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Thermochemistry is the study of the energy evolved or absorbed in chemical reactions and any physical transformations, such as melting and boiling. Thermochemistry, generally, is concerned with the energy exchange accompanying transformations, such as mixing, phase transitions, chemical reactions, and including calculations of such quantities as the heat capacity, heat of combustion, heat of formation, enthalpy, and free energy.
The measurement of heat changes is performed using calorimetry, usually an enclosed chamber within which the change to be examined occurs. The temperature of the chamber is monitored either using a thermometer or thermocouple, and the temperature plotted against time to give a graph from which fundamental quantities can be calculated. Modern calorimeters are frequently supplied with automatic devices to provide a quick read-out of information, one example being the DSC or differential scanning calorimeter.
Several thermodynamic definitions are very useful in thermochemistry. A system is the specific portion of the universe that is being studied. Everything outside the system is considered the surrounding or environment. A system may be: Isolated system - when it cannot exchange energy or matter with the surroundings, as with an insulated bomb reactor; Closed system - when it can exchange energy but not matter with the surroundings, as with a steam radiator; Open system - when it can exchange both matter and energy with the surroundings, as with a pot of boiling water.
Many chemical reactions release energy in the form of heat, light, or sound. These are exothermic reactions. Exothermic reactions may occur spontaneously and result in higher randomness or entropy (ΔS > 0) of the system. They are denoted by a negative heat flow (heat is lost to the surroundings) and decrease in enthalpy (ΔH < 0). In the lab, exothermic reactions produce heat or may even be explosive.
There are other chemical reactions that must absorb energy in order to proceed. These are endothermic reactions. Endothermic reactions cannot occur spontaneously. Work must be done in order to get these reactions to occur. When endothermic reactions absorb energy, a temperature drop is measured during the reaction. Endothermic reactions are characterized by positive heat flow (into the reaction) and an increase in enthalpy (+ΔH).




Examples of Endothermic and Exothermic Processes
Photosynthesis is an example of an endothermic chemical reaction. In this process, plants use the energy from the sun to convert carbon dioxide and water into glucose and oxygen. This reaction requires 15MJ of energy (sunlight) for every kilogram of glucose that is produced:
sunlight + 6CO2(g) + H2O(l) = C6H12O6(aq) + 6O2(g)
An example of an exothermic reaction is the mixture of sodium and chlorine to yield table salt. This reaction produces 411 kJ of energy for each mole of salt that is produced:
Na(s) + 0.5Cl2(s) = NaCl(s)
Thermochemical equations are just like other balanced equations except they also specify the heat flow for the reaction. The heat flow is listed to the right of the equation using the symbol ΔH. The most common units are kilojoules, kJ. Here are two thermochemical equations:
H2 (g) + ½ O2 (g) → H2O (l); ΔH = -285.8 kJ
HgO (s) → Hg (l) + ½ O2 (g); ΔH = +90.7 kJ
When you write thermochemical equations, be sure to keep the following points in mind:
1. Coefficients refer to the number of moles. Thus, for the first equation, -282.8 kJ is the ΔH when 1 mol of H2O (l) is formed from 1 mol H2 (g) and ½ mol O2.
2. Enthalpy changes for a phase change, so the enthalpy of a substance depends on whether is it is a solid, liquid, or gas. Be sure to specify the phase of the reactants and products using (s), (l), or (g) and be sure to look up the correct ΔH from heat of formation tables. The symbol (aq) is used for species in water (aqueous) solution.
3. The enthalpy of a substance depends upon temperature. Ideally, you should specify the temperature at which a reaction is carried out. When you look at a table of heats of formation, notice that the temperature of the ΔH is given. For homework problems, and unless otherwise specified, temperature is assumed to be 25°C. In the real world, temperature may different and thermochemical calculations can be more difficult.


