Chapter 5 – Thermochemistry Energy – the ability to do work Law of Conservation of Energy

Also called 1st Law of Thermodynamics

Energy cannot be created or destroyed

Heat lost must equal heat gained by surroundings

2 types of energy: o Kinetic – energy of motion (Ek) o Potential – stored energy (Ep)

Thermochemistry – the study of energy changes that accompany physical, chemical or nuclear changes in matter Physical Change

A change in the form of a substance in which no chemical bonds are broken

Energy is used to overcome intermolecular forces to act (ie. there are forces between molecules)

Fundamental particles remain unchanged at the molecular level

Examples: o Water vapour forms frost on a cold day, H2O(g) H2O(s) + heat o Liquid hydrogen changes to a gas at -252ºC

Chemical Change

A change in the chemical bonds between atoms by the rearrangement of atoms into new substances

New substances are formed with new chemical bonds

Examples: o Combustion of propane gas, C3H8(g) + 5 O2(g) 3 CO2(g) + 4 H2O(g) + heat o Formation of calcium hydroxides, Ca + 2 H2O H2 + Ca(OH)2 + heat

Nuclear Change

A change in the protons or neutrons in the atom, resulting in the formation of a new atom

Energy change overcomes the forces between protons and/or neutrons in the nuclei

Example:

o Uranium 238 decays,

Kinetic Energy

Sum total of vibrational, translational, and rotational energies of molecules

When molecules move and have kinetic energy, this is referred to as thermal energy of the substance

Very difficult to accurately measure this total energy, therefore, we are concerned with energy changes (not total energy) called change in enthalpy (heat content) ΔH

Textbook pg. 296 - 357

Changes in Enthalpy

All reactions release or absorb heat (enthalpy)

ΔH = ΔHproducts - ΔHreactants Heat – a term used to describe the transfer of energy between substances

Therefore, an object can posses thermal energy but not heat Exothermic Reactions

Reactions that transfer heat (q)(amount of energy transferred between substances) from the chemical system to the surroundings (may include air, clothing and metal parts)

This transfer increases the thermal energy of the particles in the surroundings which makes them move faster, thereby increasing the temperature of the surroundings

ΔH is negative (-)

Ex. a combustion reaction o 2 C2H2(g) + 5 O2(g) 4 CO2(g) + 2 H2O(g) + energy

Endothermic Reactions

Cause a transfer of heat from the surroundings to the chemical system where it is changed into potential energy, so the temperature of the chemical system does not increase

ΔH is positive (+) Open systems – systems that allow matter and energy in and out (ex. a chemical reaction that produces gas in a solution in a beaker, boiling water in a pot, lighting an explosive) Closed systems – systems that prevent the flow of matter but not energy Isolated systems – systems that prevent the flow of both energy and matter (used for most calculations of energy changes), an ideal system achieved by using a calorimeter Calorimetry – used to measure the amount of heat transferred

Typically a reaction occurs in a large insulated container of water and the change of the water’s temperature is measured

Depends on careful measurements of mass and temperature changes

Three factors must be considered: mass (m), temperature change (ΔT), and type of substance

Uses the following equation to represent the quantity of heat (q) transferred: q = mcΔT q = quantity of heat transferred m = mass that is warming up c = specific heat capacity of the substance ΔT = change in temperature, T2 – T1

Specific heat capacity (c) – amount of heat required to raise the temperature of one gram of a substance by 1ºC (or 1 Kelvin)

o Units = J/g•ºC o Note – x J/g•ºC = x kJ/kg•ºC

Sample Problem Ex. 1) When 600 mL of water in an electric kettle is heated from 20ºC to 85ºC to make a cup of tea, how much heat flows into the water? Ex. 2) What would the final temperature be if 250.0 J of heat were transferred into 10.0 g of methanol initially at 20.0ºC? Enthalpy Change (ΔH)

The energy absorbed from or released to the surroundings when a system changes from reactants to products

ΔH of the system is equal to the energy changes of the surroundings – consistent with the law of conservation of energy as energy can be converted but must remain in the same amount

