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Chapter 6

Chapter 6. Thermochemistry. Overview. Understanding Heats of Reaction Energy and its Units Heat of Reaction Enthalpy and Enthalpy Change Thermochemical Equations Stoichiometry and Heats of Reaction Measuring Heats of Reaction Uses of Heats of Reaction Hess’s Law

pascale-joseph
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Chapter 6

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  1. Chapter 6 Thermochemistry

  2. Overview • Understanding Heats of Reaction • Energy and its Units • Heat of Reaction • Enthalpy and Enthalpy Change • Thermochemical Equations • Stoichiometry and Heats of Reaction • Measuring Heats of Reaction • Uses of Heats of Reaction • Hess’s Law • Standard Enthalpies of Formation • Fuels – Foods, etc.

  3. Energy and Its Units • Thermochemistry the study of the energy changes that take place during a reaction. • Reactions generally proceed in whichever direction that will produce products with lower energy than the reactants. Heat and Energy • Heat: energy transferred from hotter to colder one. • Kinetic energy: the energy of movement of matter . Units: Joule = 1 kg*m2/s2. E.g. what is the kinetic energy of 50.0 kg person running at a velocity of 20 m/s. • Potential energy: stored energy. E.g. water at the top of a mountain, a compressed spring, a chemical bond.

  4. Energy Changes and Energy Conservation First law of Thermodynamics:Energy is neither created nor destroyed but may be converted from one form to another. Energy forms: • Thermal energy a form of kinetic energy; energy transfer results in a temperature change. • Chemical energy a form of potential energy. Energy is stored in chemical bonds and released when a compound reacts. • During reaction, energy is usually transformed from chemical to thermal energy. • First law can be written as: E = q + w where q = heat involved in the process and w = work done by or to the system. • Work can be electrical or pressure –volume

  5. Internal Energy and The First Law of Thermodynamics • Internal Energy, E, is the sum of the potential and kinetic energy of a system. • System - that part of the universe upon which we are focusing, e.g. reactions. • Surroundings - eveything else in the universe which is not the system. • State function - property depending only upon initial and final states and not upon path. • Extent of transfer of energy is: E = Efinal Einitial. • System is the reference point and a negative sign indicates that energy is flowing from the system to the surrounding. • exothermic (exo “out of”). Heat flows from system to surroundings. • endothermic (endo “into”). Heat flows from surroundings to system. • C(gr) + O2(g)  CO2(g) + 393.5 kJ Exothermic • CO2(g) + 393.5 kJ  C(gr) + O2(g) Endothermic

  6. Sign conventions • Sign of DE will depend upon the sign of q and w. DE = q + w

  7. Internal Energy and The First Law of Thermodynamics2 • The conditions of measurement must be included when discussing the total internal energy since it is related to • chemical identity of reactants and products • their temperature, pressure, and physical state. • Internal energy of a system is a state function. • State Function a property of the system which depends only in the initial and final states and is independent of the history of the system. • Several energy functions to be discussed have this property.

  8. Expansion Work Work = force acting over some distance: w =  d x F (referenced to the system). • During reactions often there is an expansion of gases against some pressure where pressure is equal to the force per unit area: or . • Work is obtained by substitution: • w =  d x F =  d x (PxA) or • w =  PV. • The first law can be restated as E = q  PV. • This equation indicates that the amount of heat involved in a reaction will be reduced by the amount of work being done for a given change in the internal energy. • E.g. Calculate the work done when during a reaction the gaseous products cause the volume to change from 22.4 L to 44.8 L against a constant pressure of 1.00 atm.

  9. Expansion Work2 • If work is performed at constant temperature, then the amount of work performed will depend upon the change in the number of moles (n) • Modifications of the ideal gas law (PV = nRT where n = # mol and R = 8.3145 J/molK) lead to an alternative way of determining work. PV = nRT • The presence of solids and liquids need not be considered since the molar volume of either a solid or liquid is about 1000x smaller than the molar volume of a gas. E.g. determine the work performed during the combustion of methane at 1.00 atm and 298.15 K.: CH4(g) + 2O2(g)  CO2(g) +2H2O(l)

  10. Energy and Enthalpy • From the first law: q = E + PV. • With no change in volume the equation simplifies to qV = E. • At constant pressure: qP = E + PV. • There are times when both volume and pressure can change; the heat involved in the reaction is then a more complicated function of E. • Enthalpy: the heat output at constant pressure. H = E + PV. • In general, H = E + PV + VP. • At constant pressure, a change in enthalpy is given by: H = E + PV = qP. • Normally, H and E are fairly close to each other in magnitude. In the combustion of propane (see book), E = 2043 kJ, H = 2041 kJ and w = PV = 2kJ.

