Glossary
Calorimetry
The science of measuring the heat of chemical reactions or physical changes, typically using a calorimeter.
Example:
A student performs calorimetry to determine the specific heat of an unknown metal by measuring the temperature change of water it heats.
Condensation
The phase transition from a gas to a liquid, occurring as energy is released and intermolecular forces reform.
Example:
When humid air touches a cold glass, water vapor undergoes condensation, forming liquid droplets on the surface.
Cooling Curves
Graphs that show how the temperature of a substance changes as heat energy is continuously removed over time.
Example:
Observing water as it freezes into ice would produce a cooling curve with a plateau at 0°C.
Critical Point
The point on a phase diagram representing the critical temperature and critical pressure, beyond which a distinct liquid phase no longer exists, and the substance becomes a supercritical fluid.
Example:
Above the critical point for water, there's no clear distinction between liquid water and steam; it's a single supercritical fluid.
Endothermic
A process that absorbs heat energy from its surroundings.
Example:
The melting of ice into water is an endothermic process, as it requires energy input to break intermolecular forces.
Exothermic
A process that releases heat energy into its surroundings.
Example:
The freezing of water into ice is an exothermic process, as it releases energy to the environment.
Freezing
The phase transition from a liquid to a solid, occurring at the freezing point as energy is released and a more ordered structure forms.
Example:
Placing liquid water in a freezer causes it to undergo freezing, turning into solid ice.
Heat of Combustion (ΔH_comb)
The heat released when one mole of a substance undergoes complete combustion with oxygen under standard conditions.
Example:
The heat of combustion of methane is a large negative value, indicating that burning natural gas releases a significant amount of energy.
Heat of Formation (ΔH_f°)
The change in enthalpy when one mole of a compound is formed from its constituent elements in their standard states.
Example:
The heat of formation of carbon dioxide is the enthalpy change when solid carbon and gaseous oxygen combine to form one mole of CO2 gas.
Heat of Fusion (H_f)
The amount of energy required to change a unit mass (or mole) of a substance from a solid to a liquid at its melting point.
Example:
To calculate the energy needed to melt a block of ice, one would use the heat of fusion for water.
Heat of Vaporization (H_v)
The amount of energy required to change a unit mass (or mole) of a substance from a liquid to a gas at its boiling point.
Example:
The high heat of vaporization of water is why sweating is an effective cooling mechanism for the human body.
Heating Curves
Graphs that show how the temperature of a substance changes as heat energy is continuously added over time.
Example:
A chemist plots the temperature of ice as it's heated, observing a heating curve with flat segments at 0°C and 100°C.
Hess's Law
A principle stating that the total enthalpy change for a chemical reaction is the same, regardless of the pathway taken, as long as the initial and final conditions are the same.
Example:
Using Hess's Law, one can calculate the enthalpy change for a complex reaction by summing the enthalpy changes of a series of simpler reactions.
Intermolecular Forces (IMFs)
Attractive forces that exist between molecules, influencing a substance's physical properties like boiling point and melting point.
Example:
Hydrogen bonding, a strong type of intermolecular force, explains why water has a relatively high boiling point compared to other similar-sized molecules.
Melting
The phase transition from a solid to a liquid, occurring at the melting point as energy is absorbed to overcome intermolecular forces.
Example:
When an ice cube is left out, it undergoes melting as it absorbs heat from the room, turning into liquid water.
Phase Diagrams
Graphs that illustrate the conditions of temperature and pressure at which different phases (solid, liquid, gas) of a substance can exist in equilibrium.
Example:
A phase diagram for CO2 shows that at room temperature and atmospheric pressure, it exists as a gas, but at lower temperatures, it can be a solid (dry ice).
Plateaus (Phase Change Zones)
Flat regions on a heating or cooling curve where the temperature remains constant despite continuous energy transfer, indicating a phase change is occurring.
Example:
During the boiling of water, the temperature stays at 100°C, forming a plateau on the heating curve as liquid turns to gas.
Specific Heat (capacity)
The amount of heat energy required to raise the temperature of one gram of a substance by one degree Celsius (or Kelvin).
Example:
Water has a high specific heat, meaning it takes a lot of energy to change its temperature, which helps regulate Earth's climate.
Thermochemistry
The branch of chemistry concerned with the study of the heat changes that accompany chemical reactions and physical transformations.
Example:
Understanding bond energies and enthalpy changes falls under the study of thermochemistry.
Triple Point
A unique point on a phase diagram where all three phases (solid, liquid, and gas) of a substance coexist in thermodynamic equilibrium.
Example:
At the triple point of water, ice, liquid water, and water vapor can all be present simultaneously.
Unit Conversions
The process of changing a measurement from one unit to another (e.g., grams to moles, Joules to kilojoules) while maintaining its value.
Example:
Before using a specific heat value in J/g°C, a student might need to perform unit conversions if the given mass is in kilograms.
Vaporizing
The phase transition from a liquid to a gas, occurring at the boiling point as energy is absorbed to completely overcome intermolecular forces.
Example:
When water boils on a stove, it is vaporizing into steam, a process that requires significant energy.
q = mcΔT
An equation used to calculate the heat (q) absorbed or released by a substance when its temperature (ΔT) changes, involving its mass (m) and specific heat capacity (c).
Example:
To find out how much energy is needed to warm a cup of coffee, you would use the formula q = mcΔT.