Chemistry & Energy

1. Chemistry & Energy Unit 4

2. Excerpt from The Web of Life by John H. Storer • “Air, rock, water, and sunlight – these are the four sources from which come all living things and their environment. On the bare sands of the desert, the sun’s rays strike in tiny units of energy moving with atomic speed. Some of them we can feel as heat or see as light. These speeding units impart some of their energy to the dead sands, which temporarily store it in the form of heat, but when the sun sinks, this newly acquired energy is radiated back into the air and lost. The sand becomes as cold and dead as ever. But chlorophyll in the leaves of green plants exists as an agent for garnering these units of solar energy. It makes of the green leaf a laboratory in which nature creates food for living creatures and carries on unceasingly the magic of building life. • Like the sand, a field of grass absorbs the sun’s rays; but when night comes the grass does not give back this newly gained energy. In its green laboratory, the chlorophyll blends the sun’s captured radiance together with elements taken from the air, the water, and the soil, and builds these dead materials into organized living form to make new blades of grass. • This grass is cool and quiet, giving no hint of the sunlight stored within its framework. But dry it out and touch a match to it. The blades of grass – these tiny bits of organized gas and sunlight – blaze up with flame hot enough to kill a man. All of that fierce heat is merely a release of the same energy that the cool, moist plants have been quietly gathering from the sunlight and storing for later use. • If the grass is not burned, the energy will remain stored within its substance. If it is eaten by an animal, its life force is transferred with it into the body of the animal to sustain the spark that we call life.”

3. The Concept of Energy: • All physical and chemical changes are accompanied by changes in energy. •  Energy is defined as the ability to do work or to transfer heat. • Forms of Energy: • Energy can be classified as potential or kinetic.  • The total energy an object possesses Etot = K.E. + P.E.

4. POTENTIAL ENERGY • Potential energy is the energy that results from an object’s position. It is stored energy that can be converted to kinetic energy. • EXAMPLES OF POTENTIAL ENERGY: • Chemical potential energy: resulting from attractions among electrons and atomic nuclei in molecules. This is the energy associated with the forces within (Intramolecular) and between (Intermolecular) molecules. Chemicals store energy in bonds (Intramolecular forces). • The stronger the bond = more energy stored • Gravitational potential energy: such as that of a ball held well above the floor or that of water at the top of a waterfall.

5. KINETIC ENERGY • An object has kinetic energy because it is moving. Kinetic energy can be converted into potential energy. • EXAMPLES OF KINETIC ENERGY • Thermal kinetic energy: atoms, molecules, or ions (atoms with a net charge) in motion at the submicroscopic level—All matter has thermal kinetic energy because, according to the kinetic-molecular theory, the submicroscopic particles of matter are in constant motion. • Mechanical kinetic energy of macroscopic objects: moving baseball or automobile • Electrical kinetic energy: electrons moving through a conductor • Radiant kinetic energy: electrons transitioning between energy levels with the atoms producing light • Sound kinetic energy: corresponds to the compression and expansion of the spaces between molecules.

6. Let’s try some….. • When gasoline is burned in a car engine, its chemical potential energy is converted into other forms of energy. • Name three. • What kind of energy conversion takes place in a toaster? • What kind of energy conversion takes place in a hair dryer?

7. UNITS OF ENERGY • The SI unit for energy is Joule, J. This unit is named in honor of James P. Joule, a British scientist who devoted much of his life to investigating energy in all of its forms. He determined that work (or energy) could be measured in terms of the heat it could produce. • Another energy unit specifically used when talking about heat energy is the calorie (cal). 1 calorie = 4.184 J • One calorie represents a small quantity of heat, thus the kilocalorie (kcal) is preferred. A dietary calorie, represented Calorie (Cal), has the same value as kcal. 1000 cal = 1 kcal = 1 Cal

8. Definition of Energy Units

11. Conversion of Energy and its Conservation: • Law of Conservation of energy states that, “Energy cannot be created nor destroyed, just transferred.” • The total amount or energy in the universe is CONSTANT. • Is mass and energy related? E = mc2 E = Energy m = mass c = speed of light (2.998 x 108 m/s) • This equation corresponds to when a body or system releases energy, the body or system decreases in mass. Likewise, when a body or system absorbs energy, the body or system increases in mass. • Usually, these changes in mass go undetected, unless we discuss reactions that involve nuclei of atoms that occur in nuclear reactors or an atomic bomb. YES!!!

