Test: Thermochemistry (Chapter 7)

Multiple Choice Questions

1. Thermochemistry primarily studies:

a) Reaction rates

b) Equilibrium positions

c) Heat changes accompanying chemical reactions

d) Electrical properties of reactions
2. A reaction that releases heat to the surroundings is called:

a) Endothermic

b) Exothermic

c) Isothermic

d) Adiabatic
3. A reaction that absorbs heat from the surroundings is called:

a) Endothermic

b) Exothermic

c) Isothermic

d) Adiabatic
4. In an exothermic reaction, the temperature of the system initially:

a) Decreases

b) Remains constant

c) Increases

d) Becomes zero
5. The standard unit for heat changes in the SI system is:

a) Calorie

b) Watt

c) Joule (or Kilojoule)

d) Degree Celsius
6. The formation of water from H₂(g) and O₂(g) (H₂ + ½O₂ -> H₂O) is an example of an ___ reaction.

a) Endothermic

b) Exothermic

c) Equilibrium

d) Slow
7. The decomposition of water into H₂(g) and O₂(g) (H₂O -> H₂ + ½O₂) is an example of an ___ reaction.

a) Endothermic

b) Exothermic

c) Spontaneous

d) Fast
8. Thermochemistry is primarily based on which law of thermodynamics?

a) Zeroth Law

b) First Law

c) Second Law

d) Third Law
9. A process that takes place on its own without outside assistance is termed:

a) Reversible

b) Non-spontaneous

c) Spontaneous

d) Equilibrium
10. Which of the following is generally a spontaneous process?

a) Water flowing uphill

b) Heat transfer from cold to hot

c) Dissolving sugar in water

d) Decomposition of water at room temp
11. Burning coal requires initiation (a spark) but then proceeds on its own. This is considered:

a) Non-spontaneous

b) Spontaneous

c) Reversible

d) Endothermic
12. A non-spontaneous process:

a) Occurs rapidly

b) Releases energy

c) Does not take place on its own

d) Is always exothermic
13. Can all spontaneous reactions be predicted based solely on whether they are exothermic or endothermic?

a) Yes, only exothermic are spontaneous

b) Yes, only endothermic are spontaneous

c) No, energy change alone is not sufficient

d) Yes, spontaneity depends only on ΔH
14. The dissolution of NH₄Cl in water is a spontaneous process that is:

a) Exothermic

b) Endothermic

c) Neutral (ΔH=0)

d) Very rapid
15. In thermochemistry, the part of the universe under study is called the:

a) Surroundings

b) Boundary

c) System

d) Universe
16. Everything outside the system is known as the:

a) Boundary

b) System

c) Universe

d) Surroundings
17. A property of a system that depends only on its initial and final states, not the path taken, is called a:

a) Path function

b) Process variable

c) State function

d) System property
18. Which of the following is generally considered a state function?

a) Heat (q)

b) Work (w)

c) Internal Energy (E)

d) Path taken
19. The change in a state function (like ΔV or ΔT) depends on:

a) The path taken

b) The rate of change

c) Only the initial and final states

d) The surroundings
20. The internal energy (E) of a system represents the sum of:

a) Only kinetic energies of particles

b) Only potential energies of particles

c) Kinetic and potential energies of particles

d) Heat and work
21. Kinetic energy of molecules includes which types of motion?

a) Translational only

b) Rotational only

c) Vibrational only

d) Translational, rotational, and vibrational
22. Potential energy within a system includes:

a) Only chemical bond energy

b) Only intermolecular forces

c) All types of bonds and intermolecular forces

d) Only nuclear energy
23. Can the absolute value of internal energy (E) be easily measured?

a) Yes, using a calorimeter

b) Yes, using a thermometer

c) No, only the change (ΔE) can often be measured

d) Yes, it is always zero
24. Heat (q) is defined as energy transfer due to:

a) Work done

b) A temperature difference

c) A pressure difference

d) A chemical reaction
25. Work (w) done by a system involving gas expansion against external pressure P is given by:

a) PΔV

b) -PΔV

c) ΔV/P

d) P/ΔV
26. According to the sign convention used, 'q' is positive when:

a) Heat is absorbed by the system

b) Heat is released by the system

c) Work is done by the system

d) Work is done on the system
27. According to the sign convention used, 'w' is positive when:

