- A **system** consists of reactants and products being studied, often represented as a chemical equation.
- The **surroundings**/**environment** are all matter outside of the system capable of absorbing or releasing energy.
- **Open** systems allow **energy and matter** to move in and out of the system.
- **Closed** systems allow only **energy** to move in and out of the system.
- **Isolated** systems do not allow energy or matter to move in and out of the system. This is an ideal but unrealistic scenario.
### Changes
As **breaking bonds requires energy** and **forming bonds releases energy**:
!!! definition
- An **endothermic** reaction overall requires energy.
- An **exothermic** reaction overall releases energy.
**Physical** changes such as state changes or dissolving substances may release or require energy depending on the energy of intermolecular bonds being broken and formed.
!!! example
- Ice melting requires energy to break the stronger bonds in a solid.
- Dissolving salt in water breaks the intermolecular bonds holding the salt together but regains it all by forming new bonds with the water.
**Chemical** changes all involve breaking old bonds to form new bonds. Depending on the energy required/released in breaking/forming those bonds, the reaction may end up endothermic or exothermic. Regardless, all reactions need a small initial **activation energy** to begin.
Represented as $H$ in joules, enthalpy represents the total energy in a system. Absolute enthalpy is not measurable, so change in enthalpy ($\Delta H$) is often used instead. The magnitude of enthalpy change is dependent on the type of change and quantity of substance that is changing.
A **negative** $\Delta H$ indicates that energy has left the system and so is an **exothermic** reaction.
In a balanced chemical equation, change in enthalpy is written to the right after the product.
$$
a + b \to c\ \ \Delta H = x\text{ kJ}
$$
!!! example
Energy is required for the decomposition of water so its enthalpy is positive.
$\Delta H$ can also be included in a balanced thermochemical equation as a reactant or product instead of listed at the end. In this case, it is always positive and its sign determines whether it is a reactant or product.
$$
a + b + x\text{ kJ} \to c
$$
!!! example
Using the same formula as in the previous example:
The **molar enthalpy of reaction** $\Delta H_x$ expresses the change in enthalpy when exactly one mole of the substance is involved in the reaction.
!!! example
The molar enthalpy of combustion (also known as the **heat of combustion**) of ethanol is $\pu{-1367 kJ/mol}$, indicating that every one mole of ethanol combusted releases 1367 kilojoules of energy.
The **standard molar enthalpy of reaction** $\Delta H^\theta_x$ is the molar enthalpy of reaction when initial and final conditions of the reaction are at standard atmospheric temperature and pressure (SATP, 25°C @ 100 kPA). Therefore, the activation energy, energy released/required during the reaction, and energy released/required following the reaction to return to SATP are all included.
!!! warning
This includes energy required for some substances to change state, such as water vapour from combustion cooling to 25°C.
### Energy profiles
Also known as **reaction profiles**, energy profiles are a visual representation of the change in chemical potential energy of the system.
- Absolute enthalpy ($H$) is placed on the y-axis while the reaction progress (time, sort of) is placed on the x-axis.
- A horizontal line representing the enthalpy before the change is placed at the beginning labelled with the reactants.
- A horizontal line representing the enthalpy after the change is placed at the end labelled with the products.
- The change in enthalpy is labelled with an arrow in the direction of the change with its value if known.
- A hump shows the reaction in progress (even exothermic reactions require some activation energy).
$\Delta H_B$, also known as **bond association energies**, the enthalpy of a bond type (e.g., $\ce{C-H}$) is the energy required to break **1 mol** of that bond type when the reactants and products are **gaseous** so energy is not lost from state changes. Compared of other methods of determining reaction enthalpy, this method is less accurate due to the other compounds affecting bond strength and thus enthalpy on a per-molecule basis.
The change in enthalpy of a reaction can be approximated by considering the bonds broken and formed:
$$\Delta H = \sum n\Delta H_B\text{reactants} - \sum n\Delta H_B\text{products}$$
A basic calorimeter uses a lid and insulation to keep matter in and minimise energy changes with its surroundings. A thermometer is used to measure the temperature change of the water, and a stirrer is common to ensure accurate thermometer readings. The reactants are placed in water to react.
It is assumed that all the heat lost/gained by the reaction is gained/lost from the water.
$$-q_\ce{H2O}=\Delta H_\text{reaction}$$
In the event that reactants cannot be placed in water to react (e.g., combustion), a **bomb calorimeter** is used, which contains a metal sealed box submerged in the waterfilled with reactant and oxygen. A circuit leads into the box to start the reaction with a spark.
!!! warning
Assumptions in calorimetry:
- All energy released/absorbed from the system goes to/from the surroundings of the calorimeter (water). This usually needs to be corrected for in bomb calorimeters by measuring the heat capacity and mass of the metal box inside the calorimeter as well.
- No energy is transferred outside the calorimeter — the insulation should work properly.
- The calorimeter itself does not absorb or release energy — this is not a good assumption but can be compensated for.
- A dilute aqueous solution is assumed to have the same density and specific heat capacity as water — this assumption is best when the solute is diluted close to 1 mol/L.
### Measuring calorimeters
Instead of recording the temperature of the calorimeter at any one point, a range of temperatures over time per trial should be plotted to obtain a curve. As calorimeters are not perfect and absorb/release energy, it will generate a graph that peaks and slowly returns to ambient temperature. To remedy this, the line returning the temperature to normal should be **linearly regressed** and extrapolated to the reaction start time to obtain a more accurate peak temperature.
Hess's law asserts that the change in enthalpy works like displacement - so long as the products and reactants are the same, any reaction with any number of intermediate steps will result in the same change in enthalpy.
$$\Delta H = \sum \Delta H \text{ of intermediate reactions}$$
