Glossary
Bond Enthalpies
The average energy required to break one mole of a specific type of bond in the gaseous state, used to estimate the enthalpy change of a reaction.
Example:
Calculating the sum of bond enthalpies for reactants and products helps predict whether a reaction will be endothermic or exothermic.
Calorimeter
A device used in calorimetry to measure the heat changes of a reaction by isolating the system and measuring the temperature change of a known mass of water or other substance.
Example:
A simple coffee-cup calorimeter can be used in a lab to measure the heat of dissolution of a salt.
Calorimetry
The experimental technique used to measure the heat released or absorbed during a chemical reaction or physical process.
Example:
Scientists use calorimetry to determine the caloric content of food by burning it and measuring the heat produced.
Endothermic Processes
Processes that absorb heat energy from the surroundings, leading to an increase in the system's energy and a positive change in enthalpy (ΔH > 0).
Example:
An instant cold pack feels chilly because the chemical reaction inside is an endothermic process, drawing heat from your skin.
Enthalpies of Formation (ΔH°f)
The heat change that occurs when one mole of a compound is formed from its constituent elements in their standard states under standard conditions.
Example:
The standard enthalpy of formation for elemental oxygen (O2) is zero, as it is already in its most stable form.
Enthalpy (ΔH)
A thermodynamic property representing the total heat content of a system at constant pressure; its change (ΔH) indicates the heat absorbed or released during a chemical reaction.
Example:
A negative enthalpy change for a reaction, like the combustion of methane, signifies that it is an exothermic process.
Exothermic Processes
Processes that release heat energy to the surroundings, resulting in a decrease in the system's energy and a negative change in enthalpy (ΔH < 0).
Example:
A burning candle is an exothermic process because it releases light and heat into the room.
First Law of Thermodynamics
States that energy cannot be created or destroyed, only transferred or converted from one form to another; it is also known as the law of conservation of energy.
Example:
When a car burns gasoline, the chemical energy isn't destroyed but converted into kinetic energy and heat, illustrating the First Law of Thermodynamics.
Heat Capacity
The amount of heat energy required to raise the temperature of a given quantity of a substance by one degree Celsius (or Kelvin).
Example:
Water has a high heat capacity, which is why it takes a lot of energy to boil water, but it also stays warm for a long time.
Hess's Law
States that if a reaction can be expressed as the sum of a series of steps, then the enthalpy change for the overall reaction is the sum of the enthalpy changes for each step.
Example:
Using Hess's Law, you can calculate the enthalpy change for a complex reaction by combining the known enthalpy changes of simpler, related reactions.
State Function
A property of a system that depends only on its current state, not on the path taken to reach that state.
Example:
Enthalpy is a state function because the total heat change for a reaction is the same whether it occurs in one step or multiple steps.
Thermal Equilibrium
The state reached when two objects or systems in contact have exchanged heat until they are at the same temperature, and no net heat transfer occurs between them.
Example:
If you put a warm can of soda into a cooler full of ice, they will eventually reach thermal equilibrium when the soda is as cold as the melted ice water.
Thermodynamics
The branch of chemistry and physics that studies energy and its transformations, particularly how heat and work relate to chemical reactions and physical changes.
Example:
Understanding thermodynamics helps predict if a reaction will spontaneously occur and how much energy it will release or absorb.