Worksheetschapter 6-1
Total questions: 85
Worksheet time: 43mins
Which statement best defines thermochemistry?
The study of biological reactions inside cells
The study of the relationship between chemistry and energy (heat)
The measurement of electric potential in solutions
The analysis of reaction rates only
Energy is defined in this lesson as:
A property that always increases in chemical reactions
The capacity to do work and/or transfer heat
The total mass of a system
The amount of substance in moles
According to the material, which option correctly describes work (w)?
Random molecular motion within a system
A force acting through a distance
Energy stored in chemical bonds only
The flow of energy from cold to hot
Heat (q) is described as:
Energy associated with height above the ground
The flow of energy from hotter temperature to colder temperature
A force acting on a mass
Energy due to the composition of a compound
Which statement about energy transfer is emphasized in the curling images showing stones contacting each other?
Energy can only be created during motion
Energy can be transferred from one object to another by work or by heat
Energy transfer requires chemical change
Energy transfer occurs only at absolute zero
Which type of energy is associated with the motion of an object?
Kinetic energy
Thermal energy
Potential energy
Chemical energy
Thermal energy is best described as energy associated with:
The position of electrons and nuclei
The temperature of an object
The composition of a mixture
The work done by a force
Chemical energy, as described in the lesson, is:
A form of kinetic energy due to temperature
A form of potential energy associated with the relative positions of electrons and nuclei in atoms and molecules
Only the energy released in nuclear reactions
Energy that cannot be converted to other forms
A ball compressed on a spring is labeled “Mechanical potential energy,” and when released it shows “Kinetic energy.” What principle does this illustrate?
Entropy must decrease in all processes
Energy is created during motion
Law of Conservation of Energy: energy changes form but is not created or destroyed
Temperature remains constant during motion
Which statement aligns with the Law of Conservation of Energy as presented with the waterfall and spring-ball images?
Energy can be destroyed if transferred as heat
Energy can neither be created nor destroyed, but it can be transferred from one object or system to another
Energy is only conserved in isolated chemical systems
Potential energy cannot convert to kinetic energy
According to the diagrams of energy gauges, what happens to the surroundings when a system loses energy?
The surroundings lose the same amount of energy
The surroundings gain the exact amount of energy lost by the system
Energy is destroyed
The total energy of system and surroundings increases
Which statement best defines a system in thermodynamic discussions of energy exchange?
Everything in the universe
The part singled out to study
Anything that cannot exchange energy
Only the surroundings of a reaction
Which option correctly distinguishes surroundings from the system?
Surroundings are the part singled out to study; the system is everything else
Surroundings are everything else with which the system can exchange energy; the system is the part singled out to study
Surroundings and system are identical
Surroundings can exchange mass but never energy with the system
Which statement aligns with the Law of Conservation of Energy as depicted by the energy gauges?
Energy can be created in the system if the surroundings lose it
Total energy is conserved; a loss by the system equals a gain by the surroundings
Energy is sometimes destroyed during transfer
Only the system’s energy changes during transfer
What is the SI unit of energy, and how is it defined dimensionally?
Calorie (cal), defined as 1 g cm s−2
Joule (J), defined as 1 kg m2 s−2
Kilowatt-hour (kWh), defined as 1000 W s
Calorie (Cal), defined as 4184 N m
Which conversion is correct based on the table of energy units?
1 calorie (cal) = 4184 J
1 kilocalorie (kcal) = 1000 J
1 kilowatt-hour (kWh) = 3.60 × 10^6 J
1 Calorie (Cal) = 4.184 J
Energy is often expressed in kJ. Which statement is accurate?
1 kJ = 100 J
1 kJ = 1000 J
1 kJ = 4184 J
1 kJ = 3.60 × 10^6 J
Which is the SI unit of energy?
calorie (cal)
joule (J)
kilowatt-hour (kWh)
Calorie (Cal)
According to the table of Energy Uses in Various Units, which unit is most appropriate for reporting household electricity usage over a day?
joule (J)
calorie (cal)
kilowatt-hour (kWh)
Calorie (Cal)
Using 1 cal = 4.184 J, what is 4.91 J in calories? Choose the closest value.
