Chapter 12 “Heat in Chemical Reactions”

1. Chapter 12 “Heat in Chemical Reactions”

2. Learning Target: Be able to calculate the amount of energy absorbed or released during a chemical or physical change. Learning Outcomes: • Know how to apply the 1st Law of Thermodynamics to chemical reactions. • Know that chemical potential energy is transferred from warm objects to cool objects in the form of heat.

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! Does the reaction in the picture release energy or absorb energy?

4. 2-1 Chemical Reactions That Involve Heat • 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)

5. 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.

6. 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 measured in Joules (J)! 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!

7. 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.

8. Learning Target: Be able to calculate the amount of energy absorbed or released during a chemical or physical change. Learning Outcomes: • Know how to apply the 1st Law of Thermodynamics to chemical reactions. • Know that chemical potential energy is transferred from warm objects to cool objects in the form of heat.

9. 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.

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

11. 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).

12. 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.

13. 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.

14. What is the enthalpy change for the following reaction? Is the reaction exothermic or endothermic? Mg + 2H2O →Mg(OH)2 + H2

15. Exothermic Reaction Diagram • CH4(g) + 2O2(g) →CO2(g) + 2H2O(g) + Heat CH4(g) + 2O2(g) ΔH = -803kJ (EXOTHERMIC; HEAT RELEASED FROM THE ‘SYSTEM’ TO THE SURROUNDINGS) ENTHALPY CO2(g) + 2H2O(g) NOTE: Hproducts IS LOWER THAN Hreactants REACTION PROGRESS (TIME)

16. 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)

17. “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°).

18. 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: • .

19. Learning Target: Be able to use thermochemical equations to determine enthalpy change (ΔH) associated with a given amount of reactant. Learning Outcomes: • Know how to calculate ΔH, and produce an enthalpy diagram. • Know how to calculate enthalpy for a given amount of reactant.

20. Warm Up • Calculate the enthalpy change (∆H) for the following reactions (You must balance the equation): a. NaCl(s)→ Na+(aq) + Cl- (aq) b. C2H5OH(l) + O2 (g)→ CO2(g) + H2O(g)

21. How much energy would be released if 50.0g of water were reacted? NH3 + H2O → NO2 + H2 ∆H° = -142.5kJ

22. Learning Target: Be able to determine the amount of standard enthalpy change (ΔH°) for a series of reactions by applying Hess’ Law. Learning Outcomes: Understand how to calculate ΔH° for a series of reactions.

23. How much energy is released if 10.0g of ethanol, C2H5OH , were burned? C2H5OH + O2 → CO2 + H2O ∆H° = -1234.7 kJ

24. 12-3 Hess’s Law 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.

25. Hess’ Law Example • Example N2 (g) + O2 (g) 2NO(g)ΔH1 = 181 kJ 2NO (g) + O2 (g) 2NO2(g) ΔH2 = -113 kJ Net equation: N2 (g) + 2 O2 (g) + 2NO (g) 2NO(g) + 2NO2(g) To write the net equation, the sum of the reactants is placed on the left hand side, and the sum of the products is placed on the right side. Substances on both sides cancel out – like algebra. In example above the next reaction is: N2 (g) + 2 O2 (g) 2NO2(g) ΔH = ΔH1 + ΔH2 = 181kJ + -113 kJ = 68kJ (endo)

26. Applying Hess’s Law • If the coefficients of an equation are multiplied by a factor, the enthalpy change the reaction is multiplied by the same factor. • If the equation is reversed, the sign of the ΔH changes also. • For example C (s) + H2O(g) CO(g) + H2(s) ΔH = 113kJ Then CO(g) + H2(s) C (s) + H2O(g) ΔH = -113kJ

27. Hess’s Law Example • Calculate the ∆H 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 • 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.

28. Calculate the ∆H for the reaction: 2S(s) + 2OF2(g) → SO2(g) + SF4(g) OF2(g) + H2O(l) → O2(g)+ 2HF(g) ∆H = -277 kJ SF4(g) + 2H2O(l) →SO2(g) + 4HF(g) ∆H = -828kJ S(s) + O2(g) →SO2(g) ∆H = -297kJ

29. 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). 12-4 Calorimetry

30. Calorimeter Experiment

31. 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.) 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. Heat equation: q = m X C X (Tf – Ti) HEAT vs. TEMPERATURE

32. A 15.75-g piece of iron absorbs 1086.75 joules of heat energy, and its temperature changes from 25°C to 175°C. Calculate the specific heat capacity of iron. • How many joules of heat are needed to raise the temperature of 10.0 g of aluminum from 22°C to 55°C, if the specific heat of aluminum is 0.90 J/g°C?

33. 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) Calorimeter Experiment

34. Calorimeter Experiment

35. 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!)

36. 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. Benjamin Thompson (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.’ James Joule 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.