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Jetzt kostenlos anmeldenEndothermic and exothermic processes are a part of everyday life. For example, when we put melted ice cream into the refrigerator, the reaction of the ice cream is to release heat into the surroundings. This transfer of heat energy from the melted ice cream into the surroundings results in the cooling of the ice cream. This is an exothermic, or heat releasing, process. Now if we transfer the cold ice cream into a freezer, even more heat energy is transferred to the surroundings resulting in the formation of hardened ice cream. The transformation of melted ice cream into hardened ice cream is also a heat-releasing or exothermic process. Yummy...
On the other hand, the refrigerant in your refrigerator causes a lowering of the temperature of the surroundings. The way that the refrigerant causes the lowering of the temperature of the surroundings is by transforming, or changing phase, from a liquid to a gas. This is a heat absorbing or endothermic process. You see, a change of phase from liquid to gas requires the absorption of heat energy from the surroundings. The engine in your refrigerator causes a liquid coolant to expand into a gas vapor, an endothermic process. This change of phase, from liquid to gas, causes cooling to take place in the freezer compartment.
In this article we will go over:
Let's start of by defining some key background terms:
There are three main types of Thermodynamic processes:
What is an endothermic and an exothermic process? How do endothermic and exothermic processes differ?
Both endothermic and exothermic processes involve the transfer of energy in the form of heat.
Let's classify the additional examples of processes as exothermic and endothermic.
1. Almost all of the machines currently in use, and the engines that drive them, are a result of the application of thermodynamics and involve endothermic and exothermic processes.
Figure 1: Thermodynamic model of an engine.
2. Chemistry labs are actually factories that manipulate heat to create substances with new properties and possibilities. Heat is a form of energy that is absorbed or released in Chemical Reactions. It is this energy, in the form of heat, that transforms one substance into another.
Figure 2: Modern chemistry laboratory
Consider the following enthalpy, or potential energy, diagrams for a set of hypothetical chemical reactions:
1. An endothermic chemical reaction:
Figure 3: Enthalpy diagram, endothermic reaction.
2. An exothermic chemical reaction:
Figure 4: Enthalpy diagram, exothermic reaction.
The above energy diagrams graph the enthalpy, or potential energies, associated with products and reactants. Reactants change to products through a reaction pathway that involves the addition of kinetic energy, in the form of heat, to the system. The enthalpy of the formation of a chemical substance can be viewed as being equivalent to the potential energy that is stored as heat within the chemical bonds of a compound. We note that:
When we look at the process of solvation, the process can be endothermic or exothermic.
Solvation - the process by which a solute (for example, table salt) dissolves into a solvent (like water) to form a solution.
Solute - the minor component of a solution.
Dissolve - to cause a solute to be incorporated into a liquid. (verb form: Dissolution)
Solvent - a substance which can dissolve a solute.
Driving Force - the Gibbs Free Energy difference, ΔG, associated with a reaction process.
Spontaneous (Spontaneous Change) - a process that occurs naturally without the input of external matter or energy into the system.
Figure 5: Model of how a solvent dissolves a solute to form a solution.
In the above figure, we note that the solute is often a salt which is held together by ionic bonds in a regularly ordered array (crystal lattice). On the other hand, within the solvent there are comparatively weaker interactions, such as Hydrogen Bonds, dipole-dipole interactions and London forces. In addition, the solvent is less ordered in some sense when compared to the solute.
The first step in the process of dissolving the solute into the solvent would involve the breaking of an ionic bond between a cation (positively charged) and an anion (negatively charged). This is accomplished by replacing each ionic bond with multiple solvent-solute interactions. This step in the solvation process involves the breaking of a strong bond by multiple weak interactions with the solvent and as a result is heat-releasing, or an exothermic reaction.
$$NaOH_{(s)} + H_2O_{(l)} \rightarrow Na^+_{(aq)} + OH^-_{(aq)} + H_2O_{(l)}$$
ii. This is a strongly exothermic process and results from the breaking of the sodium hydroxide ionic bond by multiple interactions with water:
Figure 6: Model of dissolution of sodium hydroxide in water.
iii. The overall process of the dissolution of sodium hydroxide in water can also be depicted in an enthalpy diagram:
Figure 7: Enthalpy diagram for the dissolution of sodium hydroxide.
Now, let's consider the thermochemistry for the balanced chemical equation for the dissolution of sodium hydroxide in water:
$$NaOH_{(s)} + H_2O_{(l)} \rightarrow Na^+_{(aq)} + OH^-_{(aq)} + H_2O_{(l)}$$
Or just,
$$NaOH_{(s)} \rightarrow Na^+_{(aq)} + OH^-_{(aq)}$$
Let's calculate the final temperature of the water for this reaction.
