Difficulty: Introductory | Prerequisites: Basic understanding of states of matter, temperature measurement
Heating and cooling curves plot temperature against time as a substance gains or loses energy. The slope tells you how fast the temperature changes, and a flat plateau means the substance is changing phase (melting, freezing or boiling) while the temperature holds steady. Knowing the shape of these curves lets you identify phase-change temperatures and calculate how much energy a substance absorbs or releases.
Heating curve
A graph of temperature (y-axis) versus time (x-axis) recorded while a substance is being heated at a steady rate. In simple terms, it shows what happens to the temperature as you keep adding heat.
Cooling curve
The reverse: temperature versus time as a substance loses heat to its surroundings. Think of it as a heating curve flipped, running downhill instead of uphill.
Specific heat capacity (c)
The amount of energy (in joules) needed to raise the temperature of one gram of a substance by one degree Celsius. For liquid water, c = 4.18 J/g°C. In simple terms, it measures how "stubborn" a substance is about changing temperature.
Phase change / change of state
The physical transformation from one state of matter to another (solid to liquid, liquid to gas, or the reverse). During a phase change, added or removed energy goes into breaking or forming intermolecular forces rather than changing the temperature.
Boiling point
The temperature at which a liquid's vapour pressure equals the surrounding atmospheric pressure, causing the liquid to convert to gas throughout its bulk. For water at standard pressure this is 100 °C (the lab data showed 99.9 °C, which is typical of real measurements).
Freezing point
The temperature at which a liquid solidifies. For a pure substance, this is the same temperature as the melting point. In the lab, salol's freezing point was identified as approximately 40 °C from its cooling curve plateau.
Plateau (on a heating or cooling curve)
The flat, horizontal section of the curve where temperature remains constant even though energy is still being added or removed. This is the signature of a phase change in progress. Think of it as the substance "spending" all incoming (or outgoing) energy on changing state rather than getting hotter (or cooler).
Heat of fusion (enthalpy of fusion)
The energy required to convert one gram of a solid into a liquid at its melting point, without any temperature change. For water, the accepted value is about 334 J/g.
Heat of vaporisation (enthalpy of vaporisation)
The energy required to convert one gram of a liquid into a gas at its boiling point. For water, this is about 2260 J/g, roughly seven times the heat of fusion.
Temperature change (delta T / ΔT)
The difference between the final temperature and the initial temperature of a substance: ΔT = T_final − T_initial. A positive ΔT means the substance warmed up; a negative ΔT means it cooled down.
In this lab, 100.0 mL of water (mass 100.0 g, since the density of water is 1.00 g/mL) was heated from an initial temperature of 2.2 °C. Temperature readings were taken every minute for 15 minutes.
From 0 to about 11 minutes, the temperature rose steadily from 2.2 °C to 99.2 °C. This is the liquid-phase heating region, where all the energy from the heat source goes into increasing the kinetic energy of the water molecules.
The steepest rise occurs between minutes 4 and 10 (roughly 10 °C to 87 °C), where the temperature climbs about 12 to 13 degrees per minute.
From minute 11 onward the temperature levels off at approximately 99.9 °C, forming a plateau. The water has reached its boiling point and is undergoing a phase change from liquid to gas. Energy is still being supplied, but it is being used to overcome intermolecular forces (converting liquid water to steam) rather than raising the temperature.
The recorded boiling point of 99.9 °C is very close to the accepted value of 100 °C. Small deviations are normal and can result from local atmospheric pressure, thermometer calibration or dissolved substances.
Salol (phenyl salicylate) was heated until molten, then allowed to cool. Temperature was recorded every minute for 12 minutes.
From 0 to about 6 minutes, the temperature dropped steadily from 71.1 °C to 37.2 °C. The liquid is losing thermal energy to the surroundings.
Between minutes 6 and 7 the temperature rose slightly from 37.2 °C to 41.1 °C. This brief bump happens because salol supercools: it dips below its true freezing point before crystallisation begins, and the act of crystallising releases energy (the heat of fusion) which warms the sample back up to the freezing point.
From minutes 7 through 12 the temperature holds nearly constant at about 41.0 °C. This plateau marks the freezing point: liquid and solid coexist while the substance releases its heat of fusion at a constant temperature.
The freezing point of salol, read from the plateau, is approximately 40 °C.
A plateau on a heating curve indicates boiling (or melting, if the substance starts as a solid). A plateau on a cooling curve indicates freezing (or condensation, if the substance starts as a gas). In both cases, the temperature stays constant because the energy being added or removed is going entirely into the phase change, not into changing the temperature. The length of the plateau depends on the amount of substance and the rate of energy transfer.
q = m \times c \times \Delta TWhere:
q = energy transferred (joules, J)
m = mass of the substance (grams, g)
c = specific heat capacity (J/g°C), for liquid water c = 4.18 J/g°C
ΔT = change in temperature (°C), calculated as T_final − T_initial
Worked example from the lab: heating 100.0 g of water from 2.2 °C (initial, at minute 0) to a final temperature around 97.6 °C (the temperature just before the plateau, though the worksheet uses the boiling point of 99.9 °C minus the initial 2.2 °C, giving a temperature change rounded to 95.4 °C).
q = 100.0 g × 4.18 J/g°C × 95.4 °C = 39,877.2 J ≈ 39.88 kJ
To convert joules to kilojoules, divide by 1000.
