Difficulty: Intermediate | Prerequisites: Part 1 (Systems, Processes, and State Functions).
Big Picture
Once you know what a system is and how energy is conserved (Part 1), the next question is: how do you measure the energy that moves as heat? This section covers the distinction between heat and temperature, the equations used to calculate heat transfer during temperature changes and phase changes, the two main types of calorimetry, and how phase diagrams map out where each state of matter is stable. If Part 1 is the framework, this section is the measurement toolkit.
Heat is energy transferred because of a temperature difference, measured with Q = mcΔT for temperature changes and Q = mL for phase changes. Calorimetry captures these measurements in controlled conditions. Phase diagrams show which state of matter is stable at any combination of temperature and pressure, with special points (triple point, critical point) where phases meet or merge.
Temperature
A measure of the average kinetic energy of particles in a substance. When thermal energy increases, temperature increases. Think of it as how fast the particles are jiggling on average.
Heat (Q)
The transfer of energy from one substance to another due to a temperature difference. It is a process function, not a state function. Heat always flows from warmer to colder.
Zeroth law of thermodynamics
Objects are in thermal equilibrium only when their temperatures are equal. In simple terms, if A is the same temperature as B, and B is the same as C, then A and C are also in thermal equilibrium.
Endothermic
A process where the system absorbs heat from the surroundings (ΔQ > 0). Think of it as the system pulling energy in.
Exothermic
A process where the system releases heat to the surroundings (ΔQ < 0). The system is giving energy away.
Specific heat (c)
The amount of energy required to raise the temperature of one gram of a substance by one degree Celsius. Water's specific heat is 1 cal/g·K (or 4.184 J/g·K), which is unusually high.
Heat capacity (HC)
The energy required to raise the temperature of any given amount of a substance by one degree Celsius. Calculated as HC = mc (mass times specific heat).
Calorimeter
A device that measures heat transferred during a process. Constant-pressure calorimeters work at atmospheric pressure; constant-volume (bomb) calorimeters keep volume fixed.
Enthalpy (H)
Equivalent to heat under constant pressure. It is a state function.
Latent heat (L)
The enthalpy change during a phase transition at constant temperature. The heat of fusion applies to solid-liquid transitions; the heat of vaporization applies to liquid-gas transitions.
Vapor pressure
The pressure exerted by gas molecules above a liquid at equilibrium in a closed container. It increases with temperature.
Boiling point
The temperature at which a liquid's vapor pressure equals atmospheric pressure. At this point, vaporization occurs throughout the liquid, not just at the surface.
Melting / freezing point
The temperature at which solid and liquid phases are in equilibrium. Crystalline solids have precise melting points; amorphous solids have broader ranges.
Sublimation
The direct transition from solid to gas, bypassing the liquid phase.
Deposition
The direct transition from gas to solid, the reverse of sublimation.
Triple point
The unique temperature and pressure at which all three phases (solid, liquid, gas) coexist in equilibrium.
Critical point
The temperature and pressure above which there is no distinction between liquid and gas phases. The heat of vaporization is zero at and above this point.
Heating curve
A graph of temperature versus heat added. Sloped regions show temperature changes within a single phase; flat (plateau) regions show phase changes where temperature remains constant.
Temperature measures average kinetic energy of particles. Heat is the energy transferred between substances because of a temperature difference.
You cannot say that a hot object has greater thermal energy than a cold one, because thermal energy also depends on mass and specific heat. A bathtub of lukewarm water holds more thermal energy than a thimble of boiling water.
Heat always flows from warmer to colder (zeroth law).
Joule (J) is the SI unit. Calorie (cal) is the older unit.
1 cal = 4.184 J.
Q = heat absorbed or released.
m = mass of the substance.
c = specific heat of the substance (energy per gram per degree).
ΔT = change in temperature (in Celsius or Kelvin; the size of one degree is the same in both).
Specific heat is per gram. Heat capacity (HC = mc) is for any given mass.
Water's specific heat: 1 cal/g·K = 4.184 J/g·K.
Constant-pressure calorimeter: operates at atmospheric pressure. The heat measured equals ΔH.
Constant-volume (bomb) calorimeter: volume is fixed, so W = PΔV = 0. No heat escapes to the universe (Q = 0 for the calorimeter as a whole). The heat measured equals ΔU.
Gas-liquid equilibrium:
Temperature reflects average kinetic energy, but individual molecules have a range of kinetic energies.
Evaporation (vaporization): high-energy molecules escape the liquid surface. This is endothermic, and it cools the remaining liquid. It occurs at all temperatures to some extent.
Boiling: vaporization that occurs throughout the liquid, only at or above the boiling point (where vapor pressure equals atmospheric pressure).
Condensation: gas molecules lose energy and fall into the liquid phase. Favoured by low temperature and high pressure.
In a closed container, evaporation and condensation reach equilibrium. The pressure at equilibrium is the vapor pressure, which increases with temperature.
Liquid-solid equilibrium:
Adding heat increases vibrational motion until atoms break free into the liquid phase (melting / fusion).
Removing heat causes freezing / solidification / crystallization.
Crystalline solids have sharp melting points. Amorphous solids (like glass) soften over a range.
Gas-solid equilibrium:
Sublimation: solid directly to gas (e.g. dry ice).
