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Module 3

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  1. Module 3 Heat and Energy in Chemical Reactions.

  2. Introduction: Thermochemistry Chemical reactions often involve changes in temperature, and in the heat energy that causes temperature to change. This branch of chemistry that studies heat changes is called THERMOCHEMISTRY.

  3. Some Quick Definitions • Energy: The ability to (a) do work or (b) to supply heat. • Kinetic Energy: energy associated with movement of objects or particles. • Potential Energy: energy that has not produced heat or motion yet. Potential energy may be associated with gravitation and with chemical bonds.

  4. Background Information(notes optional) Most energy on earth comes from the sun, but a small amount comes from the radioactive decay of certain atoms, and an even smaller amount from the tidal effects of the moon. Most of the energy that we experience is thermal energy, that is, one of several forms of energy related to heat

  5. (Background Info: Notes optional) • Thermal energy can be transferred in many ways: • Conduction: heat transferred between touching objects. This form of heat transfer is the most important in chemistry. • Convection: heat transferred to air molecules, which then may expand and may rise due to their reduced density. This is important in meteorology • Radiant heat: A hot object, like a heat lamp, may directly radiate heat in the form of infrared radiation. This is important in physics

  6. Heat and Temperature • Heatis a form of thermal energy that is transferred when two systems with different temperatures come into contact with each other. • Heat “flows” from one object to another. • Temperatureis a measure of the agitation of atoms or molecules in a system. • Temperature is directly related to the average kinetic energy of the molecules.

  7. Heat vs. Work When energy transfer is in the form of “work” the transfer is orderly. A whole object or group of molecules move in an orderly fashion in one direction. When energy transfer is in the form of “heat” the transfer is more random. The molecules of the heated object move more quickly in random directions. P. 128-130 Questions p. 130

  8. Law of Conservation of Energy • Energy can be neither created nor destroyed in any chemical process. • However: • Energy can be transferred from one object to another. • Eg. Collisions transfer energy between objects. • Energy can be transformed from one type of energy to another. • Eg. Potential energy can be transformed into kinetic energy, and vice versa. • Nuclear reactions can change matter into energy or vice versa

  9. Systems and Their Surroundings • System: the part of the universe on which we are focusing our attention (usually a beaker, flask or test tube) • Surroundings: Everything else in the universe. • A system may be: • Open: matter & energy easily exchanged between the system and its surroundings • Closed: matter cannot be exchanged, but some energy can escape or enter the system • Isolated: neither matter nor energy can be exchanged between system and surroundings.

  10. Calorimeters • A calorimeter is an instrument used to measure the amount of heat released or absorbed during a change. • A calorimeter is a closed system containing a fixed amount of water that can absorb heat • The ΔH can be calculated by measuring the water temperature before and after the reaction • Calorimeters can be as complicated as the “bomb” calorimeter illustrated on p. 132 or as simple as a Styrofoam cup

  11. The Calorimeter Formula Remember This? Q=mcΔT Where Q=Heat Energy m = mass of the water c = specific heat capacity of the water ΔT = the change in temperature

  12. As You Remember! • Specific Heat capacity (c) is the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius. • For water, c is 4.19 J / g°C

  13. Simple examples: copy and complete* +25°C 26187.5 J -2°C -1676 J 18°C 7542 J -6704 J 9°C 95.5g 52°C *Assume measurements are precise, c = 4.19J/g° for water and that no heat is lost in the calorimeter.

  14. Molar heat of Reaction • The molar heat of a reaction is the amount of heat released / absorbed by the reaction of one mole of the material • ΔH (Molar) = Q / n • A.K.A. molar heat of combustion, molar heat of solution, etc.

  15. Example: copy and try this • The following data was recorded for the partial combustion of wax (C25H52) in a calorimeter. Use this data to calculate the molar heat of combustion of wax. • Initial mass of wax 22.35 g • Final mass of wax 12.08 g • Volume of water in calorimeter 352.5 mL • Initial temperature of water 12.6 °C • Final temperature of water 43.5 °C

  16. Solution: part 1- heat captured • Mass of water: m=352.5 g (water) • Change of temp: ΔT= 30.9 °C • Specific heat: c= 4.19 J/g °C • Q=mc ΔT = 352.5×4.19 × 30.9 • Q = 45638.5275 Joules • But… that is not for a whole mole.

  17. Solution: part 2 – number of moles • Mass (wax) m=10.27 g • Molar mass (wax) M=352.61 g/mol • Moles of wax: n= m ÷M • n= 10.27 ÷ 352.61 • n= 0.0291 mol • The molar heat of combustion is the heat produced per mole of wax

  18. Solution: part 3 - molar heat • Q ÷n = 45638.5275 ÷ 0.0291 • Solution= 1568334.278 • Round to 3 significant digits ANSWER: • The molar heat of combustion of wax is about 1570000 Joules/mole • or 1.57x106 J/mol • or 1570 kJ/mol

  19. Heat Transfer • When two bodies (or substances) of different temperatures come into contact, heat is transferred from the warmer body (or substance) to the cooler body, until equilibrium is reached, when the temperature is the same. • At equilibrium, the heat lost (Q-loss) by the substance that was warmer, must equal the heat gained (Q-gain) by the substance which was cooler.