A heating curve for water starting from ice would show five distinct regions, though the lab began with liquid water:
Region 1: solid heating (temperature rises)
Region 2: melting plateau (temperature constant at 0 °C)
Region 3: liquid heating (temperature rises, this is what the lab measured)
Region 4: boiling plateau (temperature constant at ~100 °C, visible in the lab data at minutes 12 to 15)
Region 5: gas/steam heating (temperature rises again)
The cooling curve is the mirror image. For salol in the lab:
Liquid cooling region: temperature drops steadily
Brief supercooling dip (temperature drops below the freezing point momentarily)
Freezing plateau: temperature holds constant at the freezing point (~40 °C for salol)
Once fully solidified, the solid would continue to cool (not shown in this data set)
Students often think that temperature keeps rising while a substance boils. It does not. During a phase change, all the energy goes into breaking intermolecular bonds, and the temperature stays constant until the entire substance has changed state.
Students sometimes assume that the boiling point of water is always exactly 100.0 °C. In practice, laboratory readings routinely come in at 99 to 100 °C depending on atmospheric pressure, altitude and thermometer precision. The accepted value at standard pressure is 100 °C, but slight deviations are normal and expected.
Students may confuse the supercooling dip on a cooling curve with an error. The brief temperature drop below the freezing point before the plateau begins is a real phenomenon: the liquid can cool past its freezing point before crystallisation starts, and the energy released by crystallisation then brings the temperature back up to the true freezing point.
Students sometimes mix up heat of fusion and heat of vaporisation, or assume they are similar in magnitude. The heat of vaporisation of water (about 2260 J/g) is roughly seven times larger than its heat of fusion (about 334 J/g), because converting a liquid to a gas requires overcoming far more intermolecular attraction than converting a solid to a liquid.
Heating and cooling curves are one of the most commonly tested topics in introductory chemistry. Expect to be asked to sketch a curve, label its regions, or identify the phase change from a graph.
You will almost certainly need to use q = m × c × ΔT to calculate energy absorbed or released. Know the specific heat of water (4.18 J/g°C) from memory.
Be prepared to read a freezing or boiling point directly from a plateau on a graph. If the exam gives you a data table, plot it mentally (or on scrap paper) and look for where the temperature stops changing.
Questions often ask what is happening at the molecular level during each region. Rising temperature means increased kinetic energy. A plateau means energy is going into overcoming intermolecular forces.
This is how engineers design heat exchangers and why your car radiator has fins: the principles of specific heat and phase-change energy transfer underpin real thermal-management systems.
True or false: During boiling, the temperature of water continues to increase as long as heat is applied.
False. Temperature remains constant during a phase change.
Fill in the blank: The specific heat capacity of liquid water is ______ J/g°C.
4.18
True or false: On a cooling curve, a plateau indicates the substance is freezing.
True.
Fill in the blank: The brief dip below the freezing point on salol's cooling curve before the plateau is called ______.
Supercooling.
True or false: The mass of water in grams equals its volume in mL because water's density is 1.00 g/mL.
True (at or near room temperature and standard pressure).
Q: When water is heated and the temperature reaches a constant value forming a plateau on the graph, what does the plateau indicate?
A: The plateau indicates that the water has reached its boiling point. All the energy being added is going into the phase change (liquid to gas) rather than raising the temperature. The temperature will remain constant at the boiling point until all the liquid has evaporated or the heat source is removed.
Q: 175 g of water is heated from 15 °C to 88 °C. How many kilojoules are absorbed by the water?
A: ΔT = 88 °C − 15 °C = 73 °C. Then q = 175 g × 4.18 J/g°C × 73 °C = 53,399.5 J = 53.40 kJ.
Q: Looking at a cooling curve for salol, how would you identify the freezing point?
A: The freezing point is the temperature at which the plateau appears on the cooling curve. For salol, the temperature drops steadily, may briefly dip below the freezing point (supercooling), then levels off at approximately 40 °C. That constant-temperature region is the freezing point.
Q: Why does the temperature of salol briefly drop below its freezing point before levelling off?
A: This is supercooling. The liquid can cool past the freezing point before crystallisation begins. Once crystals start forming, the heat of fusion released by solidification warms the sample back up to the true freezing point, where the temperature then holds steady.
Q: A 50.0 g sample of water is heated from 20.0 °C to 75.0 °C. Calculate the energy absorbed in kilojoules.
A: ΔT = 75.0 − 20.0 = 55.0 °C. q = 50.0 g × 4.18 J/g°C × 55.0 °C = 11,495 J = 11.50 kJ.
This material connects directly to thermodynamics and calorimetry. The formula q = mcΔT reappears whenever you calculate energy transfer in solution chemistry, enthalpy of reaction problems or Hess's law calculations. Understanding phase changes also ties into intermolecular forces: the stronger the forces between molecules, the more energy is needed to break them, which means higher boiling points, higher heats of fusion and longer plateaus.
The concept of supercooling from the salol experiment connects to materials science and crystallography. In everyday life, supercooling is why hand warmers work: a supersaturated solution can be triggered to crystallise, releasing stored heat of fusion on demand.
Heating curve, cooling curve, phase change diagram, temperature vs time graph, specific heat capacity, specific heat of water, q = mcΔT, energy calculation chemistry, boiling point plateau, freezing point plateau, supercooling, heat of fusion, heat of vaporisation, enthalpy of fusion, latent heat, calorimetry, salol freezing point, phenyl salicylate, change of state energy, Lab 5 energy and matter, introductory chemistry lab