Certain laws or rules apply when using thermochemical equations:
1. ΔH is directly proportional to the quantity of a substance that reacts or is produced by a reaction.
Enthalpy is directly proportional to mass. Therefore, if you double the coefficients in an equation, then the value of ΔH is multiplied by two. For example:
H2 (g) + ½ O2 (g) → H2O (l); ΔH = -285.8 kJ
2 H2 (g) + O2 (g) → 2 H2O (l); ΔH = -571.6 kJ
2. ΔH for a reaction is equal in magnitude but opposite in sign to ΔH for the reverse reaction.
For example:
HgO (s) → Hg (l) + ½ O2 (g); ΔH = +90.7 kJ
Hg (l) + ½ O2 (l) → HgO (s); ΔH = -90.7 kJ
This law is commonly applied to phase changes, although it is true when you reverse any thermochemical reaction.
3. ΔH is independent of the number of steps involved.
This rule is called Hess's Law. It states that ΔH for a reaction is the same whether it occurs in one step or in a series of steps. Another way to look at it is to remember that ΔH is a state property, so it must be independent of the path of a reaction.
If Reaction (1) + Reaction (2) = Reaction (3), then ΔH3 = ΔH1 + ΔH2
The molar heat of formation of a compound (ΔHf) is equal to its enthalpy change (ΔH) when one mole of compound is formed at 25°C and 1 atm from elements in their stable form. This is a table of the heats of formation for a variety of common compounds. As you can see, most heats of formation are negative quantities, which implies that the formation of a compound from its elements usually is an exothermic process.

Compound ΔHf (kJ/mol) Compound ΔHf (kJ/mol)
AgBr(s) -99.5 C2H2(g) +226.7
AgCl(s) -127.0 C2H4(g) +52.3
AgI(s) -62.4 C2H6(g) -84.7
Ag2O(s) -30.6 C3H8(g) -103.8
Ag2S(s) -31.8 n-C4H10(g) -124.7
Al2O3(s) -1669.8 n-C5H12(l) -173.1
BaCl2(s) -860.1 C2H5OH(l) -277.6
BaCO3(s) -1218.8 CoO(s) -239.3
BaO(s) -558.1 Cr2O3(s) -1128.4
BaSO4(s) -1465.2 CuO(s) -155.2
CaCl2(s) -795.0 Cu2O(s) -166.7
CaCO3 -1207.0 CuS(s) -48.5
CaO(s) -635.5 CuSO4(s) -769.9
Ca(OH)2(s) -986.6 Fe2O3(s) -822.2
CaSO4(s) -1432.7 Fe3O4(s) -1120.9
CCl4(l) -139.5 HBr(g) -36.2
CH4(g) -74.8 HCl(g) -92.3
CHCl3(l) -131.8 HF(g) -268.6
CH3OH(l) -238.6 HI(g) +25.9
CO(g) -110.5 HNO3(l) -173.2
CO2(g) -393.5 H2O(g) -241.8
H2O(l) -285.8 NH4Cl(s) -315.4
H2O2(l) -187.6 NH4NO3(s) -365.1
H2S(g) -20.1 NO(g) +90.4
H2SO4(l) -811.3 NO2(g) +33.9
HgO(s) -90.7 NiO(s) -244.3
HgS(s) -58.2 PbBr2(s) -277.0
KBr(s) -392.2 PbCl2(s) -359.2
KCl(s) -435.9 PbO(s) -217.9
KClO3(s) -391.4 PbO2(s) -276.6
KF(s) -562.6 Pb3O4(s) -734.7
MgCl2(s) -641.8 PCl3(g) -306.4
MgCO3(s) -1113 PCl5(g) -398.9
MgO(s) -601.8 SiO2(s) -859.4
Mg(OH)2(s) -924.7 SnCl2(s) -349.8
MgSO4(s) -1278.2 SnCl4(l) -545.2
MnO(s) -384.9 SnO(s) -286.2
MnO2(s) -519.7 SnO2(s) -580.7
NaCl(s) -411.0 SO2(g) -296.1
NaF(s) -569.0 So3(g) -395.2
NaOH(s) -426.7 ZnO(s) -348.0
NH3(g) -46.2 ZnS(s) -202.9
Heat of combustion
The heat of combustion (ΔHc0) is the energy released as heat when one mole of a compound undergoes complete combustion with oxygen. The chemical reaction is typically a hydrocarbon reacting with oxygen to form carbon dioxide, water and heat. It may be expressed with the quantities:
• energy/mole of fuel (J/mol)
• energy/mass of fuel
• energy/volume of fuel
The heat of combustion is traditionally measured with a bomb calorimeter. It may also be calculated as the difference between the heat of formation (ΔfH0) of the products and reactants.
In thermodynamics and molecular chemistry, enthalpy (denoted as H, or specific enthalpy denoted as h) is a thermodynamic property of a thermodynamic system. It can be used to calculate the heat transfer during a quasistatic process taking place in a closed thermodynamic system under constant pressure (isobaric process). Change in enthalpy ΔH is frequently a more useful value than H itself. For quasistatic processes under constant pressure, ΔH is equal to the change in the internal energy of the system, plus the work that the system has done on its surroundings.[1] This means that the change in enthalpy under such conditions is the heat absorbed by a chemical reaction.

