Chapter 9: Thermochemistry

Question 1

Define heat.

Answer 1

Heat (q) is the flow of energy caused by a temperature difference.

Question 2

State the first law of thermodynamics.

Answer 2

This is the law of energy conservation, which states that the total energy of the universe is constant.

Question 3

Define internal energy.

Answer 3

The internal energy of a system is the sum of the kinetic and potential energies of all the particles that compose the system.

Question 4

What does it mean to say that internal energy is a state function?

Answer 4

It means that its value depends only on the state of the system, not on how the system arrived at that state. The value of a change in a state function is always the difference between its final and initial values.

Question 5

If the reactants of a reaction have a higher internal energy than the products, how does the energy of the system change?

Answer 5

The change in energy of the system is negative and energy flows out of the system into the surroundings.

Question 6

If the reactants of a reaction have a lower internal energy than the products, how does the energy of the system change?

Answer 6

The change in energy of the system is positive and energy flows into the system from the surroundings.

Question 7

According to the first law of thermodynamics, what is the change in the internal energy of a system equal to?

Answer 7

Change in E is equal to the sum of the heat transferred (q) and the work done (w).

Question 8

What is the difference between heat and temperature?

Answer 8

Temperature is a measure of the thermal energy within a sample of matter. Heat is the transfer of thermal energy.

Question 9

Define thermal equilibrium.

Answer 9

Thermal equilibrium is the point at which heat transfer from the system to the surroundings (or vice versa) stops because both are at the same temperature. At thermal equilibrium, there is no additional net transfer of heat.

Question 10

Define heat capacity.

Answer 10

Heat capacity is the quantity of heat required to change the temperature of a system by 1 degree Celsius.

Question 11

Define specific heat capacity.

Answer 11

Specific heat capacity is the measure of the intrinsic capacity of a substance to absorb heat and is the amount of heat required to raise the temperature of 1 gram of the substance by 1 degree Celsius.

Question 12

In what way are specific heat capacity and molar heat capacity intensive properties?

Answer 12

They depend on the kind of substance being heated, not the amount.

Question 13

What does calorimetry measure, and how does it do so?

Answer 13

Calorimetry measures the heat evolved in a chemical reaction. Calorimetry measures the thermal energy the reaction (defined as the system) and the surroundings exchange by observing the change in temperature of the surroundings.

Question 14

Define enthalpy.

Answer 14

The enthalpy (H) of a system is the sum of its internal energy and the product of its pressure and volume.

Question 15

Why is enthalpy a state function?

Answer 15

Internal energy, pressure, and volume–the components that go into calculating enthalpy–are all state functions, so enthalpy is as well.

Question 16

Compare and contrast ∆H and ∆E.

Answer 16

∆H and ∆E are similar. They both represent changes in a state function for the system, but ∆E is a measure of all the energy (heat and work) exchanged with the surroundings, while ∆H is a measure of only the heat exchanged under conditions of constant pressure.

Question 17

What is the enthalpy of reaction?

Answer 17

The enthalpy of reaction is the enthalpy change for a chemical reaction. It is an extensive property, meaning it depends on the amount of material undergoing the reaction.

Question 18

How are bomb and coffee-cup calorimetry different?

Answer 18

Bomb calorimetry occurs at constant volume and measures ∆E for a reaction. Coffee-cup calorimetry occurs at constant pressure and measures ∆H for a reaction.

Question 19

What are the three quantitative relationships between a chemical equation and ∆Hrxn?

Answer 19

1. If a chemical equation is multiplied by some factor, then ∆Hrxn is also multiplied by the same factor.
2. If a chemical equation is reversed, then ∆Hrxn changes sign.
3. If a chemical equation can be expressed as the sum of a series of steps, then ∆Hrxn for the overall equation is the sum of the heats of reactions for each step.

Question 20

State Hess’s law.

Answer 20

If a chemical equation can be expressed as the sum of a series of steps, then ∆Hrxn for the overall equation is the sum of the heats of reactions for each step.

Question 21

Is ∆H positive or negative for breaking a bond?

Answer 21

Positive.

Question 22

Is ∆H positive or negative for forming a bond?

Answer 22

Negative

Question 23

In relation to bonds, when is a reaction exothermic?

