II. THERMOCHEMISTRY
Almost every physical or chemical change either absorbs or releases energy, most often in the form of heat. Where does this heat come from, and where does it go? The science of thermochemistry attempts to answer these questions. Thermochemistry is the study of heat exchanges that occur during phase changes and chemical reactions.
II.A Heat Changes in Reactions Bunsen burners in the laboratory utilize the combustion of natural gas (methane, or CH4) to produce heat according to the following equation:
CH4(g) + 2O2(g) --> CO2(g) + 2H2O(g) + 891 kJ The heat produced in the reaction is shown as a product, on the right side of the equation. This is an example of an exothermic reaction, one in which heat is released from the reaction to the surroundings. In an exothermic reaction, heat energy can be considered a product of the reaction. In the study of thermochemistry, we must carefully define the system as the chemicals or substance involved in the reaction or physical change. Everything in the universe outside of the system is called the surroundings. Therefore, the universe is defined as follows: universe = system + surroundings In the combustion of methane, heat is released to the surroundings from the system. In an endothermic reaction, however, heat is absorbed from the surroundings and can therefore be considered a reactant, required for the reaction to proceed. The following endothermic reaction is commonly used in instant medical cold-packs for the treatment of injuries (Fig. II.A.3). A small vial of water enclosed in a larger package containing granular ammonium nitrate is broken and mixed. The ammonium nitrate dissolves and an endothermic reaction occurs: 26 kJ + NH4NO3 (s) → NH4+ (aq) + NO3- (aq) In this endothermic reaction, heat is absorbed from the surroundings. Medical cold-packs remove heat from the area of injury to provide the energy necessary for the reaction to occur. The energy change in a system during the course of a reaction is symbolized by chemists as ΔH. The H symbol is a quantity called enthalpy, the heat content of a system at a constant pressure, which is often complicated, unknown, and possibly immeasurable. However, the change in heat during a reaction can be easily be measured using a calorimeter. The change in enthalpy for a reaction is called the enthalpy (heat) of reaction (ΔHrxn). ΔHrxn is the difference between the enthalpy of the products coming out of the reaction versus the enthalpy of the reactants going into the reaction. |
Eqn. 3: ΔHrxn = ΔHproducts - ΔHreactants
II.B Sign of Enthalpy of Reaction
Recall the reaction equation for the combustion of methane:
CH4(g) + 2O2(g) --> CO2(g) + 2H2O(g) + 891 kJ
Heat is released during this combustion reaction. In other words, the reactants lose energy as they are transformed to the products; the products therefore have less energy than the reactants. In other words, ΔHreactants > ΔHproducts. If the values for ΔHreactants and ΔHproducts are substituted into Equation 2, the answer would be a negative number and the reaction equation can be re-written as follows:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ΔHrxn = -891 kJ
Enthalpy changes for exothermic reactions are always negative, because the products always contain less energy than the reactants. Similarly, the sign for ΔHrxn for endothermic equations is always positive, because the products always contain more energy than the reactants. The medical cold-pack reaction can be re-written as:
NH4NO3(s) → NH4+(aq) + NO3-(aq) ΔHrxn = +26 kJ
In both of the reactions described above, the exo- or endothermic nature of the reaction can be determined based on the sign for ΔHrxn.
CH4(g) + 2O2(g) --> CO2(g) + 2H2O(g) + 891 kJ
Heat is released during this combustion reaction. In other words, the reactants lose energy as they are transformed to the products; the products therefore have less energy than the reactants. In other words, ΔHreactants > ΔHproducts. If the values for ΔHreactants and ΔHproducts are substituted into Equation 2, the answer would be a negative number and the reaction equation can be re-written as follows:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ΔHrxn = -891 kJ
Enthalpy changes for exothermic reactions are always negative, because the products always contain less energy than the reactants. Similarly, the sign for ΔHrxn for endothermic equations is always positive, because the products always contain more energy than the reactants. The medical cold-pack reaction can be re-written as:
NH4NO3(s) → NH4+(aq) + NO3-(aq) ΔHrxn = +26 kJ
In both of the reactions described above, the exo- or endothermic nature of the reaction can be determined based on the sign for ΔHrxn.
