Thermal Energy and Calorimetry
| Website: | Young Education |
| Kurs: | Thermal Energy and Gas Laws |
| Buch: | Thermal Energy and Calorimetry |
| Gedruckt von: | Guest user |
| Datum: | Montag, 9. November 2026, 16:39 |
1. Specific Heat Capacity
Learning outcomes
- I can define specific heat capacity and explain what it represents.
- I can explain why different substances heat and cool at different rates.
- I can calculate the thermal energy gained or lost by an object using specific heat capacity.
- I can identify the factors that affect changes in temperature.
- I can apply specific heat concepts to everyday situations and scientific problems.
Specific Heat Capacity
What Is Specific Heat Capacity?
Different substances require different amounts of energy to change their temperatures.
For example, if equal masses of water and metal receive the same amount of thermal energy, the metal will usually experience a much larger temperature increase than the water.
This difference is described using specific heat capacity.
Specific heat capacity is the amount of energy required to raise the temperature of 1 kg of a substance by 1°C (or 1 K).
The symbol for specific heat capacity is:
c
Its SI unit is:
J/(kg·°C) or J/(kg·K)
A material with a high specific heat capacity requires a large amount of energy to change its temperature.
A material with a low specific heat capacity requires less energy to produce the same temperature change.
Comparing Materials
Approximate specific heat capacities include:
| Substance | Specific heat capacity, J/(kg·°C) |
|---|---|
| Water | 4200 |
| Aluminium | 900 |
| Glass | 800 |
| Iron | 450 |
| Copper | 390 |
| Lead | 130 |
These values show that water has a much higher specific heat capacity than common metals.
For example, approximately 4200 J is required to raise the temperature of 1 kg of water by 1°C.
Only about 390 J is required to raise the temperature of 1 kg of copper by the same amount.
Therefore, if equal masses receive the same amount of energy:
copper temperature changes more
water temperature changes less
Why Different Substances Heat at Different Rates
When energy is transferred into a substance, it increases the substance's internal energy.
Different materials store this energy differently because of differences in:
- particle structure
- bonding
- molecular motion
- interactions between particles
As a result, the same amount of transferred energy can produce very different temperature changes in different substances.
A high specific heat capacity means:
more energy required for each degree of temperature change
A low specific heat capacity means:
less energy required for each degree of temperature change
This is why a metal pan can become hot relatively quickly while the water inside it warms more slowly.
Calculating Energy Transfer
The thermal energy gained or lost by an object when its temperature changes can be calculated using:
Q = mcΔT
where:
- Q = thermal energy transferred, in joules (J)
- m = mass, in kilograms (kg)
- c = specific heat capacity, in J/(kg·°C)
- ΔT = temperature change, in °C or K
Temperature change is calculated using:
ΔT = T_final − T_initial
For the magnitude of the energy transferred, we usually use the magnitude of the temperature change.
The same relationship can also be rearranged:
c = Q/(mΔT)
m = Q/(cΔT)
ΔT = Q/(mc)
Exploring the Relationship
The relationship shows that temperature change depends on the energy transferred, mass and specific heat capacity.
For the same amount of energy:
larger mass → smaller temperature change
larger specific heat capacity → smaller temperature change
For the same material and mass:
more energy transferred → larger temperature change
Worked Example: Heating Water
A 2.0 kg sample of water is heated from 20°C to 30°C.
The specific heat capacity of water is:
c = 4200 J/(kg·°C)
First calculate the temperature change:
ΔT = 30 − 20
ΔT = 10°C
Now use:
Q = mcΔT
Q = 2.0 × 4200 × 10
Q = 84,000 J
Therefore:
Q = 84,000 J
or:
Q = 84 kJ
It takes 84 kJ of energy to produce this temperature increase.
Worked Example: Heating Copper
Suppose the same 84,000 J of energy is transferred to 2.0 kg of copper.
For copper:
c = 390 J/(kg·°C)
Use:
ΔT = Q/(mc)
ΔT = 84,000/(2.0 × 390)
ΔT ≈ 108°C
The copper would experience a temperature increase of approximately:
108°C
Compare this with the water, which increased by only 10°C.
This demonstrates the enormous effect of specific heat capacity.
High and Low Specific Heat Capacity
High Specific Heat Capacity
Materials with high specific heat capacity:
- require more energy to warm
- release more energy as they cool through the same temperature range
- tend to experience smaller temperature changes for a given energy transfer
Water is an important example.
Low Specific Heat Capacity
Materials with low specific heat capacity:
- require less energy to warm
- require less energy to cool through the same temperature range
- experience larger temperature changes for a given energy transfer
Many metals have relatively low specific heat capacities.
Factors Affecting Temperature Change
From:
ΔT = Q/(mc)
we can identify three major factors affecting the temperature change.
Energy Transferred
More energy produces a greater temperature change when mass and material remain constant.
larger Q → larger ΔT
Mass
A larger mass requires more energy to produce the same temperature change.
larger m → smaller ΔT
for the same energy transfer.
Specific Heat Capacity
A material with a larger specific heat capacity requires more energy for the same temperature change.
larger c → smaller ΔT
for the same mass and energy transfer.
Example: Different Masses of Water
Suppose 42,000 J is transferred into two samples of water.
Sample A
m = 1.0 kg
c = 4200 J/(kg·°C)
ΔT = 42,000/(1.0 × 4200)
ΔT = 10°C
Sample B
m = 2.0 kg
ΔT = 42,000/(2.0 × 4200)
ΔT = 5°C
Doubling the mass halves the temperature change when the energy transferred remains the same.
Water and Climate
Water's high specific heat capacity has major effects on Earth's climate.
Oceans can absorb enormous quantities of thermal energy while experiencing relatively small temperature changes.
They can also release large amounts of energy as they cool.
As a result, large bodies of water help moderate temperature changes.
Coastal areas often experience smaller temperature variations than inland areas because nearby oceans heat and cool relatively slowly.
Land generally changes temperature more rapidly.
Land and Sea Breezes
Different specific heat capacities also contribute to local wind patterns.
During the day:
- land warms relatively quickly
- water warms more slowly
- air above the warmer land heats and rises
- cooler air from over the water moves inland
This can produce a sea breeze.
At night:
- land cools relatively quickly
- water remains relatively warmer
- the temperature difference can reverse
This can contribute to a land breeze.
Specific heat capacity therefore connects directly with atmospheric convection.
Water in Heating and Cooling Systems
Water's high specific heat capacity makes it useful as a heat-transfer fluid.
Water can absorb a large amount of energy without undergoing an extreme temperature increase.
It is used in applications such as:
- heating systems
- some industrial cooling systems
- engine cooling systems
- thermal energy storage
A useful coolant should be capable of absorbing substantial thermal energy while remaining within a practical temperature range.
Cooking and Specific Heat Capacity
Specific heat capacity also affects cooking.
A metal saucepan has a relatively low specific heat capacity compared with water.
The pan therefore heats relatively quickly.
The water inside requires considerably more energy to produce the same temperature increase.
Once hot, however, the water can store a large quantity of thermal energy.
This is useful when boiling or simmering food because the water's temperature does not change extremely rapidly when relatively small amounts of energy are gained or lost.
The Human Body
Much of the human body consists of water.
Water's high specific heat capacity helps reduce rapid changes in body temperature.
A substantial amount of energy must be transferred to or from the body to produce a large temperature change.
This contributes to maintaining a relatively stable internal temperature.
The body also regulates temperature through mechanisms including:
- sweating
- blood circulation
- changes in metabolic activity
- changes in blood flow near the skin
Thermal Energy Storage
Materials with high specific heat capacities can be useful for thermal energy storage.
During periods when excess thermal energy is available, a material can absorb energy.
Later, the material can release that energy as it cools.
Water is often useful because it is:
- inexpensive
- widely available
- relatively safe
- capable of storing substantial thermal energy
Large hot-water tanks can therefore act as thermal energy stores.
Experimental Measurement of Specific Heat Capacity
Specific heat capacity can be measured experimentally.
One common investigation uses:
- a solid metal block
- an electrical heater
- a thermometer or temperature sensor
- an ammeter
- a voltmeter
- insulation
If an electrical heater supplies energy, the electrical energy transferred can be estimated using:
Q = Pt
where:
- Q = energy transferred, in J
- P = power, in W
- t = time, in s
If voltage and current are measured:
P = VI
so:
Q = VIt
The specific heat capacity can then be calculated using:
c = Q/(mΔT)
Example: Determining Specific Heat Capacity
A 1.5 kg metal block receives 27,000 J of energy.
Its temperature rises from 20°C to 60°C.
First calculate:
ΔT = 60 − 20
ΔT = 40°C
Use:
c = Q/(mΔT)
c = 27,000/(1.5 × 40)
c = 450 J/(kg·°C)
Therefore:
c = 450 J/(kg·°C)
This value is close to the specific heat capacity of iron.
Energy Loss in Experiments
In a real experiment, not all electrical energy supplied by a heater necessarily enters the sample.
Some energy may be transferred to:
- the surrounding air
- the container
- the thermometer
- connecting wires
- other equipment
This can introduce experimental error.
Insulation can reduce these unwanted transfers.
A lid may also help when heating liquids by reducing convection and evaporation from the surface.
Specific Heat Capacity vs Thermal Conductivity
Specific heat capacity should not be confused with thermal conductivity.
Specific Heat Capacity
Describes how much energy is required to change a material's temperature.
Thermal Conductivity
Describes how quickly thermal energy can move through a material.
A material can have:
- high thermal conductivity and low specific heat capacity
- low thermal conductivity and high specific heat capacity
- or other combinations
These properties describe different aspects of thermal behaviour.
Specific Heat Capacity vs Latent Heat
Specific heat capacity applies when energy transfer causes a temperature change without a phase change.
