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Thermochemistry Notes

Thermochemistry Notes. I. Thermochemistry deals with the changes in energy that accompany a chemical reaction. Energy is measured in a quantity called enthalpy , represented as H . The change in energy that accompanies a chemical reaction is represented as  H . Page 519.

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Thermochemistry Notes

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  1. Thermochemistry Notes

  2. I. Thermochemistry deals with the changes in energy that accompany a chemical reaction. Energy is measured in a quantity called enthalpy, represented as H. The change in energy that accompanies a chemical reaction is represented as H. Page 519

  3. a. The energy absorbed or released as heat in a chemical or physical change is measured in a calorimeter. In one kind of calorimeter, known quantities of reactants are sealed in a reaction chamber, which is immersed in a known quantity of water in an insulated vessel. Therefore, the energy given off (or absorbed) during the reaction is equal to the energy absorbed (or given off) by the known quantity of water. The amount of energy is determined from the temperature change of the known mass of surrounding water.

  4. Calorimeter

  5. Using the change in temperature, T, determined from calorimeter one can use the following equation to determine the quantity of energy gained or lost during the reaction or for a physical change: q = (cp)(m)(T) T in kelvin q represents the energy lost or gained (in J) m is the mass of the sample (in g) cp is the specific heat of a substance at a given temperature pg513

  6. Practice Problem: 1. How much heat energy is needed to raise the temperature of a 33.0 gram sample of aluminum from 24.0C to 100.C? • 2. Determine the specific heat of a material if a 35 g sample absorbed 96 J as it was heated from 293K to 313K. • 3. During a chemical reaction carried out in a calorimeter the temperature of water within the calorimeter raised from 24.0C to 125C. If 250. grams of water were present calculate the amount of heat energy the water gained.

  7. b. In thermochemical equations the quantity of energy released or absorbed as heat during a reaction is written and is represented by H. Example: 2H2(g) + O2(g)  2H2O(l) H= - 571.6 kJ/mol

  8. c. H can be used to determine if the reaction is exothermic or endothermic. If the H value of an equation is negative that represents an exothermic reaction. (Meaning energy is released, therefore the energy of the products would be less.) If the H value of an equation is positive that represents an endothermic reaction. Example:

  9. 2H2(g) + O2(g)  2H2O(g) H= - 483.6 kJ/mol H reactants = 1450.8 kJ/mol H products = 967.2 kJ/mol Type of Reaction: Exothermic 2H2O(g)  2H2(g) + O2(g) H= + 483.6 kJ/mol H reactants = 967.2 kJ/mol H products = 1450.8 kJ/mol Type of Reaction: Endothermic

  10. d. Hess’s law provides a method for calculating the H of a reaction from tabulated data. This law states that if two or more chemical equations are added, the H of the individual equations may also be added to find the H of the final equation. As an example of how this law operates, look at the three reactions below.

  11. (1) 2H2(g) + O2(g)  2H2O(l)H = 571.6 kJ/mol (2)2H2O2(l)  2H2(g) + 2O2(g)H = +375.6 kJ/mol (3)2H2O2(l)  2H2O(l) + O2(g)H = ? kJ/mol

  12. When adding equations 1 and 2, the 2 mol of H2(g) will cancel each other out, while only 1 mol of O2(g) will cancel. (3)2H2O2(l)  2H2O(l) + O2(g)H = ? kJ/mol

  13. Warm-up What type of energy and energy transfer do you see in this picture Thank you Ms. Bouwman (from McNeil High School) for this wonderful power point!

  14. Basic Thermochemistry Courtesy of lab-initio.com

  15. Energyis the capacity to do work • Thermal energy is the energy associated with the random motion of atoms and molecules • Chemical energy is the energy stored within the bonds of chemical substances • Nuclear energy is the energy stored within the collection of neutrons and protons in the atom • Electrical energy is the energy associated with the flow of electrons • Potential energy is the energy available by virtue of an object’s position 6.1

  16. Energy Heat (Enthalpy) Change, ΔH • Energy is the capacity to do work, and can take many forms • Potential energy is stored energy or the energy of position • Kinetic energy is the energy of motion • Thermal energy (heat) is an outward manifestation of movement at the atomic level Definition:The amount of heat energy released or absorbed during a process.

  17. Temperature = Thermal Energy 900C 400C Energy Changes in Chemical Reactions Heat is the transfer of thermal energy between two bodies that are at different temperatures. Temperature is a measure of the thermal energy. greater thermal energy 6.2

  18. Cup gets cooler while hand gets warmer Heat The flow of thermal energy from one object to another. Heat always flows from warmer to cooler objects. Ice gets warmer while hand gets cooler

  19. 3 Types of Heat Transfer • Radiation- the transfer of energy by electromagnetic waves. • Convection – Transfer of energy by currents • Conduction – Transfer of energy by touching objects

  20. 2H2(g) + O2(g) 2H2O (l) + energy H2O (g) H2O (l) + energy energy + 2HgO (s) 2Hg (l) + O2(g) energy + H2O (s) H2O (l) Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. Endothermic process is any process in which heat has to be supplied to the system from the surroundings. 6.2

  21. Exothermic Processes Processes in which energy is released as it proceeds, and surroundings become warmer Reactants  Products + energy

  22. Endothermic Processes Processes in which energy is absorbed as it proceeds, and surroundings become colder Reactants + energy  Products

  23. Water phase changes constant Temperature remains __________ during a phase change.

  24. Phase Change Diagram Processes occur by addition of energy   Processes occur by removal of energy

  25. Thermochemical Calculations

  26. Units for Measuring Heat The Joule is the SI system unit for measuring heat: The calorie is the heat required to raise the temperature of 1 gram of water by 1 Celsius degree

  27. Specific Heat The amount of heat required to raise the temperature of one gram of substance by one degree Celsius.

  28. Specific Heat (cp, sometimes s, but usually c) Things heat up or cool down at different rates. Land heats up and cools down faster than water, and aren’t we lucky for that!?

