Chapter 6: Thermochemistry

Energy Energy - the capacity to do work Work is the result of a force acting over a distance Types of Energy Kinetic Potential Law of conservation of energy - energy can not be created or destroyed, but it can be transformed

Kinetic energy Kinetic energy - Energy associated with the motion of an object Mechanical energy – the energy of moving macroscale objects Thermal energy (heat) – the energy of moving microscale objects

Potential energy Potential energy – energy associated with position or composition, stored energy Has the potential to be transformed into kinetic energy Chemical energy – relative positions of electrons and nuclei in atoms/molecules represents potential energy

Energy: System vs surroundings System - A component or set of components of interest You define the system Surroundings – everything that a system can exchange energy with, everything touching/connected to it That which “surrounds” the system

Units of energy Joule (J) - basic unit of energy in the metric system 1J =1kg •m2/s2 1kJ = 1000 J calorie (cal) – 1cal = the energy required to raise the temp of 1g of H2O by 1ºC 4.184 J = 1 cal 1000 cal = 1kcal = 1Cal Calorie (Cal) – unit associated with “food” calories Calorie = kcal = food calorie

First law of thermodynamics Thermodynamics – the study of energy and its interconversions First law of thermodynamics – the total energy of the universe is constant Law of conservation of energy is built into the definition Perpetual motion machine (a device that puts out constant energy) cannot exist

Internal energy Internal energy (E) – the sum of the kinetic and potential energies of all particles that compose a system E is a state function which is only dependent on the state of the system, not how it got there

Energy of a system in a reaction Where does lost energy go or where does gained energy come from? ………. the surroundings!

Practice A chemical reaction occurs!!! The reactants have a total energy of 75 kJ, while the products have a total energy of 45 kJ. What is the change in internal energy of this system (the compounds involved in the chemical reaction are the system)? Did the system lose or gain energy? Did the surroundings lose or gain energy?

Heat and work Heat and work can be exchanged between the surroundings and the system q = heat, the flow of energy caused by a temperature difference w = work

Heat flow Thermal equilibrium – heat is distributed to molecules in contact with them until the thermal energy of the system and surroundings is the same temperature

Quantifying heat Substances can transfer heat, but have different capacities to do so Heat capacity (C) – the quantity of heat required to change a substances temperature by 1 °C

Problems A thief breaks into a safe at a bank to steal a block of gold in the shape of a coin from Super Mario Bros. The safe keeps the gold at 14.0 °C. Once the thief grabs the gold coin in his hand, the coin eventually reaches body temperature at 37.0°C. How much heat was transferred to the gold coin which has a mass of 11.25g?

Thermal energy transfer When two substances of different temperatures reach thermal equilibrium with each other, the heat from the hotter substance is transferred to the colder substance The hotter substance cools down, as the colder substance warms up, with their final temperatures being the same 𝑞 𝑠𝑦𝑡𝑒𝑚 =− 𝑞 𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔 𝑞 𝑚𝑒𝑡𝑎𝑙 =− 𝑞 𝑤𝑎𝑡𝑒𝑟 𝑚 𝑚𝑒𝑡𝑎𝑙 𝐶 𝑚𝑒𝑡𝑎𝑙 Δ 𝑇 𝑚𝑒𝑡𝑎𝑙 =− 𝑚 𝑤𝑎𝑡𝑒𝑟 𝐶 𝑤𝑎𝑡𝑒𝑟 Δ 𝑇 𝑤𝑎𝑡𝑒𝑟

Thermal energy transfer problems With the masses and initial temperatures of the metal and water, what will the final temperature be when they both reach thermal equilibrium?

Distribute Kmetal and Kwater Collect Tfinal terms on the leftside, with other terms going to right side Factor out Tfinal Isolate Tfinal

Practice A 32.5 g block of aluminum initially at 45.8 °C is submerged into 105.6 g of water initially at 15.4 °C. What is the final temperature once the two substances reach thermal equilibrium?

Pressure-volume work – when a force causes a volume change against an external pressure If the system does work to increase the volume against an external pressure, the ΔV is positive Pressure-volume work

Using the equation Units of pressure = atm Units of volume = L Pressure x volume = L x atm You can convert units of L x atm to units of J 1 L x atm = 101.3 J

ΔE for chemical reactions If a system is at constant volume, then there is no work done If a chemical reaction does no work…. qv = the heat at constant volume

Measuring ΔErxn Calorimetry- measurement of the external energy exchanged between a reaction(system) and the surrounding by monitoring the temperature change Bomb calorimeter is a device used in calorimetry The chemical reaction(system) transfers heat to the bomb (surrounding)

Enthalpy (H) Enthalpy – the sum of the internal energy of a system and the product of its pressure and volume For a process at a constant pressure….. ΔH is the heat absorbed or evolved by the system Typically ΔE and ΔH are the same

