -543

Given the following reactions CaCO3 (s) → CaO (s) + CO2 (g) ΔH = 178.1 kJ C (s, graphite) + O2(g) → CO2(g) ΔH = -393.5 kJthe enthalpy of the reaction

CaCO3 (s) → CaO (s) + C (s, graphite) + O2 (g)

is ________ kJ.A) 215.4B) 571.6C) -215.4D) -571.6

E) 7.01 × 104

Given the following reactions H2O (l) → H2O (g) ΔH = 44.01 kJ 2H2 (g) + O2 (g) → 2H2O (g) ΔH = -483.64 kJthe enthalpy for the decomposition of liquid water into gaseous hydrogen and oxygen

2H2O (l) → 2H2 (g) + O2 (g)

is ________ kJ.A) -395.62B) -527.65C) 439.63D) 571.66

E) 527.65

Given the following reactions

N2 (g) + O2 (g) → 2NO (g) ΔH = +180.7 kJ 2NO( g) + O2 (g) → 2NO2 (g) ΔH = -113.1 kJthe enthalpy for the decomposition of nitrogen dioxide into molecular nitrogen and oxygen

2NO2 (g) → N2 (g) + 2O2 (g)

is ________ kJ.A) 67.6B) -67.6C) 293.8D) -293.8

E) 45.5

Given the following reactions N2 (g) + O2 (g) → 2NO (g) ΔH = +180.7 kJ 2NO (g) + O2(g) → 2N (g) ΔH = -113.1 kJthe enthalpy of reaction for

4NO (g) → 2NO2 (g) + N2 (g)

is ________ kJ.A) 67.6B) 45.5C) -293.8D) -45.5

E) 293.8

Given the following reactions N2 (g) + O2 (g) → 2NO (g) ΔH = +180.7 kJ 2N2O (g) → O2 (g) + 2N2 (g) ΔH = -163.2 kJthe enthalpy of reaction for

2N2O (g) → 2NO (g) + N2 (g)

is ________ kJ.A) 145.7B) 343.9C) -343.9D) 17.5

E) -145.7

The value of ΔH° for the reaction below is -186 kJ.
H2 (g) + Cl2 (g) → 2HCl (g)

The value of ΔH°f for HCl (g) is ________ kJ/mol.A) -3.72 × 102B) -1.27 × 102C) -93.0D) -186

E) +186

The value of ΔH° for the following reaction is -3351 kJ: 2Al (s) + 3O2(g) → 2Al2O3(s)The value of ΔH°f for Al2 (s) is ________ kJ.A) -3351B) -1676C) -32.86D) -16.43

E) +3351

Given the data in the table, ΔH°rxn for the reaction

Ca(OH)2 + 2H3AsO4 → Ca(H2AsO4)2 + 2H2O

is ________ kJ.

A) -744.9B) -4519C) -4219D) -130.4

E) -76.4

Given the data in the table, ΔH°rxn for the reaction

4NH3 (g) + 5O2 (g) → 4NO (g) + 6H2O (l)

is ________ kJ.

A) -1172B) -150C) -1540D) -1892

E) The ΔH°f of O2 (g) is needed for the calculation.

Given the data in the table, ΔH°rxn for the reaction
C2H5OH (l) + O2 (g) → CH3CO2H (l) + H2O (l)

is ________ kJ.

A) -79.0B) -1048.0C) -476.4D) -492.6

E) The value of ΔH°f of O2 (g) is required for the calculation.

Given the data in the table, ΔH°rxn for the reaction
3NO2 (g) + H2O (l) → 2HNO3 (aq) + NO (g)

is ________ kJ.

A) 64B) 140C) -140D) -508

E) -64

Given the data in the table, ΔH°rxn for the reaction

IF5 (g) + F2 (g) → IF7 (g)

is ________ kJ.