0.184 cal
0.879 cal
1.17 cal
20.6 cal
A nutrition label lists 2.22 Cal (food Calories). Using 1 Cal = 1000 cal and 1 cal = 4.184 J, which is the best equivalent in kilojoules?
0.532 kJ
2.22 kJ
9.28 kJ
18.6 kJ
From the energy uses table, which unit shows approximately 0.100 required to light a 100 W bulb for 1 hour?
joule (J)
calorie (cal)
Calorie (Cal)
kilowatt-hour (kWh)
Which statement best compares Calorie (Cal) and calorie (cal) as used in the table?
They are identical units.
1 Cal equals 100 cal.
1 Cal equals 1000 cal.
1 cal equals 1000 Cal.
Recall: What is the SI base unit expression equivalent to 1 joule of energy?
kg m s⁻¹
kg m² s⁻²
N m⁻¹
W s
Skill/Concept: Using 1 cal = 4.184 J, what is 250 cal in kilojoules?
0.059 kJ
1.046 kJ
4.184 kJ
59.0 kJ
Strategic Thinking: A food label lists 200 Calories (nutritional kilocalories). Using 1 Cal = 1 kcal = 4.184 kJ, which is the best estimate of the energy in joules?
8.37×10² J
8.37×10³ J
8.37×10⁴ J
8.37×10⁵ J
Which statement best defines the internal energy (U) of a system?
The energy required to create heat only
The sum of the kinetic and potential energies of all particles in the system
The work needed to move the system against external pressure only
The temperature of the system multiplied by its volume
According to the First Law expression ΔU = q + w, which combination results in an increase in internal energy?
q < 0 and w < 0
q > 0 and w < 0
q < 0 and w > 0
q > 0 and w > 0
In the sign convention shown, what does a positive value of q indicate?
The system loses thermal energy
Work is done by the system
The system gains thermal energy
Energy flows out of the system
A gas expands against a piston while absorbing heat. Which sign combination is consistent with this process?
q > 0, w > 0
q < 0, w > 0
q > 0, w < 0
q < 0, w < 0
Which quantity is a state function according to the material?
Heat (q)
Work (w)
Internal energy (U)
Pathway taken to reach the final state
Why are heat and work not considered state functions?
They depend only on the final state
They depend only on the initial state
They are pathways between states rather than properties of the state
They are constant for any process
Two different processes take a system from the same initial state to the same final state. Which statement must be true?
The total heat transferred is the same for both processes
The total work done is the same for both processes
The change in internal energy (ΔU) is the same for both processes
The temperature remains constant during both processes
Which statement best defines a state function in thermodynamics?
A property that depends only on the current state of the system, not on the path taken
A property that depends on the path taken to reach a state
A quantity that is always conserved during reactions
A variable that only applies to open systems
In the combustion of H2(g), the reactants have higher internal energy than the product H2O(l). What is the sign of ΔrU for the system?
Positive
Negative
Zero
Cannot be determined
During the combustion of H2(g), energy flow is best described as:
From surroundings to system; ΔUsys > 0
From system to surroundings; ΔUsys < 0
No net flow; ΔUsys = 0
Oscillating between system and surroundings
For the reverse reaction converting H2O(l) to H2(g) and O2(g), how does the internal energy of the system change?
It decreases and ΔrU is negative
It increases and ΔrU is positive
It remains unchanged with ΔrU = 0
It alternates between positive and negative
In the reverse reaction (H2O(l) → H2(g) + 1/2 O2(g)), the energy flow is:
Into the system from the surroundings; ΔUsys > 0
Out of the system to the surroundings; ΔUsys < 0
No energy exchange; ΔUsys = 0
Only work is done with no energy change
Which situation leads to ΔUsys being negative according to the summary?
