CHAPTER 6: PRINCIPLES OF REACTIVITY ENERGY AND CHEMICAL REACTIONS

6.0 Objectives:  Describe various forms of energy and energy transfer.  Understand the terms reactant –favored, product-favored, and thermodynamics.  Differentiate between kinetic and potential energy and know the SI unit used to measure thermal energy.  Understand the term specific heat capacity and know how to calculate amount of thermal energy transferred from one object to another.  Use heat of fusion and heat of vaporization to solve simple thermal problems.  Recognize and correctly use vocabulary related to thermodynamics: system, surroundings, endothermic, exothermic, enthalpy, first law of thermodynamics, and calorimetry.  Calculate the enthalpy change of a system in several ways: graphically, experimentally, and using Hess’s Law and the summation equation.

Homework:  Homework #1 – 13, 15, 17, 19, 21, 69, 71, 73  Energy and Specific Heat  Homework #2 – 27, 29, 31, 33, 35, 37  Phase Changes and Enthalpy  Homework #3 – 39, 41, 43, 45, 47, 79  Calorimetry  Homework #4 – 51, 53, 83, Hess’s Law Worksheet  Hess’s Law  Homework #5 – 55, 57, 61, 63, Heats of Formation worksheet  Heats of Formation

6.1ENERGY: SOME BASIC PRINCIPLES  1.Definitions  Thermodynamics  the study of the effects of work, heat, and energy on a system  “energy changes that occur during chemical reactions”  Energy  “capacity to do work”

 Kinetic Energy  “energy of motion”  KE = ½ mv 2  Includes Mechanical, Thermal, Electric, Sound  Types of Motion  translational, vibrational, rotational  Potential Energy  “energy of position or arrangement”  Examples: gravitational potential, spring potential, chemical, electrostatic

2.Law of Conservation of Energy, or First Law of Thermodynamics  “Energy cannot be created nor destroyed”  Energy of the universe is constant

3.Heat vs. Temperature  definitions  Heat and Temperature ARE NOT the same!  Heat (q)-  Total kinetic energy of the atoms of a substance; HEAT IS A FORM OF ENERGY  Temperature (T)-  a measure of the average kinetic energy of the molecules in a system  Measure temperature in degrees Kelvin

 4.System vs. Surroundings  System- what is being studied  Surroundings- everything else  5.Direction of Heat Flow – 3 Principles 1. Heat transfers from a hotter object to a cooler one 2. K.E. transfer occurs until both objects are at the same temp. 3. Heat lost by hotter object = heat gained by cooler object

 6.Exothermic vs. Endothermic processes  Exothermic- heat is transferred from the system to the surroundings  Endothermic- heat is transferred from the surroundings to the system

 7.Units of Energy  Joule (J) = SI unit of energy  kg*m 2 /s 2 = Force x distance = Work  1 KJ = 1000 J  calorie (cal)  amount of energy required to raise 1 g of H 2 O 1 o C  1 cal = J  Calorie (Cal)  1 Cal = 1000 cal = 1 kcal

6.2 SPECIFIC HEAT CAPACITY AND HEAT TRANSFER  1.Specific heat capacity, definition, units, examples, c of water  Specific heat capacity  “c” - quantity of heat required to change the temp of the system by 1 o C  Units = J/g. o C or J/g. K  c-H 2 O = J/g. o C Examples:  Which heats up faster? An empty pan or a pan full of water?

 2.Calculating heat changes:  equation  q = m c  T

 Ex6.1 How many kilojoules of heat are required to increase the temperature of 35.5g of iron from 23.6 o C to 434 o C? The specific heat of iron is 0.451J/g. K.  (6.57 KJ)

 Ex6.2 What is the specific heat of benzene if 3,450J of heat is added to a 150.0g sample of benzene and its temperature increases from 22.5 o C to 35.8 o C?  C = 1.73 J/g. K

 5.Sign conventions  q = (+)  Heat is being put into the system  Endothermic  q = (-)  Heat is being released from the system  Exothermic  6.Heat transfer – definition and equation  Heat is transferred from one substance to another  Can measure this using q = mc  T  -q lost = q gained (conservation of energy)

 Ex6.3 A g piece of lead was heated in boiling water in Salt Lake City to 94.1 o C. It was removed from the water and placed into 100.0mL of water in a Styrofoam cup. The initial temperature of the water was 18.7 o C and the final temperature of the lead and water was 26.1 o C. What is the specific heat of lead according to this data?

6.3 ENERGY AND CHANGES OF STATE  1.  H fus &  H vap q=m  H   H fus – heat required to covert a substance from a solid to a liquid at its melting point (MP)  333 J/g for H 2 O   H vap – heat required to covert a substance from a liquid to a gas at its boiling point (BP)  2256 J/g for H 2 O  c H2O solid = 2.09 J/g oC  c H2O liq. = J/g oC  c H2O gas = 1.84 J/g oC

 2.Graph – Equations, PE and KE, Phase changes

 Ex6.4 Calculate the quantity of heat needed to convert 125 g of ice at –25.0 o C to steam at 175 o C.