Deposition: gas directly to solid (e.g. frost forming).
On a heating curve, sloped segments represent temperature rising within a single phase. Use Q = mcΔT here.
Flat (plateau) segments represent phase changes. Temperature stays constant because energy is breaking or forming intermolecular forces rather than raising kinetic energy. Here ΔT = 0, so Q = mcΔT does not apply.
For phase changes, use Q = mL instead, where L is the latent heat:
Solid to liquid: use the heat of fusion. Enthalpy is positive (endothermic).
Liquid to gas: use the heat of vaporization. Also endothermic.
The reverse transitions release heat (exothermic).
A phase diagram plots pressure (y-axis) vs temperature (x-axis) and shows where each phase is stable.
General layout: solid at low temperature / high pressure (left), liquid in the middle, gas at high temperature / low pressure (right).
Lines on the diagram represent equilibrium between two phases.
Triple point: all three phases coexist in equilibrium.
Critical point: beyond this temperature and pressure, liquid and gas are indistinguishable (a supercritical fluid). Heat of vaporization is zero at and above this point.
Q = mc\Delta THeat absorbed or released during a temperature change. m = mass, c = specific heat, ΔT = change in temperature.
Q = mLHeat absorbed or released during a phase change. m = mass, L = latent heat (heat of fusion or heat of vaporization).
HC = mcHeat capacity. The energy needed to raise a given amount of substance by one degree.
Unit conversion: 1 cal = 4.184 J
Water's specific heat: 1 cal/g·K = 4.184 J/g·K
The high specific heat of water is why coastal cities have milder climates than inland ones: the ocean absorbs and releases enormous amounts of heat without large temperature swings. Evaporative cooling (sweat on your skin, a wet cloth on your forehead) works because vaporization is endothermic, pulling heat from the surface it leaves. Pressure cookers raise the boiling point by increasing pressure above atmospheric, which is a direct application of the vapor pressure concept.
Students often assume that a hotter object always contains more thermal energy than a colder one. It does not, because thermal energy also depends on mass and specific heat.
Confusing boiling with evaporation is common. Evaporation happens at the surface at any temperature; boiling happens throughout the liquid only at or above the boiling point.
During a phase change, many students try to use Q = mcΔT. This does not work because ΔT = 0 during the transition. You must use Q = mL instead.
Students sometimes think the triple point means all three phases are at the same temperature and pressure as standard conditions. The triple point is a specific, unique value for each substance.
⚠️ Know when to use Q = mcΔT (temperature change within a phase) versus Q = mL (phase change at constant temperature). Heating curve problems test exactly this distinction.
⚠️ Be able to read a phase diagram: identify the triple point, critical point, and which phase is stable in each region.
⚠️ Know the difference between constant-pressure and constant-volume calorimetry, and which thermodynamic quantity each measures (ΔH vs ΔU).
⚠️ Water's specific heat (1 cal/g·K or 4.184 J/g·K) is frequently given or required in calculations.
⚠️ Understand that vapor pressure increases with temperature, and that boiling point is where vapor pressure equals atmospheric pressure.
True or False: During a phase change, the temperature of the substance continues to rise.
Answer: False. Temperature remains constant during a phase change.
Fill in the blank: The formula for heat during a phase change is Q = ___.
Answer: Q = mL.
True or False: Evaporation occurs only at the boiling point.
Answer: False. Evaporation occurs at all temperatures at the liquid surface.
Fill in the blank: At the triple point, ___ phases coexist in equilibrium.
Answer: Three (solid, liquid, gas).
True or False: In a constant-volume calorimeter, the work done is zero.
Answer: True. W = PΔV = 0 when volume is fixed.
Q: You heat 200 g of water from 20 °C to 80 °C. The specific heat of water is 4.184 J/g·K. How much heat is absorbed?
A: Q = mcΔT = (200 g)(4.184 J/g·K)(60 K) = 50,208 J, or about 50.2 kJ.
Q: A heating curve shows a flat plateau at 100 °C for water. Why is the temperature constant even though heat is being added?
A: The energy is being used to break intermolecular forces (converting liquid to gas) rather than increasing kinetic energy. During a phase change, ΔT = 0, and you must use Q = mL.
Q: On a phase diagram, you move from low temperature/high pressure toward high temperature/low pressure. Which phases do you pass through?
A: Solid, then liquid, then gas (left to right across the diagram).
Q: What is the difference between specific heat and heat capacity?
A: Specific heat (c) is per gram: the energy to raise 1 g by 1 °C. Heat capacity (HC = mc) is for a given mass: the energy to raise that entire sample by 1 °C.
Q: Why does a bomb calorimeter measure ΔU rather than ΔH?
A: Because the volume is constant (rigid container), no PΔV work is done. At constant volume, the heat measured equals the change in internal energy, not enthalpy.
Calorimetry measurements feed directly into enthalpy calculations (Part 3), because ΔH is defined as heat at constant pressure. The phase change concepts here connect to Gibbs free energy (Part 3): whether a phase change is spontaneous depends on ΔG, which balances enthalpy and entropy. Heating curves also link to intermolecular forces from earlier in general chemistry, since the strength of those forces determines the size of the heat of fusion and vaporization.
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