EXAMPLE: When one mole of ethanol (C2H5OH) is combusted, 1,367 kJ of heat are released. Write a thermochemical equation for the reaction.
C2H5OH(l) + 7/2O2(g) → 2CO2(g) + 3H2O(l) ΔHrxn = -1,367 kJ
C2H5OH(l) + 7/2O2(g) → 2CO2(g) + 3H2O(l) ΔHrxn = -1,367 kJ
II.C Thermochemical Equations
When the term for ΔHrxn is included in the reaction expression, the result is called a thermochemical equation, a balanced chemical equation that includes all states of matter and the value of the energy change, normally expressed as the change in enthalpy (ΔHrxn). Thermochemical equations can be written in two possible formats. One includes the value for ΔHrxn stated after the balanced chemical equation. Another possible format includes the value of ΔHrxn as a reactant (in an endothermic reaction) or as a product (in an exothermic reaction). In the previous section, both correct formats are shown for the combustion of methane, an exothermic reaction. It is important to state, however, that in thermochemical equations, the coefficients always refer to the number of moles and not the number of particles.
Endothermic and exothermic reactions can be depicted in reaction diagrams, shown in Figure II.C.1. In an exothermic reaction, the products have less energy than the reactants, and vice versa for an endothermic reaction. Note that the value of ΔHrxn can be shown on the diagram as the vertical change in energy content from the reactants to the products.
Certain reactions are so common in chemistry that they have their own symbol for ΔH. For example, combustion is a reaction of a substance with oxygen that is always exothermic. The standard enthalpy of combustion (ΔHcomb) is the enthalpy change for the complete burning of one mole of a substance; because it is always exothermic, the sign for ΔHcomb is always negative. Standard enthalpy changes are carried out under standard conditions in a laboratory (1 atm and 298K) and are symbolized using a zero superscript (ΔHo). Standard enthalpies of combustion for several common fuels and other substances are listed in Table 2. The energy released during a combustion reaction can be calculated using values from Table 2. EXAMPLE PROBLEM: Your backyard barbeque uses propane (C3H8) to fuel the grill. How much heat is evolved (in Joules) when 100g of propane is used to grill a couple of burgers? This problem is solved in the VoiceThread shown below, or click here to view it directly.
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II.D Changes in State Not only chemical changes can absorb or release heat. Changes in states of matter (i.e., phase changes) will also require a flow of heat into or out of the system. Endothermic phase changes would include those changes that result in an increased energy state of the products, such as melting (fusion) or vaporization. The heat required to vaporize one mole of a substance is called the molar heat of vaporization (ΔHvap), and the heat required to melt one mole of a substance is called the molar heat of fusion (ΔHfus). Because both of these processes are endothermic, the sign for ΔH is positive. Standard molar enthalpies of fusion and vaporization for a few common substances are shown in Table 3.
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Similarly, exothermic phase changes, such as condensation or freezing (solidification), would release energy, resulting in a lower energy state than the initial phase. The same amounts of energy are released in the exothermic processes as are absorbed in the endothermic processes, such that the molar enthalpies of fusion and solidification have the same numerical value but opposite signs:
ΔHvap = - ΔHcond ΔHfus = - ΔHsolid
In fact, because any chemical reaction or phase change is in theory reversible, the sign for ΔH can be reversed for the opposite process.
The overall energy required for heating or cooling a substance to a phase change can now be calculated. Recall that any problem that involves the heating or cooling of a single phase of matter, without a phase change, must also include the specific heat of the substance to calculate the total amount of heat absorbed or released.
ΔHvap = - ΔHcond ΔHfus = - ΔHsolid
In fact, because any chemical reaction or phase change is in theory reversible, the sign for ΔH can be reversed for the opposite process.
The overall energy required for heating or cooling a substance to a phase change can now be calculated. Recall that any problem that involves the heating or cooling of a single phase of matter, without a phase change, must also include the specific heat of the substance to calculate the total amount of heat absorbed or released.
EXAMPLE PROBLEM: How much energy would be required if 30g of water were heated from room temperature (21 oC) and vaporized at the boiling temperature of water? This problem is solved in the Voicethread shown to the right, or click here to view directly.
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