Chapter 12 “Heat in Chemical Reactions”

1. Chapter 12 “Heat in Chemical Reactions” Where did all of that energy come from? T. Witherup 2006

2. OBJECTIVES • Distinguish between “exothermic” and “endothermic” chemical reactions. • Explain what is meant by enthalpy and enthalpychange. • Define standardenthalpychange and explain its use. • State and apply “Hess’s Law” to determine ΔH for a reaction. • Describe calorimetry and how it is used to determine heats of reactions. • Compare “heat” and “temperature.”

3. 2-1 Chemical Reactions That Involve Heat • What do you visualize when thinking about chemical reactions? You probably imagine… • Heat, Light, Sound • These are energy, and chemical reactions involve energy because chemical bonds are being broken and made! • Energy: the capacity to do work. • What kinds of energy are there? • Radiant (solar) • Thermal (motion of atoms & molecules) • Chemical (stored within chemical bonds) • Potential (available because of an object’s position) • Kinetic (energy of motion)

4. 12-1 Thermochemistry & Thermodynamics • Heat is the transfer of thermal energy between two objects at different temperature. • Thermochemistry is the study of heat changes that occur in chemical reactions. • ‘Thermes’ is Greek for “heat.” • Thermodynamics is the broader study of energy and work changes during such processes.

5. 12-1 Exothermic vs. Endothermic Reactions • Exothermic reactions release heat energy. • ‘Exo’ means outside. • Example: Burning natural gas (methane, CH4) to keep warm in the winter. • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + Heat • Notice that “heat” is a product! • Endothermic reactions absorb heat energy. • ‘Endo’ means inside. • Example: Decomposing HgO into mercury and oxygen. (Lavoisier used this reaction to discover oxygen.) • 2HgO(s) + Heat  2Hg(l) + O2 (g) • Now note that “heat” is a reactant!

6. 12-1 Exothermic vs. Endothermic Reactions (Examples) • U.S. Army “MREs” (“Meals ready to eat”) • Magnesium reacts with water to release heat and hydrogen gas. Write the reaction. • Mg + 2H2O  Mg(OH)2 + H2 + Heat • MREs heat the food to 60° C in 15 minutes! • EXOTHERMIC Reaction • Ammonium Nitrate (A common fertilizer) • NH4NO3 dissolves in water, causing a drop in temperature. Write the reaction. • NH4NO3 + H2O + Heat  NH41+(aq) + NO31-(aq) • NH4NO3 is used in cold packs for sports injuries. • ENDOTHERMIC Reaction • Note: NH4NO3 can also undergo an EXOTHERMIC reaction when it explodes, hence its use as a blasting agent in mining and (unfortunately) in terrorists’ weapons, such as that used in the Oklahoma City attack.

7. 12-2 Heat and Enthalpy Changes • Remember the “Law of Conservation of Energy”? • It is also the First Law of Thermodynamics: • “Energy cannot be created or destroyed, but may be converted from one form to another.” • To study thermochemistry we must be careful about the ‘system’ involved. • Open System: allows exchange of both energy and mass. • Closed System: allows the exchange of energy, but not mass. • Isolated System: allows the exchange of neither energy nor mass. • See diagrams on next slide.

8. Types of “Systems” OPEN Mass Loss or Gain CLOSED ISOLATED Energy Loss or Gain

9. A New Term: Enthalpy • Most chemical processes we use occur at normal, constant atmospheric pressure, and heat energy is either gained or lost during the chemical reaction. • Chemists use a special term, enthalpy, to describe heat changes for chemical reactions at constant pressure. • Enthalpy is the energy of a substance plus a small adjustment for pressure-volume work. • Enthalpy (H) = E + PV = Internal Energy + Pressure-Volume (Work) = Heat absorbed or released by a reaction • “Enthalpy” is from German (enthalpein, to warm).

10. Enthalpy Change (ΔH) • Enthalpy change, ΔH, for a chemical reaction that occurs at constant pressure is the heat released or absorbed in the reaction. • For the reaction Reactants  Products ΔH is just the difference between the enthalpy of the products minus the enthalpy of the reactants: ΔH = Hproducts – Hreactants (Here the Δ means “change” or “a difference.”) Think of enthalpy in terms of a savings account balance. You are just keeping track of money going into the account and the money spent vs. the net balance.

11. Some Conventions (Rules) • We use diagrams to show enthalpy changes. • For exothermic reactions, Hproducts is lower than Hreactants so ΔH is negative. • Heat is released (lostby the system). • For endothermic reactions, Hproducts is higher than Hreactants so ΔH is positive. • Heat is absorbed (gainedby the system). • See diagrams on the next slide.

12. Exothermic Reaction Diagram • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + Heat CH4(g) + 2O2(g) ΔH = -890kJ (EXOTHERMIC; HEAT RELEASED FROM THE ‘SYSTEM’ TO THE SURROUNDINGS) ENTHALPY CO2(g) + 2H2O(g) NOTE: Hproducts IS LOWER THAN Hreactants REACTION PROGRESS (TIME)

13. Endothermic Reaction Diagram • 2HgO(s) + Heat  2Hg(l) + O2 (g) 2Hg(l) + O2(g) ΔH = +181kJ (ENDOTHERMIC; HEAT ABSORBED BY THE ‘SYSTEM’ FROM THE SURROUNDINGS) ENTHALPY 2HgO(s) NOTE: Hproducts IS HIGHER THAN Hreactants REACTION PROGRESS (TIME)

14. “Standard States” • To be exact when comparing enthalpy changes chemists must define the conditions, such as pressure, temperature and phase. • The “Standard State” of a substance is its pure form at the standard pressure of 1 atmosphere and 25° C. • Standard state of oxygen is as the diatomic gas. • Standard state of carbon is graphite. • Standard state of phosphorus is P4 (white phosphorus). • Enthalpy changes measured under “standard state” conditions are called “Standard Enthalpy Changes” (ΔH°).

