Learning Objectives Define endothermic and exothermic reactions. Determine whether a reaction is endothermic or exothermic through observations, temperature changes, or an energy diagram. Describe the calculation of heat of reaction using bond energies. When physical or chemical changes occur, they are generally accompanied by a transfer of energy. The law of conservation of energy states that in any physical or chemical process, energy is neither created nor destroyed. In other words, the entire energy in the universe is conserved. In order to better understand the energy changes taking place during a reaction, we need to define two parts of the universe, called the system and the surroundings. The system is the specific portion of matter in a given space that is being studied during an experiment or an observation. The surroundings is everything in the universe that is not part of the system. In practical terms for a laboratory chemist, the system is the particular chemicals being reacted, while the surroundings is the immediate vicinity within the room. During most processes, energy is exchanged between the system and the surroundings. In the course of an endothermic process, the system gains heat from the surroundings and so the temperature of the surroundings decreases. A chemical reaction or physical change is exothermic if heat is released by the system into the surroundings. Because the surroundings is gaining heat from the system, the temperature of the surroundings increases (Figure \(\PageIndex{1}\)) . Exothermic and endothermic reactions can be thought of as having energy as either a product of the reaction or a reactant. Exothermic reactions give off energy, so energy is a product. Endothermic reactions require energy, so energy is a reactant. Definition: Enthalpy Energy comes in many forms, such as sound, light, and heat. For chemical reactions and phase changes that occur in everyday life, i.e. under standard conditions, we refer to the energy released and absorbed as the enthalpy . For our purposes enthalpy is interchangeable with the word heat, but it is a very complex quantity; a full discussion is beyond the scope of our class. The symbol for enthalpy H . However, we usually refer to enthalpy changes (going from a solid to a liquid, for example, or from reactants to products). Therefore, you will see \(\Delta H\) as the symbol for enthalpy. Phase changes , discussed in the previous section 6.1 , are also classified in a similar way. The change from gas to liquid (condensation) and liquid to solid (freezing) are exothermic . The change from solid to liquid (melting), and liquid to gas (evaporation and boiling) are endothermic . The exothermic processes release heat to the surroundings while the endothermic processes absorb heat from the surroundings. When methane gas is combusted, heat is released, making the reaction exothermic. Specifically, the combustion of \(1 \: \text{mol}\) of methane releases 890.4 kilojoules of heat energy. This information can be shown as part of the balanced equation. \[ \ce{CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l)} \hspace{0.5in} \Delta H = -890 kJ \nonumber \] Since energy is released, the reaction is exothermic . The process in the above thermochemical equation can be shown visually in the figure below. In the combustion of methane example, the enthalpy change is negative because heat is being released by the system. Therefore, the overall enthalpy of the system decreases. The heat of reaction is the enthalpy change for a chemical reaction. In the case above, the heat of reaction is \(-890.4 \: \text{kJ}\). Endothermic reactions absorb energy from the surroundings as the reaction occurs. When \(1 \: \text{mol}\) of calcium carbonate decomposes into \(1 \: \text{mol}\) of calcium oxide and \(1 \: \text{mol}\) of carbon dioxide, \(177.8 \: \text{kJ}\) of heat is absorbed. The process is shown visually in the figure above (B). The thermochemical reaction is shown below. \[ \ce{CaCO3(s) -> CaO(s) + CO2(g)} \hspace{0.5in} \Delta H = +177.8 kJ \nonumber \] Because the heat is absorbed by the system, the \(177.8 \: \text{kJ}\) is written as a reactant. The heat of reaction is positive for an endothermic reaction. Example \(\PageIndex{1}\) Label each of the following processes as endothermic or exothermic. water boiling gasoline burning ice forming on a pond Solution endothermic - you must put a pan of water on the stove and give it heat in order to get water to boil. Because you are adding heat/energy, the reaction is endothermic. exothermic - when you burn something, it feels hot to you because it is giving off heat into the surroundings. exothermic - think of ice forming in your freezer instead. You put water into the freezer, which takes heat out of the water, to get it to freeze. Because heat is being pulled out of the water, it is exothermic. Heat is leaving. Exercise \(\PageIndex{1}\) Label each of the following processes as endothermic or exothermic. water vapor condensing gold melting Answer (a) exothermic Answer (b) endothermic Example \(\PageIndex{2}\) Is each chemical reaction exothermic or endothermic? \(\ce{2H2(g) + O2(g) -> 2H2O(l)}\hspace{0.5in} \Delta H = -565 kJ \nonumber \) \(\ce{N2(g) + O2(g) -> 2NO(g)}\hspace{0.5in} \Delta H = +188 kJ \nonumber \) Solution Because energy is a product, energy is given off by the reaction. Therefore, this reaction is exothermic. Because energy is a reactant, energy is absorbed by the reaction. Therefore, this reaction is endothermic. Exercise \(\PageIndex{2}\) Is each chemical reaction