12. Energy and Chemical Reactions: • Energy changes that accompany chemical changes are generally more noticeable than those that take place during physical changes. • The substances that we begin with before the chemical reaction are referred to as the reactants. • The substances that are produced as a result of a chemical change are referred to as the products. • If a reaction takes place and the temperature of the surroundings drops, does this indicate the destruction of energy? • NO, this indicates that energy was absorbed from the surroundings into the system (reaction) and can is classified as an endothermic reaction. • What if the temperature of the surroundings increases? Was energy created? • NO, This means the system (reaction) released energy to the surroundings. This is an exothermic reaction.

13. How can this be explained????? • When two atoms chemically bond together, their bond has a specific amount of energy due to its stability (strength). The energy required to break this bond varies depending on the type of bond and the atomsthat are bonded together. • During a chemical change, the reactants begin with a certain amount of potential energy. In order to break these bonds a specific amount of energy, called the activation energy, is needed. When the products are yielded they may have a higher or lower potential energy than the reactants due, once again, to the number of bonds formed and the relative strengths of the bonds. • If the reactants have more energy than the products, then heat must be released from the reaction……. EXOTHERMIC. • The products are more stable than the reactants. • If the products require more energy than the reactants, then heat must be absorbed from the surroundings……… ENDOTHERMIC. • The reactants are more stable than the products. • The heat of reaction is the measure of the energy flow; quantity of heat released of absorbed during a reaction and is represented by q or H. (q = H at constant pressure)

14. Potential Energy Diagrams

15. POTENTIAL ENERGY DIAGRAMS

16. Heat Potential energy

17. N2 + O2 + HEAT 2 NO 2 NO Potential energy Heat N2 + O2

18. Direction • Every energy measurement has three parts. • A number (how many). • A unit (Joules or calories). • A sign to tell direction. • negative - exothermic • positive- endothermic 593.2 - 593.2 J 593.2 J

19. + - • The reaction: 2CO(g) + O2(g)  2CO2(g) is exothermic. • What is the sign of H for this reaction? • What is the sign of H for the reverse reaction? • Is the heat content of the products greater or less than that of the reactants? • For the reaction: CaCO3(s) CaO(s) + CO2(g) H = +176 KJ/mol • Is this reaction exothermic or endothermic? • b. What is the value of H for the reverse reaction? • c. How does the heat content of the products compare to that of the reactants? endothermic Less than Greater than -176 kJ/mol

20. Heat Energy and Temperature: • We have all measured temperature before using a thermometer, but how does it work? • When substances heat up, typically they expand, which explains the concept behind a thermometer. A glass (or plastic) tube is attached to a bulb. The bulb is filled with alcohol or liquid mercury. Upon heating the liquid expands and is forced from the bulb into the tube. Upon cooling, the liquid contracts and flows back into the tube. The tube is then calibrated to precisely record changes in temperature.

21. temperature scales • The three temperature scales are • Fahrenheit, F • Celsius, C • Kelvin, K. • The Celsius temperature scale is the one most commonly used for scientific work. Comparison of the Three Temperature Scales F C K Water Boils 212.00 100.00 373.15 Body Temperature 98.6 37.0 310.2 Water Freezes 32.0 0.00 273.15 Absolute Zero -459.67 -273.15 0.00

22. FAHRENHEIT: • The Fahrenheit scale was created (1714, Gabriel Fahrenheit) based on temperatures he could recreate in the lab. • The lowest temperature he could achieve was with a mixture of freezing ice, snow salt, and water. This became 0F. • The reference point for 100F is told to be based on the body temperature of a cow, around 98.6F, simply because Fahrenheit liked cows!! Except he chose a sick cow when he created the scale.

23. CELSIUS • The Celsius scale was created in 1742 by Anders Celsius. He designated the freezing and boiling points of water to be the 0C and 100C, respectively, and then divided the difference into 100graduations. This is why the Celsius scale is sometimes referred to as “Centigrade.”

24. KELVIN • The Kelvin temperature scale (1848) was derived from the observation of gases. • Theoretically, the lowest possible temperature an object can obtain is that of absolute zero. This temperature has never been reached. Scientists have come to within 1 billionth of a degree above absolute zero. • This value is the basis for Lord Kelvin’s scale. At absolute zero, all motion will cease.

