Chemistry 1A: Chapter 6 P a g e | 1 Chapter 6: Thermochemistry Homework: Read Chapter 6 Work out sample/practice exercises in the sections as you read, Keep up with MasteringChemistry Chemical Thermodynamics: Chemical thermodynamics is the study of the changes and the transfers of energy that accompany chemical and physical processes. It is VERY IMPORTANT to always identify the phases in chemical reactions. This unit also will occasionally require balancing reactions with fractions. Thermodynamics addresses 3 fundamental questions: 1) Will two or more substances spontaneously react when mixed under specified conditions or will it be nonspontaneous (will not react)? 2) What energy changes and transfers are associated with the reaction? 3) To what extent does a reaction occur (equilibrium)? (learn in Chem 1B) Systems tend toward… 1) Minimizing potential energy 2) Maximizing entropy (degrees of freedom or ways of being). Thermochemistry tells us nothing of the reaction rate. Reactions may be spontaneously favored under specified conditions, but may not react at an observable rate, or reactions may need a large initial amount of energy to initiate and reach the activation energy. Thermodynamic examples: Methane will spontaneously react and combust in oxygen, but the reaction does not initiate without a spark. Ice will spontaneously melt into water, but only if the temperature is above freezing. We will study physical and chemical processes accompanied by the transfer of heat or work. Nuclear energy, where the atom changes, is different. Chemistry 1A: Chapter 6 P a g e | 2 Example 1: Systems tend toward minimizing potential energy and maximizing entropy. Identify the side that minimizes the potential energy. Identify the side that maximizes ways of being (entropy). Assume the same temperature and atmospheric pressure for each. a) Top of waterfall or Bottom of waterfall b) 1 mole H2 molecules or 2 mole of H atoms c) H2O (g) at 100°C or H2O (l) at 100°C Nature of Energy: Energy is the capacity to do work or transfer heat. Energy (a quantity) can exist in a variety of forms and can be transformed from one form to another. It is measured by the amount of work done, usually in joules or watts. Chemistry generally uses joules. SI units of energy: 1 Joule = 1 kg m2/s2 Energy Conversions 1 calorie (cal) = 4.184 joules (J) (exact) 1 Calorie (Cal, food) = 1000 cal = 1 kcal = 4184 J 1 kilowatt-hour (kWh) = 3.60 x 106 J Some Forms of Energy: • Electrical kinetic energy associated with the flow of electrical charge • Mechanical kinetic energy associated with moving parts • Heat or thermal energy kinetic energy associated with molecular motion • Light or radiant energy kinetic energy associated with energy transitions in an atom • Nuclear potential energy in the nucleus of atoms • Chemical potential energy due to the structure of the atoms, the attachment between atoms, the atoms’ positions relative to each other in the molecule, or the molecules, relative positions in the structure Chemistry 1A: Chapter 6 P a g e | 3 Energy can be classified into two major categories, potential and kinetic. Potential Energy is the energy by virtue of its position of composition. a) For an object that has height, potential energy is the mass x gravity (9.8 m/s2) x height; Ep = mgh. b) Another form is electrostatic potential energy from interactions with charged 9 2 particles; Eel = kQ1Q2/d, where k = 8.99 x 10 J m/C and C is coulomb that is a unit of charge, Q is charge, and d is distance. Like charges repel and opposites attract. c) Potential energy also comes from the arrangement of chemical compositions. When bonds break and new ones form in a chemical reaction, the energy change generally is due to the changes in potential energy/composition. Kinetic Energy is the energy of motion. For a moving object kinetic energy is half 2 the mass x velocity squared; Ek = ½ mv Example 2: What is the potential energy of a 400 g ball on top of a building that is 30.0 m tall? Example 3: What is the kinetic energy of a 400 g ball moving at 30.0 m/s? Some Thermochemistry terms… System: the portion we study Surroundings: everything else affected by the portion we study Universe: the combined system plus surroundings. In the lab experiments our universe will be a calorimeter and its contents. Chemistry 1A: Chapter 6 P a g e | 4 Laws of Thermodynamics: Laws of thermodynamics are useful in predicting outcomes, but unlike a theory do not explain the expected behavior. Zeroth Law: Temperature Concept: Temperature measures the intensity of hotness or coldness of an object. When two objects are brought together heat always flows spontaneously from a hotter object to a colder