Endothermic vs Exothermic If the ΔH of a chemical reaction is positive, the system absorbs heat from the surroundings, this is an endothermic reaction If the ΔH of a chemical reaction is negative, the system releases heat to the surroundings, this is an exothermic reaction

How does a chemical reaction give off heat??? Chemical “potential” energy is stored in the bonds Chemical energy arises from the electrostatic forces between protons and electron that compose atoms and molecules Breaking bonds absorbs energy, while making bonds releases energy Breaking stronger bonds absorbs more energy than breaking weaker bonds

ΔH and thermochemical equations ΔHrxn – the enthalpy of reaction or heat of reaction, the heat transferred in a reaction C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) ΔHrxn = -2044 kJ When the reaction occurs with the corresponding moles from the equation, 2044kJ of heat is released from the system (exothermic reaction) 1 mol C3H8: -2044kJ or 5 mol O2: -2044kJ

Practice An LP gas tank in a home barbeque contains 13.2 kg of propane, C3H8. Calculate the heat (in kJ) associated with the complete combustion of all the propane in the tank. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) ΔHrxn = -2044 kJ

Measuring ΔHrxn 𝑞 soln = 𝑚 soln × 𝐶 s,soln ×Δ𝑇 − 𝑞 soln = 𝑞 𝑟𝑥𝑛 Can calculate ΔHrxn using a coffee-cup calorimeter By recording the change in temperature after a reaction occurs, depending on the mass and solution heat capacity, you can calculate ΔHrxn Coffee-cup calorimetry occurs at constant pressure 𝑞 soln = 𝑚 soln × 𝐶 s,soln ×Δ𝑇 − 𝑞 soln = 𝑞 𝑟𝑥𝑛 Δ 𝐻 𝑟𝑥𝑛 = 𝑞 𝑝 = 𝑞 𝑟𝑥𝑛

Practice Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g) In an experiment to determine the enthalpy change for this reaction, 0.158 g Mg metal is combined with enough HCl to make 100.0 mL of solution in a coffee cup calorimeter. The HCl is sufficiently concentrated so that the Mg completely reacts. The temperature of the solution rises from 25.6 °C to 32.8 °C as a result from the reaction. The density of the solution is 1.00 g/mL and the Cs,soln = 4.18 J/(g·°C). What is the mass of the solution? What is the qsoln? What is the qrxn? How many moles of Mg are initially present? What is the ΔHrxn in kJ/mol?

Chemical equations and ΔHrxn Rules for chemical equations and ΔHrxn 1. If a chemical equation is multiplied by some factor, then ΔHrxn is also multiplied by the same factor 2. If a chemical equation is reversed, then ΔHrxn changes sign

Chemical equations and ΔHrxn 3. If a chemical equation can be expressed by the sum of a series of steps, then ΔHrxn for the overall equation is the sum of the heats of reaction for each step Anything on both the product and reactant side cancels out This is known as Hess’s law

Practice Calculate ΔHrxn for the following reaction Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) Use the following reactions 2Fe(s) + 3/2O2(g)  Fe2O3(s) ΔH = -824.2 kJ CO(g) +1/2O2(g)  CO2(g) ΔH = -282.7 kJ

Standard enthalpies of formation ΔH°f ΔH°f is the change in enthalpy when 1 mol of product forms from its constituent elements H2(g) + ½O2(g)  H2O(g) ΔH°f = -241.8 kJ/mol ΔH°rxn can be calculated using ΔH°f of each reactant and product if all compounds are in their standard states

Standard states For a gas: when a pure gas is at a pressure of 1 atm Standard states are defined below For a gas: when a pure gas is at a pressure of 1 atm For a liquid/solid: A pure substance in its most stable form at 1 atm and 25° C For a substance in solution: a substance in a solution at a concentration of 1M

Calculating the ΔH°rxn Standard enthalpy change of a reaction ΔH°rxn – can be determined through the following equation np = moles of product nr = moles of reactant

Gathering ΔH°f values 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g) Identify the ΔHf° of the products and reactants using charts ΔHf° of NH3(g) = -45.9 kJ/mol ΔHf° of O2(g) = 0 kJ/mol ΔHf° of NO(g) = 91.3 kJ/mol ΔHf° of H2O(g) = -241.8 kJ/mol

Setting up equation for ΔH°rxn 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g) Take the ΔHf° of a product and multiply by the number of moles Then do this for each product Then take the sum of those values 4(ΔHf°NO)+ 6(ΔHf°H2O) = 4(91.3kJ/mol)+ 6(-241.8 kJ/mol) = -1085.6 kJ (products) Do the same for the reactants 4(ΔHf°NH3)+ 5(ΔHf°O2) = 4(-45.9 kJ/mol)+ 5(0 kJ/mol) = -183.6 kJ (reactants)

Finishing the calculation for ΔH°rxn Finally: products – reactants (-1085.6 kJ) – (- 183.6kJ) = -902.0 kJ = ΔH°rxn

Problem What is the ΔH°rxn for the combustion of C3H8(g)?

Chapter 6 done.