25. TEMPERATURE CONVERSIONS • The size of one degree is the same for both the Celsius and Kelvin scales, however not for Fahrenheit. To convert among these three scales: Convert the following temperatures from one unit to another. • 263 K to oF • 38 K to oF • 13 oF to oC • 1390 oC to K • 3000. oC to oF F = 1.8 (K – 273.15) + 32 F = -391 F F = 14 F

26. How to Derive Equations 0º C = 32 ºF Compare Freezing Points 0ºC 32ºF

27. 100 ºC = 212 ºF Compare Boiling Points 0ºC 100ºC 212ºF 32ºF

28. 100ºC = 212ºF 0ºC = 32ºF 100ºC = 180ºF Find the distance between the two points on both scales 0ºC 100ºC 212ºF 32ºF

29. 100ºC = 212ºF 0ºC = 32ºF 100ºC = 180ºF 1ºC = (180/100)ºF 1ºC = 1.8ºF A 1º change in Celsius corresponds to 1.8ºF 0ºC 100ºC 212ºF 32ºF

30. A 1º change in Celsius corresponds to 1.8ºF Lets set-up ordered pairs to determine the equation of the line. Whatever data set has a zero point, assign that scale as the x coordinate (0,32) (100, 212) ºF = 1.8(ºC) + b y = mx + b

31. Find the y-intercept (b) by substitution!!! ºF = 1.8(ºC) + b 212 = 1.8(100) + b 32 = 1.8(0) + b 212 - 180 = b 32 = b 32 = b ºF = 1.8(ºC) + 32

32. Water boils at 140 °X and freezes at 14 °X. Derive the relationship between °X and °C. 1ºC = (126/100)ºX A 1º change in Celsius corresponds to 1.26ºX 1ºC = 1.26ºX Lets set-up ordered pairs to determine the equation of the line. (0,14) (100, 140)

33. y = mx + b ºX = 1.26(ºC) + b Find the y-intercept (b) by substitution!!! 140 = 1.26(100) + b 14 = 1.26(0) + b 14 = b 140 - 126= b 14 = b ºX = 1.26(ºC) + 14

34. Heat and Its Measurement: • We can now convert among the three temperature scales, but what exactly is temperature? • Temperature: the average kinetic energy of a substance’s particles. • Are temperature and heat the same thing? NO!!! • Heat: form of energy that is often associated with the changing temperature of an object. An objects temperature increases because energy is transferred to it. The energy transfer, in the form of heat, happens when two objects are brought into contact, as nature attempts to equalize their temperatures, that is, to establish a thermal equilibrium.

35. A lit match burns at a temperature of 200C. • A building on fire burns at a temperature of 200C. • Which one is hotter? • Neither, they are the same temperature. • Which one produces more heat? • The burning building. • A hot cup of coffee is 85C. • A Jacuzzi is 85C. • A hot cup of coffee contains enough heat energy to probably melt one ice cube before coming to room temperature. • A Jacuzzi would melt an entire freezer full of ice before coming to room temperature. Heat is affected by both temperature and mass.

38. The Kinetic Theory of Heat and Temperature: • Temperature can be used to determine the direction of flow of energy. Energy spontaneously flows from a warmer object to a cooler one. (HEAT DOES NOT RISE). • Kinetic theory explains the flow of energy in terms of particle collisions. Because the warm object is at a higher temperature than the cool object, the average kinetic energy of its particles is greater than that of the cool object. Because the objects are in contact, the particles of the warm object collide with those of the cool object. Gradually the particles in the warm object will transfer kinetic energy to the particles in the cool object through collisions. As a result, the particles in the cool object gain kinetic energy. Eventually the average kinetic energy of the particles in both objects will be equal and the two objects will have the same temperature. • For example: You can feel that a bowl of hot soup warms the air around it. Left undisturbed, the soup and bowl will eventually cool to room temperature.

39. Measuring Changes in Heat : • In the early 1780’s Antoine Lavoisier measured the heat given off by a guinea pig. He kept the animal enclosed in a container so that its body heat would melt ice. From the amount of ice melted, he calculated the body heat produced by the animal. This was an example of an ice calorimeter. • A calorimeter is a device used to measure the energy given off or absorbed during chemical or physical changes. • There are two main types of calorimeters, a constant pressure calorimeter (coffee-cup) and a constant volume calorimeter (bomb calorimeter).