object until thermal equilibrium is reached. Think of the heat flow direction when placing your hand in ice water verses hot water. First Law: Law of Conservation of Energy: Energy cannot be created or destroyed. The energy of the universe is constant. This law works for ordinary chemical and physical processes. It was later combined with another (conservation of matter) to include nuclear reactions after E =mc2. It was changed to the Law of Conservation of Matter and Energy Example 4: What is the difference between temperature and heat? Example 5: Review your temperature conversions between °F, °C, and Kelvin. a) 98.6 °F = ? °C = ? K b) 788 K = ? °C = ? °F When you see X where X can be anything remember that it is just the change from the final minus the initial values. X = Xfinal – Xinitial Definitions: Heat is the flow of energy caused by temperature differences Enthalpy,H; the energy accounting for heat flow in constant pressure processes. Units are generally kJ/mol. When the energy is measured in J or kJ for a particular system, qp may be used. Chemistry 1A: Chapter 6 P a g e | 5 Internal Energy, E or U; the sum of all the kinetic and potential energies of all components. This includes all motions of vibration, rotation, movement through space of the object, its atoms, nuclei, and even including subatomic particles like electrons. We cannot truly measure all of this, so we generally try to solve for the difference or change in internal energy in our studies. E = Efinal – Einitial When solving for E we work with a constant volume system. Units are generally kJ/mol. When the energy is measured in J or kJ for a particular system, qv may be used. Work, w; Work is force acting over a distance, (work = force x distance), the energy used to cause an object to move. Pressure-Volume work involves the compression or expansion of gases. Work = -PV ; if using atm and liter: convert to J with R/R (101.3 J/L atm) Work = -nRT; use the R with joules: R = 8.314 J/mol K For most reactions the energy from work is quite small. Expansion work is (-); work done by the system on the surroundings. Compression work is (+); work done on the system by the surroundings Relating together; qv = qp + w , or it can be written E = H + w Endothermic: a process that absorbs heat (+) Exothermic: a process that releases heat (-) State Function: a property that depends only on the present state and not the path the system took to reach the state. In Thermochemistry state functions are generally capitalized. (Notice work is a small w and is not a state function) Examples of state functions include: E, H, V, P, T. Chemistry 1A: Chapter 6 P a g e | 6 Example 6: a) As 1.00 mole of H2O gas becomes H2O liquid at 100°C, 40.7 kJ of energy are transferred at 1.00 atm pressure. What is the enthalpy sign for this conversion? Is this process endothermic or exothermic? b) During this process at 1.00 atm pressure, the volume decreases from 30.6 liters down to 18 ml. Solve for the work in L atm and in J. Is the sign of work positive or negative? Is work done on the system by the surroundings (compression) or by the system on the surroundings (expansion) ? Example 7: Predict the sign for work, determine if the system is expanding or compressing and state if work is done on or by the system. a) 2 NO (g) + O2 (g) 2 NO2 (g) b) 2 C4H10 (g) + 13 O2 (g) 10 H2O (l) + 8 CO2 (g) Example 8: Which of the following are state functions and which are not state functions ? a) The temperature of an ice cube. b) The enthalpy of fusion for water. c) The time to travel 1-mile Chemistry 1A: Chapter 6 P a g e | 7 Enthalpies of Reaction: Thermochemical Equations with Enthalpy: 1) Always include phases (s, l, g, aq) of substances. Energy varies with phase. 2) Assume temperature is constant for reactants at the beginning and products at the end, even though the temperature may change during the reaction. H is the energy required to return a system to the starting temperature at the completion of the reaction. 3) Reactions may be balanced using fractions and not whole numbers. 4) Enthalpy (H) is an extensive property. When a reaction is doubled, the enthalpy doubles. 5) When a reaction is reversed the enthalpy changes its sign. 6) Change in enthalpy may be experimentally determined using calorimeters. 7) The ° indicates standard conditions (1 atm and 25°C). 8) Scientists arbitrarily agreed to set a zero Hf° defined as the enthalpy of formation of elements in their most common form under standard state conditions. All other Hf° values are calculated from these 9) Enthalpy of formation Hf°) is the formation of 1 mole of a compound from its elements in their most common form under standard conditions.