Chemistry: Atoms First Second Edition Julia Burdge & Jason Overby Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 10 Energy Changes in Chemical Reactions M. Stacey Thomson Pasco-Hernando State College

Energy and Energy Changes The system is a part of the universe that is of specific interest. The surroundings constitute the rest of the universe outside the system. The system is usually defined as the substances involved in chemical and physical changes. System Surroundings Universe = System + Surroundings 10.1

Energy and Energy Changes Thermochemistry is the study of heat (the transfer of thermal energy) in chemical reactions. Heat is the transfer of thermal energy. Heat is either absorbed or released during a process. The SI energy unit is a Joule, J. Often calories are used and 1 calorie is the amount of heat required to raise 1 g of water by 1°C.  1 cal = J  calorie is not the same as a nutritional calorie (Cal)  1 Cal=1000 cal Surroundings heat

Energy and Energy Changes An exothermic process occurs when heat is transferred from the system to the surroundings. “Feels hot!” Surroundings Universe = System + Surroundings heat System 2H 2 (g) + O 2 (g)2H 2 O(l) + energy

Energy and Energy Changes An endothermic process occurs when heat is transferred from the surroundings to the system. “Feels cold” energy + 2HgO(s)2Hg(l) + O 2 (g) Surroundings Universe = System + Surroundings heat System hill.com/sites/ /student_view0/chapter10/animations.html# hill.com/sites/ /student_view0/chapter10/animations.html#

Introduction to Thermodynamics Thermodynamics is the study of the interconversion of heat and other kinds of energy. In thermodynamics, there are three types of systems: An open system can exchange mass and energy with the surroundings. A closed system allows the transfer of energy but not mass. An isolated system does not exchange either mass or energy with its surroundings. 10.2

The First Law of Thermodynamics The first law of thermodynamics states that energy can be converted from one form to another, but cannot be created or destroyed. ΔU is the change in the internal energy. “sys” and “surr” denote system and surroundings, respectively. ΔU = U f – U i ; the difference in the energies of the initial and final states. ΔU sys + ΔU surr = 0 ΔU sys = –ΔU surr

Work and Heat The overall change in the system’s internal energy is given by: q is heat  q is positive for an endothermic process (heat absorbed by the system)  q is negative for an exothermic process (heat released by the system) w is work  w is positive for work done on the system  w is negative for work done by the system ΔU = q + w

Work and Heat ΔU = q + w

Worked Example 10.1 Strategy Combine the two contributions to internal energy using ΔU = q + w and the sign conventions for q and w. Text Practice: Calculate the overall change in internal energy, ΔU, (in joules) for a system that absorbs 188 J of heat and does 141 J of work on its surroundings.

Enthalpy The thermodynamic function of a system called enthalpy (H) is defined by the equation: H = U + PV A note about SI units: Pressure:pascal; 1Pa = 1 kg/(m. s 2 ) Volume:cubic meters; m 3 PV:1kg/(m. s 2 ) x m 3 = 1(kg. m 2 )/s 2 = 1 J Enthalpy:joules U, P, V, and H are all state functions. 10.3

Enthalpy and Enthalpy Changes The enthalpy of reaction (ΔH) is the difference between the enthalpies of the products and the enthalpies of the reactants: Assumes reactions in the lab occur at constant pressure ΔH > 0 (positive)endothermic process ΔH < 0 (negative)exothermic process ΔH = H(products) – H(reactants)

Thermochemical Equations Concepts to consider:  What is the system?  What are the surroundings?  ΔH > 0endothermic H 2 O(s)H 2 O(l) ΔH = kJ/mol

Thermochemical Equations Concepts to consider:  What is the system?  What are the surroundings?  ΔH < 0exothermic CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) ΔH = −890.4 kJ/mol

Thermochemical Equations Enthalpy is an extensive property. Extensive properties are dependent on the amount of matter involved. H 2 O(l) → H 2 O(g)ΔH = +44 kJ/mol 2H 2 O(l) → 2H 2 O(g)ΔH = +88 kJ/mol Units refer to mole of reaction as written Double the amount of matterDouble the enthalpy

Thermochemical Equations The following guidelines are useful when considering thermochemical equations: 1)Always specify the physical states of reactants and products because they help determine the actual enthapy changes. CH 4 (g) + 2O 2 (g)ΔH = −802.4 kJ/molCO 2 (g) + 2H 2 O(g) CH 4 (g) + 2O 2 (g)ΔH = kJ/molCO 2 (g) + 2H 2 O(l) different statesdifferent enthalpies

Thermochemical Equations The following guidelines are useful when considering thermochemical equations: 2)When multiplying an equation by a factor (n), multiply the ΔH value by same factor. CH 4 (g) + 2O 2 (g) ΔH = − kJ/mol CO 2 (g) + 2H 2 O(g) 2CH 4 (g) + 4O 2 (g) ΔH = − kJ/mol 2CO 2 (g) + 4H 2 O(g) 3)Reversing an equation changes the sign but not the magnitude of ΔH. CH 4 (g) + 2O 2 (g) ΔH = − kJ/mol CO 2 (g) + 2H 2 O(g) ΔH = kJ/molCH 4 (g) + 2O 2 (g)

Worked Example 10.3 Given the thermochemical equation for photosynthesis, 6H 2 O(l) + 6CO 2 (g) → C 6 H 12 O 6 (s) + 6O 2 (g)ΔH = kJ/mol calculate the solar energy required to produce 75.0 g of C 6 H 12 O 6.