During melting or boiling, energy can be transferred without changing temperature.
That energy is associated with latent heat rather than specific heat capacity.
For example:
heating liquid water from 20°C to 80°C → specific heat capacity
boiling water at constant temperature → latent heat
This distinction is essential when analyzing heating and cooling processes involving phase changes.
Heating Curves
A heating curve shows how temperature changes as energy is supplied.
During regions where the substance remains in one state:
energy transferred → temperature increases
Specific heat capacity determines how steeply temperature changes.
During a phase change:
energy transferred → state changes
while the temperature may remain approximately constant.
This shows why energy transfer and temperature change are not always the same thing.
Common Mistakes
Mistake: "Specific heat capacity is the amount of heat an object contains."
Specific heat capacity is a property describing how much energy is required to change the temperature of a unit mass.
Mistake: "A high specific heat capacity means the material heats quickly."
It means more energy is required to change its temperature, so under comparable heating conditions its temperature generally changes more slowly.
Mistake: "Q = mcT."
The equation uses the change in temperature, not simply the final temperature:
Q = mcΔT
Mistake: "Temperature change is final temperature plus initial temperature."
Use:
ΔT = T_final − T_initial
Mistake: "Mass can be entered in grams."
When c is given in J/(kg·°C), mass must be converted to kilograms.
Mistake: "Specific heat capacity and thermal conductivity mean the same thing."
Specific heat capacity describes energy required for temperature change. Thermal conductivity describes the rate at which energy moves through a material.
Key Terms
Specific heat capacity — energy required to raise the temperature of 1 kg of a substance by 1°C or 1 K.
Thermal energy transfer — transfer of energy caused by thermal processes.
Temperature change — difference between final and initial temperatures.
Mass — amount of matter in an object, measured in kilograms for SI calculations.
Joule (J) — SI unit of energy.
Thermal conductivity — measure of how effectively a material conducts thermal energy.
Thermal equilibrium — condition in which systems have the same temperature and there is no net heat transfer between them.
Thermal energy storage — storing energy by increasing the internal energy of a material for later use.
Latent heat — energy transferred during a phase change without an accompanying temperature change.
Heating curve — graph showing how the temperature or state of a substance changes as energy is supplied.
Key Takeaways
- Specific heat capacity measures how much energy is required to change the temperature of a substance.
- Its common SI unit is J/(kg·°C) or J/(kg·K).
- Materials with high specific heat capacity require more energy for the same temperature change.
- Water has a particularly high specific heat capacity compared with many common materials.
- Thermal energy transferred during a temperature change can be calculated using Q = mcΔT.
- Temperature change depends on the energy transferred, mass and specific heat capacity.
- Increasing mass reduces the temperature change produced by a fixed amount of energy.
- Increasing specific heat capacity also reduces the temperature change produced by a fixed amount of energy.
- Water's high specific heat capacity helps moderate coastal climates and makes water useful in heating and cooling systems.
- Specific heat capacity and thermal conductivity describe different properties.
- During a phase change, energy can be transferred without changing temperature, so latent heat must be considered instead.
- Real experiments must account for unwanted energy transfers to the surroundings.
Check Your Understanding
- Define specific heat capacity.
- What does it mean if a substance has a high specific heat capacity?
- Why does water usually warm more slowly than an equal mass of copper receiving energy at the same rate?
- Calculate the energy required to heat 3.0 kg of water from 20°C to 40°C. Use c = 4200 J/(kg·°C).
- A 2.0 kg metal block receives 18,000 J and its temperature increases by 20°C. Calculate its specific heat capacity.
- Two equal masses receive the same amount of energy. Material A has c = 400 J/(kg·°C), while material B has c = 800 J/(kg·°C). Which experiences the larger temperature change? Explain.
- Explain why oceans help reduce extreme temperature changes in coastal areas.
- Why is water useful as a coolant?
- Explain the difference between specific heat capacity and thermal conductivity.
- A 500 g aluminium block has c = 900 J/(kg·°C). Calculate the energy needed to increase its temperature by 30°C.
- Why might an experimental value for specific heat capacity differ from the accepted value?
- Explain why Q = mcΔT cannot by itself describe the energy required to boil water that is already at its boiling point.
2. Latent Heat and Phase Changes
Learning outcomes
- I can distinguish between sensible heating and latent heating.
- I can explain why temperature remains constant during a phase change.
- I can describe the processes of melting, freezing, boiling, condensation, sublimation, and deposition.
- I can calculate the energy required for a phase change using latent heat.
- I can relate evaporation and condensation to latent heat transfer.
Latent Heat and Phase Changes
Sensible Heating and Latent Heating
When thermal energy is transferred to or from a substance, two different things can happen:
- its temperature can change
- its state can change
These processes are described as sensible heating and latent heating.
Sensible heating occurs when energy transfer changes the temperature of a substance without changing its state.
For example:
20°C liquid water → 60°C liquid water
The water remains liquid, but its temperature increases.
The energy transferred can be calculated using:
Q = mcΔT
Latent heating occurs when energy is transferred during a change of state without changing the temperature.
For example:
0°C ice → 0°C liquid water
Energy is absorbed, but the temperature remains constant while the ice melts.
Why Temperature Remains Constant During a Phase Change
Temperature is related to the average kinetic energy of particles.
During ordinary heating within a single state:
energy added → average kinetic energy increases → temperature increases
During a phase change, however, the transferred energy is used mainly to change the potential energy associated with interactions between particles.
For example, when a solid melts, energy is required to overcome some of the forces holding particles in their fixed arrangement.
Therefore:
energy is transferred
but:
average kinetic energy does not increase significantly
so:
temperature remains constant during the phase change for a pure substance at constant pressure.
Once the phase change is complete, further energy transfer can again increase the temperature.
Changes of State
Matter can change between solid, liquid and gas.
The six major changes are:
| Phase change | Change |
|---|---|
| Melting | Solid → Liquid |
| Freezing | Liquid → Solid |
| Vaporization | Liquid → Gas |
| Condensation | Gas → Liquid |
| Sublimation | Solid → Gas |
| Deposition | Gas → Solid |
The complete cycle can be explored here:
Some changes require energy to be absorbed, while others release energy to the surroundings.
Melting
Melting is the change:
solid → liquid
During melting, the substance absorbs energy.
In a solid, particles vibrate around relatively fixed positions. As energy is supplied, particle motion increases until the melting temperature is reached.
During melting itself, additional energy helps overcome interactions maintaining the solid structure.
The particles become able to move past one another, forming a liquid.
For pure water at standard atmospheric pressure, melting occurs at approximately:
0°C
During the melting process:
temperature remains approximately 0°C until the ice has melted
even though energy continues to enter the system.
Freezing
Freezing is the reverse process:
liquid → solid
During freezing, energy is transferred from the substance to the surroundings.
Particles become more organized into the structure of the solid.
For pure water at standard atmospheric pressure:
liquid water at 0°C → ice at 0°C
The temperature remains approximately constant while freezing occurs.
This means freezing water releases energy even though its temperature is not decreasing during the actual phase change.
Boiling
Boiling is a form of vaporization in which:
liquid → gas
throughout the liquid.
At the boiling point, bubbles of vapor form inside the liquid and rise to the surface.
Energy supplied during boiling is used to separate particles and move them into the gas phase.
At standard atmospheric pressure, pure water boils at approximately:
100°C
Once boiling begins, continuing to supply energy does not necessarily make the liquid water hotter.
Instead, the energy converts more liquid into vapor.
Only after all the liquid has vaporized can additional heating increase the temperature of the vapor.
Condensation
Condensation is the reverse of vaporization:
gas → liquid
During condensation, particles move closer together and energy is released to the surroundings.
For example, water vapor can condense into droplets on a cold surface.
The water vapor releases latent heat as it condenses.
This is important in:
- cloud formation
- weather systems
- refrigerators
- air conditioners
- distillation
Sublimation
Sublimation occurs when a substance changes directly from:
solid → gas
without first becoming liquid.
Examples include:
- dry ice
- some solid air fresheners
- iodine under suitable conditions
- ice and snow slowly disappearing in cold, dry conditions
During sublimation, the substance absorbs energy.
Dry ice is solid carbon dioxide. At ordinary atmospheric pressure, it sublimates rather than melting into liquid carbon dioxide.
Deposition
Deposition is the reverse of sublimation:
gas → solid
without passing through the liquid state.
During deposition, energy is released.
A familiar example is frost.
Under suitable conditions, water vapor in the air can form ice directly on a cold surface.
water vapor → ice
Energy Absorbed and Released
The six phase changes can be divided according to the direction of energy transfer.
Energy Absorbed
These changes require the substance to gain energy:
solid → liquid
melting
liquid → gas
vaporization
solid → gas
sublimation
The particles become less strongly bound and generally more widely separated.
Energy Released
These changes release energy:
gas → liquid
condensation
liquid → solid
freezing
gas → solid
deposition
The particles form stronger interactions as they move into more condensed states.
What Is Latent Heat?
The word latent means hidden.
Latent heat refers to energy transferred during a phase change without a corresponding temperature change.
The energy required for a phase change depends on:
- the mass of the substance
- the substance involved
- the type of phase change
The relationship is:
Q = mL
where:
- Q = energy transferred, in J
- m = mass, in kg
- L = specific latent heat, in J/kg
Specific Latent Heat
Specific latent heat is the energy required to change the state of 1 kg of a substance without changing its temperature.
There are different types of specific latent heat.
Specific Latent Heat of Fusion
The specific latent heat of fusion is associated with:
solid ↔ liquid
For water:
L_f ≈ 334,000 J/kg
This means approximately 334,000 J is required to melt 1 kg of ice at 0°C without changing its temperature.