exothermic or endothermic? \(\ce{H2(g) + F2(g) -> 2HF(g)}\hspace{0.5in} \Delta H = -544 kJ \nonumber \) \(\ce{2C(s) + H2(g) -> C2H2(g)}\hspace{0.5in} \Delta H = +22 kJ \nonumber \) Answer a. The energy (544 kJ) is produced, hence the reaction is exothermic b. The energy (22 kJ) is supplied or absorbed to react, hence, the reaction is endothermic Bond Energy Atoms bond together to form compounds because in doing so they attain lower energies than they possess as individual atoms. A quantity of energy, equal to the difference between the energies of the bonded atoms and the energies of the separated atoms, is released, usually as heat. That is, the bonded atoms have a lower energy than the individual atoms do. When atoms combine to make a compound, energy is always given off, and the compound has a lower overall energy. When a chemical reaction occurs, molecular bonds are broken and other bonds are formed to make different molecules. For example, the bonds of two water molecules are broken to form hydrogen and oxygen. \[ \ce{2H2O \rightarrow 2H2 + O2} \nonumber \] Energy is always required to break a bond, which is known as bond energy. While the concept may seem simple, bond energy serves a very important purpose in describing the structure and characteristics of a molecule. Energy is always required to break a bond; it is an endothermic process. Energy is release when a bond is made; it is an exothermic process. Bond Bond Energy (kJ/mol) C–H 413.0 C–O 358.0 C=O* 745.0 C–N 305.0 C–C 347.0 C=C 614.0 C≡C 839.0 N–H 391.0 O—H 467.0 H–H 432.0 H—Cl 427.0 H—I 295.0 *C=O (CO2) = 799 NaN Note: kJ or kJ/mol? You might notice that the enthalpy units in Table 1 are in kJ/mol, but up to this point we've just been using kJ. Why? It's because bond energies, and enthalpies, are measured like everything else in chemistry—in terms of moles of atoms or molecules. This means that the bond energy for a C-H bond, listed in Table 1 as 413 kJ/mol, has been measured for 1 mol of C-H bonds. When a chemical reaction occurs, the atoms in the reactants rearrange their chemical bonds to make products. The new arrangement of bonds does not have the same total energy as the bonds in the reactants. Therefore, when chemical reactions occur, there will always be an accompanying energy change . The enthalpy change for a given reaction can be calculated using the bond energy values from Table \(\PageIndex{1}\). In this process, one adds energy to the reaction to break bonds, and extracts energy for the bonds that are formed. \[\Delta H = \sum (\text{bonds broken}) - \sum (\text{bonds formed}) \nonumber \] Example \(\PageIndex{3}\): Generation of Hydrogen Iodide What is the enthalpy change for this reaction and is it endothermic or exothermic? \[\ce{H2(g) + I2(g) \rightarrow 2HI(g)} \nonumber \] Solution First look at the equation and identify which bonds exist in the reactants. one H-H bond and one I-I bond Now do the same for the products two H-I bonds Then identify the bond energies of these bonds from the table above: H-H bonds: 436 kJ/mol I-I bonds: 151 kJ/mol The sum of enthalpies on the reaction side is: 436 kJ/mole + 151 kJ/mole = 587 kJ/mol. This is how much energy is needed to break the bonds on the reactant side. Then we look at the bond formation which is on the product side: 2 mol H-I bonds: 297 kJ/mol The sum of enthalpies on the product side is: 2 x 297 kJ/mol= 594 kJ/mol This is how much energy is released when the bonds on the product side are formed. Remember we use the equation: \[\Delta H = \sum (\text{bonds broken}) - \sum (\text{bonds formed}) \nonumber \] The net change of the reaction is therefore 587 kJ/mol - 594 kJ/mol= -7 kJ/mol. Since the enthalpy is negative, the reaction is exothermic . Exercise \(\PageIndex{3}\) Using the bond energies given in the chart above, find the enthalpy change for the thermal decomposition of water: \[ \ce{2H2O(g) \rightarrow 2H2(g) + O2(g)} \nonumber \] Is the reaction written above exothermic or endothermic? Explain. Answer The enthalpy change deals with breaking 4 mole of O-H bonds (there are two water molecules, and one water molecule has two O-H bonds) and the formation of 1 mole of O-O bonds and two moles of H-H bonds (Table \(\PageIndex{1}\)). The sum of the energies required to break the bonds on the reactants side is 4 x 460 kJ/mol = 1840 kJ/mol. The sum of the energies released to form the bonds on the products side is 2 moles of H-H bonds = 2 x 436.4 kJ/mol = 872.8 kJ/mol 1 moles of O=O bond = 1 x 498.7 kJ/mol = 498.7 kJ/mol which is an output (released) energy = 872.8 kJ/mol + 498.7 kJ/mol = 1371.5 kJ/mol. Remember we use the equation: \[\Delta H = \sum (\text{bonds broken}) - \sum (\text{bonds formed}) \nonumber \] to determine the enthalpy of the reaction. Therefore, total energy difference is 1840 kJ/mol – 1371.5 kJ/mol = 469 kJ/mol, which indicates that the reaction is endothermic and that 469 kJ of heat is needed to be supplied to carry out this reaction. Summary Chemical processes are labeled as exothermic or endothermic based on whether they give off or absorb energy, respectively. Atoms are held together by a certain amount of energy called bond energy. Energy is released to generate bonds, which is why the enthalpy change for breaking bonds is positive. Energy is required to break bonds. Atoms are much happier when they are "married" and release energy because it is easier and more stable to be in a relationship (e.g., to generate octet electronic configurations ). The enthalpy change is negative because the system is releasing energy when forming bond.