Enthalpy of Formation
Enthalpy of formation refers to the heat change that occurs when one mole of a substance is formed from its elements in their standard states. It is a measure of the stability of a compound and is often used to calculate the heat of reaction in chemical processes. The enthalpy of formation values are useful in understanding and predicting the energy changes in chemical reactions.
7 Key excerpts on "Enthalpy of Formation"
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...The standard state of an element or a compound is defined as the most stable form of the physical state that exists at 1 atmosphere pressure and 298 K. For the chemical reaction that results in the formation of a molecule directly from its elements, the standard enthalpy is called the standard Enthalpy of Formation or standard heat of formation (Δ H ° f). Other types of standard enthalpies include the standard enthalpy of neutralization (Δ H ° neut), which is the change in enthalpy that occurs when an acid and a base undergo a neutralization reaction to form one mole of water and a salt, and standard enthalpy of combustion (Δ H ° comb), which is the enthalpy change that occurs when one mole of the compound is burned completely in oxygen. Here we will focus mainly on the standard enthalpies of formation. Standard enthalpies of formation are very useful in determining whether a compound is stable at the standard room temperature of 298 K. They are also useful in determining the enthalpies of formation for unknown reactions or reactions that are not easy to measure quantitatively. The standard Enthalpy of Formation is defined as the change in enthalpy that occurs during the formation of one mole of a compound from its elements, with all substances in their standard states. So, the chemical equation describing the formation must be written with all species in their standard states and the stoichiometry of the reaction must be set so that that one mole of product is produced. For example, the reaction of solid carbon with oxygen to form carbon monoxide is normally written as; 2C s + O 2 g → 2CO g However, the stoichiometry of this equation shows two moles of carbon monoxide being produced...
eBook - ePub- Atef Korchef(Author)
- 2022(Publication Date)
- CRC Press(Publisher)
...They release energy. 13. A phase change occurs when we cross the lines or the curves on the phase diagram. The temperature remains constant during a phase change (Figure 5.9) and energy is used to overcome attractive forces between molecules. 14. For a chemical reaction, when Δ H rxn < 0 the reaction is exothermic, when Δ H rxn > 0 the reaction is endothermic and when Δ H rxn = 0 the reaction is athermic (no heat is gained or lost). 15. The standard Enthalpy of Formation, designated by Δ H f °, is the change in enthalpy when one mole of a substance is formed under standard conditions (P = 1 atm and T = 25°C) from its pure elements under the same standard conditions. Conventionally, the standard Enthalpy of Formation of a pure element in its most stable form is zero. 16. Bond breaking requires energy. The energy required to break one mole of that chemical bond is called the bond enthalpy or bond dissociation enthalpy. It is also a measure of the bond strength. Bond enthalpy values are always positive since bond breaking is an endothermic process. However, bond making is an exothermic process (it releases energy). Note that a chemical reaction can be described as the breaking of bonds in the reactants and the making of bonds in the products. Thus, if the bond enthalpies of the reactants and products are known, we can calculate the standard enthalpy of the reaction, Δ H rxn ° : Δ H rxn ° = ∑ Δ H broken bonds − ∑ Δ H formed bonds 17. Hess’s law states that, if a chemical reaction is carried out in a series of steps, the enthalpy change ΔH for the overall reaction is equal to the sum of the enthalpy changes for the individual steps...
eBook - ePub- Warren C. Strahle, William A. Sirignano, William A. Sirignano(Authors)
- 2020(Publication Date)
- Routledge(Publisher)
...In the above hydrogen—fluorine example the substances, in fact, behave as ideal gases so that we may write at the 1 bar, 298 K condition The Q P comes out as a difference between the enthalpies of the product and the elements from which it is formed. In principle, all compounds can be formed by reactions by their constituent elements. In thermodynamics we never know absolute values of energy variables, but that is not important because only changes in these variables are dealt with in thermodynamics. To facilitate numerical work, we define a reference set of substances and their thermodynamic properties. We define reference chemicals as the elements in the form most abundantly found in nature when they are found alone. So, for example, oxygen in the air is found as the diatomic molecule O 2 in the gas phase. The reference state for oxygen is the gaseous diatomic molecule. Similarly, reference states for fluorine, hydrogen, and nitrogen are gaseous F 2, H 2, and N 2 respectively, not F, H, and N. By contrast, the reference state for carbon is C(s), or graphite. We define formation reactions for a compound or atom as the reaction that forms one mole of the substance from the elements or element in their reference state. The heat of formation for a substance is the standard state enthalpy change for a reaction forming the substance from its reference state elements. Moreover, for thermodynamic reference the heat of formation of the elements in their standard states is defined as zero. The heat of formation may be defined at any temperature and for species i. It is denoted as In the above hydrogen—fluorine reaction it is clear that it was a formation reaction at 298 K, since hydrogen and fluorine behave perfectly at the stated conditions...