a) Heat is absorbed by the system

b) Heat is released by the system

c) Work is done by the system

d) Work is done on the system
28. The First Law of Thermodynamics is essentially a statement of:

a) Entropy increase

b) Conservation of energy

c) Ideal gas behavior

d) Rate of reaction
29. The mathematical expression for the First Law of Thermodynamics is:

a) ΔE = q * w

b) ΔE = q / w

c) ΔE = q - w

d) ΔE = q + w
30. If a process occurs at constant volume (ΔV = 0), then ΔE equals:

a) q (heat absorbed/released)

b) w (work done)

c) 0

d) ΔH
31. Enthalpy (H) is defined as:

a) E - PV

b) E + PV

c) PV / E

d) E / PV
32. Is Enthalpy (H) a state function?

a) Yes

b) No

c) Only at constant pressure

d) Only for gases
33. The change in enthalpy (ΔH) is related to the change in internal energy (ΔE) by:

a) ΔH = ΔE - PΔV

b) ΔH = ΔE + PΔV (at constant P)

c) ΔH = ΔE / PΔV

d) ΔH = PΔV - ΔE
34. For reactions involving only liquids and solids, where volume changes are negligible (ΔV ≈ 0):

a) ΔH >> ΔE

b) ΔH << ΔE

c) ΔH ≈ ΔE

d) ΔH = 0
35. At constant pressure, the change in enthalpy (ΔH) is equal to:

a) w (work done)

b) ΔE

c) qp (heat absorbed/released at constant P)

d) 0
36. Why is enthalpy (ΔH) often more convenient to work with than internal energy (ΔE) in chemistry?

a) ΔE is always zero

b) Most reactions are carried out at constant volume

c) Most reactions are carried out at constant pressure (open to atmosphere)

d) ΔH is easier to calculate
37. For the reaction 2H₂(g) + O₂(g) -> 2H₂O(g) at 100°C, Δn (change in moles of gas) is:

a) 0

b) 1

c) -1

d) -2
38. If ΔH for a reaction is -242.2 kJ/mol and PΔV is -1.55 kJ/mol, what is ΔE?

a) -243.75 kJ/mol

b) -240.65 kJ/mol

c) -242.2 kJ/mol

d) +240.65 kJ/mol
39. Standard conditions for thermochemical measurements are typically defined as:

a) 0°C and 1 atm

b) 25°C and 1 atm

c) 100°C and 1 atm

d) 0 K and 0 atm
40. Standard enthalpy of reaction (ΔH°) refers to the enthalpy change when reactants in their standard states form products in their ___ states, according to the balanced equation.

a) Gaseous

b) Liquid

c) Standard

d) Most stable
41. Standard enthalpy of formation (ΔH°f) is the enthalpy change when ___ mole(s) of a compound is formed from its ___ in their standard states.

a) one, constituent atoms

b) two, elements

c) one, elements

d) two, constituent atoms
42. The standard enthalpy of formation (ΔH°f) of any element in its standard state is:

a) Always positive

b) Always negative

c) Zero

d) Equal to its ΔH°c
43. Standard enthalpy of atomization (ΔH°at) is the enthalpy change when one mole of ___ are formed from an element in its standard state.

a) Gaseous molecules

b) Liquid atoms

c) Gaseous ions

d) Gaseous atoms
44. The process described by ½H₂(g) -> H(g) represents the enthalpy of ___ for hydrogen.

a) Formation

b) Combustion

c) Atomization

d) Neutralization
45. Standard enthalpy of neutralization (ΔH°n) is the heat evolved when one mole of ___ reacts with one mole of ___ to form one mole of water.

a) Acid, Base

b) H⁺ ions, OH⁻ ions

c) Salt, Water

d) Oxidant, Reductant
46. The enthalpy of neutralization for any strong acid reacting with a strong base is approximately:

a) 0 kJ/mol

b) +57.4 kJ/mol

c) -57.4 kJ/mol

d) -285.8 kJ/mol
47. Why is the enthalpy of neutralization constant for strong acid-strong base reactions?

a) All salts formed are soluble

b) The only net reaction is H⁺(aq) + OH⁻(aq) -> H₂O(l)

c) Strong acids and bases don't ionize

d) The reactions are endothermic
48. Standard enthalpy of combustion (ΔH°c) is the heat evolved when one mole of a substance is completely burnt in excess ___ under standard conditions.