Calorimetry Calorimetry is the measurement of heat changes. Heat changes are measured in a device called a calorimeter. The specific heat (s) of a substance is the amount of heat required to raise the temperature of 1 g of the substance by 1°C. 10.4

Specific Heat and Heat Capacity The heat capacity (C) is the amount of heat required to raise the temperature of an object by 1°C. The “object” may be a given quantity of a particular substance. Specific heat capacity has units of J/(g °C) Heat capacity has units of J/°C Specific heat capacity of water heat capacity of 1 kg water

Specific Heat and Heat Capacity The heat associated with a temperature change may be calculated: m is the mass. s is the specific heat. ΔT is the change in temperature (ΔT = T final – T initial ). C is the heat capacity. q = msΔTq = CΔT

Worked Example 10.4 Calculate the amount of heat (in kJ) required to heat 255 g of water from 25.2°C to 90.5°C.

Calorimetry A coffee-cup calorimeter may be used to measure the heat exchange for a variety of reactions at constant pressure:

Calorimetry Concepts to consider for coffee-cup calorimetry: q P = ΔH System:reactants and products (the reaction) Surroundings:water in the calorimeter For an exothermic reaction:  the system loses heat  the surroundings gain (absorb) heat q sys = −msΔT The minus sign is used to keep sign conventions consistent. q surr = msΔTq sys = −q surr

Worked Example 10.5 A metal pellet with a mass of g, originally at 88.4°C, is dropped into 125 g of water originally at 25.1°C. The final temperature of both pellet and the water is 31.3°C. Calculate the heat capacity C (in J/°C) of the pellet. hill.com/sites/ /student_view0/chapter10/animations.html# hill.com/sites/ /student_view0/chapter10/animations.html#

Constant-Volume Calorimetry Constant volume calorimetry is carried out in a device known as a constant-volume bomb. q cal = −q rxn A constant-volume calorimeter is an isolated system. Bomb calorimeters are typically used to determine heats of combustion.

Constant-Volume Calorimetry q rxn = −C cal ΔT q cal = C cal ΔT q rxn = −q cal To calculate q cal, the heat capacity of the calorimeter must be known.

Worked Example 10.6 A chocolate chip cookie weighing 7.25 g is burned in a bomb calorimeter to determine its energy content. The heat capacity of the calorimeter is kJ/°C. During the combustion, the temperature of the water in the calorimeter increases by 3.90°C. Calculate the energy content (in kJ/g) of the cookie.

Study Guide for Sections DAY 26, Terms to know: Sections system, surroundings, thermochemistry, thermal energy, exothermic, endothermic, open system, closed system, isolated system, first law of thermodynamics (law of conservation of energy), enthalpy, calorimetry, specific heat, heat capacity, heat of reaction, coffee cup calorimeter, bomb calorimeter DAY 26, Specific outcomes and skills that may be tested on exam 4: Sections Be able to describe potential energy changes for the system, kinetic energy and temperature changes for the surroundings, and bond stability changes for the system in any endothermic or exothermic system. Given either change in energy for the system or surroundings, be able to calculate the other Given two out of three of any of the following quantities, be able to calculate the third: change in heat for the system, change in work for the system, overall change in energy for the system Given initial heat and final heat, be able to calculate the change in enthalpy for a system Be able to determine whether a process is endo or exothermic based on the mathematical sign for the change in enthalpy Be able to calculate how the enthalpy change will be different if a chemical equation is reversed or multiplied by a certain stoichiometric coefficient Given two of the following quantities, be able to calculate the third: temperature change, heat change, heat capacity Given three of the following quantities, be able to calculate the fourth: temperature change, heat change, mass, specific heat Be able to describe how to set up a coffee cup or bomb calorimeter to calculate changes in energy based on measured changes in temperature and energy exchanges between system and surroundings

Extra Practice Problems for Sections Complete these problems outside of class until you are confident you have learned the SKILLS in this section outlined on the study guide and we will review some of them next class period VC10.3 VC

Prep for Day 27 Must Watch videos: (energy calculations, Tyler DeWitt) (calculations with energy, crash course chemistry) (Energy, UC-Berkeley, lessons and ) (Bond enthalpy, TheChemistrySolution) (colligative properties, Isaacs) (colligative properties II, Isaacs) Other helpful videos: lectures/lecture-16/ lectures/lecture-16/ (MIT, start at 29 minute mark) (UV-Irvine, lecture 6 and the first 30 minutes of lecture 7) (Bond enthalpy, Annette Tomory) (UC-Irvine) Read Sections ,