Standard enthalpy changes
Standard enthalpy changes describe the change in enthalpy observed in the constituents of a thermodynamic system when going between different states under standard conditions. The standard enthalpy change of vaporization, for example gives the enthalpy change when going from liquid to gas. These enthalpies are reversible; the enthalpy change of going from gas to liquid is the negative of the enthalpy change of vaporization. A common standard enthalpy change is the standard enthalpy change of formation, which has been determined for a large number of substances. The enthalpy change of any reaction under any conditions can be computed, given the standard enthalpy change of formation of all of the reactants and products.
Chemical Properties
• Standard enthalpy change of reaction
Standard enthalpy change of formation, defined as the enthalpy change observed in a constituent of a thermodynamic system when, one mole of a compound is formed from its elementary antecedents under standard conditions. The standard enthalpy of formation "standard heat of formation" of a compound is the change of enthalpy that accompanies the formation of 1 mole of a substance in its standard state from its constituent elements in their standard states (the most stable form of the element at 1 bar of pressure and the specified temperature, usually 298.15 K or 25 degrees Celsius). Its symbol is ΔHfO or ΔfHO.
A similar type of enthalpy change, known as the standard enthalpy change of hydrogenation is defined as the enthalpy change observed when 1 mol of an unsaturated compound reacts with an excess of hydrogen to become fully saturated, all elements within the reaction being within their standard states.
For example, the standard enthalpy of formation of carbon dioxide would be the enthalpy of the following reaction under the conditions above:
C(s,graphite) + O2(g) → CO2(g)
The standard enthalpy change of formation is measured in units of energy per amount of substance. Most are defined in kilojoules per mole, or kJ mol−1, but can also be measured in calories per mole, joules per mole or kilocalories per gram (any combination of these units conforming to the energy per mass or amount guideline). In physics the energy per particle is often expressed in electronvolts which corresponds to about 100 kJ mol−1.
All elements in their standard states (oxygen gas, solid carbon in the form of graphite, etc.) have a standard enthalpy of formation of zero, as there is no change involved in their formation.
The standard enthalpy change of formation is used in thermochemistry to find the standard enthalpy change of reaction. This is done by subtracting the sum of the standard enthalpies of formation of the reactants from the sum of the standard enthalpies of formation of the products, as shown in the equation below.
ΔHreactionO = ΣΔHfO (Products) - ΣΔHfO (Reactants)
The standard enthalpy of formation is equivalent to the sum of many separate processes included in the Born-Haber cycle of synthesis reactions. For example, to calculate the standard enthalpy of formation of sodium chloride, we use the following reaction:
Na(s) + (1/2)Cl2(g) → NaCl(s)