Reactants have lower internal energy than products
Reactants and products have equal internal energy
Reactants have higher internal energy than products
The reaction occurs at constant temperature
According to the summary, when reactants have a lower internal energy than the products:
Energy flows out of the system and ΔUsys is negative
Energy flows into the system and ΔUsys is positive
There is no energy flow and ΔUsys is zero
Energy flows cyclically between system and surroundings
Which pair correctly matches the sign of ΔUsys with the direction of energy flow for an exothermic internal energy change shown?
ΔUsys > 0; energy into surroundings
ΔUsys < 0; energy into surroundings
ΔUsys = 0; no flow
ΔUsys > 0; energy into system
Consider two reactions: A has reactants at lower internal energy than its products; B has reactants at higher internal energy than its products. Which statement is correct?
Reaction A: ΔUsys < 0; Reaction B: ΔUsys > 0
Reaction A: ΔUsys > 0; Reaction B: ΔUsys < 0
Both reactions have ΔUsys = 0
Signs of ΔUsys cannot be inferred from relative energies
Which statement correctly describes a negative change in internal energy (ΔUsys < 0) for a system?
The system absorbs heat and has work done on it.
Energy flows out of the system to the surroundings.
Both heat and work are transferred into the system.
The reactants have lower energy than the products.
In which situation is ΔUsys positive?
Reactants have higher internal energy than products, and energy is released.
Products have higher internal energy than reactants, and energy flows into the system.
Heat is released and work is done by the system.
Energy flows out of the system to the surroundings.
A reaction results in products with lower internal energy than reactants. What can you conclude about ΔUsys and the direction of energy flow?
ΔUsys is positive; energy flows into the system.
ΔUsys is negative; energy flows out of the system.
ΔUsys is zero; no energy transfer occurs.
ΔUsys sign cannot be determined from this information.
For a closed system, the first law of thermodynamics is expressed as ΔU = q + w. If 250 J of heat is absorbed by the system and the surroundings do 150 J of work on the system, what is ΔU? Use the sign convention: heat absorbed by the system is positive q; work done on the system is positive w.
−400 J
−100 J
+100 J
+400 J
+250 J
A gas expands against a constant external pressure, doing 85 J of work on the surroundings, while it releases 40 J of heat to the surroundings. Using ΔU = q + w with the sign convention that heat released is negative q and work done by the system is negative w, what is the change in internal energy?
−125 J
−45 J
+45 J
+125 J
Which statement best defines heat capacity (C) for a system?
The mass of a system per degree Celsius
The amount of heat required to change the system’s temperature by 1°C
The average kinetic energy of particles in a system
The temperature at which a system reaches equilibrium
According to the relationship between heat and temperature change, which equation correctly relates the variables?
q = C × ΔT
q = ΔT/C
q = m/ΔT
q = C/ΔT
Which statement describes temperature in this context?
A form of energy transferred due to a temperature difference
A measure that changes only when mass changes
Directly proportional to heat capacity only
Independent of any heat transfer
Specific heat capacity (Cs) is defined as the amount of heat required to raise the temperature of which quantity by 1°C?
1 gram of the substance
1 kilogram of the substance
1 mole of the substance
The entire system regardless of mass
What are the units of specific heat capacity Cs as given?
J g⁻¹ °C
J g °C⁻¹
J g⁻¹ °C⁻¹
J mol⁻¹ °C⁻¹
Which formula correctly computes the heat q absorbed or released by a sample when its temperature changes by ΔT?
q = m × Cs × ΔT
q = m/Cs × ΔT
q = Cs/ΔT
q = m × ΔT/Cs
Molar heat capacity (Cm) is the amount of heat required to raise the temperature of what amount of substance by 1°C?
1 gram
1 kilogram
1 mole
Any mass
Which element listed has the highest specific heat capacity according to the table?
Lead (0.128 J g⁻¹ °C⁻¹)
Silver (0.235 J g⁻¹ °C⁻¹)
Copper (0.385 J g⁻¹ °C⁻¹)
Aluminum (0.903 J g⁻¹ °C⁻¹)
A 50.0 g sample of copper (Cs = 0.385 J g⁻¹ °C⁻¹) is heated causing a temperature increase of 10.0°C. What heat q is absorbed?
