Enthalpy. Enthalpy. Enthalpy. Enthalpy. E = q + w. Internal Energy at Constant Volume SYSTEM. heat transfer in (endothermic), +q

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1 heat transfer in (endothermic), +q heat transfer out (exothermic), -q SYSTEM E = q + w w transfer in (+w) w transfer out (-w) Internal Energy at Constant Volume E = KE + PE ΔE = q + w Because most systems, especially those not involving gases, undergo little to no change in volume at the constant pressure of the atmosphere, the work that is done on or by the chemical system is small. When the change in volume of a system is zero, work energy lost or gained is also zero: w = -P external ΔV So at constant volume, the change in internal energy is equal to heat gained or lost: Since the volume change accompanying many reactions is close to zero, another thermodynamic quantity is often used, called enthalpy. is the thermodynamic quantity equivalent to the total heat content of a system. is defined as the internal energy plus the product of pressure and volume: H = E + PV ΔE = q v H = E + PV When the system changes at constant pressure, the change in enthalpy, H, is This can be written H = (E + PV) H = E + P V In this equation, the change in total heat content of the system is equal to the sum of E, thought as the total the energy required to make the system, and P V, thought as the total energy needed to make room for the system. When at constant pressure, the H value assumes no work energy is lost or gained other than expansion work that makes room for the system or from the atmosphere collapsing. Since E = q + w and w = P V, we can substitute into the change in enthalpy expression: H = E + P V H = (q + w ) + P V H = (q + ( P V)) + P V H = q p

2 vs. Internal Energy At constant pressure the observed heat gained or lost by the system is the change in enthalpy H = q p At constant volume the observed heat gained or lost by the system is the change in internal energy : ΔE = q v Given the volume change of many reactions is close to zero, the difference between H and ΔE is small; making H a satisfactory measure of energy changes during most chemical reactions. Types of Changes changes ( H) can be observed in one of three modes: 1) Heat is transferred from system into surroundings, or visa versa, with no physical or chemical change. 2) Heat is transferred as the system or the surroundings undergoes a physical change (phase change, dissociation, etc ) 3) Heat is transferred between the system (reactants and products) and the surroundings as a chemical change occurs. Important Slide! Change of a Chemical Reaction All chemical reactions either release or absorb heat, even if we can not observe it directly. Reactants Products + Energy (exo) Reactants + Energy Products (endo) The change in enthalpy can be described by: ΔH = H f H i The total enthalpy of the products minus the total enthalpy of the reactants describes the total change in internal heat energy of a chemical reaction. The enthalpy change for any given reaction is called the reaction enthalpy (ΔH rxn ) Example: The combustion of hydrogen gas with oxygen to produce water produces kj of heat energy Since energy was produced, the reaction is said to be exothermic. 2H 2(g) + O 2(g) 2H 2 O (g) kJ Endothermicity and Exothermicity A process is endothermic when H is positive. A process is exothermic when H is negative.

3 Thermochemical Equations Two ways to express energy transfer in thermochemical equations: 2SO 2(g) + O 2(g) 2SO 3(g) kj 2SO 2(g) + O 2(g) 2SO 3(g) H = kj Notice, when the enthalpy is represented after the equation the sign changed; the reaction is exothermic Thermochemical Equations Write the thermochemical equations for the combustion of methane The heat released or absorbed by a reaction depends on the physical states of the reactants (s, l or g) and on the physical environment (T and P). Observe the difference in the following: CH 4(g) + 2O 2 (g) CO 2(g) + 2 H 2 O (g) ΔH = -802 kj CH 4(g) + 2O 2 (g) CO 2(g) + 2 H 2 O (l) ΔH = -890 kj The Truth about 1. is an extensive property. 2. H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction. 3. H for a reaction depends on the state of the products and the state of the reactants. Because thermochemical reactions are affected by physical states and environmental conditions, standard conditions must be defined to compare enthalpies of different reactions. Standard Molar of a Reaction ΔH o rxn = the enthalpy of a reaction in which all reactants and products are in their standard states and 1 atm pressure. The standard state refers to the physical state of a pure substance at 1 atm. So, which reaction indicates the enthalpy of the reaction under standard state conditions? CH 4(g) + 2O 2 (g) CO 2(g) + 2 H 2 O (g) ΔH = -802 kj CH 4(g) + 2O 2 (g) CO 2(g) + 2 H 2 O (l) ΔH = -890 kj

4 It is understood that the coefficients of a balanced chemical reaction describes the molar ratio of reactants and products. The coefficients for a chemical reaction also describe the ratio of heat produced/consumed by any chemical reaction. Notice in the standard state combustion of methane: CH 4(g) + 2O 2 (g) CO 2(g) + 2 H 2 O (l) ΔH = -890 kj For every one mole of methane combusted, 890 kj of energy are produced. What about: 1 mol water? 4 mol oxygen? 1. In the thermite reaction, you observed the production of liquid iron from iron (III) oxide and aluminum metal. If the reaction produces kj per mole iron (III) oxide, how much energy is evolved if 0 grams of aluminum are reacted completely? For any reaction, an energy diagram (enthalpy diagram) can be drawn: For the reaction: H 2(g) + O 2(g) 2H 2 O (g) kJ The enthalpy diagram for the reaction is: H H 2(g) + O 2(g) 2H 2 O (g) ΔH < 0 (exo) = kj H 2(g) + O 2(g) ΔH = kj 2H 2 O (g) 2. Write the correct thermochemical equation for a reaction that requires 21KJ of energy to combine xenon gas with fluorine gas producing xenon tetraflouride. Then generate an enthalpy diagram for the reaction. Change of a Physical Change changes accompany physical state changes as well as chemical reactions. T ( o C) time 4 H total q i i 1

5 Energy processes: 1. Within a phase ΔH = q p 2. Between phases ΔH rxn = ΔH vap ΔH rxn = ΔH fus Heat of Vaporization = The heat energy required to convert 1 mol of a substance from its liquid state to its gaseous state. ΔH vap(h2o) = kj/mol Heat of Fusion = The heat energy required to convert one mol of a substance from its solid state to its liquid state. ΔH fus(h2o) = 6.01 kj/mol Remember however, the heat of vaporization and heat of fusion only reflect the energy change that accompanies the conversion of a substance between two phases. They do not represent the energy required to raise the temperature to the temperature of the phase change. T ( o C) H total q i i 1 time 1 q = sm ΔT 2 ΔH fus 3 q = sm ΔT 4 ΔH vap q = sm ΔT 1. How much total heat is needed to raise the temperature of g H 2 O from -1 o C to 12 o C? Given: ΔH vap of H 2 O = kj/mol ΔH fus of H 2 O = 6.01 kj/mol s ice = 2.06 J/gK s water vap = 2.03 J/gK