Specific Latent Heat of Vaporization
The specific latent heat of vaporization is associated with:
liquid ↔ gas
For water near its normal boiling point:
L_v ≈ 2,260,000 J/kg
A much larger amount of energy is required to vaporize water than to melt the same mass.
Worked Example: Melting Ice
Calculate the energy required to melt 2.0 kg of ice at 0°C.
For water:
L_f = 334,000 J/kg
Use:
Q = mL
Q = 2.0 × 334,000
Q = 668,000 J
Therefore:
Q = 668,000 J
or:
Q = 668 kJ
Notice that the ice remains at approximately 0°C during the melting process.
Worked Example: Boiling Water
Calculate the energy required to convert 0.50 kg of water at 100°C completely into steam at 100°C.
Use:
L_v = 2,260,000 J/kg
Q = mL
Q = 0.50 × 2,260,000
Q = 1,130,000 J
Therefore:
Q = 1.13 × 10⁶ J
or:
Q = 1.13 MJ
This large value explains why converting liquid water into vapor requires substantial energy.
Worked Example: Finding Mass
A heater transfers 167,000 J of energy to ice at its melting point.
If:
L_f = 334,000 J/kg
calculate the mass of ice melted.
Rearrange:
m = Q/L
Substitute:
m = 167,000 / 334,000
m = 0.50 kg
Therefore, 0.50 kg of ice melts.
Heating Curves
A heating curve shows how the temperature of a substance changes as energy is supplied.
A typical heating curve contains sloping and flat regions.
Sloping Regions
The substance remains in one state.
Energy increases the average kinetic energy of the particles.
Therefore:
temperature increases
This is sensible heating.
The energy can be calculated using:
Q = mcΔT
Flat Regions
A phase change is occurring.
Energy changes the arrangement and interactions of particles rather than increasing their average kinetic energy.
Therefore:
temperature remains constant
This is latent heating.
The energy can be calculated using:
Q = mL
Multi-Stage Heating Problems
Some problems require both specific heat capacity and latent heat.
Suppose ice at −10°C is heated until it becomes liquid water at 20°C.
There are three stages.
Heating the Ice
ice at −10°C → ice at 0°C
Use:
Q₁ = mc_iceΔT
Melting the Ice
ice at 0°C → water at 0°C
Use:
Q₂ = mL_f
Heating the Water
water at 0°C → water at 20°C
Use:
Q₃ = mc_waterΔT
The total energy is:
Q_total = Q₁ + Q₂ + Q₃
This approach is essential whenever a problem crosses one or more phase changes.
Evaporation and Latent Heat
Evaporation occurs when energetic particles escape from the surface of a liquid.
To escape, these particles require energy.
This energy is associated with the latent heat of vaporization.
When higher-energy particles escape:
- energy leaves the liquid
- the average kinetic energy of the remaining particles decreases
- the liquid cools
This is evaporative cooling.
Sweating and Latent Heat
Sweating is an important application of latent heat.
Sweat spreads across the skin and evaporates.
The phase change is:
liquid water → water vapor
Energy is required for this change.
That energy is transferred partly from the skin.
Therefore:
skin loses thermal energy → skin cools
The effectiveness of sweating depends on whether the sweat actually evaporates.
Sweat that simply runs off the skin provides much less evaporative cooling.
Why Humidity Matters
Evaporation occurs more slowly when the surrounding air already contains a large amount of water vapor.
Therefore:
high humidity → slower evaporation → less cooling
In dry air, sweat can evaporate more rapidly.
This allows more latent heat to be transferred away from the body.
Wind can also increase evaporation by removing humid air from close to the skin.
Condensation and Latent Heat
Condensation does the opposite of evaporation.
During:
gas → liquid
latent heat is released.
This is extremely important in the atmosphere.
When water vapor condenses to form cloud droplets, energy is released into the surrounding air.
This released energy can contribute to atmospheric convection and the development of powerful weather systems.
Latent heat therefore plays an important role in Earth's weather and climate.
Refrigeration and Air Conditioning
Refrigerators and air conditioners deliberately use evaporation and condensation.
A refrigerant circulates through the system.
Inside the cooled region:
liquid → gas
The refrigerant evaporates and absorbs latent heat.
Outside the cooled region:
gas → liquid
The refrigerant condenses and releases latent heat.
A compressor and other components allow this cycle to operate continuously.
This allows thermal energy to be transferred from a cooler region to warmer surroundings using external energy.
Steam and Thermal Energy
Steam can transfer very large amounts of energy because water has a high specific latent heat of vaporization.
When steam condenses:
steam → liquid water
a large quantity of latent heat is released.
This makes steam useful in:
- industrial heating
- power stations
- sterilization
- some heating systems
It also makes steam burns potentially serious because condensation can transfer substantial energy to the skin.
Melting Ice as a Cooling Method
Ice is effective at keeping food and drinks cold for two reasons.
First, ice is cold.
More importantly, melting ice absorbs substantial latent heat.
At approximately 0°C:
ice + energy → liquid water
The energy used to melt the ice does not initially raise its temperature.
Instead, it changes the state.
This allows ice to absorb a considerable amount of energy while remaining close to 0°C.
Sensible Heat vs Latent Heat
| Sensible Heating | Latent Heating |
|---|---|
| Temperature changes | Temperature remains constant during the phase change |
| State remains the same | State changes |
| Mainly changes particle kinetic energy | Mainly changes particle potential energy/interactions |
| Use Q = mcΔT | Use Q = mL |
| Example: water 20°C → 60°C | Example: ice 0°C → water 0°C |
Both involve energy transfer, but the effect of that energy is different.
Common Mistakes
Mistake: "If energy is being added, temperature must increase."
During a phase change, energy can be absorbed while temperature remains constant.
Mistake: "Particles stop moving during a phase change."
Particles continue moving. The transferred energy is primarily changing their interactions and arrangement.
Mistake: "Latent heat is stored heat that cannot be measured."
Latent heat refers to energy transferred during a phase change without a temperature change.
Mistake: "Boiling and evaporation are identical."
Both involve liquid becoming gas, but boiling occurs throughout the liquid while evaporation occurs at the surface.
Mistake: "Condensation absorbs latent heat."
Condensation releases latent heat.
Mistake: "Use Q = mcΔT during melting."
During the phase change itself, use:
Q = mL
Mistake: "Temperature difference is needed in Q = mL."
There is no ΔT in the latent heat equation because the temperature remains constant during the ideal phase change.
Key Terms
Sensible heating — energy transfer that produces a temperature change without a phase change.
Latent heating — energy transfer associated with a phase change without a temperature change.
Phase change — physical change from one state of matter to another.
Melting — change from solid to liquid.
Freezing — change from liquid to solid.
Vaporization — change from liquid to gas.
Boiling — vaporization occurring throughout a liquid.
Evaporation — vaporization occurring at the surface of a liquid.
Condensation — change from gas to liquid.
Sublimation — direct change from solid to gas.
Deposition — direct change from gas to solid.
Specific latent heat — energy required per kilogram to change the state of a substance without changing its temperature.
Specific latent heat of fusion — energy per kilogram associated with the solid–liquid phase change.
Specific latent heat of vaporization — energy per kilogram associated with the liquid–gas phase change.
Heating curve — graph showing how temperature and state change as energy is supplied.
Key Takeaways
- Sensible heating changes temperature without changing state.
- Latent heating changes state without changing temperature during the phase change.
- Sensible heating calculations use Q = mcΔT.
- Phase-change calculations use Q = mL.
- Temperature remains constant during an ideal phase change because energy changes particle interactions and potential energy rather than average kinetic energy.
- Melting, vaporization and sublimation require energy to be absorbed.
- Freezing, condensation and deposition release energy.
- Specific latent heat is measured in J/kg.
- Water has a large specific latent heat of vaporization, so vaporizing or condensing water transfers substantial amounts of energy.
- Evaporation removes energy from a liquid and can produce cooling.
- Sweating uses evaporation and latent heat to cool the body.
- Condensation releases latent heat and is important in cloud formation and weather.
- Refrigerators and air conditioners use evaporation and condensation to transfer thermal energy.
- Multi-stage heating problems may require both Q = mcΔT and Q = mL.
Check Your Understanding
- Explain the difference between sensible heating and latent heating.
- Why does the temperature of melting ice remain approximately constant even while energy is being supplied?
- Identify the phase change: solid → gas.
- Which phase changes release energy to the surroundings?
- Calculate the energy required to melt 0.75 kg of ice at 0°C if L_f = 334,000 J/kg.
- Calculate the energy required to vaporize 0.20 kg of water at 100°C if L_v = 2,260,000 J/kg.
- Explain why evaporation causes cooling.
- Why is sweating less effective at cooling the body when humidity is high?
- Explain why condensation releases energy to the surroundings.
- Why can steam transfer a large amount of energy when it condenses?
- Explain why ice can keep a drink close to 0°C for a considerable time while it melts.
- A sample of ice starts below 0°C and is heated until it becomes liquid water above 0°C. Identify the stages that require Q = mcΔT and the stage that requires Q = mL.
3. Heating and Cooling Curves
Learning outcomes
- I can interpret heating and cooling curves.
- I can identify regions where temperature changes and where phase changes occur.
- I can explain how particle motion changes throughout a heating or cooling process.
- I can relate energy input to changes in temperature and state.
- I can predict the behavior of substances during heating and cooling.
Heating and Cooling Curves
What Is a Heating Curve?
A heating curve is a graph showing how the temperature of a substance changes as thermal energy is added.
A typical heating curve follows a substance as it changes through:
solid → liquid → gas
The graph usually contains two types of regions:
- sloping regions, where temperature changes
- flat regions, where a phase change occurs
Heating curves are useful because they show that adding energy does not always cause the temperature to increase.