19.3 J
193 J
0.77 J
500 J
Two equal-mass metal blocks, aluminum (Cs = 0.903 J g⁻¹ °C⁻¹) and lead (Cs = 0.128 J g⁻¹ °C⁻¹), each receive the same amount of heat. Which block undergoes the larger temperature increase, and why?
Aluminum, because higher Cs means larger ΔT for the same q
Aluminum, because lower Cs means larger ΔT for the same q
Lead, because lower Cs means larger ΔT for the same q
Lead, because higher Cs means larger ΔT for the same q
According to a specific heat capacity reference table, which type of data would you look up to calculate heat absorbed by a substance during a temperature change?
Molar mass
Density
Specific heat capacity
Thermal conductivity
A reference table lists specific heat capacities for elements, compounds, and materials. Which statement best interprets these listings for problem solving?
Use element values only; compounds share the same values as their constituent elements.
Choose the specific heat that matches the exact substance state or composition given in the problem.
Any value in the table can be used interchangeably because specific heat is universal.
Use the highest value in the table to avoid underestimating heat.
You need to compute heat absorbed by a 48.1 g copper sample heated from 20.9°C to 41.1°C. Which formula from specific heat capacity concepts should you apply with the tabulated value for copper?
q = m·c·ΔT
q = m/ΔT
q = c/ΔT
q = m·ΔT/c
A 50.0 g sample of a metal is placed into 100.0 g of water. The metal cools while the water warms until they reach the same temperature. Which statement best represents the heat exchange between the two substances?
q_metal = q_water
q_metal = - q_water
q_system = q_surroundings
Heat flows from the colder water to the hotter metal
Which statement best describes the direction of heat flow between a hot object and a cold object when they are placed in contact?
Heat flows from the cold object to the hot object until both reach the cold object's temperature.
Heat flows from the hot object to the cold object until thermal equilibrium is reached.
No heat flows because energy is conserved.
Heat flows randomly between the objects with no net direction.
A closed system is cooled by its surroundings. If q_sys represents the heat gained by the system and q_surr the heat gained by the surroundings, which relationship must hold?
q_sys = q_surr
q_sys = -q_surr
q_sys > 0 and q_surr > 0
q_sys + q_surr = +q_sys
A 200 g metal block at 90°C is placed into 300 g of water at 20°C in an insulated container until they reach the same final temperature. Which reasoning correctly predicts the sign of q for the metal block?
q_metal is positive because the metal loses heat to the cooler water.
q_metal is negative because the metal loses heat to the cooler water.
q_metal is positive because the system is closed.
q_metal is zero because the container is insulated.
Which expression represents pressure–volume work for a constant external pressure?
w = PΔV
w = -PΔV
w = ΔV/P
w = -ΔV/P
During compression of a gas in a piston where the volume decreases, what is the sign of ΔV and of the work done on the system?
ΔV positive; work negative
ΔV negative; work positive
ΔV negative; work negative
ΔV zero; work zero
A balloon expands against a constant external pressure. Which statement best describes the direction of energy transfer as work?
Work is done on the system and energy enters the gas.
Work is done by the system and energy leaves the gas as work.
No work occurs because pressure is constant.
Work depends only on temperature, not on volume change.
For a piston system initially at volume Vi and finally at volume Vf, which definition of ΔV is used to compute pressure–volume work?
ΔV = Vi − Vf
ΔV = Vf − Vi
ΔV = Vi/Vf
ΔV = Vf/Vi
The volume of a balloon increases from 0.111 L to 1.24 L against a constant external pressure of 1.05 bar. Using w = −PΔV and 1 L·bar = 100 J, what is the work (in J) done by the balloon?
−118 J
+118 J
+113 J
−113 J
In which scenario would the work calculated by w = −PΔV be zero for a gas under a constant external pressure?
The gas is compressed.
The gas expands.
The gas volume does not change.
