5 - Energetics

• Energetics is the study of heat changes in chemical reactions.
• Heat is a form of energy.
• Total energy is always conserved and can only be transferred.
• Heat energy always flows from a higher temperature object to a lower temperature object.
• The Kelvin scale is based on kinetic energy, so 0K means that there is no kinetic movement at all.
• Enthalpy (H) is the amount of heat energy contained in a substance.
• Enthalpy is stored in the chemical bonds as potential energy.
• When substances react, the total enthalpy of a system cannot be measured due to loss of heat, but it is possible to measure the difference in the enthalpy between the reactants and products.
• Enthalpy is denoted as H, however heat change is denoted as ΔH.
• The enthalpy change for chemical reactions is denoted kJ mol - 1.
• The reaction mixture is called the system (the chemical reaction), which gives heat to or takes heat from the surroundings (anything around the system).
• Chemical reactions that involve transfer of heat between system and the surroundings are described as exothermic and endothermic.
• In exothermic reactions, heat is released to the surroundings because more heat energy is released than what is added, causing the overall heat energy to be released from the system, causing the surroundings to become hotter.
• Exothermic reactions have negative ΔH values because heat is released, thus enthalpy decreases, -ΔH.
• In an exothermic reaction, the products are more stable than the reactants as they have a lower enthalpy, meaning the reaction is downhill in terms of heat energy.
• Exothermic reactions release energy (as heat).
• Calorimetry is the process of measuring the amount of heat released or absorbed during a chemical reaction.
• When the reaction occurs, it is either going to give off or take in heat.
• The change in heat can be measured by observing the temperature of water, as water can serve as the surroundings.
• There are many ways to perform calorimetry, the most common is to use a bomb shell.
• Calorimetry is performed many different ways, but when using calorimetry several assumptions are made due to the potential for heat loss and incomplete combustion to lead to systematic errors in experimental results.
• These assumptions include that all the heat is transferred to the water, that all the solution has dissolved, the mass of the water remained constant, there was an unlimited oxygen source, and the experiment was under standard conditions.
• To compensate for heat lost by the water, the temperature rise that would have taken place had the reaction been instantaneous can be calculated by taking 𝑇5 − 𝑇6.
• Hess’s Law states that regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes.
• This means the enthalpy change going from A to B is the same whether the reaction proceeds directly to A or whether it goes via A to C then B.
• Average bond enthalpy is the energy required to break one mole of the same type of bond in the gaseous state averaged over a variety of similar compounds.
• Energy is released through the formation of chemical bonds.
• Bond forming: Removing heat brings atoms closer together, forming bonds.
• Releases energy, ∆𝐻 negative therefore exothermic.
• Energy is required when breaking chemical bonds.
• Bond breaking: Adding heat separates atoms, breaks bonds.
• Requires energy, ∆𝐻 positive therefore endothermic.
• The enthalpy of reaction (∆𝐻 $) can be calculated using bond enthalpies.
• Since there are average bond enthalpies, the calculation gives only an approximate result.
• Examples of exothermic reactions include bond forming, where removing heat brings atoms closer together, forming bonds, and gas-liquid-solid reactions, where heat is removed, these are exothermic reactions.
• Thermochemical equations give the balanced equation with an enthalpy charge.
• State symbols must be shown, as ∆H depends on the state of the reactants or products.
• Heat changes (enthalpy) can be calculated from the temperature changes: ΔH = 𝑚𝑐 ∆𝑇.
• If the temperature of the compound increased, then the reaction is endothermic and ΔH must be positive.
• If the temperature of the compound decreased, then the reaction is exothermic and ΔH must be negative.