Standard enthalpy change of formation Born-Haber diagram for lithium fluoride.
1. The standard enthalpy of atomization of solid sodium
2. The first ionization energy of gaseous sodium
3. The standard enthalpy of atomization of chlorine gas
4. The electron affinity of chlorine atoms
5. The lattice enthalpy of sodium chloride
The sum of all these values will give the standard enthalpy of formation of sodium chloride.
Additionally, applying Hess's Law shows that the sum of the individual reactions corresponding to the enthalpy change of formation for each substance in the reaction is equal to the enthalpy change of the overall reaction, regardless of the number of steps or intermediate reactions involved. In the example above the standard enthalpy change of formation for sodium chloride is equal to the sum of the standard enthalpy change of formation for each of the steps involved in the process. This is especially useful for very long reactions with many intermediate steps and compounds.
Chemists may use standard enthalpies of formation for a reaction that is hypothetical. For instance carbon and hydrogen will not directly react to form methane, yet the standard enthalpy of formation for methane is determined to be -74.8 kJ mol−1 from using other known standard enthalpies of reaction with Hess's law. That it is negative shows that the reaction, if it were to proceed, would be exothermic; that is, it is enthalpically more stable than hydrogen gas and carbon.
• Standard enthalpy change of combustion. Defined as the enthalpy change observed in a constituent of a thermodynamic system, when one mole of a substance combusts completely with oxygen under standard conditions. The standard enthalpy of combustion is the enthalpy change when one mole of a substance completely reacts under standard thermodynamic conditions (although experimental values are usually obtained under different conditions and subsequently adjusted). By definition, combustion reactions are generally strongly exothermic and so enthalpies of combustion are generally strongly negative.
It is commonly denoted as ΔHc0. Enthalpies of combustion are typically measured using bomb calorimetry, and have units of energy (typically kJ); strictly speaking, the enthalpy change per mole of substance combusted is the standard molar enthalpy of combustion (which typically would have units of Kj.
• Standard enthalpy change of hydrogenation. Defined as the enthalpy change observed in a constituent of a thermodynamic system, when one mole of an unsaturated compound reacts completely with an excess of hydrogen under standard conditions to form a saturated compound. The standard enthalpy of formation "standard heat of formation" of a compound is the change of enthalpy that accompanies the formation of 1 mole of a substance in its standard state from its constituent elements in their standard states (the most stable form of the element at 1 bar of pressure and the specified temperature, usually 298.15 K or 25 degrees Celsius). Its symbol is ΔHfO or ΔfHO.
A similar type of enthalpy change, known as the standard enthalpy change of hydrogenation is defined as the enthalpy change observed when 1 mol of an unsaturated compound reacts with an excess of hydrogen to become fully saturated, all elements within the reaction being within their standard states.
• Standard enthalpy change of atomization, defined as the enthalpy change required to atomize one mole of compound completely under standard conditions.
Physical Properties
• Standard enthalpy change of solution, defined as the enthalpy change observed in a constituent of a thermodynamic system, when one mole of an solute is dissolved completely in an excess of solvent under standard conditions.
• Standard enthalpy change of fusion, defined as the enthalpy change required to completely change the state of one mole of substance between solid and liquid states under standard conditions.
• Standard enthalpy change of vapourization, defined as the enthalpy change required to completely change the state of one mole of substance between liquid and gaseous states under standard conditions.
• Standard enthalpy change of sublimation, defined as the enthalpy change required to completely change the state of one mole of substance between solid and gaseous states under standard conditions.
• Standard enthalpy change of denaturation, defined as the enthalpy change required to denature one mole of compound under standard conditions.
• Lattice enthalpy, defined as the enthalpy required to separate one mole of an ionic compound into separated gaseous ions to an infinite distance apart (meaning no force of attraction) under standard conditions
Hess's law
Hess's law is a relationship from physical chemistry named for Germain Hess, a Swiss-born Russian chemist and doctor. The law is based on the principle of conservation of energy and the path independence of energy changes. Hess's law can be used to predict energy changes that are not easily measured.
The law states that the energy change for any chemical or physical process is independent of the pathway or number of steps required to complete the process. In other words, an energy change is path independent, only the initial and final states being of importance. This path independence is true for all state functions.
Hess's law allows the enthalpy change (ΔH) for a reaction to be calculated even when it cannot be measured directly. This is accomplished by performing arithmetic operations on chemical equations and known ΔH values. Chemical equations may be multiplied (or divided) by a whole number. When an equation is multiplied by a constant, its ΔH must be multiplied by the same number as well. If an equation is reversed, ΔH for the reaction must also be reversed (i.e. -ΔH).
Addition of chemical equations can lead to a net equation. If enthalpy change is included for each equation and added, the result will be the enthalpy change for the net equation. If the net enthalpy change is negative (ΔHnet < 0), the reaction will be exothermic and is more likely to be spontaneous; positive ΔH values correspond to endothermic reactions. Note that entropy also plays an important role in determining spontaneity, so some reactions with a positive enthalpy change are nevertheless spontaneous.
Hess's Law says that enthalpy changes are additive. Thus the ΔH for a single reaction can be calculated from the difference between the heat of formation of the products minus the heat of formation of the reactants.
BOND ENERGY
Bond energy (enthalpy) is the energy (enthalpy) required per mole of gaseous compound to break a particular bond to produce gaseous fragments.
 Bond energies (enthalpies) are positive.
ΔH is positive.
 Breaking bonds is an endothermic reaction.
 Bond formation releases energy.
ΔH is negative.Bond formation is an exothermic reaction.
 Bond energies (enthalpies) can be used to indicate how stable a compound is or how easy it is break a particular bond. The more energy that is required to break a bond, the more stable the compound will be. A larger bond energy implies the bond is harder to break, so the compound will be more stable.
 Bond energies (enthalpies) can be used to estimate the heat (enthalpy) of a reaction.
ΔHo(reaction) = sum of the bond energies of bonds being broken - sum of the bond energies of the bonds being formed.
ΔHo(reaction) = ΔH(reactant bonds broken) - ΔH(product bonds formed)
Calculating heat (enthalpy) of reaction, ΔHo, from bond energies of reactants and products:
 Write the balanced chemical equation, with all reactants and products in the gaseous state.
If a reactant or product is NOT in the gaseous state, you will need to use Hess's Law to include the relevant energy (enthalpy) for the change of state.
 Write the general equation for the heat (enthalpy) of reaction:
ΔHo(reaction) = ΔH(reactant bonds broken) - ΔH(product bonds formed)
 Substitute bond energy values into the equation and solve for ΔHo(reaction)
Bond Energy and Chemical Stability
Breaking chemical bonds requires energy. The more energy that is required to break a bond, the more chemically stable the compound will be.
When comparing 2 compounds, the compound containing the bond with the lowest bond energy will be the least stable compound, regardless of the strength of the other bonds present in the compound.
Single Bond Energies (kJ mol-1)
S Si I Br Cl F O N C H