Reading the Axes
A heating curve normally has:
Vertical axis: temperature
Horizontal axis: energy added or heating time
If the heater supplies energy at a constant rate, time can be used as an indication of the amount of energy transferred.
Moving from left to right means that more energy is being supplied to the substance.
The Five Main Regions
A complete heating curve often contains five regions:
solid warming → melting → liquid warming → boiling → gas warming
These regions alternate between sloping and flat sections.
| Region | State | Temperature | What happens to energy? |
|---|---|---|---|
| Solid warming | Solid | Increases | Particle kinetic energy increases |
| Melting | Solid + liquid | Constant | Potential energy increases |
| Liquid warming | Liquid | Increases | Particle kinetic energy increases |
| Boiling | Liquid + gas | Constant | Potential energy increases |
| Gas warming | Gas | Increases | Particle kinetic energy increases |
Understanding this pattern makes heating curves much easier to interpret.
Solid Warming
At the beginning of a typical heating curve, the substance is completely solid.
Particles in a solid:
- are closely packed
- remain in relatively fixed positions
- vibrate around those positions
As thermal energy is supplied, the particles vibrate more vigorously.
Their average kinetic energy increases.
Therefore:
energy added → kinetic energy increases → temperature increases
The graph slopes upward.
No phase change occurs during this section.
Reaching the Melting Point
Eventually, the solid reaches its melting point.
At this temperature, additional energy begins changing the substance from solid to liquid.
The important observation is:
energy continues to enter the substance, but its temperature remains approximately constant.
The heating curve becomes flat.
The Melting Plateau
During melting:
solid → liquid
Both solid and liquid are present.
The transferred energy is used mainly to overcome some of the attractive interactions maintaining the solid structure.
Instead of substantially increasing average particle kinetic energy, the energy increases the potential energy associated with particle arrangement.
Therefore:
average kinetic energy remains approximately constant
so:
temperature remains constant
This flat section is called a plateau.
Latent Heat During Melting
The energy transferred during melting is called latent heat.
The energy required can be calculated using:
Q = mL_f
where:
- Q = energy transferred, in J
- m = mass, in kg
- L_f = specific latent heat of fusion, in J/kg
No ΔT appears in this equation because the temperature does not change during the ideal phase change.
Liquid Warming
Once all the solid has melted, the substance is completely liquid.
Additional energy now increases the average kinetic energy of the liquid particles.
The particles:
- move faster
- collide more frequently and energetically
- have greater average kinetic energy
Therefore the temperature begins increasing again.
The heating curve slopes upward.
Energy transferred during this region can be calculated using:
Q = mcΔT
Reaching the Boiling Point
Eventually, the liquid reaches its boiling point.
Additional energy now begins converting the liquid into gas.
The temperature again stops increasing while the phase change takes place.
The graph enters another flat region.
The Boiling Plateau
During boiling:
liquid → gas
Both liquid and gas are present.
Energy continues to enter the system, but it is used mainly to overcome attractive interactions between particles.
The particles become much more widely separated as they enter the gas phase.
Therefore:
energy added → potential energy increases
while:
average kinetic energy remains approximately constant → temperature remains constant
Latent Heat During Boiling
The energy required to vaporize a substance can be calculated using:
Q = mL_v
where:
- Q = energy transferred
- m = mass
- L_v = specific latent heat of vaporization
For water, the specific latent heat of vaporization is much larger than its specific latent heat of fusion.
This means considerably more energy is required to vaporize 1 kg of water than to melt 1 kg of ice at their respective phase-change temperatures.
Gas Warming
Once all the liquid has become gas, additional energy again increases particle kinetic energy.
Gas particles:
- move faster
- have greater average kinetic energy
- collide more energetically
The gas temperature therefore increases.
The heating curve once again slopes upward.
Slopes and Plateaus
A simple rule helps when interpreting heating curves:
Sloping line
Temperature is changing.
The substance remains primarily in one state.
Use:
Q = mcΔT
Flat line
A phase change is occurring.
Temperature remains approximately constant.
Use:
Q = mL
So:
slope = temperature change
plateau = phase change
Particle Motion Along a Heating Curve
Heating curves can also be interpreted using the particle model.
Solid warming
Particles vibrate faster.
Kinetic energy increases.
Melting
Particles begin leaving their fixed arrangement.
Potential energy increases.
Liquid warming
Particles move faster.
Kinetic energy increases.
Boiling
Particles separate substantially and enter the gas phase.
Potential energy increases.
Gas warming
Particles move increasingly rapidly.
Kinetic energy increases.
This gives an important relationship:
temperature change → mainly kinetic energy change
phase change → mainly potential energy change
A Heating Curve for Water
Consider water being heated from ice below 0°C to steam above 100°C at standard atmospheric pressure.
The sequence is:
ice below 0°C
Temperature rises.
↓
ice + water at approximately 0°C
Melting occurs while temperature remains approximately constant.
↓
liquid water from approximately 0°C to 100°C
Temperature rises.
↓
water + steam at approximately 100°C
Boiling occurs while temperature remains approximately constant.
↓
steam above 100°C
Temperature rises.
This assumes approximately standard atmospheric pressure.
Energy Calculations Across a Heating Curve
A complete heating problem may require several calculations.
Suppose ice begins at −20°C and ends as steam above 100°C.
The total energy cannot be found using one single Q = mcΔT calculation.
Instead, divide the process into stages.
Heating the Ice
Q₁ = mc_iceΔT
Melting the Ice
Q₂ = mL_f
Heating the Liquid Water
Q₃ = mc_waterΔT
Vaporizing the Water
Q₄ = mL_v
Heating the Steam
Q₅ = mc_steamΔT
Then:
Q_total = Q₁ + Q₂ + Q₃ + Q₄ + Q₅
Each section of the graph represents a different physical process.
What Does the Length of a Plateau Mean?
If energy is being supplied at a constant rate, a longer plateau means more energy is required for that phase change.
For water, the boiling plateau is typically much longer than the melting plateau under comparable conditions because:
L_v > L_f
Therefore, vaporizing water requires considerably more energy per kilogram than melting ice.
The horizontal length of a heating curve must be interpreted carefully, however. It represents energy directly only if the horizontal axis is energy, or if energy is being supplied at a constant rate when the axis is time.
What Does the Slope Tell Us?
During a region without a phase change:
Q = mcΔT
Rearranging:
ΔT = Q/(mc)
For the same mass and the same rate of energy input, a substance with a lower specific heat capacity experiences a greater temperature change.
Its heating curve will therefore rise more steeply.
A substance with a higher specific heat capacity experiences a smaller temperature change for the same energy input.
Its curve will be less steep.
Therefore, the slope can provide information about the thermal properties of a substance.
What Is a Cooling Curve?
A cooling curve shows what happens as thermal energy is transferred from a substance to its surroundings.
The processes on a cooling curve are essentially the reverse of those on a heating curve.
For a substance beginning as a gas:
gas → liquid → solid
The temperature decreases during single-state regions and remains approximately constant during phase changes.
Gas Cooling
Initially, the substance may be entirely gaseous.
As energy leaves the gas:
- particles move more slowly
- average kinetic energy decreases
- temperature decreases
The graph slopes downward.
Eventually, the gas reaches its condensation temperature.
Condensation Plateau
During condensation:
gas → liquid
Energy continues to leave the substance.
However, the temperature remains approximately constant while the phase change occurs.
Particles move closer together and form stronger intermolecular attractions.
Energy is released to the surroundings.
This is the reverse of vaporization.
Vaporization absorbs latent heat.
Condensation releases latent heat.
Liquid Cooling
Once condensation is complete, the substance is entirely liquid.
As more energy leaves:
- particles move more slowly
- average kinetic energy decreases
- temperature decreases
The cooling curve slopes downward again.
Eventually, the freezing point is reached.
Freezing Plateau
During freezing:
liquid → solid
Energy continues to leave the substance, but its temperature remains approximately constant.
Particles become arranged into the more ordered structure of the solid.
Latent heat is released to the surroundings.
The process is the reverse of melting.
Melting absorbs latent heat.
Freezing releases latent heat.
Solid Cooling
Once freezing is complete, the substance is entirely solid.
Further energy loss reduces the vibrational motion of the particles.
Average kinetic energy decreases.
Therefore:
temperature decreases
The graph slopes downward again.
Heating vs Cooling
| Heating | Cooling |
|---|---|
| Energy enters substance | Energy leaves substance |
| Solid warms | Gas cools |
| Melting absorbs energy | Condensation releases energy |
| Liquid warms | Liquid cools |
| Boiling absorbs energy | Freezing releases energy |
| Gas warms | Solid cools |
The same phase-change temperatures appear on ideal heating and cooling curves for a pure substance at the same pressure.
Predicting What Happens
Heating and cooling curves allow us to predict the state of a substance.
Suppose a substance has:
melting point = 40°C
boiling point = 120°C
Then, under the specified pressure:
Below 40°C:
solid
At 40°C while melting:
solid + liquid
Between 40°C and 120°C:
liquid
At 120°C while boiling:
liquid + gas
Above 120°C:
gas
This information can often be determined directly from a heating curve.
Example: Identifying a Phase Change
Suppose a substance is heated continuously.
Its temperature increases from 20°C to 65°C.
It then remains at 65°C for several minutes even though energy continues to be supplied.
Afterward, the temperature begins increasing again.
The flat section indicates a phase change.
If the substance began as a solid, then 65°C is its approximate melting point under those conditions.
During the plateau:
solid + liquid are present
Example: Predicting Particle Behaviour
Suppose a liquid is heated but has not yet reached its boiling point.