ΔHsystem = ±ǀqsurroundingsǀ = mcΔT Example – Consider the following reaction: Zn(s) + 2 HCl(aq) H2(g) + ZnCl2(aq)

Through collisions, the kinetic energy is transferred to particles in the surroundings

We can observe and measure this transfer of energy of measuring the increase in temperature of the surroundings (including solvent water molecules, the flask, the air surrounding the flask)

Molar Enthalpy (ΔHx)

The enthalpy change associated with a physical, chemical, or nuclear change involving one mole of a substance, where the x represents the type of change occurring

Enthalpy changes for exothermic reactions are give a negative sign, ΔH = (-)

Enthalpy changes for endothermic reactions are given a positive sign, ΔH = (+) Representing Enthalpy Change (ΔH)

Molar Enthalpy of Vaporization

A physical change (liquid to gas)

Obtained empirically and listed in reference books

Enthalpy Changes for Any Amount of a Substance

The amount of energy involved in an enthalpy change depends on the quantity of matter undergoing that change – do not always have 1 mol of the substance

To calculate an enthalpy change ΔH for some amount of a substance other than a mole, use the following formula: ΔH = nΔHx Where n is the number of moles you have of a substance, and ΔHx is an enthalpy value obtained from a reference source

Calorimetry and Physical Changes

Molar enthalpies are determined by studying energy changes in an isolated system (calorimeter)

Analysis of energy changes is based on the law of conservation of energy – the total energy change of the chemical system is equal to the total energy change of the surroundings ΔHsystem = ±ǀqsurroundingsǀ

There are 3 assumptions made in calorimetry: o No heat is transferred to the outside environment o The calorimeter doesn’t absorb or release any heat o The dilute aqueous solution has c = 4.18 J/(g•ºC) and d = 1.00 g/mL (pure water)

VERY IMPORTANT – when ΔHsystem = (+) (endothermic), qsurroundings will be (-)(exothermic) and VICE VERSA

o If the system gains energy, the surroundings must lose energy Sample Problems Ex. 1) Freon-12 has a molar mass 120.91 g/mol. The molar enthalpy of vaporization for Freon-12 is 34.99 kJ/mol. If 500.0 g of the refrigerant is vaporized, what is the expected enthalpy change ΔH? Ex. 2) Ethylene glycol has a molar enthalpy of vaporization of 55.8 kJ/mol. What amount of ethylene glycol would vaporize while absorbing 200.0 kJ of heat?

Ex. 3) In a calorimetry experiment, 7.46 g of potassium chloride is dissolved in 100.0 mL (100.0 g) of water at an initial temperature of 24.1ºC. The final temperature of the solution is 20.0ºC. What is the molar enthalpy of the solution of potassium chloride? Ex. 4) What mass of lithium chloride must have dissolved if the temperature of 200.0 g of water increased by 6.0ºC? The molar enthalpy of solution of lithium chloride is -37 kJ/mol. (Be careful of units!)

Representing Enthalpy Change 1. Thermochemical Equations with Energy Terms To represent thermochemical equations with energy terms, you will first need to form your chemical equation. Next, you must write your energy terms on the reactant side if your reaction if endothermic, or on the product side if your reaction is exothermic. Example: Endothermic: H2O + 285.8 kJ H2 + ½ O2 Exothermic: Mg + ½ O2 MgO + 601.6 kJ Sample Problem Write a thermochemical equation to represent the exothermic reaction that occurs when 2 moles of butane burns in excess oxygen. The molar enthalpy of combustion is -2871 kJ/mol. 2. Thermochemical Equations with ΔH Values The second way to describe the enthalpy change in the reaction is to write a balanced chemical equation and then the ΔH value beside it. Make sure that the ΔH value is given the correct sign – positive (+) for endothermic, negative (-) for exothermic. Example: CH3OH(l) + 3/2 O2(g) CO2(g) + 2 H2O(g) ΔH = -726 kJ Sample Problem Sulfur dioxide and oxygen react to form sulfur trioxide. The molar enthalpy for the combustion of sulfur dioxide in the reaction is -98.9 kJ/mol SO2. What is the enthalpy change for this reaction?