a) Nitrogen

b) Hydrogen

c) Oxygen

d) Chlorine
49. Enthalpy of combustion values are always:

a) Positive (endothermic)

b) Negative (exothermic)

c) Zero

d) Equal to enthalpy of formation
50. Standard enthalpy of solution (ΔH°sol) is the heat change when one mole of a substance dissolves in enough solvent so that:

a) The solution becomes saturated

b) The solution boils

c) Further dilution causes no more heat change

d) The concentration is exactly 1 M
51. The dissolution of ammonium chloride (NH₄Cl) in water is an ___ process (ΔH°sol = +16.2 kJ/mol).

a) Exothermic

b) Endothermic

c) Isothermic

d) Adiabatic
52. A device used to measure heat changes in chemical reactions is called a:

a) Thermometer

b) Barometer

c) Calorimeter

d) Spectrometer
53. The heat absorbed or released (q) in calorimetry is calculated using q = m × s × ΔT, where 's' represents:

a) Mass

b) Specific heat

c) Temperature change

d) Volume
54. The heat capacity (C) of a calorimeter system is the product of:

a) Mass and ΔT

b) Specific heat and ΔT

c) Mass and specific heat

d) Heat and ΔT
55. A bomb calorimeter is typically used to measure:

a) Enthalpy of neutralization

b) Enthalpy of solution

c) Enthalpy of combustion

d) Enthalpy of formation
56. Why is oxygen supplied under high pressure (e.g., 20 atm) in a bomb calorimeter?

a) To keep the volume constant

b) To ensure complete combustion

c) To cool the system

d) To measure pressure change
57. In a bomb calorimeter experiment, the heat change (q) is calculated using q = c × ΔT, where 'c' is the:

a) Specific heat of the sample

b) Specific heat of water

c) Heat capacity of the calorimeter system

d) Heat capacity of the bomb
58. Bomb calorimetry measures heat changes at constant:

a) Pressure (ΔH)

b) Temperature

c) Volume (ΔE)

d) Entropy
59. Hess's Law states that the total enthalpy change for a reaction:

a) Depends on the catalyst used

b) Depends on the number of steps

c) Is independent of the path taken

d) Is always zero
60. Hess's Law is essentially an application of the:

a) Law of Mass Action

b) First Law of Thermodynamics

c) Second Law of Thermodynamics

d) Ideal Gas Law
61. According to Hess's Law, if A -> D directly has ΔH, and A -> B -> C -> D has steps ΔH₁, ΔH₂, ΔH₃, then:

a) ΔH = ΔH₁ + ΔH₂ + ΔH₃

b) ΔH = ΔH₁ - ΔH₂ - ΔH₃

c) ΔH = ΔH₃ - ΔH₂ - ΔH₁

d) ΔH = ΔH₁ * ΔH₂ * ΔH₃
62. Hess's Law is particularly useful for calculating enthalpy changes that are:

a) Very large

b) Very small

c) Difficult or impossible to measure directly

d) Measured using a bomb calorimeter
63. If C(s) + O₂(g) -> CO₂(g) has ΔH = -393 kJ and CO(g) + ½O₂(g) -> CO₂(g) has ΔH = -283 kJ, what is ΔH for C(s) + ½O₂(g) -> CO(g)?

a) -676 kJ

b) +110 kJ

c) -110 kJ

d) +676 kJ
64. The Born-Haber cycle is an application of Hess's Law used primarily to calculate:

a) Enthalpy of combustion

b) Enthalpy of formation

c) Lattice energy of ionic compounds

d) Bond energy
65. Lattice energy (ΔH°Latt) is the enthalpy change for the process:

a) M(s) + X₂(g) -> MX(s)

b) M⁺(g) + X⁻(g) -> MX(s)

c) MX(s) -> M(g) + X(g)

d) MX(s) -> M⁺(aq) + X⁻(aq)
66. Can lattice energy be measured directly in a calorimeter?

a) Yes, easily

b) Yes, using a bomb calorimeter

c) No, it is calculated indirectly

d) Yes, using a glass calorimeter
67. Which enthalpy change corresponds to Na(s) -> Na(g)?

a) Ionization energy

b) Enthalpy of formation

c) Enthalpy of atomization

d) Lattice energy
68. Which enthalpy change corresponds to Na(g) -> Na⁺(g) + e⁻?

