When you boil water for tea or light a candle, heat moves between the system and its surroundings. This exchange is explained by enthalpy, a way to measure the total heat content of a system.
Chemists use the enthalpy formula, $H = U + PV$, to calculate heat changes in reactions or phase transitions. In this guide, you’ll learn about enthalpy change, standard enthalpy, and how to use these concepts to solve real-world chemistry problems.
Enthalpy Change: Quick Summary
Do you just need the basics? Here’s a simple explanation of what enthalpy is:
🟠 Enthalpy ($H$): The total heat content of a system, calculated as $H = U + PV$, where $U$ is internal energy, and $PV$ is pressure-volume work.
🟠 Enthalpy Change ($\Delta H$): Measures heat absorbed or released during a process at constant pressure, such as exothermic ($\Delta H < 0$) or endothermic ($\Delta H > 0$) reactions.
🟠 Standard Enthalpy ($\Delta H^\circ$): Heat change measured under standard conditions ($298\,\text{K};\ 1\,\text{bar}$) for reactions, phase changes, or formations.
🟠 Phase Transition Enthalpies: Melting, vaporization, and sublimation require energy, expressed as $\Delta H_\text{mol}$, $\Delta H_\text{vap}$, and $\Delta H_\text{sub}$.
🟠 Reaction enthalpy $\Delta H_\text{reaction}$: Reaction enthalpy shows the heat released or absorbed during a chemical reaction at constant pressure.
What is Enthalpy?
Enthalpy ($H$) is the total heat content of a system. It combines the internal energy ($U$) of the system with the energy linked to its pressure and volume ($PV$). The formula for enthalpy is straightforward:
$H = U + PV$
Enthalpy is a state function, which means it depends only on the initial and final states of the system, not on the steps taken to get there. This makes it a consistent way to measure energy changes in reactions and processes.
In constant-pressure processes like chemical reactions, the change in enthalpy ($\Delta H$) equals the heat absorbed or released. For instance, in a combustion reaction, the heat released can be calculated directly from $\Delta H$.
You experience enthalpy changes in daily life. When you boil water, heat is added to increase its temperature and eventually change its phase from liquid to gas. Melting ice or warming soup involves similar energy exchanges.
Enthalpy Change ($\Delta H$)
Enthalpy change ($\Delta H$) measures the heat absorbed or released during a process at constant pressure. It shows whether energy flows into or out of a system.
In exothermic reactions ($\Delta H < 0$), heat is released to the surroundings. A good example is the combustion of methane:
$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O , (\Delta H < 0)$
This reaction generates heat, warming its surroundings.
In endothermic reactions ($\Delta H > 0$), heat is absorbed. For instance, melting ice requires heat from the surroundings:
$H_2O(s) \rightarrow H_2O(l) , (\Delta H > 0)$
Here’s a quick comparison of the two processes:
| Type | Description | $\Delta H$ | Example |
| Exothermic | Releases heat to surroundings | $\Delta H < 0$ | Combustion, freezing water |
| Endothermic | Absorbs heat from surroundings | $\Delta H > 0$ | Melting ice, boiling water |
Reaction Enthalpy and Enthalpy of Reaction ($\Delta H_{\text{reaction}}$)
Reaction enthalpy ($\Delta H_{\text{reaction}}$) measures the heat change during a chemical reaction at constant pressure. It helps calculate how much energy is released or absorbed when reactants become products.
You can calculate reaction enthalpy using this formula:
$\Delta H_{\text{reaction}} = \sum \Delta H_f (\text{products}) – \sum \Delta H_f (\text{reactants})$
Here, $\Delta H_f$ stands for the standard enthalpy of formation, which is the heat change when one mole of a compound forms from its elements in their standard states.
Take the combustion of methane as an example:
$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$
Using the formula:
$\Delta H_{\text{reaction}} = \left[ \Delta H_f (CO_2) + 2 \Delta H_f (H_2O) \right] – \left[ \Delta H_f (CH_4) + 2 \Delta H_f (O_2) \right]$
With these values:
$\Delta H_f (CO_2) = -393.5 \, \text{kJ/mol}$, $\Delta H_f (H_2O) = -285.8 \, \text{kJ/mol}$, $\Delta H_f (CH_4) = -74.8 \, \text{kJ/mol}$, and $\Delta H_f (O_2) = 0 \, \text{kJ/mol}$, the calculation becomes:
$\Delta H_{\text{reaction}} = \left[ -393.5 + 2(-285.8) \right] – \left[ -74.8 + 2(0) \right]$
$\Delta H_{\text{reaction}} = -965.1 \, \text{kJ/mol}$
Burning one mole of methane releases $965.1 \, \text{kJ}$ of energy.