3. Molar Enthalpies of Reaction This is a method used to describe the energy change associated with the reaction of one mole of a substance. Example: CO(g) + 2 H2(g) CH3OH ΔHCH3OH = -128.6 kJ/mol CH3OH Sample Problem Write an equation whose energy change is the molar enthalpy of combustion of propanol (C3H7OH). 4. Potential Energy Diagram Using a graphical representation of the energy transferred during a physical or chemical change. The vertical axis represents the potential energy of the system, and the horizontal axis represents the reaction progress. Exothermic Reactions – products have less potential energy than reactants because energy is released to surroundings as products form Endothermic Reactions – products have more potential energy than the reactants because energy is absorbed from the surroundings

Hess’s Law Hess’s Law states that the value of ΔH for any reaction can be written in steps equal to the sum of the values of ΔH for each of the individual steps. ΔHtarget = ΔH1 + ΔH2 +ΔH3 + ... = Rules for chemical equation and enthalpy change:

1. If a chemical equation is reversed, then the sign of ΔH changes 2. If the coefficients of a chemical equation are altered by multiplying or dividing by a constant

factor, then the ΔH is altered in the same way Sample Problem Ex. 1) What is the enthalpy change for the formation of one mole of butane (C4H10) gas from its elements? The reaction is:

4 C + 5 H2 C4H10 ΔHº = ? The following known equations, determined by calorimetry, are provided:

(1) C4H10 + 13/2 O2 4 CO2 + 5 H2O ΔHº1 = -2657.4 kJ (2) C + O2 CO2 ΔHº2 = -393.5 kJ (3) 2 H2 + O2 2H2O ΔHº3 = -483.6 kJ

Ex. 2) How much energy can be obtained from the roasting of 50 kg of zinc sulfide ore?

ZnS + 3/2 O2 ZnO + SO2 You are given the following chemical equations:

(1) ZnO Zn + ½ O2 ΔHº1 = 350.5 kJ (2) S + O2 SO2 ΔHº2 = -296.8 kJ (3) ZnS Zn + S ΔHº3 = 206 kJ

Standard Enthalpies of Formation Standard enthalpy of formation is the quantity of energy associated with the formation of one mole of a substance from its elements in their standard states. For instance, the standard enthalpy of formation of carbon dioxide is:

C(s) + O2(g) CO2

We can use the standard enthalpies of formation to find the enthalpy change of our target compound by finding the sum of the standard enthalpies of formation for our product and subtracting standard enthalpies of formation for our reactant. The equation is given by: ΔH = You can find standard enthalpies of formation from a reference table (Appendix C6 in your textbook). Please note that the standard enthalpy of formation for an element in its standard state is zero (ex. O2(g), H2(g), Fe(s), etc.) Sample Problem Ex. 1) The main component in natural gas used in home heating is methane. What is the molar enthalpy of combustion of methane fuel?

CH4(g) + O2(g) CO2(g) + 2 H2O(l)

Ex. 2) The standard enthalpy of combustion of benzene to carbon dioxide and liquid water is -3273 kJ/mol. What is the standard enthalpy of formation of benzene?

C6H6(g) + 15/2 O2(g) 6 CO2(g) + 3 H2O(l)

Multi-Step Energy Calculations Using Standard Enthalpies of Formation In solving multi-step energy calculations we need to use two key relationships:

1. enthalpy change in a system = heat transferred to/from the surroundings ΔH = q = mcΔT

2. ΔH = nΔHreactant Sample Problem Ex. 1) When octane burns in an automobile, heat is released to the air, the metal in the car engine, and the majority to the cooling system – ethylene glycol. What mass of octane is completely burned to cause the heating of 20 kg of aqueous ethylene glycol from 10ºC to 70ºC? The specific heat capacity of aqueous ethylene glycol is 3.5 J/(g•ºC). Assume that water is produced as a gas and that all the heat flows into the coolant.

Ex. 2) If 3.2 g of propane burns, what temperature change will be observed if all the heat from combustion transfers into 4 kg of water?