standard enthalpy of formation

Examples of standard enthalpy of formation in the following topics:

• Standard Enthalpy of Reaction

  • In order to calculate the standard enthalpy of a reaction, we can sum up the standard enthalpies of formation of the reactants and subtract this from the sum of the standard enthalpies of formation of the products.
  • In order to calculate the standard enthalpy of reaction, we need to look up the standard enthalpies of formation for each of the reactants and products involved in the reaction.
  • Note that because it exists in its standard state, the standard enthalpy of formation for oxygen gas is 0 kJ/mol.
  • Next, we sum up our standard enthalpies of formation.
  • A calculation of standard enthalpy of reaction (∆H°rxn) from standard heats of formation (∆H°f)

• Standard States and Standard Enthalpy Changes

  • The standard enthalpy of formation refers to the enthalpy change when one mole of a compound is formed from its elements.
  • The standard enthalpy of formation, or standard heat of formation, of a compound is the change in enthalpy that accompanies the formation of one mole of the compound from its elements in their standard states.
  • For example, the standard enthalpy of formation for carbon dioxide would be the change in enthalpy for the following reaction:
  • Note that standard enthalpies of formation are always given in units of kJ/mol of the compound formed.
  • Graphite is the most stable state of carbon and is used in thermochemistry to define the heat of formation of carbon compounds.

• Bond Enthalpy

  • The total enthalpy, H, of a system cannot be measured directly.
  • Generally, a positive change in enthalpy is required to break a bond, while a negative change in enthalpy is accompanied by the formation of a bond.
  • In other words, breaking a bond is an endothermic process, while the formation of bonds is exothermic.
  • Bond enthalpy, also known as bond dissociation energy, is defined as the standard enthalpy change when a bond is cleaved by homolysis, with reactants and products of the homolysis reaction at 0 K (absolute zero).
  • Describe the changes in enthalpy accompanying the breaking or formation of a bond

• Change in Enthalpy

  • Enthalpy (H) is a measure of the total energy of a thermodynamic system.
  • By absorbing heat, the temperature, and thus the enthalpy of a substance increases.
  • Hess's law states that the standard enthalpy change of the overall reaction is the sum of the enthalpy change of all the intermediate reactions that make up the overall reaction.
  • This lesson introduces Enthalpy and the energy of chemical bonding.
  • We discuss where the energy in chemical bonds comes from in terms of internal energy and enthalpy, as well as how to approximate the overall heat of reaction using bond enthalpies.

• Hess's Law

  • Hess's Law sums the changes in enthalpy for a series of intermediate reaction steps to find the overall change in enthalpy for a reaction.
  • This law states that if a reaction takes place in several steps, then the standard reaction enthalpy for the overall reaction is equal to the sum of the standard enthalpies of the intermediate reaction steps, assuming each step takes place at the same temperature.
  • However, because we know the standard enthalpy change for the oxidation for these two substances, it is possible to calculate the enthalpy change for this reaction using Hess's law.
  • First it looks at combining reactions according to Hess's law and their heats of reaction, and then it discusses using standard heats of formation of the reactants and products to find the overall heat of reaction.
  • By Hess's law, the net change in enthalpy of the overall reaction is equal to the sum of the changes in enthalpy for each intermediate transformation: ΔH = ΔH1+ΔH2+ΔH3.

• Standard Free Energy Changes

  • The standard Gibbs Free Energy is calculated using the free energy of formation of each component of a reaction at standard pressure.
  • These same definitions apply to standard enthalpies and internal energies.
  • To accomplish this, combine the standard enthalpy and the standard entropy of a substance to get the standard free energy of a reaction:
  • Standard Gibbs free energies of formation are normally found directly from tables.
  • The standard Gibbs free energy of formation of a compound is the change of Gibbs free energy that accompanies the formation of 1 mole of that substance from its component elements, at their standard states.

• Heat of Solution

  • Heat of solution refers to the change in enthalpy when a solute is dissolved into a solvent.
  • The heat of solution, also referred to the enthalpy of solution or enthalpy of dissolution, is the enthalpy change associated with the dissolution of a solute in a solvent at constant pressure, resulting in infinite dilution.
  • The heat of solution, like all enthalpy changes, is expressed in kJ/mol for a reaction taking place at standard conditions (298.15 K and 1 bar).
  • The heat of solution can be regarded as the sum of the enthalpy changes of three intermediate steps:
  • Solute-solvent attractive bond formation (the exothermic step in the process of solvation) is indicated by dashed lines.

• Internal Energy and Enthalpy

  • The enthalpy of reaction measures the heat released/absorbed by a reaction that occurs at constant pressure.
  • To correct for this, we introduce the concept of enthalpy, which is much more commonly used by chemists.
  • The enthalpy of reaction is defined as the internal energy of the reaction system, plus the product of pressure and volume.
  • Due to this relation, the change in enthalpy is often referred to simply as the "heat of reaction."
  • An explanation of why enthalpy can be viewed as "heat content" in a constant pressure system.

• Thermochemical Equations

  • Thermochemical equations are chemical equations which include the enthalpy change of the reaction, $\Delta H_{rxn}$ .
  • Enthalpy (H) is a measure of the energy in a system, and the change in enthalpy is denoted by $\Delta H$.
  • Since enthalpy is a state function, the value of $\Delta H$ is independent of the path taken by the reactions to reach the products.
  • Values of $\Delta H$ can be determined experimentally under standard conditions.
  • A thermochemical equation is a balanced stoichiometric chemical equation which includes the enthalpy change.

• Changes in Temperature

  • Reactions can be classified by their enthalpies of reaction.
  • Reactions with positive enthalpies—those that absorb heat from their surroundings—are known as endothermic.
  • In contrast, reactions with negative enthalpies—those that release heat into their surroundings—are known as exothermic.
  • Le Chatelier's Principle predicts that the addition of products or the removal of reactants from a system will reverse the direction of a reaction, while the addition of reactants or the removal of products from a system will push the reaction towards the formation of products.
  • Endothermic reactions, on the other hand, will be shifted towards product formation as heat is removed from the reaction's surrounding environment.