In this section, we will delve into the concept of standard enthalpy change, a fundamental aspect in the study of chemical reactions and thermodynamics. Understanding how energy changes during reactions is crucial for predicting reaction behaviour and manipulating conditions for desired outcomes.
Definition and Units
Standard Enthalpy Change (ΔH⦵) is defined as the heat change that results when molar quantities of reactants as stated in a balanced chemical equation react under standard conditions of 100 kPa and 298 K.
- Units: The SI unit for enthalpy change is the Joule (J), but it is commonly expressed in kJ mol(-1) in the context of chemical reactions.
Calculating Standard Enthalpy Change
Using Heat Capacity and Temperature Change
The equation Q = mcΔT is used to calculate the heat change in a reaction, where:
- Q is the heat energy absorbed or released,
- m is the mass of the substance,
- c is the specific heat capacity, and
- ΔT is the change in temperature.
To find the standard enthalpy change, the equation ΔH = − Q/n can be applied, where:
- ΔH is the enthalpy change,
- Q is the heat change,
- n is the amount of moles.
Image courtesy of the science hive
It is crucial to pay attention to the sign of ΔH:
- A negative ΔH indicates an exothermic reaction, where heat is released.
- A positive ΔH indicates an endothermic reaction, where heat is absorbed.
Experimental Investigation
- Combustion Reactions: Often carried out in a calorimeter to ensure heat exchange is measured accurately. For example, burning a known mass of a hydrocarbon and measuring the temperature change in water surrounding the combustion chamber.
- Discrepancies in Measurements: There may be differences between theoretical ΔH values and those measured experimentally due to heat losses to the surroundings, incomplete combustion, or other experimental errors. It is important to acknowledge and minimise these where possible.
Energy Profiles and Reaction Stability
Energy profiles provide a visual representation of the energy changes during a reaction, plotting potential energy against the reaction progress.
Endothermic Reactions
- Reactants are more stable: They have lower potential energy.
- Products have higher potential energy: This is because energy is absorbed from the surroundings.
- Profile Appearance: The energy of products is higher than that of the reactants, with an uphill curve from reactants to products.
Image courtesy of Brazosport College
Exothermic Reactions
- Reactants have higher potential energy: Indicating they are less stable.
- Energy is released to the surroundings: This makes the products more stable with lower potential energy.
- Profile Appearance: A downhill curve from reactants to products, reflecting the release of energy.
A graphical representation of exothermic reaction.
Image courtesy of Slashme
Axes Labelling
- Y-axis: Potential Energy
- X-axis: Reaction Progress
Combustion Reactions and Bonding in N2
Combustion reactions are typically exothermic, releasing energy in the form of heat and/or light.
- Bond Breaking and Making: Requires energy to break bonds in reactants and releases energy when new bonds form in products. For combustion, energy released in forming bonds is greater than that required to break bonds, resulting in an overall release of energy.
- Specific Look at N2 Bonding: Nitrogen gas (N2) has a triple bond, one of the strongest types of covalent bonds. Breaking this bond requires a substantial amount of energy. However, in combustion reactions involving nitrogen, such as the formation of NO or N2O, the energy released in forming new bonds is still greater than that required to break the N≡N bond, contributing to the exothermic nature of these reactions.
In summary, understanding standard enthalpy change and how it varies in different types of reactions is essential in predicting and controlling chemical processes. Energy profiles and the stability of reactants and products provide additional layers of insight, particularly in the context of combustion and the role of bonding in energy changes.
FAQ
In an exothermic reaction, the released energy primarily increases the kinetic energy of the surrounding particles, often leading to a rise in temperature. This released energy can be utilised in various ways depending on the context of the reaction. For instance, in combustion engines, the energy released during fuel combustion is converted into mechanical work to propel vehicles. In residential heating systems, the energy from exothermic reactions is used to warm living spaces. In biological systems, exothermic reactions contribute to maintaining body temperature and powering cellular processes.
Discrepancies in calorimetry experiments can be minimised through careful experimental design and execution. Using a well-calibrated and well-insulated calorimeter helps reduce heat loss to the surroundings, improving accuracy. Ensuring complete combustion of the sample and thorough mixing during the reaction helps achieve uniform temperature distribution, leading to more accurate temperature measurements. Minimising sources of systematic error, such as heat capacities of the calorimeter components, and using precise measuring instruments also contribute to reducing discrepancies. Additionally, conducting multiple trials and averaging the results can help mitigate the impact of random errors.
The standard enthalpy change of formation for elements in their standard states is defined as zero to provide a reference point for comparing the enthalpy changes of different substances. This convention simplifies calculations and allows chemists to predict the enthalpy change of a reaction based on the enthalpies of formation of the reactants and products. By setting the standard enthalpy of formation for elements in their most stable form at 298 K and 100 kPa to zero, it creates a consistent and straightforward way to determine the relative stability and reactivity of different compounds.
The standard conditions for measuring enthalpy change are 100 kPa of pressure and a temperature of 298 K (25°C). These conditions are chosen to provide a consistent basis for comparing the enthalpies of different reactions. Using standard conditions ensures that the values obtained are reproducible and can be reliably used in thermodynamic calculations, databases, and reference materials. This standardisation is essential for scientists and engineers worldwide to have a common reference point, facilitating communication and collaboration across different disciplines and industries.
Measuring standard enthalpy changes is crucial for understanding the energy changes associated with chemical reactions, which has implications in various fields such as industry, environmental science, and biochemistry. Knowing the standard enthalpy change enables chemists to predict the feasibility of reactions, design energy-efficient processes, and assess the environmental impact of reactions. For instance, in industrial settings, reactions with large negative enthalpy changes are preferred as they release energy, potentially lowering operational costs. In contrast, understanding reactions with positive enthalpy changes is vital for processes that store energy, such as in the formation of biological molecules.
Practice Questions
Standard enthalpy change (ΔH⦵) refers to the heat change that occurs when molar quantities of reactants, as indicated in a balanced chemical equation, react under standard conditions, specifically at 100 kPa and 298 K. Experimentally, ΔH⦵ can be determined using a calorimeter, a device that measures the amount of heat absorbed or released during a chemical reaction. In the context of a combustion reaction, a known mass of a substance, such as a hydrocarbon, is burned inside the calorimeter. The heat produced or consumed during the reaction leads to a temperature change in a known volume of water surrounding the combustion chamber. Using the formula Q = mcΔT, where Q is the heat change, m is the mass of the water, c is the specific heat capacity of water, and ΔT is the change in temperature, one can determine the heat change of the reaction.
Combustion reactions are typically exothermic because the energy released during the formation of new bonds in the products is greater than the energy required to break the bonds in the reactants. In chemical reactions, energy is always required to break existing bonds, a process that is endothermic. However, when new bonds form in the products, energy is released, which is an exothermic process. In combustion reactions, such as the burning of hydrocarbons in oxygen to produce carbon dioxide and water, the bonds formed in the carbon dioxide and water molecules release more energy than the energy needed to break the bonds in the hydrocarbon and oxygen molecules. As a result, the overall reaction releases energy, making it exothermic.