A) 1801B) -1801C) 121D) -121

E) -101

Given the data in the table, ΔH° for the reaction
2CO (g) + O2 (g) → 2CO2 (g)

is ________ kJ.

A) -566.4B) -283.2C) 283.2D) -677.0

E) The ΔH°f of O2 (g) is needed for the calculation.

The value of ΔH° for the following reaction is 177.8 kJ. The value of Δ for CaO(s) is ________ kJ/mol.
CaCO3 (s) → CaO (s) + CO2 (g)

A) -1600B) -813.4C) -635.5D) 813.4

E) 177.8

Given the data in the table, ΔH°rxn for the reaction
2Ag2S (s) + O2 (g) → 2Ag2O (s) + 2S (s)

is ________ kJ.

A) -1.6B) +1.6C) -3.2D) +3.2

E) The ΔH°f of S (s) and of O2 (g) are needed for the calculation.

Given the data in the table, ΔH°rxn for the reaction
Ag2O (s) + H2S (g) → Ag2S (s) + H2O (l)

is ________ kJ.

A) -267B) -370C) -202D) -308

E) More data are needed to complete the calculation.

Given the data in the table, ΔH°rxn for the reaction
2SO2 (g) + O2 (g) → 2SO3 (g)

is ________ kJ.

A) -99B) 99C) -198D) 198

E) The ΔH°f of O2 (g) is needed for the calculation.

Given the data in the table, ΔH°rxn for the reaction
SO3 (g) + H2O (l) → H2SO4 (l)

is ________ kJ.

A) -132B) 1496C) 704D) -704

E) -2.16 × 103

Given the data in the table, ΔH°rxn for the reaction
3Cl2 (g) + PH3 (g) → PCl3 (g) + 3HCl (g)

is ________ kJ.

A) -385.77B) -570.37C) 570.37D) 385.77

E) The ΔH°f of Cl2 (g) is needed for the calculation.

100

Given the data in the table, ΔH°rxn for the reaction
PCl3 (g) + 3HCl (g) → 3Cl2 (g) + PH3 (g)

is ________ kJ.

A) -570.37B) -385.77C) 570.37D) 385.77

E) The ΔH°f of Cl2 (g) is needed for the calculation.

101

Given the data in the table and ΔH°rxn for the reaction
SO2Cl2 (g) + 2H2O (l) → H2SO4 (l) + 2HCl (g) ΔH° = -62 kJ

ΔH°f of HCl (g) is ________ kJ/mol.

A) -184B) 60C) -92D) 30

E) Insufficient data are given.

102

A 19.5 g candy bar contains 8% protein, 33% fat, and 18% carbohydrate. The respective fuel values for protein, fat, and carbohydrate are 17, 38, and 17 kJ/g, respectively. What is the fuel value (kJ) for this piece of candy?A) 241B) 331C) 27D) 60.0

E) 17.0

103

A 3.00 L pitcher of sweetened ice tea contains 600. g of sugar. Assuming that the sugar is the only fuel source, what is the fuel value (in kJ) of a 250. mL serving? The respective fuel values for protein, fat, and carbohydrate are 17, 38, and 17 kJ/g, respectively.
A) 8.50 × 102 kJ
B) 10.2 × 104 kJ
C) 2.55 × 103 kJD) 38 kJ

E) 17 kJ

104

A typical fast food meal consists of a burger, fries, and a soft-drink and contains 58.0 grams of fat, 39.0 grams of protein, and 177 grams of carbohydrate. If jogging burns 950.0 kJ/hour, how many minutes would it take to completely burn off the meal? The respective fuel values for protein, fat, and carbohydrate are 17, 38, and 17 kJ/g, respectively.A) 6.19B) 208C) 371D) 17.3

E) 9.70

105

A slice of cake contains 29.0 grams of fat, 9.0 grams of protein, and 77 grams of carbohydrate. If swimming burns 1000.0 kJ/hour, how many minutes would it take to completely burn off the slice of cake? The respective fuel values for protein, fat, and carbohydrate are 17, 38, and 17 kJ/g, respectively.A) 154B) 2.56C) 23.4D) 117

