9.5 KiB
HL Chemistry - 2
The course code for this page is SCH4UZ.
Thermal concepts
!!! definition - 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.
!!! info Acid-base reactions are always exothermic.
Specific heat capacity
Please see SL Physics 1#3.1 - Thermal concepts for more information.
Enthalpy
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. \[\ce{H2O_{(l)} -> H2_{(g)} + 1/2 O2_{(g)}\ \ \Delta H = +280 kJ}\]
\(\Delta H\) 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: \[\ce{H2O_{(l)} + 285.5 kJ -> H2_{(g)} + 1/2 O2_{(g)}}\]
(Standard) Molar enthalpy of reaction
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.
\(\Delta H_\text{combustion} = \ce{-1367
kJ/mol C2H6O}\)
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).
Bond enthalpies
\(\Delta H_B\), 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: \[\Delta H = \sum n\Delta H_B\text{reactants} - \sum n\Delta H_B\text{products}\]
Calorimetry
!!! definition - A calorimeter measures changes in energy.
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.
(Source:
Kognity)
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
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. \[\Delta H = \sum \Delta H \text{ of intermediate reactions}\]
Formation equations
A formation equation is a balanced chemical equation where exactly one mole of product and its reactants in elemental form are in their diatomic state — -gens are diatomic, phosphorus is \(\ce{P4}\), sulfur is \(\ce{S8}\), and at SATP (25°C, 100 kPa).
!!! info Fractions are permitted as coefficients on the reactant side to get exactly one mole of product.
!!! example \[\ce{6C_{(s)} + 6H2_{(g)} + 3O2_{(g)} -> C6H12O6}\] \[\ce{2C_{(s)} + 3/2 H2_{(g)} + 1/2 Cl2_{(g)} -> C2H3Cl_{(g)}}\]
The standard enthalpy of formation \(\Delta H^\theta_f is the energy change from the formation of one mole of its substance from its elements in their standard states. It can be determined by subtracting the sum of the enthalpy of each element/compound on the reactant side and adding those on the product side.\)\(\Delta H = \sum n\Delta H\text{ products} - \sum n\Delta H\text{ reactants}\)$
!!! warning It is assumed that there is no state change that would affect enthalpy when calculating standard enthalpy of formation.
Enthalpy cycles
Enthalpy cycles are a visual representation of Hess’s law. It is used to show that the energy is the same from initial reactants to a product regardless of any intermediate steps.
!!! example \(\Delta H_1 = \Delta H_2 +
\Delta H_3\). Note that both arrows point to the intermediate
product.
(Source:
Kognity)