a) Ionization energy

b) Enthalpy of atomization

c) Electron affinity

d) Lattice energy
69. Which enthalpy change corresponds to ½Cl₂(g) -> Cl(g)?

a) Ionization energy

b) Enthalpy of formation

c) Enthalpy of atomization

d) Electron affinity
70. Which enthalpy change corresponds to Cl(g) + e⁻ -> Cl⁻(g)?

a) Ionization energy

b) Enthalpy of atomization

c) Electron affinity

d) Lattice energy
71. In the Born-Haber cycle for NaCl, ΔH°f (NaCl) equals:

a) ΔH°Latt + ΔH°x

b) ΔH°Latt - ΔH°x

c) ΔH°x - ΔH°Latt

d) ΔH°Latt * ΔH°x

Short Answer Questions

Provide your answers in the text boxes below.

1. Define thermochemistry.
2. Differentiate between exothermic and endothermic reactions, giving one example of each.
3. What happens to the temperature of the system and surroundings during an exothermic reaction?
4. What happens to the temperature of the system and surroundings during an endothermic reaction?
5. What are the common SI units for measuring heat changes?
6. Define a spontaneous process.
7. Give two examples of spontaneous processes.
8. Define a non-spontaneous process.
9. Give two examples of non-spontaneous processes.
10. Explain why stating that 'all exothermic reactions are spontaneous' is incorrect. Give an example.
11. Define system, surroundings, and boundary in thermodynamics.
12. What is meant by the 'state' of a system?
13. Define 'state function' and give three examples.
14. Explain why properties like heat (q) and work (w) are not state functions.
15. Define the internal energy (E) of a system.
16. What types of energy contribute to the internal energy?
17. State the First Law of Thermodynamics (Law of Conservation of Energy).
18. Write the mathematical equation for the First Law of Thermodynamics and define each term (ΔE, q, w).
19. Explain the sign conventions for heat (q) and work (w).
20. What is pressure-volume work? Give the formula.
21. Show that at constant volume, the heat change (qv) is equal to the change in internal energy (ΔE).
22. Define enthalpy (H).
23. Write the mathematical relationship between enthalpy (H), internal energy (E), pressure (P), and volume (V).
24. Derive the relationship ΔH = ΔE + PΔV for a process occurring at constant pressure.
25. Show that at constant pressure, the heat change (qp) is equal to the change in enthalpy (ΔH).
26. Under what conditions is ΔH approximately equal to ΔE?
27. Why is ΔH generally preferred over ΔE in chemical reactions?
28. Define standard enthalpy of reaction (ΔH°). What are standard conditions?
29. Define standard enthalpy of formation (ΔH°f).
30. What is the standard enthalpy of formation of an element in its standard state?
31. Define standard enthalpy of atomization (ΔH°at). Give an example.
32. Define standard enthalpy of neutralization (ΔH°n).
33. Why is the enthalpy of neutralization approximately -57.4 kJ/mol for all strong acid-strong base reactions?
34. Define standard enthalpy of combustion (ΔH°c).
35. Are enthalpies of combustion typically positive or negative? Why?
36. Define standard enthalpy of solution (ΔH°sol).
37. What is a calorimeter?
38. Describe the basic components of a simple glass calorimeter.
39. How is the heat change (q) calculated in glass calorimetry using q = m × s × ΔT?
40. What is meant by the 'heat capacity' (C) of a calorimeter?
41. Describe the main components and setup of a bomb calorimeter.
42. What type of enthalpy change is usually measured using a bomb calorimeter?
43. How is the heat change (q) calculated in bomb calorimetry using q = c × ΔT?
44. Does a bomb calorimeter measure ΔE or ΔH directly? Explain.
45. State Hess's Law of Constant Heat Summation.
46. Explain how Hess's Law is a consequence of the First Law of Thermodynamics.
47. Give an example of how Hess's Law can be used to calculate an enthalpy change that is difficult to measure directly (e.g., ΔH°f of CO).
48. Show mathematically why the enthalpy change for a cyclic process is zero according to Hess's Law.
49. What is the main application of the Born-Haber cycle?
50. Define lattice energy (ΔH°Latt).
51. Draw a labelled Born-Haber cycle for the formation of NaCl(s) from Na(s) and Cl₂(g).
52. List the individual enthalpy steps involved in the Born-Haber cycle for NaCl.
53. Explain how the lattice energy of NaCl can be calculated using the Born-Haber cycle and Hess's Law.