15. Using Enthalpy Changes How much heat will be released when 6.440g of sulfur reacts with excess oxygen by the reaction: 2S (s) + 3O2 (g)  2SO3 (g) ΔH° = -791.4kJ Solution: 1 mol S -791.4kJ 6.440g S X X = -79.46 kJ 32.07g S 2 mol S • See sample problems (p. 386) & 12-2 worksheets.

16. 12-3 Hess’s Law • Consider the energy changes that occur as two groups of adventurers climb a mountain using different trails. • Group 1 makes the difficult climb straight up the face of the mountain. • Group 2 takes the gently sloping trail that winds slowly up the mountain. • What is their relative Potential Energy (PE) at the base camp and at the summit? • Hess’s Law: If a series of chemical reactions are added together, the enthalpy change for the netreaction will be the sum of the enthalpy changes for the individual steps.

17. Hess’s Law Example • Calculate the standard enthalpy change for the reaction 2C(s) + O2(g)  2CO(g) from • C(s) + O2(g)  CO2(g) ΔH = -393.5 kJ • 2CO(g) + O2(g)  2CO2(g) ΔH = -566.0 kJ • (Ans.: -221.0 kJ) • See also sample problems (p390-391) & 12-3PP sheet. • The trick is to simply combine equations so that only the relevant parts are present (everything else cancels out). Remember, whatever you do to the equations, you must do to the value of ΔH for the reaction.

18. 12-4 Calorimetry • What happens during chemical reactions? • During an exothermic reaction, the surroundings gain heat from the ‘system.’ • During an endothermic reaction, the surroundings lose heat to the ‘system.’ • Calorimetry is the study of heat flow and heat measurement during chemical reactions. • Calorimetry experiments use accurate measurements of temperature changes to determine the enthalpy changes (heat flow).

19. Heat vs. Temperature • Heat Capacity: the amount of heat needed to raise the temperature of an object by 1 degree Celsius (1° C). • Specific Heat: the heat capacity of one gram of a substance. • Specific Heat of water = 4.184 J/g.C° • To raise the temperature of one gram of liquid water by 1 degree C requires 4.184 J of heat energy, or 1 calorie (cal.) • Recall: 1,000 calories = 1 kilocalorie (1 kcal) = 1 Calorie (the ‘food’ unit) • Note that heat and temperature are related, but different, concepts. • We detect heat changes by measuring temperature changes. • A small temperature change does not mean a small amount of heat transfers! (Consider melting a lake!) • Use specific heat and mass to measure heat changes.

20. A Calorimetry Experiment • Calorimeter: a well insulated container filled with a known mass of water, a way to conduct a chemical reaction, a stirrer and a thermometer. • Measure the initial temperature of the water (Ti). • Conduct the chemical reaction. • Measure the temperature periodically. • Determine the final temperature (Tf). • Note that all heat released by a reaction is gained by the water in the calorimeter, and vice versa! • Mathematically we may say: (qrxn = -qsur) • Calculate the quantity of heat absorbed by the water (qsur) from the mass of water, the Specific Heat of water and the temperature change (Tf – Ti). (And viceversa.) qsur = m X C X (Tf – Ti) • The sign of qsur indicates the direction of heat flow for the reaction. • Exothermic reaction, heat goes to the water:(-) sign. • Endothermic reaction, heat goes from the water:(+) sign.

21. Calorimetry Example • When 1.008g of hydrogen gas reacts with chlorine gas in a calorimeter containing 500.00g of water, the temperature rises from 25.00°C to 29.24°C. Calculate ΔH for the process H2(g) + Cl2(g)  2HCl(g) ΔH = ? • Use qsur = m X C X (Tf – Ti) • qsur = (500.00g) X (4.184J/g.C°) X (29.24 - 25.00 C°) = 8870J • qsur = 8870J = -qrxn (An exothermic reaction.) • Therefore qrxn = -8870J. • 1.008g H2 X (1 mol H2/2.00g H2) = 0.504 mol H2, which released this -8870J of energy. • So for the overall reaction: (-8870J/0.502 mol H2) = ( -17,700J/1mol H2) Or -17.7kJ/mol H2

22. Fuel for Our Body • Our body is like a chemical factory, taking in raw materials of various kinds to make new products that keep us alive and help us grow. • These reactions involve bond-breaking and bond-making, and they meet the energy needs of our body. • Carbohydrates and fats are major food sources, and they release lots of energy (exothermic reactions). • Glucose releases -2803kJ/mol C6H12O6 • See Table on page 399. • “Calories Count” Project • Recall 1000 cal = 1 Kcal = 1 Cal (the food kind!)

23. 12-5 What is Heat? • “Caloric Theory” • Heat was thought of as an invisible, weightless fluid capable of flowing from a hot object to a colder one. • BenjaminThompson (Count Rumford) showed that friction heat between objects was continuous as long as two objects were rubbed together. • This was not consistent with ‘Caloric Theory.’ • JamesJoule used a paddlewheel experiment to convert mechanical energy into heat, thus showing there is a ‘mechanical equivalent of heat.’ • Modern “Kinetic Theory” • Heat is defined as the transfer of kinetic energy from a hotter object to a colder one.

24. OBJECTIVES • Distinguish between “exothermic” and “endothermic” chemical reactions. • Explain what is meant by enthalpy and enthalpychange. • Define standardenthalpychange and explain its use. • State and apply “Hess’s Law” to determine ΔH for a reaction. • Describe calorimetry and how it is used to determine heats of reactions. • Compare “heat” and “temperature.”