H 339 339 299 366 432 563 463 391 413 436

C 259 290 240 276 328 441 351 292 348

N 200 270 161

O 369 203 185 139

F 541 258 237 254 153

Cl 250 359 210 219 243

Br 289 178 193

I 213 151

Si 227 177

S 213

Example of Bond Energy and Chemical Stability
The energy required to break the O - H bond in water, H2O(g), is 463 kJ mol-1
The energy required to break the O - O bond in hydrogen peroxide, H2O2(g), is 139 kJ mol-1.
Hydrogen peroxide is less stable than water because it contains the O - O bond which is easier to break.
Hydrogen peroxide will be more chemically reactive than water.
Calculating Heat (Enthalpy) of Reaction Using Bond Energy
Bond energies (enthalpies) can be used to estimate the heat (enthalpy) of a reaction.
ΔHo(reaction) = sum of the bond energies of bonds being broken - sum of the bond energies of the bonds being formed.
Or,
ΔHo(reaction) = ΔH(reactant bonds broken) - ΔH(product bonds formed)



Example of Using Bond Energies to Calculate Heat (enthalpy) of Reaction
Use the bond energies provided in the table above to calculate the heat (enthalpy) of reaction, Ho, for the reaction:
CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g)
 Write the balanced chemical equation, with all reactants and products in the gaseous state.
CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g)
 Write the general equation for the heat (enthalpy) of reaction:
ΔHo(reaction) = ΔH(reactant bonds broken) - ΔH(product bonds formed)
 Substitute bond energy values into the equation and solve for ΔHo(reaction)


Bonds Broken Bonds Formed

bond type bond energy bond type bond energy
4 x C - H 4 x 413 = 1652 4 x C - Cl 4 x 328 = 1312
4 x Cl - Cl 4 x 243 = 972 4 x H - Cl 4 x 432 = 1728
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ΔH(reactant bonds broken) = 2624 ΔH(product bonds formed) = 3040
ΔHo(reaction) = 2624 - 3040 = -416 kJ mol-1