  29. Specific heat is the amount of heat required to raise the temperature of 1 kg (but in Chem we use g) of a material by one degree (C or K, they’re the same size). Cp water = 4184 J / kg C (“holds” its heat) Cp sand = 664 J / kg C (less E to change) This is why land heats up quickly during the day and cools quickly at night and why water takes longer.

  30. Calculations Involving Specific Heat OR cp = Specific Heat Q = Heat lost or gained T = Temperature change m = Mass

  31. The amount of heat required to raise the temperature of one gram of substance by one degree Celsius. Specific Heat

  32. The specific heat (s) of a substance is the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius. The heat capacity (C) of a substance is the amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius. C = ms Heat (q) absorbed or released: q = msDt q = CDt Dt = tfinal - tinitial 6.4

  33. Specific Heat Capacity If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al? heat gain/lose = q = (c)(mass)(∆T) where ∆T = Tfinal - Tinitial q = (0.897 J/g•K)(25.0 g)(37 - 310)K q = - 6120 J Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al.

  34. How much heat is given off when an 869 g iron bar cools from 940C to 50C? s of Fe = 0.444 J/g • 0C Dt = tfinal – tinitial = 50C – 940C = -890C q = msDt = 869 g x 0.444 J/g • 0C x –890C = -34,000 J 6.4

  35. Calorimetry

  36. Calorimetry The amount of heat absorbed or released during a physical or chemical change can be measured, usually by the change in temperature of a known quantity of water in a calorimeter.

  37. Constant-Volume Calorimetry qsys = qwater + qbomb + qrxn qsys = 0 qrxn = - (qwater + qbomb) qwater = msDt qbomb = CbombDt Reaction at Constant V No heat enters or leaves! 6.4

  38. Constant-Pressure Calorimetry qsys = qwater + qcal + qrxn qsys = 0 qrxn = - (qwater + qcal) qwater = msDt qcal = CcalDt Reaction at Constant P No heat enters or leaves! 6.4

  39. BOOM! Combustible material ignited at constant volume! This heats up the “bomb”, which heats up the water surrounding it… First, some heat from reaction warms the water, which we know the mass of and “c” for… qwater = (c)(water mass)(∆T) THEN, some heat from reaction warms “bomb,” which has a known specific heat for the entire apparatus (typically), so we don’t need the mass… qbomb = (heat capacity, J/K)(∆T) Total heat evolved = qtotal = qwater + qbomb

  40. Practice A sample of iron metal is added to 75.00 grams of water originally at 35.0°C in a calorimeter. The final temperature of the metal and water in the calorimeter is measured to be 95.0°C. a) Describe the transfer of heat energy that occurs in the calorimeter. (b) Assuming no heat is lost to the outside, how many joules of heat energy are transferred?

  41. Enthalpy (H)

  42. Phase Change Diagram D B A C E

  43. Changing Phase • Heat of Fusion - heat change for freezing and melting • Heat of Vaporization – heat change for condensation or evaporation For Water: Heat fusion = 340 J/g Heat vaporization = 2,300 J/g Heat = (mass)(heat of fusion or vaporization) How many joules of heat are necessary to melt 500g of ice at its freezing point? = 500g * 340J/g = 170,000 J or 170KJ

  44. Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. DH = H (products) – H (reactants) DH = heat given off or absorbed during a reaction Hproducts < Hreactants Hproducts > Hreactants DH < 0 DH > 0 6.3

  45. ∆Hfo, standard molar enthalpy of formation ∆Hfo = Enthalpy change when 1 mol of compound is formed from the corresponding elements under standard conditions H2(g) + 1/2 O2(g) --> H2O(g) ∆Hfo (H2O, g)= -241.8 kJ/mol

  46. Decomposition • ∆Hfomay also be used to calculate the decomposition of something • If … H2(g) + 1/2 O2(g) --> H2O(g) ∆Hf˚ = - 242 kJ/mol Then… H2O(g) --> H2(g) + 1/2 O2(g) ∆Hf˚ = + 242 kJ/mol

  47. Enthalpy Values • Depend on how the reaction is written and on phases of reactants and products… H2(g) + 1/2 O2(g) --> H2O(g) ∆H˚ = -242 kJ 2 H2(g) + O2(g) --> 2 H2O(g) ∆H˚ = -484 kJ H2O(g) ---> H2(g) + 1/2 O2(g) ∆H˚ = +242 kJ H2(g) + 1/2 O2(g) --> H2O(liquid) ∆H˚ = -286 kJ

  48. Huh? So what’s that mean? To convert 1 mol of water to 1 mol each of H2 and CO requires 131 kJ of energy. Since delta H is positive, the “water gas” reaction is ENDOthermic

  49. 6.4

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