Standard Enthalpy ($\Delta H^\circ$)
Standard enthalpy ($\Delta H^\circ$) represents the heat change during a process measured under standard conditions: $298 \, \text{K}$ (room temperature) and $1 \, \text{bar}$ pressure. These conditions provide a baseline for comparing heat changes in reactions and processes.
Types of Standard Enthalpy
1. Standard Enthalpy of Formation ($\Delta H_f^\circ$):
This is the heat change when one mole of a compound forms from its elements in their most stable states. For example:
$C(s) + O_2(g) \rightarrow CO_2(g) , (\Delta H_f^\circ = -393.5 \, \text{kJ/mol})$
2. Standard Enthalpy of Combustion ($\Delta H_c^\circ$):
This measures the heat released when one mole of a substance burns completely in oxygen. For example:
$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O , (\Delta H_c^\circ = -890.3 \, \text{kJ/mol})$
3. Standard Enthalpy of Vaporization ($\Delta H_\text{vap}^\circ$):
This is the energy needed to turn one mole of a liquid into a gas. For water:
$H_2O(l) \rightarrow H_2O(g) , (\Delta H_\text{vap}^\circ = 40.7 \, \text{kJ/mol})$
4. Standard Enthalpy of Melting ($\Delta H_\text{mol}^\circ$):
This measures the heat required to melt one mole of a solid into a liquid. For ice:
$H_2O(s) \rightarrow H_2O(l) , (\Delta H_\text{mol}^\circ = 6.01 \, \text{kJ/mol})$
| Substance | $\Delta H_f^\circ$ (kJ/mol) | $\Delta H_c^\circ$ (kJ/mol) | $\Delta H_\text{vap}^\circ$ (kJ/mol) | $\Delta H_\text{mol}^\circ$ (kJ/mol) |
| Water ($H_2O$) | $-285.8$ | N/A | $40.7$ | $6.01$ |
| Methane ($CH_4$) | $-74.8$ | $-890.3$ | N/A | N/A |
Enthalpy in Phase Transitions
Phase transitions involve changes in a substance’s physical state, such as melting, boiling, or sublimation. These processes occur at a constant temperature, where the added or removed energy is used to break or form intermolecular bonds, not to change kinetic energy. Each type of transition has a specific enthalpy value.
Enthalpy of Melting
Melting occurs when a solid transitions into a liquid. This process requires energy to break intermolecular forces without increasing the temperature. The enthalpy of melting ($\Delta H_\text{mol}$) represents the energy required to melt one mole of a solid under standard conditions.
Example: Melting Ice
To melt $1 \, \text{mol}$ of ice at $0^\circ \text{C}$, the enthalpy of melting is:
$\Delta H_\text{mol} = 6.01 \, \text{kJ/mol}$
Calculation Example:
If $2 \, \text{mol}$ of ice melts, the total energy required is:
$\Delta H = n \cdot \Delta H_\text{mol}$
$\Delta H = 2 \, \text{mol} \cdot 6.01 \, \text{kJ/mol}$
$\Delta H = 12.02 \, \text{kJ}$
Enthalpy in Vaporization
Vaporization is the phase change from liquid to gas. It requires energy to separate particles far enough to overcome intermolecular forces. The enthalpy of vaporization ($\Delta H_\text{vap}$) measures the heat required to vaporize one mole of liquid under standard conditions.
For water:
$\Delta H_\text{vap} = 40.7 \, \text{kJ/mol}$
Example Calculation:
To vaporize $0.5 \, \text{mol}$ of water:
$\Delta H = n \cdot \Delta H_\text{vap}$
$\Delta H = 0.5 \, \text{mol} \cdot 40.7 \, \text{kJ/mol}$
$\Delta H = 20.35 \, \text{kJ}$
Enthalpy in Sublimation
Sublimation is the direct transition from solid to gas. The enthalpy of sublimation ($\Delta H_\text{sub}$) combines melting and vaporization energies:
$\Delta H_\text{sub} = \Delta H_\text{mol} + \Delta H_\text{vap}$
For water:
$\Delta H_\text{sub} = 6.01 \, \text{kJ/mol} + 40.7 \, \text{kJ/mol}$
$\Delta H_\text{sub} = 46.71 \, \text{kJ/mol}$
Example Calculation:
To sublime $3 \ \, \text{mol}$ of ice:
$\Delta H = n \cdot \Delta H_\text{sub}$
$\Delta H = 3 \, \text{mol} \cdot 46.71 \, \text{kJ/mol}$
$\Delta H = 140.13 \, \text{kJ}$
Calculations with Enthalpy
Accurate enthalpy calculations help measure heat transfer and energy changes in chemical processes. Let’s explore examples involving temperature changes, reaction enthalpy, and phase transitions.