  11. Enthalpies of Physical and Chemical Change • Enthalpies of Physical Change: • Heating a substance increases the temperature; the amount of heat absorbed is proportional to the heat capacity of the species being heated. • Amount of energy absorbed during phase change is proportional to the heat of phase change. • Sum the heats in each portion of the curve to determine overall heat. • The heat for converting a solid directly to a gas is called the heat of sublimation and is equal to the sum of the heats of fusion and vaporization at the same temperature.

  12. Enthalpies of Chemical Change • H is an extensive property – its value depends upon the amount of reactants. • H is attached to the chemical equation to indicate the amount of heat involved in the reaction. E.g. the combustion of methane: • CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H =  890kJ • 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(l) H =  1780kJ E.g.2 determine the amount of heat that would be evolved when 150 g of methane is burnt. • Reversing reaction changes the sign of the heat. • CO2(g) + 2H2O(l)  CH4(g) + 2O2(g) H = +890kJ.

  13. Calorimetry and Heat Capacity • Calorimeter = a device that measures the change in the heat content or internal energy. • Atmospheric pressure • Bomb calorimeter • Heat capacity = the amount of heat absorbed by a substance to raise the temperature by a given amount. • Heat transferal to a substance like a solid or a liquid, causes a change in temperature that is proportional to the amount of heat involved. • H2O absorbs 4.18 J for every gram and °C • Al absorbs 0.902 J for every gram and °C • The amount of heat absorbed is directly proportional to amount of absorbing species: where s = specific heat capacity, C = molar heat capacity and T = Tfinal Tinitial.

  14. Calorimetry and Heat Capacity2 • Energy change from any source such as reactions or phase change can be measured with heat capacity. E.g. How much heat is required to heat 500.0 g of water from 20.0°C to 100.0°C. • The enthalpy change in the system is the negative of the heat of the calorimeter. • E.g. exothermic reactions gives off heat to calorimeter. H =  qcalorimeter. E.g.2 When 2.00 g of ethanol was burned, all of the reaction energy was used to heat water in a calorimeter. Determine H for the reaction if the temperature of 200.0 g of water increased from 25.0°C to 89.0°C. • Heat capacity of a whole calorimeter is used for complicated calorimeters such as the bomb calorimeter. E.g. 800.0 J of heat caused the temperature of a calorimeter was found to increase by 2.0 K. In some other reaction, the temperature of the calorimeter was found to increase by 5.0 K. Calculate the heat of the reaction.

  15. Hess’s Law • Hess’s law when a reaction at constant temperature and pressure can be written as the summation of a series of reactions, the enthalpy change, H, of the reaction is equal to the summation of the H’s of the individual reactions. E.g. determine the heat of formation of NO2(g): ½ N2(g) + O2(g)  NO2 = ? • Forming NO2(g) from N2(g) can be thought of as 2 step process:

  16. Hess’s Law2 • Missing steps in a sequence can be determined using Hess’s law. E.g. determine the heat for methanol decomposition to its elements from the heat of combustion and the other given reactions. Heat of combustion is CH3OH(g) + O2(g)  CO2(g) + H2O(l) H = 726.4 kJ. DH° = DH1° + DH2° + DH3° or DH1° = -726.4 + 393.51 +571.66 = +238.77 kJ

  17. Standard Heats of Formation • Standard state the pure form of a substance at 1 atm usually at 25°C. • Standard reaction enthalpies, H°, difference in enthalpy between products and reactants of a reaction each in their standard states. • Standard heat (enthalpy) of formation the standard reaction enthalpy per mol for the synthesis of a compound from its elements. • Since reaction enthalpy depends upon conditions of experiment, it is usually reported at some reference condition, H° • Most tables present enthalpy data in its standard state and as the heat of formation. • E.g. (HCl) is 92.3 kJ and the reaction is: • ½H2(g) + ½Cl2(g)  HCl(g) = 92.3 kJ • H of pure elements in their most stable form under standard conditions is defined as zero. E.g. Na(g), Na(s); C(g), C(gr), C(d). • of elements in another form often given. • Na(s)  Na(g) H° = 107.8 kJ/mol. Also called the enthalpy of sublimation.

  18. Calculations with Heat of Formation • H° of a reaction can be obtained from of all reactants and products. E.g. Determine the heat of combustion of ethanol, CH3CH2OH, from heats of formation in the book. Solution: CH3CH2OH + 3O2 2CO2(g) + 3H2O(l) DHc°= ? • For any general reaction such as: aA + bB  cC + dD,

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