Thermochemistry

Topic 5 Thermochemistry

Words you need to know: • Thermochemistry • Thermodynamics • Energy • Heat • calorie/Calorie • Joule • Energy transfers occur btwn the system and its surroundings

calorie v. Calorie • calorie is the amount of heat energy needed to raise the temperature of 1 g of water 1oC • Calorie is 1000 calories • Food energy is measured in Calories • Fats are 9 Calories/g • Carbs and proteins are 4 Calories/g

1st Law of Thermodynamics • aka Law of conservation of energy • Energy in the universe is constant, cannot be created or destroyed • Energy can be converted to different forms

Exothermic Processes • System transfers heat to its surroundings • Temperature of the surroundings increases • “feels hot” • Potential energy of the system is converted to heat energy that is released (system does not “cool off”)

Endothermic Processes • System absorbs heat from the surroundings • Temperature of the surroundings decreases • “feels cool” • Heat energy from the surroundings is converted into potential energy in the system

Practice • Imagine an ice cube melting in your hand. Is the melting of ice endothermic or exothermic? Explain using the terms system and surroundings • Imagine warming your hands near a campfire. Is a campfire an endothermic or exothermic process? Explain.

Enthalpy (H) • Property of a system that explains heat flow between the system and its surroundings (constant P) • State function (middle steps don’t matter, just the beginning and end of a reaction)

Enthalpy changes (DH) • Defined as heat absorbed by the system during a physical or chemical change • DH is positive for endothermic rxns (because heat is absorbed by the system) • DH is negative for exothermic rxns (because heat is released by the system) • Expressed in kJ or kJ/mol • Magnitude of DH is directly proportional to moles of reactants and products

Types of Enthalpy Changes • DHrxn = heat of reaction of any chemical reaction • DHcomb = heat of combustion for combustion reactions (rxns with O2) only • DHfus = heat of fusion when a solid melts

More Types • DHvap = heat of vaporization when a liquid vaporizes • DHBDE = bond dissociation energy or the heat required to break a bond • DHf = heat of formation or the heat change when a compound is formed from its elements • DHsols = heat of solution or the heat change when a solute dissolves in a solvent

Practice How many kJ of heat are absorbed when 25.0 g of methane burn in air? Methane has a DHcomb of -802 kJ/mol.

Reversible Reactions • If DH is positive for the forward reaction, then it will be equal in size but opposite in sign for the reverse reaction H2O (s)  H2O (l) DH = +6.0 kJ/mol H2O (l)  H2O (s) DH = -6.0 kJ/mol

Keep in mind… • The sign for DH is positive or negative depending on the direction of energy flow • The sign does NOT indicate a positive or negative value for energy • What scientific law requires that the magnitude of the heat change for a forward and reverse reactions be the same with opposite signs? Explain.

DH and Phase Changes • DH for phase changes from solid to liquid and liquid to gas are ALWAYS positive (endothermic/absorb heat) • DH for phase changes from gas to liquid and liquid to solid are ALWAYS negative (exothermic/release heat) • Why is this? What is the sign of DH for the process of sublimation? Deposition?

Calorimetry • aka. measurement of heat flow • Calorimeter measure heat flow • Heat capacity (C) = amount of heat required to raise the temperature of any object 1oC. Expressed in J/K or J/oC • We will more often use specific heat capacity (Cp) which is the capacity of 1 g of a substance • Water has a Cp of 1 cal/gK or 4.184 J/gK

Explain • Why is the Cp of water 1 cal/gK? What is the conversion between calories and joules?

Example What is the molar heat of combustion of liquid ethanol if the combustion of 9.03 grams of ethanol causes a calorimeter to increase in temperature by 3.54 K? The heat capacity of the calorimeter is 75.8 kJ/K.

Hess’s Law • If a reaction is carried out in a series of steps the DH of the overall reaction is equal to the sum of the DH’s for each individual step. • Useful for determining DH for reactions that are difficult to measure directly, like sulfur trioxide…

Hess’s Law with SO3 The overall reaction is 2S (s) + 3O2 (g)  2SO3 (g) The reaction occurs in 2 measurable steps S (s) + O2 (g)  SO2 (g) DHrxn = -269.9 kJ 2SO2 (g) + O2 (g)  2SO3 (g) DHrxn = -196.6 kJ To get the total DH for the reaction, manipulate the equation steps like an algebraic equation. Whatever you do to the reaction, you must also do to DH.

Enthalpies of Formation • Reaction that produces 1 mole of a substance from its constituent elements in their most stable thermodynamic state • To form 1 mole of HI, the equation looks like this: ½ H2 + ½ I2 HI DH = +25.94 kJ

Standard Heat of Formation (DHof) • Heat absorbed when 1 mole of a substance is formed from its elements in their standard states at 25oC and 1 atm. • There is a BIG table in the back of your book listing standard heats of formation.

Example Write the thermochemical equation associated with the standard heat of formation of AlCl3. What is the DHof for this equation?

Things to know… • The table of standard heats of formation includes elements, ions, and compounds. • The DHof of pure elements is always zero.

One last thing… • Hess’s law allows us to calculate the DHorxn for just about any reaction. • Breaking the overall reaction into the formation reactions for both products and reactants and then putting them all together like this: DHorxn = SnDHof products – SmDHof reactants

Example Calculate the standard enthalpy change for the combustion of 1 mole of liquid ethanol. Tip: carefully watch the signs of DHof.

Another Example What is: • The enthalpy of sublimation of solid calcium? • The heat of solution of gaseous ammonia? • The bond dissociation energy of hydrogen gas? • The heat change when gaseous bromine condenses to a liquid? Write chemical equations to illustrate your answers.

Thermochemistry