As energy enters:
- particle average kinetic energy increases
- particles move faster
- temperature increases
Once boiling begins:
- the average kinetic energy remains approximately constant
- particles continue separating into the gas phase
- potential energy increases
- temperature remains approximately constant
This explains why a temperature graph changes from a slope to a plateau.
Evaporation vs a Heating-Curve Plateau
Evaporation should not automatically be confused with the boiling plateau.
Evaporation can occur from the surface of a liquid at temperatures below its boiling point.
Boiling occurs throughout the liquid at its boiling point for the given pressure.
A standard heating curve usually shows the boiling phase change, not ordinary surface evaporation.
Heating Curves and Pressure
Melting and boiling points can depend on pressure.
This is especially important for boiling.
At lower atmospheric pressure, liquids generally boil at lower temperatures.
For example, water boils below 100°C at sufficiently high altitudes.
At higher pressure, its boiling temperature increases.
Therefore, a heating curve should be interpreted with the pressure conditions in mind.
Heating Curves in Everyday Life
Heating and cooling curves help explain many familiar processes.
Melting Ice
Ice can absorb considerable energy while remaining close to 0°C as it melts.
Boiling Water
Water can continue absorbing energy at its boiling point without becoming significantly hotter while liquid water remains.
Freezing Food
Energy must be removed not only to cool food to its freezing point but also to complete the phase change.
Refrigeration
Refrigerants repeatedly evaporate and condense, absorbing and releasing latent heat.
Industrial Processing
Engineers use heating and cooling curves when controlling:
- metal processing
- food production
- chemical manufacturing
- refrigeration
- material testing
Common Mistakes
Mistake: "Every part of a heating curve slopes upward."
During phase changes, temperature remains approximately constant and the curve forms a plateau.
Mistake: "No energy is being transferred during a flat section."
Energy is still being transferred. It is changing the state rather than the temperature.
Mistake: "Particles stop moving during a plateau."
Particles continue moving. Their average kinetic energy remains approximately constant during the ideal phase change.
Mistake: "Temperature measures all the internal energy of a substance."
Temperature is related to average kinetic energy. Internal energy also includes potential energy associated with particle interactions.
Mistake: "Use Q = mcΔT everywhere on a heating curve."
Use Q = mcΔT when temperature changes and Q = mL during a phase change.
Mistake: "A flat line always means melting."
It could represent melting, boiling, freezing or condensation depending on whether the substance is being heated or cooled and which states are present.
Key Terms
Heating curve — graph showing how temperature and state change as energy is supplied to a substance.
Cooling curve — graph showing how temperature and state change as energy is removed.
Plateau — flat region of a heating or cooling curve where a phase change occurs at approximately constant temperature.
Melting point — temperature at which solid and liquid phases coexist during melting or freezing under specified pressure.
Boiling point — temperature at which boiling occurs at a specified external pressure.
Kinetic energy — energy associated with particle motion.
Potential energy — energy associated with particle positions and interactions.
Specific heat capacity — energy required to raise the temperature of 1 kg of a substance by 1°C or 1 K.
Specific latent heat — energy transferred per kilogram during a phase change without a temperature change.
Latent heat of fusion — energy associated with the solid–liquid phase change.
Latent heat of vaporization — energy associated with the liquid–gas phase change.
Condensation — change from gas to liquid.
Freezing — change from liquid to solid.
Key Takeaways
- Heating and cooling curves show how temperature and state change as energy is transferred.
- Sloping sections represent temperature changes within a single state.
- Flat sections represent phase changes.
- During a sloping section, average particle kinetic energy changes.
- During an ideal phase change, temperature and average kinetic energy remain approximately constant while particle potential energy changes.
- Melting and boiling absorb latent heat.
- Freezing and condensation release latent heat.
- Use Q = mcΔT when temperature changes without a phase change.
- Use Q = mL during a phase change.
- A longer plateau represents a larger energy transfer if the horizontal scale represents energy or energy is supplied at a constant rate.
- The slope of a heating curve can be related to specific heat capacity when mass and heating rate are controlled.
- Heating and cooling processes are essentially reverses of one another.
- Heating curves can be used to identify melting points, boiling points and the physical state of a substance.
- A complete heating or cooling problem may require several separate energy calculations.
Check Your Understanding
- What does a sloping section of a heating curve represent?
- What does a flat section represent?
- Why does temperature remain approximately constant while a solid melts?
- Describe how particle motion changes while a solid is heated toward its melting point.
- What happens to particle potential energy during melting?
- Why does the temperature begin rising again after melting is complete?
- What two states are present during the boiling plateau?
- Explain the difference between what happens to energy during a sloping region and during a plateau.
- Which equation should be used during a temperature-changing region: Q = mcΔT or Q = mL?
- Which equation should be used during a phase change?
- Describe what happens to particles as a gas cools and condenses.
- A substance is heated at a constant rate and remains at 80°C for several minutes before its temperature begins rising again. Explain what is happening during the constant-temperature region.
4. Calorimetry
Learning outcomes
- I can explain the purpose of calorimetry.
- I can use the principle of conservation of energy in thermal systems.
- I can calculate unknown temperatures or energy transfers using calorimetry.
- I can identify sources of experimental error in calorimetry investigations.
- I can interpret calorimetry results in practical situations.
Calorimetry
What Is Calorimetry?
Calorimetry is the measurement of energy transferred as heat during physical or chemical processes.
A device used to make these measurements is called a calorimeter.
Calorimetry can be used to investigate:
- specific heat capacities
- energy released by fuels
- energy contained in foods
- energy changes during chemical reactions
- final temperatures when substances are mixed
- latent heat during phase changes
The central idea is the conservation of energy:
Energy cannot be created or destroyed; it can only be transferred or transformed.
Energy Conservation in Calorimetry
Consider a hot metal block placed into cooler water.
Thermal energy flows:
hot metal → cooler water
The metal loses energy while the water gains energy.
In an ideal insulated system:
energy lost by hot object = energy gained by cold object
This relationship forms the basis of many calorimetry calculations.
Another useful way to express this is:
Q_lost + Q_gained = 0
when signs are included.
In many school calculations, however, we simply equate the magnitudes:
Q_lost = Q_gained
Reaching Thermal Equilibrium
When two objects at different temperatures are placed in thermal contact, energy transfers from the warmer object to the cooler object.
Eventually they reach the same temperature.
This condition is called thermal equilibrium.
For example:
hot metal + cool water → common final temperature
At thermal equilibrium, there is no longer any net heat transfer between them.
The final temperature should lie between the two initial temperatures if there are no additional energy sources or phase changes.
The Main Calorimetry Equation
When a substance changes temperature without changing state, the energy transferred is:
Q = mcΔT
where:
- Q = energy transferred, in J
- m = mass, in kg
- c = specific heat capacity, in J/(kg·°C)
- ΔT = temperature change, in °C or K
Remember:
ΔT = T_final − T_initial
When comparing energy lost and gained, it is often easier to use the magnitude of each temperature change.
A Simple Mixing Experiment
Suppose hot water is mixed with cold water inside an insulated container.
The hot water transfers energy to the cold water.
Ideally:
energy lost by hot water = energy gained by cold water
Using Q = mcΔT:
m_hot c_hot ΔT_hot = m_cold c_cold ΔT_cold
If both samples are water, their specific heat capacities are equal and can cancel:
m_hot ΔT_hot = m_cold ΔT_cold
This makes the final temperature relatively easy to calculate.
Worked Example: Mixing Water
100 g of water at 80°C is mixed with 200 g of water at 20°C.
Assume no energy is transferred to the surroundings or container.
Let the final temperature be T_f.
Because both substances are water:
energy lost by hot water = energy gained by cold water
Convert masses:
100 g = 0.100 kg
200 g = 0.200 kg
Set up the equation:
0.100c(80 − T_f) = 0.200c(T_f − 20)
The c values cancel:
0.100(80 − T_f) = 0.200(T_f − 20)
Expand:
8 − 0.1T_f = 0.2T_f − 4
12 = 0.3T_f
T_f = 40°C
Therefore, the final equilibrium temperature is:
40°C
Checking the Answer
A useful check is to compare the calculated final temperature with the initial temperatures.
Initially:
hot water = 80°C
cold water = 20°C
Calculated:
final temperature = 40°C
Since:
20°C < 40°C < 80°C
the answer is physically reasonable.
If the calculation had produced 100°C or −10°C, something would probably be wrong under the stated assumptions.
Different Materials
Calorimetry becomes more interesting when different substances are involved.
Suppose a hot metal block is placed in cooler water.
Because metal and water have different specific heat capacities, we use:
m_metal c_metal ΔT_metal = m_water c_water ΔT_water
If the metal's specific heat capacity is unknown, the equation can be rearranged to determine it.
This is a common laboratory method for measuring specific heat capacity.
Measuring the Specific Heat Capacity of a Metal
A typical experiment involves:
- measuring the mass of a metal sample
- heating the metal to a known temperature
- measuring the mass and initial temperature of water
- quickly placing the hot metal into the water
- insulating the container
- measuring the final equilibrium temperature
The energy lost by the metal is assumed to equal the energy gained by the water.
Worked Example: Finding Specific Heat Capacity
A 0.20 kg metal block at 100°C is placed into 0.50 kg of water at 20°C.
The final temperature is 25°C.
For water:
c_water = 4200 J/(kg·°C)
First calculate the energy gained by the water.
ΔT_water = 25 − 20
ΔT_water = 5°C
Use:
Q = mcΔT
Q_water = 0.50 × 4200 × 5
Q_water = 10,500 J
Assuming no energy is lost:
Q_metal = 10,500 J
The metal cools from 100°C to 25°C:
ΔT_metal = 100 − 25
ΔT_metal = 75°C
Now:
10,500 = 0.20 × c_metal × 75
10,500 = 15c_metal
c_metal = 700 J/(kg·°C)
Calorimeters
A good calorimeter attempts to minimize unwanted energy transfer between the experiment and the surroundings.