Answer 23

When weak bonds break and strong bonds form.

Question 24

In relation to bonds, when is a reaction endothermic?

Answer 24

When strong bonds break and weak bonds form.

Question 25

What is the standard state of a gas?

Answer 25

The pure gas at a pressure of exactly 1 atm.

Question 26

What is the standard state of a liquid or solid?

Answer 26

The pure substance in its most stable form at a pressure of 1 atm and at the temperature of interest (often taken to be 25 C)

Question 27

What is the standard state of a substance in solution?

Answer 27

The standard state of a substance in solution is a concentration of exactly 1 M.

Question 28

What is standard enthalpy change (∆H˚)?

Answer 28

The change in enthalpy for a process when all reactants and products are in their standard states. The degree sign indicates standard states.

Question 29

What is the standard enthalpy of formation (∆H˚f) for a pure compound?

Answer 29

The change in enthalpy when 1 mol of the compound forms from its constituent elements in their standard states.

Question 30

What is the standard enthalpy of formation (∆H˚f) for a pure element in its standard state?

Answer 30

∆H˚f = 0

Question 31

How is standard enthalpy of formation similar to sea level?

Answer 31

Assigning the value of zero to the standard enthalpy of formation for an element in its standard state is the equivalent of assigning an altitude of zero to sea level. Once we assume sea level is zero, we can measure all subsequent changes in altitude relative to sea level. Similarly, we can measure all changes in enthalpy relative to those of pure elements in their standard states.

Question 32

What does the negative of the standard enthalpy of formation correspond to?

Answer 32

The decomposition of a compound into its constituent elements in their standard states. The formation a compound corresponds to the positive enthalpy of formation.

Question 33

How do you calculate ∆H˚rxn?

Answer 33

Subtract the enthalpies of formation of the reactants multiplied by their stoichiometric coefficients from the enthalpies of formation of the products multiplied by their stoichiometric coefficients.

Question 34

What is the Born-Haber cycle?

Answer 34

A hypothetical series of steps that represents the formation of an ionic compound from its constituent elements.

Question 35

Why does the magnitude of the lattice energy decrease as you move down the column of lattice energies for alkali metal chlorides?

Answer 35

As the ionic radii increase as we move down the column, the ions cannot get as close to each other and therefore do not release as much energy as when the lattice forms.

Question 36

How do lattice energies change with increasing ionic radius?

Answer 36

They become less exothermic (less negative) with increasing ionic radius.

Question 37

How do lattice energies change with increasing magnitude of ionic charge?

Answer 37

Lattice energies become more exothermic (more negative) with increasing magnitude of ionic charge.

Question 38

What types of energy are considered kinetic?

Answer 38

Thermal, mechanical, sound, radiant, and electrical.

Question 39

What types of energy are considered potential?

Answer 39

Gravitational, nuclear, elastic, electrical, and chemical.

Question 40

Why does Einstein’s equation E = mc2 not apply to the chemistry we are doing?

Answer 40

This equation reveals that energy and mass can be interconverted, which is important for nuclear reactions. For standard chemical reactions, however, mass and energy are converted independently.

Question 41

What is the definition of a Joule?

Answer 41

A Joule is the energy required to move a 1 kg mass by 1 meter.

Question 42

What is the definition of a calorie?

Answer 42

1 calorie is the energy required to raise the temperature of 1 gram of water by 1˚C.

Question 43

Why does the value for ∆Hrxn only apply to one specific reaction and its balanced chemical equation?

Answer 43

The heat (or enthalpy) of reaction is an extensive property, meaning it is mass-dependent. Therefore, its value is only true for a specific reaction.

Question 44

Define lattice energy.

Answer 44

Lattice energy is the energy associated with gaseous ions coalescing into a solid lattice of alternating cations and anions.

Question 45

List the steps of the Born-Haber cycle.

Answer 45

1. Sublimating of the metal (solid to gas)
2. Breaking of nonmetal bond (if diatomic)
3. Ionization of the gas phase metal (ionization energy)
4. Reduction of gas phase nonmetal (electron affinity)
5. Lattice energy

Question 46

Do the charges or the bond length have a more significant effect on a comparison of the lattice energy of two ionic compounds?

Answer 46

The charges will have a more significant effect than bond length.