Heat Transfer in Chemical Reactions
Heat transfer during a temperature change uses the formula:
$q = m c \Delta T$
Where:
- $q$ = heat absorbed or released (J)
- $m$ = mass of the substance (g)
- $c$ = specific heat capacity ($\text{J/g}^\circ\text{C}$)
- $\Delta T$ = change in temperature ($^\circ \text{C}$)
Example: Heating water
Calculate the heat absorbed by $100 \ \, \text{g}$ of water heated from $25^\circ \text{C}$ to $100^\circ \text{C}$, with $c = 4.18 \, \text{J/g}^\circ \text{C}$:
$q = m c \Delta T$
$q = 100 \cdot 4.18 \cdot (100 – 25)$
$q = 100 \cdot 4.18 \cdot 75$
$q = 31,350 \, \text{J}$ or $31.35 \, \text{kJ}$
Reaction Enthalpy Calculations
Reaction enthalpy is calculated using:
$\Delta H_{\text{reaction}} = \sum \Delta H_f (\text{products}) – \sum \Delta H_f (\text{reactants})$
Example: Combustion of methane
For the reaction $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$:
- $\Delta H_f (CH_4) = -74.8 \, \text{kJ/mol}$
- $\Delta H_f (CO_2) = -393.5 \, \text{kJ/mol}$
- $\Delta H_f (H_2O) = -285.8 \, \text{kJ/mol}$
$\Delta H_{\text{reaction}} = [-393.5 + 2(-285.8)] – [-74.8 + 2(0)]$
$\Delta H_{\text{reaction}} = -965.1 + 74.8$
$\Delta H_{\text{reaction}} = -890.3 \, \text{kJ/mol}$
Phase Transition Enthalpy Calculations
Enthalpy changes during phase transitions involve:
$\Delta H_{\text{total}} = n \cdot (\Delta H_\text{mol} + \Delta H_\text{vap})$
Example: Freezing and boiling water
For $2\ \, \text{mol}$ of water:
- $\Delta H_\text{mol} = 6.01 \, \text{kJ/mol}$
- $\Delta H_\text{vap} = 40.7 \, \text{kJ/mol}$
$\Delta H_{\text{total}} = 2 \cdot (6.01 + 40.7)$
$\Delta H_{\text{total}} = 2 \cdot 46.71$
$\Delta H_{\text{total}} = 93.42 \, \text{kJ}$
Advance Your Knowledge in Enthalpy
Are you struggling with reaction enthalpy or enthalpy change? You can find more useful topics in our Chemistry blogs. Or find a tutor, who can explain it in a way that clicks for you.
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Enthalpy: Frequently Asked Questions
1. What is enthalpy in chemistry?
Enthalpy is the total heat content of a system, calculated using $H = U + PV$, where $U$ is internal energy, $P$ is pressure, and $V$ is volume.
2. What does $\Delta H$ represent?
$\Delta H$ represents the enthalpy change, measuring the heat absorbed or released during a process at constant pressure.
3. How do exothermic and endothermic reactions differ?
Exothermic reactions release heat and have $\Delta H < 0$, while endothermic reactions absorb heat with $\Delta H > 0$.
4. What is standard enthalpy ($\Delta H^\circ$)?
Standard enthalpy refers to enthalpy change measured under standard conditions of $298 \, \text{K}$ and $1 \, \text{bar}$.
5. How is reaction enthalpy ($\Delta H_{\text{reaction}}$) calculated?
Reaction enthalpy is calculated using $\Delta H_{\text{reaction}} = \sum \Delta H_f (\text{products}) – \sum \Delta H_f (\text{reactants})$.
6. What is enthalpy of melting ($\Delta H_\text{mol}$)?
Enthalpy of melting is the heat required to convert one mole of a solid into a liquid at constant temperature and pressure.
7. How does temperature affect enthalpy?
Higher temperatures increase molecular interactions and internal energy, raising the system’s enthalpy.
8. What is enthalpy of vaporization ($\Delta H_\text{vap}$)?
The enthalpy of vaporization is the energy needed to transform one mole of liquid into a gas at constant temperature and pressure.
Sources:
1. Geeks for Geeks
2. Britannica
3. Wikipedia
