5.3 Enthalpy
3 min read•june 24, 2024
The governs energy changes in chemical reactions. It states that energy can't be created or destroyed, only converted between forms. This principle helps us understand how heat flows during reactions and introduces the concept of .
Enthalpy measures a system's heat content and changes during reactions. As a , it depends only on the current state, not the path taken. and allow us to calculate and predict enthalpy changes for various reactions.
First Law of Thermodynamics and Enthalpy
First law of thermodynamics in reactions
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- States energy cannot be created or destroyed, only converted from one form to another (thermal, chemical, mechanical, electrical)
- In a closed system, change in internal energy () equals heat () added to system minus work () done by system:
- At constant pressure, heat absorbed or released by system equals change in enthalpy ()
- Enthalpy measures total heat content of a system
- Changes in enthalpy during reactions indicate amount of heat absorbed (, ) or released (, )
- The amount of heat required to change a system's temperature is related to its
Enthalpy as a State Function
Enthalpy as state function
- Enthalpy () is a , its value depends only on current state of system, not path taken to reach that state
- Change in enthalpy () is independent of route taken between initial and final states
- In , enthalpy changes provide information about heat absorbed or released during reactions
- Exothermic reactions release heat (), endothermic reactions absorb heat ()
- Enthalpy is an extensive property, its value depends on amount of substance present (doubling amount doubles enthalpy)
Thermochemical Equations and Enthalpy Changes
Notation for thermochemical equations
- Thermochemical equations are balanced chemical equations that include enthalpy change for reaction
- Enthalpy change is written as a term in equation, typically on right side
- () is enthalpy change when reactants and products are in standard states (1 atm, 25°C, 1 M for solutions)
- Symbol denotes standard enthalpy of reaction
- Example: (combustion of methane)
Calculation of enthalpy changes
- () is enthalpy change when one mole of compound is formed from its constituent elements in standard states
- By convention, standard enthalpy of formation for elements in standard states is zero
- Standard enthalpy of reaction can be calculated using standard enthalpies of formation of reactants and products:
- Enthalpy changes can also be determined experimentally using
- Heat absorbed or released by system is measured and used to calculate enthalpy change
- can be used to estimate enthalpy changes in gas-phase reactions
Application of Hess's law
- Hess's law states overall enthalpy change for a reaction is independent of route taken between initial and final states
- Enthalpy change for a reaction can be calculated by summing enthalpy changes of individual steps that add up to overall reaction
- Allows determination of enthalpy changes for reactions that cannot be directly measured or are difficult to carry out experimentally
- By combining known enthalpy changes of simpler reactions, enthalpy change for a more complex reaction can be calculated
- Example: If reaction A has and reaction B has , and sum of reactions A and B yields reaction C, then
Thermodynamic Considerations
Spontaneity of reactions
- Enthalpy is not the only factor determining whether a reaction occurs spontaneously
- Entropy, a measure of disorder in a system, also plays a crucial role in determining reaction spontaneity
- combines both enthalpy and entropy to predict the spontaneity of a reaction under specific conditions
Key Terms to Review (34)
$ ext{Δ H}_{ ext{rxn}}^{ ext{°}}$: $ ext{Δ H}_{ ext{rxn}}^{ ext{°}}$ represents the standard enthalpy change, or the amount of heat energy released or absorbed during a chemical reaction under standard conditions of temperature and pressure. It is a fundamental concept in thermochemistry that quantifies the energetic changes associated with chemical transformations.
$ ext{Δ H}_f^{ ext{°}}$: $ ext{Δ H}_f^{ ext{°}}$ is the standard enthalpy of formation, which represents the change in enthalpy that occurs when one mole of a compound is formed from its constituent elements in their standard states at a specified temperature, usually 25°C and 1 atm pressure.
$
abla U$: $
abla U$ represents the change in internal energy of a system. It is a fundamental concept in thermodynamics that describes the energy transferred between a system and its surroundings during a process, without consideration of the work done or heat exchanged.
$
Delta H$: $
Delta H$ is the change in enthalpy, a thermodynamic property that represents the total energy released or absorbed during a chemical process or physical transformation at constant pressure. It is a measure of the heat energy exchanged between a system and its surroundings, and is a crucial concept in understanding chemical reactions, phase changes, and spontaneity of processes.
Biofuel: Biofuel is a type of energy source derived from organic materials, such as plant and animal waste. It can be used as an alternative to fossil fuels for generating heat and power.
Bond Enthalpy: Bond enthalpy, also known as bond dissociation energy, is the amount of energy required to break a particular chemical bond in a gaseous molecule. It represents the strength of the bond and is an important factor in determining the stability and reactivity of chemical compounds.
Calorimetry: Calorimetry is the measurement of heat transfer in chemical reactions or physical changes. It is used to determine the enthalpy changes of reactions.
Calorimetry: Calorimetry is the scientific process of measuring the heat energy released or absorbed during a chemical or physical process. It provides a way to quantify the energy changes that occur in a system, which is essential for understanding thermodynamic principles.
Endothermic: Endothermic refers to a process or reaction that absorbs heat from the surrounding environment. This means that the system undergoing the endothermic process requires an input of energy in the form of heat in order to proceed. Endothermic processes are central to understanding various topics in chemistry, including energy basics, enthalpy, dissolution, equilibrium, and free energy.