E) 262

106

A 26.9 g rock rolls down the hill at a speed of 81.9 m/s . What is the kinetic energy of the rock?A) 90.2 JB) 145 JC) 0.950 JD) 90200 J

E) 1450 J

107

A 23.2 g piece of space debris is traveling at 81.9 m/s. What is the kinetic energy of the space debris?A) 145 JB) 1450 JC) 0.950 JD) 77800 J

E) 77.8 J

108

At what velocity (m/s) must a object be moving in order to possess a kinetic energy of 1.0J?A) 0.35 m/sB) 2.8 m/sC) 0.13 m/sD) 0.031 m/s

E) 0.016 m/s

109

At what velocity (m/s) must a 417.3 g object be moving in order to possess a kinetic energy of 3.2J?A) 0.12 m/sB) 26 m/sC) 0.015 m/sD) 0.0038 m/s

E) 0.00024 m/s

110

When work is done on a system, w will be a ________ value.A) positiveB) negativeC) very largeD) very small

E) There is not enough information given to determine the answer.

111

The value of ΔE for a system that performs 139 kJ of work on its surroundings and gains of heat is ________ kJ.A) -85B) 193C) 7506D) 85

E) -193

112

The value of ΔE for a system that performs 19 kJ of work on its surroundings and loses of heat is ________ kJ.A) -28B) 28C) 171D) 10

E) -10

113

Calculate the work (kJ) done during a reaction in which the internal volume expands from to against an outside pressure of 2.5atm.A) -7.3 kJB) 17 kJC) 7.3 kJD) -17 kJ

E) 0 kJ; No work is done.

114

Calculate the work (kJ) done during a reaction in which the internal volume expands from to against a vacuum (an outside pressure of 0 atm).A) 0; kJ No work is done.B) 3.6 kJC) -3.6 kJD) 6.5 kJ

E) -6.5 kJ

115

Calculate the work (kJ) done during a reaction in which the internal volume contracts from to against an outside pressure of 4.4 atm.A) 31 kJB) 43 kJC) -31 kJD) -43 kJ

E) 0 kJ; No work is done.

116

The value of ΔE for a system that performs 151 kJ of work on its surroundings and loses 79 kJ of heat is ________ kJ.A) +230.B) -230.C) +72D) -72

E) -151

117

Calculate the value of ΔE in joules for a system that loses 115 J of heat and has 150 J of work performed on it by the surroundings.A) -115 JB) -35 JC) +35 JD) +265 J

E) -265 J

118

The value of ΔH° for the reaction below is -72 kJ. ________ kJ of heat are released when 5.5 mol of HBr is formed in this reaction.

H2 (g) + Br2 (g) → 2HBr (g)

A) 144B) 72C) 0.44D) 198

E) -72

119

The value of ΔH° for the reaction below is -126 kJ. The amount of heat that is released by the reaction of 10.0 g of Na2O2 with water is ________ kJ.

2Na2O2 (s) + 2H2O (l) → 4NaOH (s) + O2 (g)

A) 8.08B) 16.2C) 67.5D) 32.3

E) -126

120

The value of ΔH° for the reaction below is -482 kJ. Calculate the heat (kJ) released to the surroundings when 10.0 g of CO (g) reacts completely.
2CO (g) + O2 (g) → 2CO2 (g)

A) 2410 kJB) 172 kJC) 86.0 kJD) 482 kJ

E) -482 kJ

121

In the presence of excess oxygen, methane gas burns in a constant-pressure system to yield carbon dioxide and water:

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) △H = -890.0 kJ

Calculate the value of q (kJ) in this exothermic reaction when 1.80 g of methane is combusted at constant pressure.A) -100.1 kJB) 0.0324 kJC) -0.0100 kJD) 30.9 kJ