6.4FIRST LAW OF THERMODYNAMICS  1.  E = q + w  1 st Law of Thermo:  “energy change for a system is the sum of heat transferred and the work transferred to or from the system”  E = internal energy  Sum of the potential and kinetic energy of all atoms & molecules

 2.Enthalpy  Enthalpy: H = E + PV (sum of Internal E and work)  Work is represented by PV  Like Internal E, enthalpy cannot be directly measured  Enthalpy Change:  H =  E + P  V  Heat of Rxn at constant pressure (most rxns open to atmosphere)  Standard Temperature and Pressure (STP)  1 atm, 273 K (0 o C)

 3.State functions  Depends only on present conditions (doesn’t matter how it got there) State FunctionPath Function Independent of path taken to establish property or value. Dependent on path taken to establish property or value. Can integrate using final and initial values. Need multiple integrals and limits of integration in order to integrate. Multiple steps result in same value. Multiple steps result in different value. Based on established state of system (temperature, pressure, amount, and identity of system). Based on how state of system was established. Normally represented by an uppercase letter. 1 Normally represented by a lowercase letter. 1

6.5 ENTHALPY CHANGES FOR CHEMICAL RXNS 1.Measuring  H  a. Graphically

 b. Experimentally  Determine how much heat is evolved or absorbed for a given reaction  Calorimetry (next topic)  c. Mathematically   H rxn = Energy (Products) - Energy (Reactants)

 2. Principles of Enthalpy  Identity and state are important  H 2 O (l) is different from H 2 O (g)  Different energies for the different states  Positive or negative values  (-) = Exothermic; heat added to products side  (+) = Endothermic; heat added to reactants side  Stoichiometric quantity  Coefficients balance the chemical equation  May not always have whole numbers or reduced coefficients  Changing the coefficients changes the  H!

 Reverse reactions  Reactions are cyclic in nature  If Rxn is reversed, must change sign  Ex. 2HgO (s) --> 2Hg (l) + O 2 (g)  H= kJ 2Hg (l) + O 2 (g) --> 2HgO (s)  H= kJ  3.Thermochemical equation –  H 2 (g) + ½ O 2 (g)  H 2 O (l)  H = kJ

 Ex6.5 When 5.00g of sucrose, C 12 H 22 O 11, is burned to produce gaseous carbon dioxide and liquid water, 82 kJ of energy are produced. Write the correct thermochemical equation for this reaction.

6.6CALORIMETRY  1.Definition and assumptions  Calorimetry - Technique used to measure a quantity of heat based on conservation of energy  Heat flows from the hotter to the cooler object resulting in a temp change  Heat lost by hot solids is gained by water in calorimeter  q = mc  T

 2.Coffee cup calorimeter, equation q H2O = -q object q object = amt. Of heat given off by substance as it cools to a common T with H 2 O cm∆T = cm ∆T q H2Oor Soln. = -q rxn q H2O = amt. Of heat given off by rxn q rxn = cm ∆T

 Ex6.6 A student placed 0.500g of Mg(s) in a coffee cup calorimeter along with 100.0mL of 1.00M HCl. The reaction produces aqueous magnesium chloride and hydrogen gas. The temperature of the solution rises from 22.2 to 44.8 o C. What is the enthalpy of this reaction? Assume that the specific heat capacity of the solution is 4.2 J/g K and the density of the solution is 1.00g/mL.

 Advantages and disadvantages:  Easy to do; simple instrumentation; relatively robust/accurate  Sig figs are limited to tools used; some heat is lost to cup

 4.Bomb Calorimeter  q rxn + q bomb + q water = 0  Initiate rxn by ignition wire  Treat whole metal assembly as one Sp. Heat capacity  Use for reactions involving gases  Ex. Combustion of Gasoline (Octane)  q bomb = C bomb  T  q water = cm  T  q rxn = -(q bomb +q water )

 Ex6.7 When a 1.00g sample of naphthalene, C 10 H 8, is burned in a bomb calorimeter with a heat capacity of 13.24kJ/K and 1.20kg of water, the temperature rises from to o C. Write the thermochemical equation for this reaction.

6.7 HESS’S LAW  1.Statement and usage  “If a reaction is the sum of two or more other reactions,  H for the overall process is the sum of the  H values of those reactions”

 2. Ex6.8 From the following information calculate the enthalpy of the formation of one mole of SO 3 from its elements. S(s) + O 2 (g)  SO 2 (g)  H = kJ 2SO 2 (g) + O 2 (g)  2SO 3 (g)  H = kJ

 3. Ex6.9 Use the following equations and a Hess’s Law process to determine the enthalpy of formation of 1.00 mole of NO(g) from its reactants. N 2 (g) + 3H 2 (g)  2NH 3 (g)  H = -91.8kJ 4NH 3 (g) + 5O 2 (g)  4NO (g) + 6H 2 O(g)  H = kJ H 2 (g) + 1/2O 2 (g)  H 2 O(g)  H = kJ

 1.Definition  “Enthalpy change for the formation of 1 mole of a compound directly from its elements in their standard states (1 atm; 25 o C)”   H f o

 2.Examples  CO 2 (g)  C (s) + O 2 (g)  CO 2 (g)  H f o = kJ  N 2 (g)  H f o = 0 kJ!!  3.Reference book values  See Reference booklet for Standard enthalpy of formation values (  H f o )

 4.Summation equation   H rxn =  H o products -  H o reactants

 5. Ex6.10 Use the summation equation to determine the enthalpy of the following reaction: 4NH 3 (g) + 5O 2 (g)  4NO(g) + 6H 2 O(g)

 Ex6.11 Use the table of heats for formation and the equation in example 6.7 to calculate the standard heat of formation for naphthalene, C 10 H 8.example 6.7