40. In a coffee cup calorimeter, the reaction takes place in the water. • A coffee cup calorimeter is great for measuring heat flow in a solution, but it can't be used for reactions which involve gases, since they would escape from the cup. • The coffee cup calorimeter can't be used for high temperature reactions, either, since these would melt the cup. A bomb calorimeter works in the same manner as a coffee cup calorimeter, with one big difference. • A bomb calorimeter is used to measure heat flows for gases and high temperature reactions. • In a bomb calorimeter, the reaction takes place in a sealed metal container, which is placed in the water in an insulated container. Heat flow from the reaction crosses the walls of the sealed container to the water. • The temperature difference of the water is measured, just as it was for a coffee cup calorimeter. Analysis of the heat flow is a bit more complex than it was for the coffee cup calorimeter because the heat flow into the metal parts of the calorimeter must be taken into account:

41. Bomb Calorimeter

42. CALORIMETRY • We will use the coffee-cup calorimeter in the lab. A bomb calorimeter is the preferred calorimeter to measure the heat released in combustion reactions. Both types of calorimeters contain water, which is used to measure the heat absorbed or discharged. • Some substances require little heat to cause a change in their temperature. Other substances require a great deal of heat to cause a change in their temperature. • For example: One gram of liquid water requires 4.184 joules of heat to cause a temperature change of 1C. It only takes 0.902 joules to raise the temperature of one gram of aluminum 1C. • The heat needed to raise the temperature of one gram of a substance by 1C is called the specific heat (Cp) of a substance. • Every substance has its own specific heat (intensive physical property). • The heat required to raise the temperature of one gram of water 1C is 4.184 J or 1 cal.

43. The Specific Heat of water (Cp) is 4.184 J/gC or 1 cal/gC. • Specific heats can be used in calculations involving the change in temperature of a specific mass of a substance. Further, the amount of energy transferred can be calculated from the relationship. q = mCT • q = heat gained or lost by substance • m = mass in grams • C = specific heat • T = change in temperature (T2 – T1) or (Tfinal – Tinitial) • According to the Law of Conservation of Energy, in a calorimeter containing water, any heat gained by the substance is lost by the water. Likewise, any heat lost by the substance is gained by the water. This transfer of energy takes place between these two quantities of matter that are at different temperatures until the two reach the same (equilibrium)temperature.

44. The heat released or absorbed by water is calculated using the following equations: q = mCT or Note: Any heat lost by the substance (-q) equals the heat gained by the water (+q) and vice versa. • Notice that throughout our discussion of calorimetry, we have assumed that the calorimeter itself does not absorb any heat. We have also assumed that no heat escapes from the calorimeter. Neither of these assumptions is completely true. However, we will continue to make these assumptions in order to simplify our calculations. The error from actual losses of heat should be considered in any laboratory exercise using calorimeters.

45. EXAMPLE: • Phosphorus trichloride, PCl3, is a compound used in the manufacture of pesticides and gasoline additives. How much heat is required to raise the temperature of 96.7 g of PCl3 from 31.7°C to 69.2°C. The specific heat of PCl3 is 0.874 J/g°C. q = mCT

46. Examples • The specific heat of graphite is 0.71 J/g°C. Calculate the energy needed to raise the temperature of 75 kg of graphite from 294 K to 348 K. q = mCT Because the size of a Celsius degree is equal to the size of a Kelvin, the difference in each scale would be equivalent

47. EXAMPLE: • An Oreo® is burned in a bomb calorimeter containing 5.0 L of water originally at 25.0 °C. After igniting the Oreo®, the water temperature increases to 30.6°C. How many Calories of energy did the Oreo® contain? q = mCT What is Given? Volume of water = 5.0 L T1 = 25.o °C What happened to the temperature of the water? How does the heat gained by the water compare to the heat lost by the Oreo? Why? T2 = 30.6 °C Law of Conservation of Energy!!!

48. -qOreo = qH2O qH2O = mH2OCH2OTH2O Volume of water = 5.0 L If we calculate the heat gained by the water, this necessarily equals the heat lost by the cookie T1 = 25.o °C How can we find the mass of the water? T2 = 30.6 °C TH2O = 5.6 °C TH2O = Density = mass / volume CH2O= CH2O= 1 cal/g °C

49. mass = Density * volume Putting it all together…. qH2O = mH2OCH2OTH2O

50. EXAMPLE: • A reaction changes the temperature of a calorimeter with 500. mL of water from 10.°C to 50.°C. What is the ΔH for the reaction? q= mCT