E) -1.00 × 105 kJ

122

Hydrogen peroxide decomposes to water and oxygen at constant pressure (△H = -196 kJ). What is the value of q (kJ) for this reaction when 4.60 g of hydrogen peroxide decomposes at constant pressure?A) -26.5 kJB) -0.0189 kJC) 1.25 kJ

D) -2.65 × 104 kJ

E) -13.3 kJ

123

The combustion of titanium with oxygen produces titanium dioxide:

Ti (s) + O2(g) → TiO2 (s)

When 0.610 g of titanium is combusted in a bomb calorimeter, the temperature of the calorimeter increases from 25.00 °C to 50.50 °C. In a separate experiment, the heat capacity of the calorimeter is measured to be 9.84 kJ/K. The heat of reaction for the combustion of a mole of Ti in this calorimeter is ________ kJ/mol.A) 2.09B) 4.14C) -311D) -0.154

E) -1.98 × 104

124

A sample of aluminum metal absorbs 11.2 J of heat, upon which the temperature of the sample increases from 23.2 °C to 30.5 °C. Since the specific heat capacity of aluminum is 0.90 J/g-K, the mass of the sample is ________ g.A) 72B) 1.7C) 10.D) 65

E) 7.3

125

A sample of calcium carbonate [CaCO3 (s)] absorbs of heat, upon which the temperature of the sample increases from 20.8 °C to 27.3 °C. If the specific heat of calcium carbonate is what is the mass (in grams) of the sample?A) 7.6 gB) 5.1 gC) -7.6 gD) 0.13 g

E) 5.3 g

126

How many joules of heat are absorbed when the temperature of a 13.9 g sample of CaCO3 (s) increases from 21.7 °C to 33.3 °C? Specific heat of calcium carbonate is 0.82 J/g-K.A) 130 JB) 0.68 JC) s-130 JD) -0.68 J

E) 9.5 J

127

An 6.11 g sample of calcium carbonate [CaCO3 (s)] absorbs of heat, upon which the temperature of the sample increases from 19.2 °C to 35.9 °C. What is the specific heat of calcium carbonate?A) 0.82 J/g-KB) -0.82 J/g-KC) 31 J/g-KD) 230 J/g-K

E) 8600 J/g-K

128

A sample of iron absorbs 81.0 J of heat, upon which the temperature of the sample increases from 19.7 °C to 28.2 °C. If the specific heat of iron is 0.450 J/g-K, what is the mass (in grams) of the sample?A) 21.2 gB) 4.29 gC) -21.2 gD) 0.0472 g

E) 3.83 g

129

The temperature of a 35.1 g sample of iron increases from 24.6 °C to 31.8 °C If the specific heat of iron is 0.450 J/g-K, how many joules of heat are absorbed?A) 115 JB) 0.0936 JC) -115 JD) 0.722 J

E) 3.29 J

130

A 22.9 g sample of iron absorbs 155 J of heat, upon which the temperature of the sample increases from 23.9 °C to 38.9 °C. What is the specific heat of iron?A) 0.451 J/g-KB) -0.451 J/g-KC) 237 J/g-KD) 102 J/g-K

E) 53,200 J/g-K

131

The specific heat capacity of liquid water is 4.18 J/g-K. How many joules of heat are needed to raise the temperature of 7.25 g of water from 20.0 °C to 44.1 °C?A) 41.8 JB) 730 J

C) 1.94 × 103 J

D) 2.39 × 10-2 J
E) 66.8 J

132

The specific heat capacity of methane gas is 2.20 J/g-K. How many joules of heat are needed to raise the temperature of 7.25 g of methane from 22.0 °C to 57.0 °C?A) 115 JB) 558 J