Simple calorimeters may use:
- insulated foam cups
- lids
- thermometers
- temperature probes
- insulating material
More advanced calorimeters use carefully controlled and highly insulated systems.
No real calorimeter is perfectly insulated, so experimental results always contain some uncertainty.
Why Foam Cups Are Useful
Foam contains many small pockets of trapped gas.
Both the foam and trapped gas are poor thermal conductors.
The small pockets also restrict convection.
Therefore, a foam cup reduces energy transfer between the contents and surroundings.
A lid further reduces:
- convection
- evaporation
- exchange of warm or cool air
This helps the experiment behave more like an isolated thermal system.
Calorimetry and Chemical Reactions
Calorimetry can also measure energy changes during chemical reactions.
Suppose two solutions react inside an insulated cup.
If the reaction releases energy:
chemical energy → thermal energy
The temperature of the solution rises.
If the reaction absorbs energy from the surroundings, the temperature may decrease.
By measuring the temperature change, the energy transferred to or from the solution can be estimated.
Exothermic Reactions
An exothermic reaction transfers energy to the surroundings.
In a simple calorimeter:
reaction releases energy → solution gains energy → temperature rises
Examples can include:
- combustion
- some neutralization reactions
- some oxidation reactions
The energy gained by the solution can be estimated using:
Q = mcΔT
where m and c refer to the solution being heated.
Endothermic Reactions
An endothermic reaction absorbs energy from the surroundings.
In a calorimeter:
reaction absorbs energy → solution loses thermal energy → temperature decreases
Some dissolving processes and chemical reactions are endothermic.
Calorimetry can therefore help determine whether a process releases or absorbs energy.
Measuring Energy in Food
Food contains chemical energy.
A simple experiment can estimate the energy released when food burns.
The food is burned underneath a known mass of water.
Energy released by the food is transferred to the water.
The water's temperature increase is measured.
Then:
Q_water = m_water c_water ΔT
This provides an estimate of the energy transferred from the burning food.
Worked Example: Food Calorimetry
A burning food sample heats 0.20 kg of water from 20°C to 45°C.
For water:
c = 4200 J/(kg·°C)
First:
ΔT = 45 − 20
ΔT = 25°C
Then:
Q = mcΔT
Q = 0.20 × 4200 × 25
Q = 21,000 J
Therefore, the water gained:
21,000 J
or:
21 kJ
If the food sample lost 0.50 g of mass while burning, the measured energy released per gram would be:
21 kJ ÷ 0.50 g = 42 kJ/g
However, this simple experiment usually underestimates the actual energy content of the food.
Why Simple Food Calorimetry Is Inaccurate
Not all the energy released by the burning food reaches the water.
Energy can be transferred to:
- surrounding air
- the container
- the thermometer
- the clamp and equipment
- nearby surfaces
The food may also undergo incomplete combustion.
As a result:
energy measured by water < total energy released by food
A simple classroom food calorimeter therefore usually produces a value lower than the accepted energy content.
Bomb Calorimeters
A bomb calorimeter is a more sophisticated instrument used to measure energy released during combustion.
The sample burns in a sealed chamber surrounded by water.
The system is carefully insulated.
Bomb calorimeters provide more accurate measurements because:
- combustion can be more complete
- energy transfer is more carefully controlled
- the calorimeter itself can be calibrated
- heat losses are minimized
Calorimetry and Phase Changes
Calorimetry can also involve phase changes.
If ice is placed into warm water, energy transfers from the water to the ice.
Some energy may:
- warm the ice to its melting point
- melt the ice
- warm the resulting liquid water
Different equations may therefore be required.
For temperature changes:
Q = mcΔT
For phase changes:
Q = mL
Energy conservation still applies:
total energy lost = total energy gained
Example: Ice in Warm Water
Suppose ice at 0°C is placed into warm water.
The warm water cools.
The ice absorbs energy and melts.
If all the ice melts:
energy lost by warm water = energy used to melt ice + energy used to warm melted ice
This might require:
Q_water = Q_melting + Q_warming
or:
m_water c_water ΔT = m_ice L_f + m_ice c_water ΔT
Calorimetry problems can therefore combine specific heat capacity and latent heat.
Finding an Unknown Final Temperature
Many calorimetry problems ask for the final equilibrium temperature.
The general strategy is:
Step 1: Identify which object loses energy.
Step 2: Identify which object gains energy.
Step 3: Write an energy equation.
Step 4: Substitute Q = mcΔT where appropriate.
Step 5: Solve for the unknown final temperature.
Step 6: Check that the answer is physically reasonable.
A useful starting equation is:
energy lost = energy gained
Sources of Experimental Error
Real calorimetry experiments are not perfectly isolated.
Several sources of error can affect the results.
Energy Transfer to the Surroundings
Energy may escape through:
- conduction
- convection
- radiation
Better insulation reduces these losses.
Energy Absorbed by the Calorimeter
The cup, thermometer and other equipment also absorb energy.
Ignoring this can produce inaccurate results.
Evaporation
Evaporation removes energy from liquids and may change their mass.
A lid can reduce this effect.
Temperature Measurement
Possible problems include:
- thermometer uncertainty
- slow sensor response
- incorrect reading
- measuring before equilibrium is reached
Incomplete Mixing
Different regions of a liquid may have different temperatures.
Stirring gently helps produce a more uniform temperature.
Delays
Energy may be lost while transferring a heated object into the calorimeter.
Performing the transfer quickly reduces this problem.
Improving a Calorimetry Experiment
Accuracy can be improved by:
- using better insulation
- using a lid
- transferring samples quickly
- stirring consistently
- using accurate temperature probes
- recording temperature continuously
- accounting for the heat capacity of the calorimeter
- reducing evaporation
- repeating measurements
- averaging reliable trials
A good experimental design attempts to make the real system behave as closely as possible to the ideal model used in the calculations.
Calorimeter Heat Capacity
In more accurate experiments, the calorimeter itself cannot be ignored.
The calorimeter may gain or lose energy according to:
Q_cal = C_calΔT
where:
- Q_cal = energy transferred to the calorimeter
- C_cal = heat capacity of the calorimeter, in J/°C
- ΔT = temperature change
The energy equation might then become:
Q_hot = Q_cold + Q_cal
This produces a more realistic model.
Interpreting Experimental Results
Calorimetry results should not simply be accepted without evaluation.
Suppose the accepted specific heat capacity of a material is:
900 J/(kg·°C)
but an experiment gives:
760 J/(kg·°C)
The difference suggests that the assumptions of the experiment were not completely satisfied.
Possible explanations include:
- energy lost to the surroundings
- inaccurate temperature measurements
- inaccurate mass measurements
- energy absorbed by equipment
- insufficient time to reach equilibrium
Scientists should consider both the calculated value and the quality of the experimental evidence.
Percentage Error
Experimental values can be compared with accepted values using percentage error:
percentage error = |experimental − accepted| / accepted × 100%
For example:
experimental value = 760 J/(kg·°C)
accepted value = 900 J/(kg·°C)
Difference:
|760 − 900| = 140
Percentage error:
140/900 × 100%
≈ 15.6%
This provides a useful measure of how far the experimental result is from the accepted value.
Practical Applications of Calorimetry
Calorimetry is used in many scientific and technological fields.
Food Science
Determining the energy released by foods.
Chemistry
Measuring energy changes during reactions.
Materials Science
Determining specific heat capacities and thermal properties.
Engineering
Evaluating heating, cooling and energy-storage systems.
Environmental Science
Studying thermal properties of materials and systems.
Fuel Testing
Comparing the energy released by different fuels.
Calorimetry therefore connects thermal physics with chemistry, biology, engineering and environmental science.
Common Mistakes
Mistake: "The hot object loses temperature."
An object loses energy. Its temperature decreases as a result.
Mistake: "Energy lost and gained are always exactly equal in a real experiment."
Total energy is conserved, but some may transfer to the container or surroundings rather than the object being measured.
Mistake: "The final temperature can be higher than both starting temperatures."
For simple mixing without an additional energy source or reaction, the equilibrium temperature should lie between the initial temperatures.
Mistake: "Use grams directly in Q = mcΔT."
If c is given in J/(kg·°C), convert mass to kilograms.
Mistake: "The calorimeter does not absorb energy."
Real calorimeters have heat capacity and may need to be included in accurate calculations.
Mistake: "A result different from the accepted value means the experiment failed."
Experimental differences can provide useful information about uncertainty, heat loss and limitations in the procedure.
Key Terms
Calorimetry — measurement of energy transferred as heat during physical or chemical processes.
Calorimeter — device used to measure thermal energy transfers.
Conservation of energy — principle stating that energy cannot be created or destroyed, only transferred or transformed.
Thermal equilibrium — state in which objects have the same temperature and there is no net heat transfer between them.
Specific heat capacity — energy required to raise the temperature of 1 kg of a substance by 1°C or 1 K.
Heat capacity — energy required to raise the temperature of an entire object by 1°C or 1 K.
Exothermic — describing a process that transfers energy to the surroundings.
Endothermic — describing a process that absorbs energy from the surroundings.
Latent heat — energy transferred during a phase change without a temperature change.
Experimental error — difference between an experimental result and the accepted or actual value resulting from limitations in measurement or method.
Percentage error — comparison between experimental and accepted values expressed as a percentage.
Key Takeaways
- Calorimetry measures energy transfers in thermal systems.
- Calorimetry is based on the conservation of energy.
- In an ideal insulated system, energy lost = energy gained.
- Objects exchanging energy eventually reach thermal equilibrium.