Given that the molar heat of solution (ΔHsolution ) of sodium hydroxide (NaOH) is ΔHsolution = -44.51 kJ/mol, we want to calculate the final temperature of the water (Tf ) contained in an insulated container, after 45 grams of NaOH is dissolved into it.
i. Let the initial temperature of the water be 20.0°C and a the volume of the water be 1L. The formula for the temperature difference (ΔT ) for the dissolution of NaOH in water is:
$$\Delta T=\frac{\Delta H_{solution\,NaOH}}{C_{p\,(H_2O)}*m_{tot}}$$
Where:
First, we calculate the molar heat of solution for 45 grams of sodium hydroxide:
$$\Delta H_{solution\,NaOH}=(45\,g\,NaOH)*\frac{1\,mol\,NaOH}{40.00\frac{g}{mol}}*\frac{-44.51\,kj}{1\,mol\,NaOH}*\frac{1000\,J}{1\,kJ}=-5.564x10^4\,J$$
Where the molar mass of sodium hydroxide is 40.00 g/mol.
Notice, that the molar heat of solution for this reaction (ΔHsolution, NaOH = -55.64 kJ/mol) is negative, which shows that the reaction is exothermic.
ii. Now we insert this value of the molar heat of solution for 45 grams of sodium hydroxide into the temperature difference formula:
$$\Delta T=\frac{5.564x10^4\,J}{(4.18\frac{J}{g\,^\circ C})*(1000\,g\,H_2O+45\,g\,NaOH)}=12.7^\circ C$$
The thermochemistry data2 is included in the table below:
| Compound | ΔH°f [kJ/mol] | ΔS° [kJ/mol] |
| Ba(OH)2 · 8H2O (s) | -3345 | 0.427 |
| NH4Cl (s) | -314 | 0.095 |
| NH3 (g) | -46 | 0.192 |
| H2O (l) | -286 | 0.070 |
| BaCl2 (s) | -859 | 0.124 |
$$\Delta H^\circ =[2*\Delta H^\circ_{f\,(g)\,NH_3}+10*\Delta H^\circ_{f\,(l)\,H_2O}+\Delta H^\circ_{f\,(s)\,BaCl_2}]-[\Delta H^\circ_{f\,(s)\,Ba(OH)_2*8H_2O}+2*\Delta H^\circ_{f\,(s)\,NH_4Cl}]$$
Then, inserting the table values:
$$\Delta H^\circ=[2*(-46\frac{kJ}{mol})+10*(-286\frac{kJ}{mol})+(-859\frac{kJ}{mol})]-[(-3345\frac{kJ}{mol})+2*(-314\frac{kJ}{mol}]=162\frac{kJ}{mol}$$
Notice, that the positive value for the enthalpy of this reaction, ΔH° = +162 kJ/mol, shows that the reaction is endothermic.
ii. Calculation of the reaction Entropy, ΔS° :
$$\Delta S^\circ = [2*\Delta S^\circ_{(g)\,NH_3}+10*\Delta S^\circ_{(l)\,H_2O}+\Delta S^\circ_{(s)\,BaCl_2}]-[\Delta S^\circ_{(s)\,Ba(OH)_2*8H_2O}+2*\Delta S^\circ_{(s)\,NH_4Cl}]$$
Then, inserting the table values:
$$\Delta S^\circ = [2*(0.192\frac{kJ}{K*mol}+10*(0.070\frac{kJ}{K*mol})+(0.124\frac{kJ}{K*mol})]-[(0.427\frac{kJ}{K*mol})+2*(0.095\frac{kJ}{K*mol})]=0.591\frac{kJ}{K*mol}$$
iii. Lastly, we calculate the Gibbs Free Energy difference, ΔGº :
$$\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ = 162\frac{kJ}{mol}-(298\,K*0.591\frac{kJ}{K*mol})=-14.12\frac{kJ}{mol}$$
Notice, that the negative value for the Gibbs free energy difference of this reaction, ΔG° = -14.12 kJ/mol, shows that the reaction is spontaneous, or favorable.
Can a process be exothermic and endothermic? Answer: Yes, a process can contain a series of chemical reactions that are individually either endothermic or exothermic. For example, the biochemical process of glycolysis includes at least 12 different reactions, a few of which are endothermic with the majority of the other reactions in process being exothermic.
Endothermic processes are those which absorb energy as heat. Exothermic processes are those which release energy as heat.
Yes, thermodynamic processes can be both exothermic and endothermic at different points in the process.
Chemical reactions can involve a sequence of endothermic and exothermic processes, but a chemical reaction cannot be both endothermic and exothermic at the same time.
Endothermic processes are those which absorb energy as heat. Exothermic processes are those which release energy as heat.
Solvation processes which absorb energy as heat are endothermic. Solvation processes which release energy as heat are exothermic processes.
Flashcards in Endothermic and Exothermic Processes11
Start learningWhat are endothermic and exothermic processes?
Endothermic processes are those which absorb energy as heat. Exothermic processes are those which release energy as heat.
Can a process be both exothermic and endothermic?
Yes, thermodynamic processes can be both exothermic and endothermic at different points in the process.
Can a reaction be both exothermic and endothermic?
Chemical reactions can involve a sequence of endothermic and exothermic processes, but a chemical reaction cannot be both endothermic and exothermic at the same time.
How do endothermic and exothermic processes differ?
Endothermic processes are those which absorb energy as heat. Exothermic processes are those which release energy as heat.
Why are some solvation processes endothermic and others exothermic?
Solvation processes which absorb energy as heat are endothermic. Solvation processes which release energy as heat are exothermic processes.
A negative value for the Gibbs free energy difference indicates what?
A spontaneous reaction.
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