Unformatted text preview: Thermochemistry Notes Thermochemistry: The study of energy changes related to chemical reactions. Work and Energy Energy is defined as the capacity to do work Potential energy (such as mgh) Kinetic energy (1/2 mv2) SI unit of energy is the joule (J) Work is defined as applying a force F through a distance d W = Fd A force of 1N applied through a distance of 1m does 1J of work (Units: kg m2/s2 or Nm) Note: Work and energy have the same units = joules Thermochemistry Definitions System: The portion of the universe under study. Surroundings: Everything else besides the system. Interaction: Exchange of energy and or matter between the system and its surroundings. Systems: Open system: Exchanges both matter and energy with its surroundings. Closed system: Exchanges only energy with its surroundings. Isolated system: Exchanges neither energy nor matter with its surroundings. Internal Energy (E) = Total energy of the system Kinetic (thermal) energy: Translational Rotational Vibrational Potential energy ‐Nuclear forces ‐Electrostatic attractions (Chemical energy) ‐Intramolecular ‐ Bonds between atoms ‐Intermolecular ‐ Attractions betweens molecules Heat (q) = energy transfer (kinetic energy) Caused by temperature differences System and surroundings reach thermal equilibrium A system does not contain heat. Energy (in the form of thermal kinetic energy) is transferred from the hotter region to the colder region until thermal equilibrium is established. Work (w) Transfer of energy between a system and its surroundings. Usually due to some physical change to the system. A system does not contain work. Focus is on pressure‐volume (P‐V) work: Work done when gases are expanded or compressed. P = F/A F = PA Work (w) = force (F) x distance (h) = P x A x h (A)(h) = V Work (w) = PV (V = Vf – Vi) A negative is added as a sign convention to show that when a system loses energy (i.e. is transferred from the system to the surroundings) the work is negative. Therefore w = ‐PV 1st Law of Thermodynamics Internal energy (E) is a state function. E = Efinal – Einitial A state function only depends on the current conditions of the system and not upon how that state was reached. Cannot measure absolute internal energy Internal energy is a state function Depends on the present state of the system and not on its history Heat (q) and work (w) are not state function (They are not contained in the system and exist as a consequence of a change in the system) For a given change, different combinations of heat and work can lead to the same E. Functions of state are reversible. Law of Conservation of Energy: In a physical or chemical change, energy can be exchanged between a system and its surroundings, but no energy can be created or destroyed. First Law of Thermodynamics: E = q + w Heat absorbed q>0 Increases by the system energy of Work done on the system w>0 the system Heat given off q<0 Decreases by the system energy of Work done by the system w<0 the system So “positive” is measured relative to the system. Heats of Reaction and Enthalpy Change, H Exothermic Reaction
Chemical energy Thermal energy Endothermic Reaction
Thermal energy Chemical energy Isolated System System Temperatur
e increases System Temperatur
e decreases Open/Closed System Heat given off to surroundings (q<0 for system) Heat absorbed from surroundings (q>0 for system) Heat of Reaction (qrxn): The quantity of heat exchanged between the reaction system and its surroundings for a reaction occurring at a fixed temperature. (Assume only pressure‐volume type of work for the following examples) For reactions that take place at constant volume: w = ‐PV = 0 (no volume change) E = q + w = qv qv = E For reactions that take place at constant pressure: E = q + w = qP ‐ PV qP = E + PV Enthalpy Most reactions are carried out at constant pressure in open vessels. Enthalpy (H) E + PV qP = enthalpy change for a reaction = H qP = H = E + PV Properties of Enthalpy 1. Enthalpy is an extensive property (depends upon the amount present) 2. Enthalpy is a state function (E, P and V are state functions) 3. Enthalpy changes have unique values (H = qP) An enthalpy diagram can be used to illustrate the enthalpy change of a reaction. The enthalpy, H, is plotted on the Y‐axis and the progression of the reaction on the X‐axis. H is the value that is determined. Calorimetry Measurement of heat energy transfer Heat Capacity (C) Quantity of heat required to changed the temperature of a system by 1oC (or 1K) C = q/T Molar Heat Capacity/Specific Heat Molar heat capacity = C/n, where C is the heat capacity and n is the moles. Specific heat = C/m = q/(mT), where m is the mass in grams and T = Tf ‐ Ti 1calorie = 1 cal = amount of energy required to raise the temperature of 1g of water by 1oC (at 1 atm). 1cal = 4.184J 1kilocalorie = 1Cal = 1,000cal The specific heat of water is 4.184J/goC Enthalpy, Specific heats and Chemical reactions: Using a Styrofoam cup calorimeter: All the heat lost by the hot solid is gained by the water in the cup. The Styrofoam acts as a good insulator and has a low specific heat. If a reaction takes place inside a insulated calorimeter qrxn = ‐qcalorimeter If this reaction takes place under constant pressure qrxn = qP Bomb Calorimeter For combustion reactions and other reactions involving gases. The chamber of the bomb calorimeter keeps the system at a constant volume. E = qv = qrxn = ‐qcalorim where qcalorim = heat capacity of calorimeter x T Note that this is not necessarily the enthalpy H. Enthalpy is measured at constant pressure. (i.e. qp = H = E + PV) Hess’s Law of Heat Summation The heat of a reaction is constant, whether the reaction is carried out directly in one step or indirectly through a number of steps. Reversing an equation: Changes the sign of H Multiplying coefficients in an equation: Multiplies the value of H by the same amount. Multiplying factors may be fractional. Standard Enthalpy of Formation Scale of relative enthalpies. Standard state of solid or liquid: Pure element or compound at 1atm and the temperature of interest (often 25oC) Standard state of gas: Pure gas substance acting ideal at 1atm and the temperature of interest. Standard enthalpy of reaction (Ho) is the enthalpy change for standard state reactants yielding standard state products. Standard enthalpy of formation (heat of formation) (Hof) = enthalpy change in creating 1 mole of a substance at standard state conditions from its elements at standard states and in its reference form. The standard enthalpy of formation of a pure element in its reference form is 0. (ex. Carbon as graphite. Graphite is the reference form of carbon.) To determine the standard enthalpy of a reaction from standard heats of formation: Ho = p x Hof(products) ‐ r x Hof(reactants) and refers to the stoichiometric coefficients of the reactants and products. Ionic Reactions in Solutions: Since cations and anions cannot be formed in isolation, you cannot directly extend enthalpy to the formation of single ions. H+(aq) is arbitrarily assigned an enthalpy of formation of zero. The formation of other ions are measured relative to this. ...
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