- For temperature changes, use Q = mcΔT.
- For phase changes, use Q = mL.
- Calorimetry can determine unknown temperatures, specific heat capacities and energy transfers.
- Chemical reactions can be studied by measuring the temperature changes they produce.
- Food calorimetry estimates the chemical energy released by food.
- Real calorimeters are not perfectly isolated.
- Energy transferred to the container and surroundings can affect experimental results.
- Insulation, lids, rapid transfer and accurate temperature measurements improve calorimetry experiments.
- More accurate calculations may include the heat capacity of the calorimeter itself.
- Calorimetry has applications in physics, chemistry, food science and engineering.
Check Your Understanding
- What is the purpose of calorimetry?
- Explain how conservation of energy is used in a calorimetry experiment.
- A hot metal block is placed into cooler water. Which object gains energy and which loses energy?
- Why should the final equilibrium temperature normally lie between the two initial temperatures?
- Calculate the energy required to heat 0.50 kg of water from 20°C to 35°C. Use c = 4200 J/(kg·°C).
- A hot object loses 12,000 J in an ideal calorimeter. How much energy does the cooler substance gain?
- Explain why foam cups are often used in simple calorimetry experiments.
- Why can ignoring the calorimeter's own heat capacity introduce error?
- A food sample heats 0.10 kg of water by 30°C. Calculate the energy gained by the water.
- Explain why a simple food-calorimetry experiment usually underestimates the energy released by the food.
- Give three ways a calorimetry experiment could be made more accurate.
- A student obtains a specific heat capacity substantially different from the accepted value. Explain two experimental factors that could have caused the difference.
5. Thermal Energy Applications
Learning outcomes
- I can explain how thermal energy principles are used in cooking, heating, and cooling systems.
- I can describe how engineers use materials with different thermal properties.
- I can evaluate methods of improving thermal efficiency.
- I can explain the importance of thermal energy in natural and industrial processes.
- I can apply thermal energy concepts to solve real-world problems.
Thermal Energy Applications
Thermal Energy in Everyday Life
Thermal energy principles are involved whenever energy is transferred because of a temperature difference.
They are important in:
- cooking
- home heating
- refrigeration
- air conditioning
- transportation
- electronics
- manufacturing
- power generation
- weather and climate
- biological systems
Engineers often face one of two challenges:
increase heat transfer when it is useful
or
reduce heat transfer when it is unwanted
Understanding conduction, convection, radiation, specific heat capacity, latent heat and insulation allows systems to be designed for these different purposes.
Controlling Thermal Energy Transfer
There are three major methods of heat transfer.
Conduction
Energy transfers through direct interactions between particles.
It is particularly important in solids.
Convection
Energy is transferred through the movement of fluids such as liquids and gases.
Radiation
Energy is transferred by electromagnetic waves and does not require matter.
Most practical thermal systems involve more than one method at the same time.
For example, when cooking soup:
stove → conduction through pan → convection through soup
while radiation and convection also transfer energy from the stove and pan to the surroundings.
Choosing Materials
Engineers select materials according to their thermal properties.
Important properties include:
- thermal conductivity
- specific heat capacity
- melting point
- boiling point
- ability to reflect or absorb radiation
- thermal expansion
- resistance to high temperatures
The best material depends on what the system is designed to accomplish.
Thermal Conductors
Materials with high thermal conductivity transfer energy rapidly.
Metals such as:
- copper
- aluminium
- steel
are widely used when efficient conduction is required.
Applications include:
- saucepans
- heat exchangers
- radiators
- heat sinks
- boilers
- engine components
Copper has excellent thermal conductivity, but it can be more expensive and heavier than aluminium.
Engineers therefore consider more than just thermal performance.
Thermal Insulators
Materials with low thermal conductivity are used when energy transfer should be reduced.
Examples include:
- foam
- fibreglass
- mineral wool
- wood
- plastics
- cork
- trapped air
Applications include:
- building insulation
- refrigerator walls
- cooler boxes
- oven gloves
- saucepan handles
- thermal clothing
An insulating material does not completely stop heat transfer. It reduces the rate at which energy is transferred.
Cooking
Cooking involves transferring thermal energy into food.
Different cooking methods use different combinations of conduction, convection and radiation.
Frying
Energy transfers:
stove → pan → food
Conduction is particularly important.
Boiling
Energy conducts through the saucepan and then convection currents distribute energy through the water.
Baking
Hot air transfers energy by convection while hot oven surfaces transfer energy by radiation.
Conduction transfers energy through the food itself.
Grilling
Infrared radiation from the heating element or flame is particularly important.
Why Cooking Pans Use Metals
The base of a saucepan should transfer thermal energy efficiently.
Metals are useful because they generally have:
- high thermal conductivity
- high temperature resistance
- good mechanical strength
Aluminium and copper are particularly good thermal conductors.
However, the handle has a different purpose.
A metal handle could conduct energy toward the user's hand.
Handles are therefore often made from materials such as:
- plastic
- wood
- silicone
One device can therefore deliberately contain both conductors and insulators.
Convection Ovens
A convection oven uses a fan to circulate hot air.
Moving air increases the rate of energy transfer between the air and food.
This can:
- produce more even heating
- reduce cooking time
- reduce temperature differences inside the oven
This is an example of forced convection.
The same principle is used in many engineering cooling systems.
Induction Cooking
Induction cooktops use changing electromagnetic fields to produce electrical currents in suitable cookware.
Electrical resistance within the cookware produces heating.
This can be efficient because energy is produced directly in the cookware rather than first heating a conventional heating element and then transferring energy to the pan.
However, induction cooking requires compatible cookware containing suitable magnetic materials.
This illustrates an important engineering idea:
thermal efficiency often depends on both the energy-transfer method and the materials selected.
Heating Buildings
Heating systems transfer energy into buildings to maintain comfortable indoor temperatures.
Common systems include:
- radiators
- underfloor heating
- heat pumps
- electric heaters
- furnaces
A radiator heats nearby air.
The warmer air becomes less dense and rises.
Cooler air moves toward the radiator and is heated.
This establishes convection currents.
Radiators also transfer energy by radiation.
Underfloor Heating
Underfloor heating distributes energy across a large surface area.
The floor warms the air immediately above it.
The warmer air rises and produces convection.
The large heated area can provide relatively uniform room temperatures.
It can sometimes operate effectively using lower water temperatures than traditional radiators, depending on the system and building.
Heat Pumps
A heat pump transfers thermal energy from one location to another.
It can move energy:
outside → inside
for heating
or:
inside → outside
for cooling.
Heat pumps use a refrigerant that undergoes evaporation and condensation.
During evaporation, the refrigerant absorbs latent heat.
During condensation, it releases latent heat.
Electrical energy is required to operate the compressor and other components.
A heat pump can deliver more thermal energy to a building than the electrical energy it consumes because it moves existing thermal energy rather than converting all of its input electricity directly into heat.
Refrigerators
A refrigerator also uses a heat-pump cycle.
Its purpose is to transfer energy:
inside refrigerator → surrounding room
A refrigerant circulates through the system.
Inside the refrigerator, the refrigerant evaporates and absorbs energy.
Outside, it condenses and releases energy.
This explains why the coils at the back or underneath a refrigerator can become warm.
The refrigerator does not "create cold."
Instead, it removes thermal energy from its interior.
Air Conditioning
An air conditioner works using similar principles.
Energy is absorbed from indoor air and transferred outside.
The system involves:
- evaporation
- condensation
- compression
- convection
- conduction
Fans increase convection across the heat exchangers.
The indoor evaporator becomes cool and absorbs energy.
The outdoor condenser becomes warm and releases energy.
Insulation and Thermal Efficiency
Heating and cooling systems become less efficient when large amounts of thermal energy are transferred unintentionally.
Building insulation reduces these unwanted transfers.
Useful methods include:
- roof insulation
- wall insulation
- double or triple glazing
- draught proofing
- reflective surfaces
- insulated doors
- appropriate shading
The best combination depends on the climate and building design.
Double Glazing
Double-glazed windows contain two panes separated by a sealed gap.
The gap may contain air or a gas such as argon.
The trapped gas:
- conducts energy poorly
- prevents large convection currents from developing
Some windows also use special coatings to reduce radiative energy transfer.
Triple glazing uses another pane and another insulating layer.
Thermal Efficiency
Thermal efficiency describes how effectively a system uses or controls thermal energy for its intended purpose.
For many energy-conversion systems:
efficiency = useful energy output / total energy input × 100%
For example, suppose a heating system receives:
10,000 J
and transfers:
8,500 J
usefully into the intended space.
Efficiency = 8500 / 10,000 × 100%
Efficiency = 85%
The remaining energy has been transferred elsewhere.
Improving Thermal Efficiency
Thermal efficiency can be improved by reducing unwanted energy transfers.
Possible methods include:
- increasing insulation
- reducing air leakage
- using reflective surfaces
- recovering waste heat
- improving heat exchangers
- controlling temperatures more accurately
- reducing unnecessary heating or cooling
- maintaining equipment properly
However, improvements should also be evaluated in terms of:
- cost
- safety
- durability
- environmental impact
- maintenance
- expected energy savings
The most expensive solution is not automatically the best solution.
Heat Exchangers
A heat exchanger transfers thermal energy between fluids while usually keeping the fluids physically separate.
Examples are found in:
- car radiators
- refrigerators
- air conditioners
- power stations
- industrial chemical plants
- building heating systems
Effective heat exchangers often use:
- highly conductive materials
- thin walls
- large surface areas
- moving fluids
- large temperature differences
These features increase the rate of heat transfer.
Surface Area and Cooling
Increasing surface area can increase heat transfer.
This is why radiators and heat sinks often contain many thin fins.