  20. Heat Transfer Heat moves from Hot to Cold -Q1 = Q2 = Q loss Q gain Hot Body Equilibrium Cool Body Until they are the same temperature

  21. If Q=mcΔT and -Q1 = Q2 Then -m1c1ΔT1 = m2c2ΔT2

  22. Example: Copy and try this! • A calorimeter contains 230.0g of water at 25.0°C. A 200.0g sample of copper at 47.0 °C is placed inside the calorimeter. Calculate the final temperature of this system. (the specific heat capacity of copper is 0.390 J/g °C)

  23. Solution • Q=mcΔT • For water: Qgain=230 × 4.19 ×ΔTwater • For copper: Qloss=200 × 0.39 ×ΔTCu • Since -Q1 = Q2 this means that: • -(200 × 0.39 ×ΔTCu)= 230x4.19x ΔTwater • Let x represent the final temperature • So ΔTwater = (x–25) °C • And ΔTCu = (x– 47)°C • We can substitute:

  24. m(copper) c (copper) ΔT (copper) m (water) c (water) ΔT (water) • -200 g× 0.39 j/g°C× (x-47)°C = 230 g× 4.19 j/g°C× (x-25) °C • -78 g•j/g°C(x - 47) °C = 963.7 g•j/g°C(x-25) • -78x + 3666 = 963.7x - 24092.5 • -963.7x - 78x = -3666 - 24092.5 • 1041.7 x = 27758.5 • 26.647 (round to 3 sig. dig.) • Answer: 26.6°C

  25. This is true when liquids are mixed too! • When liquids are mixed, the heat lost by the warm liquid must be the same as the heat gained by the cooler one. • In this case, not only do the substances transfer heat to each other, but they will mix as well.

  26. Example: copy and try • A mixture is made of 100.0 mL of water at 90.0°C and 100.0 mL of water at 25.0°C. What will the final temperature be? M(hot)=100g m(cold)=100g C(hot)=4.19 j/g c(cold)=4.19 j/g Ti(hot)=90 Ti(cold)=25 Tf=x Tf=x -ΔT(hot) =Tf –Ti(hot) ΔTcold =Tf –Ti(cold) Data: The same

  27. Solution: part 1 • Q=mcΔT so… Qgain=mcoldx ccoldx ΔTcold -Qloss = mhot x chot x ΔThot • For cold water, Qgain=100 × 4.19 ×ΔTcold • For hot water Qloss=100 × 4.19 ×ΔThot • Since –Q(lost) = Q(gained) this means that • -100 × 4.19 ×ΔThot = 100 × 4.19 ×ΔTcold • ΔTcold = (x-25)let x= final temp • -ΔThot= (x-90) • Note: since ΔTcold = ΔThotthere is a short cut you can use, we’ll look at that later.

  28. Solution: part 2 • -100 × 4.19 (x-90) =100 × 4.19 (x-25) • –(x– 90)=x -25 • x-25=-x+90 • x + x =90 + 25 • 2 x = 115 • x = 57.5 • Answer: the final temperature is 57.5 °C • Of course, in this case you could have done the problem more simply, since you knew you had equal amounts of the same substance (ie. same mass & specific heat capacity) , you could have just averaged the temperature! Warning: don't try that if the amounts or substances are Different!

  29. Homework Questions: • What will the final temperature be if you mix 100.0 g. of water at 20.0°C and 40.0 g. of water at 80.0 °C? • What is the molar heat of combustion of methane (CH4)if burning 1.00 gram of it in a calorimeter raises the temperature of 800.0g of water from 27.0 °C to 42.0 °C? • What will the final temperature be if 50.0g of copper (c=0.39 J/g °) at 80.0 °C is dropped into 200.0 mL of 20.0°C water?

  30. Review Concepts • There are three types of Change • Physical change: does not change the composition for example a change of state • Chemical change: Changes the composition, examples: effervescence, decomposition, change of colour, precipitation, combustion. • Nuclear change: Changes in the atom nucleii, examples: formation of isotopes, radioactive decay, nuclear fission, nuclear fusion

  31. Remember:Kinetic Energy of particles: Reminder • Particles (ie. Molecules) can have 3 types of motion, giving them kinetic energy • Vibrational kinetic energy • Rotational kinetic energy • Translational kinetic energy

  32. State & Kinetic Energy Reminder • Solids exhibit only vibrational energy • Virtually none of their kinetic energy comes from translation or rotation. • Liquids exhibit mostly rotational energy • A small portion of their kinetic energy can come from vibration or translation. • Gases exhibit mostly translational energy • a tiny portion of their kinetic energy can come from vibration or rotation.

  33. Melting a Pure Solid andBoiling a Pure Liquid • As you slowly add thermal energy to a pure solid, its temperature will rise as the molecular vibration increases. • There will come a point, however, where heating the solid does not increase the temperature. Instead, the increased energy is used to change the type of motion, making the molecules tumble or rotate. This is called the melting point. • The solid becomes a liquid, and the temperature again rises with increased thermal energy, until the liquid begins to boil or evaporate.