C) 1.26 × 103 J

D) 8.67 × 10-3 J
E) 41.3 J

133

The specific heat capacity of liquid mercury is 0.14 J/g-K. How many joules of heat are needed to raise the temperature of 6.00 g of mercury from 25.1 °C to 65.3 °C?
A) 1.7 × 103 JB) 34 JC) 76 J

D) 5.8 × 10-4 J

E) 2.2 J

134

How much heat is required to raise the temperature of a 1.15 kg piece of copper metal from 25.0 °C to 77.5 °C? The specific heat capacity of solid copper metal is 0.385 J/g-K.
A) 2.32 × 104 J
B) 1.57 × 105 JC) 23.2 J

D) 6.38 × 10-6 J

E) 0.00638 J

135

A 4.50-g sample of liquid water at 25.0 °C is heated by the addition of 133 J of energy. The final temperature of the water is ________ °C. The specific heat capacity of liquid water is 4.18 J/g-K.A) 149B) 25.1C) -17.9D) 32.1

E) 7.07

136

A 10.1 g sample of NaOH is dissolved in 250.0 g of water in a coffee-cup calorimeter. The temperature increases from 23.0 °C to ________°C. Specific heat of liquid water is 4.18 J/g-K and ΔH for the dissolution of sodium hydroxide in water is 44.4 kJ/mol. A) 35.2B) 24.0C) 33.7D) 33.3

E) 40.2

137

A 50.0-g sample of liquid water at 25.0 °C is mixed with 23.0 g of water at 79.0 °C. The final temperature of the water is ________ °C. A) 123B) 27.3C) 52.0D) 231

E) 42.0

138

A 5.00-g sample of copper metal at 25.0 °C is heated by the addition of 133 J of energy. The final temperature of the copper is ________ °C. The specific heat capacity of copper is A) 35.1B) 25.0C) 45.0D) 95.0

E) 70.0

139

The temperature of a 24.3 g sample of gold increases from 23.7 °C to 31.5 °C. If the specific heat of gold is 0.129 J/g-K, how many joules of heat are absorbed?A) 24.5 JB) 0.0414 JC) -24.5 JD) 0.293 J

E) 1.01 J

140

What is the enthalpy change (in kJ) of a chemical reaction that raises the temperature of 250.0 mL of solution having a density of 1.25 g/mL by 3.33 °C? (The specific heat of the solution is 3.74 J/g-K.)A) -7.43 kJB) -12.51 kJC) 8.20 kJD) -3.89 kJ

E) 6.51 kJ

141

An 8 oz. bottle of energy drink contains 6.0 g of protein, 2.0 g of fat, and 16.3 g of carbohydrate. The fuel value of this energy drink bottle is ________ kJ. The fuel values for protein, fat, and carbohydrate are 17, 38, and 17 kJ/g, respectively.A) 520B) 280C) 720D) 460

E) 72

142

________ is defined as the energy used to move an object against a force.

143

The ΔHvap of water is 40.7 kJ at 100 °C. How much liquid water in grams can be converted to vapor if 5950 J of heat are absorbed?

144

The ΔHvap of water is 40.7 kJ at 100 °C. How much heat energy is required to convert 15.0 grams of liquid water to vapor.

145

When 0.800 grams of NaOH is dissolved in 100.0 grams of water, the temperature of the solution increases from 25.00 °C to 27.06 °C. The amount of heat absorbed by the water is ________ J. (The specific heat of water is 4.18 J/g-°C.)

146

The ΔHrxn for the combustion of methane is -890.0 kJ. How much heat energy (kJ) is released if 82.1 grams of methane are burned in an excess amount of oxygen?

147

The ________ of a reaction is the enthalpy change when all reactants and products are at 1 atm pressure and a specific temperature.

148

Coal contains hydrocarbons of high molecular weight as well as compounds containing ________, sulfur, or nitrogen.

149

Work equals mass times distance.

150

One joule equals 1 kg-m2/s2.

151

Energy units include watts, volts, and newtons.

152

The primary component of natural gas is methane.