eBook - ePub- Adrian Dingle, Derrick C. Wood(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...Also called ΔH° f. Enthalpy Change for a Reaction Remember, when an element is formed from itself, there is no change. The standard Enthalpy of Formation of elements in their standard states is zero. III. Hess’s Law A. Manipulating Chemical Equations 1. Hess’s Law states that the overall enthalpy change in a reaction is the sum of all the reactions for the process and is independent of the route taken. i. Rule 1: If you reverse the reactions, then change the sign of ΔH. For example, ii. Rule 2: If you multiply the reaction by a coefficient, multiply the value of ΔH by the same coefficient. For example, iii. Rule 1 and 2 can be combined. For example, if the first reaction is tripled and reversed, Practice Questions 1. Based on the equation below, what is the enthalpy change when 56.4 g of C 2 H 5 OH(l) decomposes? 2. What is the standard enthalpy change, Δ H ° rxn, for the reaction below? The enthalpies of formation: 3. Given the data below, calculate the enthalpy change for the decomposition of phosphorous trichloride. Answers 1. 2. Enthalpy Change for a Reaction = Enthalpy Change for a Reaction 3. Reverse equation #2: Multiply equation #3 by 4: Sum the new equations and cancel out 4PCl 5 (g) and 4Cl 2 (g) from both sides: Chapter 21 Inter and Intra Forces and Physical and Chemical Change I. Chemical and Physical Change A. Differences at the Particulate Level 1. When weak intermolecular forces (between molecules) are broken or formed, physical changes take place. 2. When strong intra bonds (chemical bonds within compounds) are broken or formed, chemical changes take place. 3. Some changes (e.g.,. dissolving an ionic salt in water) involve both intermolecular and chemical bonds changing, and, as such, can be classified as chemical and/or physical changes. 4. In large molecules (particularly those encountered in biochemistry), intermolecular forces can occur between different parts of the same molecule...
- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...ΔH° represents the standard change in enthalpy; ΔS° represents the standard change in entropy; and ΔG° represents the standard change in free energy. Standard conditions are a very specific set of conditions: • The temperature is 25°C, or 298 K. • Gases are at 1 atmosphere of pressure. • Solutions are 1 mole/liter (M) in concentration. • Liquids and solids are pure. TEST TIP Do not confuse thermodynamic standard conditions, where the temperature is 298 K, with standard temperature and pressure (STP) for gases, where the temperature is 273 K. Heats of Formation The standard heat of formation refers to the energy associated with forming a compound from its elements in their standard states. The for an element at standard conditions is zero, because by definition it requires no energy to create an element from itself. In addition to calorimetry, values can be used to determine the enthalpy of a reaction based on the equation: EXAMPLE: Using the following values, determine the amount of energy released when sugar (C 12 H 22 O 11) is burned to produce gaseous carbon dioxide and water vapor. Compound ΔHf° (kJ/mol) C 12 H 22 O 11(s) –2,221 H 2 O (g) –241.81 Co 2(g) –393.5 SOLUTION: The Enthalpy of Formation for oxygen gas is not given because it is zero (all elements in their standard states are zero). TEST TIP If you encounter a problem where the values are not given for one or more substances in a reaction, do not dismay. They are probably elements in their standards states, which have a value of zero. Hess’s Law In addition to calorimetry and heats of formation, Hess’s law is a third way to determine the ΔH° of a reaction. Hess’s law takes advantage of the fact that enthalpy is a state function: you can find the ΔH° for a reaction by adding up the enthalpy values from multiple steps so that their sum results in the enthalpy change for the overall chemical reaction...
eBook - ePubFoundations for Teaching Chemistry
Chemical Knowledge for Teaching
- Keith S. Taber(Author)
- 2019(Publication Date)
- Routledge(Publisher)
...The energy released as solvation occurs does not compensate for the energy needed to break bonds (e.g., between ions in the salt). Yet the process is still energetically favourable (the total free energy decreases) as the total number of ways that the energy quanta in the system can be distributed across the species in the mixture is far greater than the number of options in the separate solid and liquid before dissolving. That is, the increase in entropy (ΔS) (T) is sufficient to give an overall negative free energy change (ΔG) at that temperature despite an increase in enthalpy (heat entering the system, ΔH) as ΔG = ΔH – TΔS. Students at the highest levels of school chemistry sometimes have to study some basic thermodynamics and so appreciate the nature and significance of entropy to the overall free energy changes in various processes. It is probably sensible to ignore entropy in introductory teaching, and this will be less problematic when students appreciate that chemistry is understood in terms of models, which involve simplifications and so will sometimes fail to explain all examples or details of a phenomena (see Chapter 3). Chemical reactions Chemistry is the science that studies the properties of different substances and, in particular, their chemical properties: which substances react with which other substances, under what conditions, and what new substances are produced. Like many central ideas in chemistry, chemical reactions are understood at two levels of description (see Chapter 4)...
eBook - ePub- Antonio Blanco, Gustavo Blanco(Authors)
- 2017(Publication Date)
- Academic Press(Publisher)
...The magnitude of this “heat of combustion” depends on the molecular structure of the substance and the remaining energy content in the products formed. If a mole (342 g) of common sugar (sucrose) is oxidized to CO 2 and H 2 O, 1350 kcal (5648 kJ) are produced (the potential energy of sucrose is represented by the bonds and configuration of the molecule; the energy content of the products CO 2 and H 2 O is much lower). Heat from the combustion of a substance is the maximum energy that can be obtained from complete oxidation of the elements of that substance. When the reaction takes place at a constant pressure, the change in heat produced is called enthalpy change and is symbolized by the notation ∆ H. It has a negative sign when heat is released. Enthalpy (H) is the caloric energy released or consumed in a system at constant temperature and pressure. When the volume is not changed, that is, when no work is done, the energy change (∆ E) equals the change in enthalpy (∆ H). Heat is the simplest manner by which energy can be applied to produce work. Typical examples are represented by engines and machines. However, in biological systems which operate at constant temperature and pressure, heat change is not a source of energy. Second Law of Thermodynamics The first law of thermodynamics allows us to understand and determine the energy changes that occur during a chemical reaction, but cannot predict its direction. One might think that a reaction always proceeds in the direction in which ∆ H is negative (exothermic) and does not occur when ∆ H is positive (endothermic). But this is not the case because ∆ H is not the only factor that determines the direction of a reaction...