The external pressure is zero but volume changes.
A sealed cylinder with a movable piston undergoes compression at constant external pressure. Which change would increase the magnitude of the work done on the gas?
Decreasing the external pressure
Increasing the magnitude of the volume decrease
Allowing the gas to expand instead
Holding volume constant
A balloon expands against a constant external pressure. Which expression correctly gives the work done by the system during this expansion?
w = +PΔV
w = −PΔV
w = −Δn/RT
w = +ΔnRT
For an ideal-gas reaction at constant temperature, which expression gives the pressure–volume work per mole of reaction?
w = −PΔV
w = −ΔnRT
w = q − ΔU
w = +ΔnRT
Consider the reaction: C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l). Based on the change in moles of gas, which statement is correct about work direction at 298 K?
Δn > 0, work done by the system
Δn = 0, no pressure–volume work
Δn < 0, work done on the system
Cannot be determined without ΔU
For the gas-phase reaction C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) at 298 K, what is Δn for gases and the sign of w? Use w = −ΔnRT.
Δn = −3; w > 0 (on system)
Δn = 0; w = 0
Δn = +1; w < 0 (by system)
Δn = −1; w > 0 (on system)
Calculate the pressure–volume work (in kJ mol−1) for reaction (a) C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l) at 298 K using w = −ΔnRT with R = 8.314 J mol−1 K−1. Treat only gas species in Δn.
−2.48 kJ mol−1
+2.48 kJ mol−1
+7.44 kJ mol−1
−7.44 kJ mol−1
In constant-volume calorimetry, which statement is correct?
ΔU = q − PΔV; at constant V, ΔU = qv
At constant V, w ≠ 0 because ΔV ≠ 0
The system cannot exchange heat with surroundings
qv equals enthalpy change ΔH
Which statement best describes a bomb calorimeter used in constant volume calorimetry?
A device where a sample burns at constant pressure in an open container
A device where a sample burns in a sealed chamber with constant volume
A device that measures enthalpy change directly from gas expansion work
A device that cools a sample to determine its freezing point
In a bomb calorimeter, which expression correctly relates the internal energy change of reaction to the calorimeter temperature change?
ΔrU = Ccal × ΔT
ΔrU = −Ccal × ΔT
ΔrU = −ΔT/Ccal
ΔrU = qcal
Why does heat released by the reaction in a bomb calorimeter equal the change in internal energy (ΔU)?
Because the reaction occurs at constant pressure with volume work allowed
Because no heat is exchanged with the surroundings
Because the reaction occurs at constant volume so w = 0 and qv = ΔU
Because the calorimeter heat capacity is zero
Which sign convention is correct for the relationship between heat of reaction and heat absorbed by the calorimeter in a bomb calorimeter?
qr = qcal
qr = −qcal
qr = 1/qcal
qr = qcal × ΔT
A 1.52 g sample of glucose (C6H12O6) combusts in a bomb calorimeter. Temperature rises from 20.11°C to 23.64°C. The calorimeter heat capacity Ccal is 6.71 kJ °C−1. What is qr for the reaction?
+23.6 kJ
+6.71 kJ
−23.6 kJ
−6.71 kJ
Using the same experiment, what is ΔrU per mole of glucose? (Molar mass of glucose ≈ 180.16 g mol−1.)
−2.79 kJ mol−1
−279 kJ mol−1
−2.79×103 kJ mol−1
−1.55×103kJmol−1
To report ΔU per mole from a bomb calorimetry experiment, which operation is required after obtaining qr for the sample?
Multiply qr by the number of moles of reactant
Divide qr by the number of moles of reactant
Add Ccal to qr
Subtract ΔT from qr
In a bomb calorimeter experiment for a combustion reaction, which quantity is directly obtained from the measured heat change when the calorimeter’s total heat capacity is known and the volume is constant?
Change in enthalpy (ΔH) of the reaction
Change in internal energy (ΔU) of the reaction
Work done by expanding gases (PΔV)
Heat capacity of the reactants