  34. Phase Diagram Boiling Gas Translation Melting Liquid Mostly Rotation Rotation Translation Temperature (°C or K )  Vibration & Rotation Solid Vibrational Thermal Energy Added (joules or kilojoules) 

  35. Info from a phase diagram(diagram shown in endothermic direction, increasing absorption of heat) 100 Specific Heat capacity (c) of the solid is the inverse of the slope, ie: Run over Rise 30÷60 = 0.5 J/g° 80 Boiling Point (56°C) 60 40 Melting Point (18°C) 20 0 Temperature °C  -20 -40 Heat of Fusion (ΔHfus=10 J/g) Heat of vaporization (ΔH vap=15 J/g) 10 20 30 40 50 60 70 80 90 Energy Added (Joules / gram) 

  36. Info from a phase diagram(diagram shown in Exothermic direction, increasing release of heat) 100 80 Condensation Point (56°C) 60 40 Freezing Point(18°C) 20 Temperature °C 0 -20 -40 Heat of condensation (ΔHcond=15 J/g) Heat of solidification (ΔH=10 J/g) 10 20 30 40 50 60 70 80 90 Energy Released (Joules / gram) 

  37. Heat Diagrams for MixturesSo far we have only shown heat diagrams for pure substances. For mixtures the “plateaus” are less clear, and the melting and boiling points less clearly defined. gas Third fractional distillation point Second fractional distillation point foamy First fractional distillation point Sometimes the difference in boiling points for different substances in a liquid mixture can be used to separate them by fractional distillation. Melting Range Boiling Range liquid slushy solid

  38. Assignments • Textbook Reading pp. 125-136 • Textbook Questions pp. 145, #1-19

  39. Module 3, Lesson 2 Endothermic and Exothermic Reactions

  40. Exothermic Reactions • Exo=outer, Therm=heat • Exothermic reactions are chemical reactions which release heat energy. • You can recognize exothermic reactions because the products are hotter than the reactants. • Example: burning wood is an exothermic reaction.

  41. Endothermic Reactions • Endo = inner; Therm = heat • Endothermic reactions absorb heat from their surroundings. • You can recognize endothermic reactions because the products become colder than the reactants were. • Examples: instant “ice packs”,

  42. More About Energy • Potential Energy (EP) is stored energy. In chemistry it is usually stored as chemical bonds. • Kinetic Energy (EK) is energy of motion. In chemistry it usually revealed by temperature, caused by moving molecules. • Heat Energy (Q) is energy transferred from one body to another due to a difference in temperature between the bodies

  43. Types of Heat Change in Chemistry • Enthalpy (H): The amount of total energy in a substance, most of it is in the form of potential or “hidden” heat energy. (Detailed discussion will follow) • Heat of Reaction (ΔH): amount of energy absorbed or released during a reaction. It represents an amount that the enthalpy has changed. • In addition to chemical reactions, energy can be absorbed or released in other changes. The symbol ΔH can also be used for these.

  44. Variations on ΔH(enthalpy can change in many types of reaction) • Heat of Formation (ΔHF): the amount of energy absorbed/released when compound is made from its elements. • Heat of Dissolution (ΔHd): The amount of energy absorbed/released when a solute dissolves. • Heat of Neutralization (ΔHn): The amount of energy absorbed/released when a solute dissolves. • Heat of Combustion (ΔHcombustion): amount of energy released when a material burns. • Heat of Fusion (melting) or Solidification (freezing) (ΔHf)=-(ΔHs) : amount of energy absorbed when a solid melts or released when a liquid freezes • Heat of Vaporization or Condensation (ΔHv)=-(ΔHcondensation) : amount of energy absorbed when a liquid evaporates or released when a gas condenses

  45. Recapping Enthalpy • Enthalpy is the total heat content of a substance, including the energy that was stored in the bonds of the substance during its formation. • Enthalpy is mostly potential energy. • Enthalpy cannot be measured directly, but it can be calculated by the amount of energy released/absorbed during reactions. • The Heat of Reaction (ΔH) or enthalpy change is the difference between the Heats of Formation (ΔHF) of the products and the reactants Next

  46. Enthalpy Diagram: Exothermic Reactants ΔH is Negative (ΔH<0) Potential Energy (Enthalpy) Products Progress of Reaction (Time)

  47. Enthalpy Diagram: Endothermic Products ΔH is Positive (ΔH>0) Potential Energy (Enthalpy) Reactants Progress of Reaction (Time)

  48. Some books show Enthalpy Graphs like this: Exothermic ΔH ΔH Endothermic It means the same as the ones before!

  49. Another Type of Enthalpy GraphThis is the type shows what happens during a reaction. We will examine this type in more detail in the next module. Activation Energy Enthalpy ΔH Reaction progress (time)

  50. Trick Questions • Some physical processes are tricky to classify as exothermic or endothermic. The following are guides: Dissolving can be EXOTHERMIC or ENDOTHERMIC depending on the solute.