The fins provide a large surface area in contact with the surrounding air.
This increases energy transfer by convection.
Radiation may also increase because more surface is exposed.
Cooling Electronics
Computers and electronic equipment generate thermal energy during operation.
Excessive temperature can damage components or reduce their performance.
A processor may use:
processor → thermal paste → heat sink → air
Thermal paste improves thermal contact between surfaces.
The heat sink conducts energy away from the processor.
A fan moves air across the fins and increases convection.
High-performance systems may also use liquid cooling.
Car Engine Cooling
Combustion engines generate large amounts of thermal energy.
If this energy is not removed, engine components can overheat.
A typical cooling system uses:
- coolant
- a pump
- a radiator
- a fan
- a thermostat
Coolant absorbs energy from the engine and carries it to the radiator.
The radiator transfers the energy to the surrounding air.
Its large surface area and thin metal fins increase heat transfer.
Specific Heat Capacity in Engineering
The specific heat capacity of a material affects how quickly its temperature changes.
Materials with high specific heat capacity can absorb substantial amounts of energy without experiencing extremely large temperature changes.
Water is particularly useful because of its high specific heat capacity.
Applications include:
- engine cooling
- industrial cooling
- heating systems
- thermal energy storage
For a temperature change:
Q = mcΔT
This equation allows engineers to determine how much energy a material can absorb or release.
Thermal Energy Storage
Thermal energy does not always need to be used immediately.
It can be stored for later use.
For example, hot-water tanks can store energy.
Because water has a high specific heat capacity, a large tank can store considerable thermal energy.
Suppose 200 kg of water is heated by 30°C.
Using:
c = 4200 J/(kg·°C)
Q = mcΔT
Q = 200 × 4200 × 30
Q = 25,200,000 J
Therefore:
Q = 25.2 MJ
This demonstrates why water can be useful for thermal energy storage.
Latent Heat in Technology
Phase changes can transfer large quantities of energy without large temperature changes.
This is used in:
- refrigerators
- air conditioners
- heat pumps
- thermal storage systems
- industrial cooling systems
During evaporation:
energy is absorbed
During condensation:
energy is released
The energy involved is:
Q = mL
Phase-change materials can therefore be useful when a system needs to absorb or release substantial energy within a narrow temperature range.
Thermal Energy in Power Stations
Many power stations use thermal energy to produce electricity.
A simplified sequence is:
energy source → thermal energy → steam → turbine → generator → electrical energy
Water is heated to produce high-temperature steam.
The steam drives a turbine.
The turbine drives a generator.
After leaving the turbine, the steam is cooled and condensed so that the water can be reused.
Heat exchangers and cooling systems are therefore essential parts of many power stations.
Waste Heat
Industrial processes often produce thermal energy that would otherwise be transferred to the environment.
This waste heat may sometimes be recovered.
Possible uses include:
- heating buildings
- heating water
- preheating industrial materials
- generating additional electricity
- supplying energy to another industrial process
Recovering waste heat can improve the overall efficiency of a system.
Thermal Energy in Nature
Thermal energy is also fundamental to natural processes.
Solar radiation heats Earth's surface unevenly.
This produces temperature and density differences in the atmosphere and oceans.
These differences contribute to:
- convection currents
- winds
- ocean circulation
- weather systems
- evaporation
- cloud formation
Thermal energy therefore plays an important role in Earth's weather and climate.
Evaporation and the Water Cycle
Solar energy causes water to evaporate from:
- oceans
- lakes
- rivers
- soil
- plants
Evaporation absorbs latent heat.
Water vapor can later rise and cool.
When it condenses:
gas → liquid
latent heat is released into the atmosphere.
This energy can contribute to convection and weather-system development.
Thermal Energy in Living Organisms
Living organisms must also manage thermal energy.
Humans regulate body temperature using processes including:
- sweating
- changes in blood flow
- convection
- radiation
- evaporation
Sweating is particularly effective because evaporation requires latent heat.
Energy transferred from the skin allows sweat to evaporate, cooling the body.
Animals use many other thermal adaptations, including:
- fur
- feathers
- body fat
- panting
- behavioral changes
Real-World Problem: Keeping a Drink Cold
Suppose you want to design a container that keeps a cold drink cold for several hours.
You should reduce all major methods of heat transfer.
Reduce Conduction
Use low-conductivity materials.
Reduce Convection
Use sealed spaces or a vacuum.
Reduce Radiation
Use reflective surfaces.
Reduce Energy Transfer Through the Opening
Use an insulated lid and narrow opening.
A vacuum flask combines several of these methods.
This illustrates an important engineering principle:
good thermal design often controls several heat-transfer mechanisms simultaneously.
Real-World Problem: Cooling a Computer
Suppose a processor is overheating.
Possible improvements include:
- using a more conductive heat sink
- increasing fin surface area
- improving thermal contact
- adding a fan
- increasing airflow
- using liquid cooling
- reducing unnecessary power consumption
Simply adding insulation would make the problem worse because the goal is to remove thermal energy.
Thermal engineering always begins by identifying the desired direction of energy transfer.
Real-World Problem: Making a Building More Efficient
Suppose a building requires excessive energy for air conditioning.
Possible improvements include:
- improving roof insulation
- installing better windows
- reducing air leakage
- using reflective roofing
- shading windows
- improving air-conditioning efficiency
The best solution should be evaluated using both thermal performance and practical factors.
For example, an improvement costing $5,000 that saves $1,000 per year may be more attractive than one costing $30,000 that saves only slightly more energy.
Evaluating Thermal Solutions
When evaluating a thermal technology, ask:
Does it improve the required heat transfer?
How much energy does it save?
How much does it cost?
How long will it last?
Is it safe?
What maintenance is required?
What environmental effects does it have?
A technically effective solution may not always be the most practical solution.
Engineering involves balancing these competing factors.
Common Mistakes
Mistake: "Thermal efficiency always means reducing heat transfer."
Some systems, such as radiators and heat sinks, are designed to increase heat transfer.
Mistake: "A refrigerator creates cold."
A refrigerator transfers thermal energy from its interior to the surrounding environment.
Mistake: "Insulation completely stops heat transfer."
Insulation reduces the rate of heat transfer.
Mistake: "Only one method of heat transfer occurs in a system."
Most practical systems involve conduction, convection and radiation simultaneously.
Mistake: "Metals are always the best materials."
Metals are useful when high thermal conductivity is needed, but insulators are better when energy transfer should be reduced.
Mistake: "Efficiency can exceed 100%."
For an energy-conversion efficiency defined as useful output divided by total input, it cannot exceed 100%. Heat-pump performance is commonly described using a different measure, the coefficient of performance, which can be greater than 1.
Key Terms
Thermal energy — part of a system's internal energy associated with the microscopic motion and interactions of its particles.
Thermal conductivity — measure of how effectively a material conducts thermal energy.
Thermal insulator — material that reduces the rate of energy transfer by conduction.
Thermal efficiency — measure of how effectively a system uses or controls energy for its intended purpose.
Heat exchanger — device that transfers thermal energy between fluids, usually without allowing them to mix.
Forced convection — convection produced or increased by a fan, pump or similar device.
Heat sink — device designed to transfer thermal energy away from a component.
Refrigerant — substance circulated through a refrigeration system to absorb and release energy, often through phase changes.
Heat pump — device that uses external energy to transfer thermal energy from a cooler region to a warmer region.
Specific heat capacity — energy required to raise the temperature of 1 kg of a substance by 1°C or 1 K.
Latent heat — energy transferred during a phase change without a temperature change.
Thermal energy storage — storing energy in a material for later use.
Waste heat — thermal energy produced by a process that is not part of its intended useful output.
Thermal equilibrium — condition in which systems are at the same temperature and there is no net heat transfer between them.
Key Takeaways
- Thermal energy principles are essential in cooking, heating, cooling, engineering, industry and natural systems.
- Engineers choose materials according to properties such as thermal conductivity, specific heat capacity and temperature resistance.
- Conductors are useful when rapid heat transfer is required.
- Insulators are useful when unwanted heat transfer must be reduced.
- Cooking usually involves combinations of conduction, convection and radiation.
- Refrigerators and air conditioners transfer thermal energy rather than "creating cold."
- Heat pumps move thermal energy and can provide efficient building heating and cooling.
- Building insulation reduces unwanted energy transfers and can reduce energy consumption.
- Heat exchangers are designed to transfer thermal energy efficiently between fluids.
- Heat sinks and radiators use conductive materials and large surface areas to increase cooling.
- Water's high specific heat capacity makes it useful for cooling and thermal energy storage.
- Evaporation and condensation allow large amounts of energy to be transferred through latent heat.
- Industrial systems can improve efficiency by recovering waste heat.
- Thermal energy drives important natural processes including atmospheric convection, evaporation, condensation and weather.
- Good thermal engineering requires balancing energy performance, cost, safety, durability and environmental impact.
Check Your Understanding
- Why are metals commonly used in saucepan bases?
- Why are saucepan handles often made from plastic or wood?
- Explain how a convection oven increases the rate of energy transfer to food.
- Explain how a refrigerator transfers thermal energy from its interior to the room.
- Why does insulation reduce the energy required to heat or cool a building?
- Give two design features that can increase the effectiveness of a heat exchanger.
- Explain why water is useful in many cooling systems.
- Calculate the energy required to heat 50 kg of water by 20°C. Use c = 4200 J/(kg·°C).
- Explain how evaporation and condensation are used in an air-conditioning system.
- Why do computer heat sinks contain many thin fins?
- Give two ways waste heat from an industrial process could be put to useful purposes.
- A building is overheating because of strong sunlight and poor insulation. Suggest three changes that could reduce the cooling energy required, and explain the thermal principle behind each.