Endothermic process: An endothermic process is a chemical reaction or physical change that absorbs heat energy from its surroundings. These processes result in a decrease in the temperature of the surrounding environment.
Enthalpy: Enthalpy is a measure of the total energy of a thermodynamic system, including both the internal energy of the system and the work done by or on the system due to changes in pressure and volume. It is a key concept in understanding the energy changes that occur during chemical reactions and phase changes.
Enthalpy (H): Enthalpy (H) is the total heat content of a system at constant pressure. It is a thermodynamic property that includes internal energy and the product of pressure and volume.
Enthalpy change (ΔH): Enthalpy change ($\Delta H$) is the amount of heat absorbed or released by a system at constant pressure during a chemical reaction. It indicates whether a reaction is endothermic (absorbs heat) or exothermic (releases heat).
Exothermic: Exothermic refers to a chemical reaction or process that releases energy in the form of heat to the surrounding environment. These reactions produce more energy than they consume, resulting in a net release of heat.
Exothermic process: An exothermic process is a chemical reaction or physical change that releases heat to its surroundings. This release of energy usually results in an increase in the temperature of the surroundings.
Expansion work: Expansion work is the work done by a system when it expands against an external pressure. It often occurs during processes where gases are involved, such as in chemical reactions or physical changes.
First law of thermodynamics: The first law of thermodynamics states that energy cannot be created or destroyed, only transferred or converted from one form to another. In chemistry, it is often formulated as $\Delta U = Q - W$, where $\Delta U$ is the change in internal energy, $Q$ is the heat added to the system, and $W$ is the work done by the system.
First Law of Thermodynamics: The First Law of Thermodynamics states that energy can be converted from one form to another, but it cannot be created or destroyed. It is the principle of conservation of energy, which says that the total energy of an isolated system is constant and energy can neither be created nor destroyed, but can only be transformed or transferred from one form to another.
Gibbs Free Energy: Gibbs free energy is a thermodynamic property that combines the concepts of enthalpy and entropy to determine the spontaneity and feasibility of a chemical process. It is a crucial factor in understanding the driving forces behind chemical reactions and phase changes.
Gibbs free energy (G): Gibbs free energy (G) is a thermodynamic potential that measures the maximum reversible work obtainable from a system at constant temperature and pressure. It is used to predict the direction of chemical reactions.
Heat Capacity: Heat capacity is a measure of the amount of energy required to raise the temperature of a substance by one degree. It is a fundamental property that describes how much heat a material can absorb or release without undergoing a phase change.
Heat capacity (C): Heat capacity (C) is the amount of heat energy required to raise the temperature of a substance by one degree Celsius. It is an extensive property dependent on the quantity of the substance.
Hess’s law: Hess's law states that the total enthalpy change for a chemical reaction is the same, no matter how many steps the reaction is carried out in. It relies on the fact that enthalpy is a state function.
Hess's Law: Hess's law is a fundamental principle in thermochemistry that states the change in enthalpy (heat energy) for a chemical reaction is independent of the path taken, but rather depends only on the initial and final states of the reaction. This law allows for the calculation of enthalpy changes for complex reactions by breaking them down into simpler, well-understood steps.
Hydrocarbons: Hydrocarbons are organic compounds consisting entirely of hydrogen and carbon atoms. They are the primary components of fossil fuels like coal, petroleum, and natural gas.
Internal energy (U): Internal energy (U) is the total energy contained within a system, encompassing both kinetic and potential energies of the particles. It is a state function, meaning its value depends only on the current state of the system, not on how that state was reached.
Standard Enthalpy of Formation: The standard enthalpy of formation is the change in enthalpy that occurs when one mole of a compound is formed from its constituent elements in their standard states. It represents the energy released or absorbed during the formation of a compound from its elements under standard conditions of temperature and pressure.
Standard Enthalpy of Reaction: The standard enthalpy of reaction is the change in enthalpy that occurs during a chemical reaction under standard conditions, which are typically defined as a temperature of 25°C (298.15 K) and a pressure of 1 atm (101.325 kPa). It represents the energy released or absorbed when reactants are converted into products and is an important concept in understanding the thermodynamics of chemical processes.
Standard state: The standard state of a substance is its pure form at a specified temperature (usually 298 K) and pressure (1 bar). It serves as a reference point for thermodynamic calculations.
State function: A state function is a property of a system that depends only on its current state and not on the path taken to reach that state. Examples include enthalpy, internal energy, and entropy.
State Function: A state function is a thermodynamic property that depends only on the initial and final states of a system, and not on the path taken to get from the initial to the final state. It is a variable that can be used to describe the condition or state of a system, such as its energy, volume, or temperature, without regard to how the system reached that state.
Thermochemical Equations: Thermochemical equations are a type of chemical equation that not only represent the reactants and products of a chemical reaction, but also include information about the energy changes that occur during the reaction. They provide a comprehensive way to describe the energetics of a chemical process.
Thermochemistry: Thermochemistry is the study of the heat energy involved in chemical reactions and changes of state. It focuses on how energy is absorbed or released during these processes.
Thermochemistry: Thermochemistry is the study of the energy changes that occur during chemical reactions and phase changes. It focuses on the relationship between chemical processes and the associated transfer of heat energy.