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Energy & Chemical Change

Energy & Chemical Change. Mr. Solsman Chapter 15. Chemistry is the study of matter, its composition and changes, and the energy effects that accompany these changes. Chapter 15 deals with energy and how energy changes relate to chemical reactions. Why are energy factors important?.

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Energy & Chemical Change

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  1. Energy & Chemical Change Mr. Solsman Chapter 15

  2. Chemistry is the study of matter, its composition and changes, and the energy effects that accompany these changes. • Chapter 15 deals with energy and how energy changes relate to chemical reactions.

  3. Why are energy factors important? • Can we power cars using water as fuel? • Energy factors are critical for the safe operation of chemical production facilities (and lab experiments). • Energy factors are important in determining if a process can be operated at a profit.

  4. The Nature of Energy • Energy is the ability to do work or produce heat. • Two Classifications - Kinetic & Potential • Kinetic Energy is the Energy of Motion • Potential Energy is Stored Energy • can be a result of position or composition. • we will be most concerned with potential energy in the form of composition.

  5. The Law of Conservation of Energy • Energy can neither be created or destroyed. • Energy can only be converted from one form to another.

  6. Kinetic Energy – the energy of motion • Depends on velocity • Depends on mass Kinetic energy = ½ mv2

  7. Positional Potential Energy • Depends on mass • Depends on distance • Depends on an attractive force Potential Energy ↔ mad

  8. Compositional Potential Energy • Depends on the chemical bonds in a compound. • A chemical reaction can store energy or release energy. • Burning wood is an example of releasing energy. • Plant growth is an example of storing energy.

  9. Energy is a State Function • State functions are those which depend only on the current state of a system and not on it’s history. • This concept is critical to the content of this chapter.

  10. Temperature and Heat • Temperature is a measure of how hot something is. • (Well, duh!) • Temperature is related to how fast the particles of a system are moving.

  11. Temperature • Temperature is a measure of the average kinetic energy of particles in the system. • As kinetic energy is added to a system, the molecules move faster and the average kinetic energy increases – we say the temperature goes up.

  12. HEAT • Heat is NOT the same as temperature. • Heat is the FLOW of energy from one object to another. • Heat ALWAYS flows from hot to cold. • Heat flow does not always raise or lower the temperature, it can change the kinetic OR the potential energy.

  13. TYPES OF ENERGY CHANGES • Exothermic changes – changes where energy is released. • Endothermic changes – changes where energy is absorbed.

  14. Terminology • The SYSTEM is what we are focusing our attention on. • The SURROUNDINGS is the rest of the universe.

  15. Units of Heat • The calorie is defined as the amount of heat needed to raise the temperature of one gram of water by one ºC. • The Calorie (food calorie) is equal to 1000 calories.

  16. Units of Energy • The SI unit of energy is the joule. • 1 calorie = 4.184 joules • How many joules is 65 calories? • How many calories is 457 joules?

  17. Heat Calculations • How much energy (in joules) does it take to heat 25 grams of water from 22 ºC to 45 ºC?

  18. Heat Calculations • Q = c x m x ∆T (where ∆ means “change in”) • ∆T = TFinal – TInitial • m is the mass • c is called the specific heat capacity. The units of c are J / g ºC

  19. Typical Calculation • How many joules are required to heat 1.50 kg of iron from 15 ºC to 37 ºC? • Q = c x m x ∆T = 0.45 J/g ºC x 1.50 kg x 1000 g/kg x 22 ºC = 15000 J

  20. WOW!!!! • How many joules are required to heat 1.50 kg of water from 15 ºC to 37 ºC? • Q = c x m x ∆T = 4.2 J/g ºC x 1.50 kg x 1000 g/kg x 22 ºC = 140000 J

  21. Compare this to the value of 15,000 J for iron. Why are these values so different? • Specific Heat Capacity

  22. Specific Heat Capacity

  23. Problem • If it required 54 J to heat 15 grams of a sample from 45 ºC to 60. ºC, what is the specific heat capacity of the material? • Which of the materials on the specific heat capacity slide would the material most likely be?

  24. Specific Heat Capacity

  25. Chemical Energy and the Universe • Thermochemistryis the study of heat changes that accompany chemical reactions and phase changes. 4Fe(s) + 3O2(g) → 2FeO3(s) + 1625 kJ

  26. The system is the specific part of the universe that contains the reaction or process you wish to study. The surroundingsare everything else other than the system in the universe. • The universeis defined as the system plus the surroundings.

  27. Enthalpyis the heat content of a system at constant pressure. • Enthalpy (heat) of reactionis the change in enthalpy during a reaction symbolized as ΔHrxn. ΔHrxn = Hfinal– Hinitial ΔHrxn = Hproducts– Hreactants

  28. Enthalpy changes for exothermic reactions are always negative. • Enthalpy changes for endothermic reactions are always positive.

  29. Cold-pack reaction: • 27 kJ + NH4NO3(s) → NH4+ (aq) + NO3- (aq)

  30. STOICHIOMETRY THERMOCHEMICAL REACTION – A REACTION THAT INCLUDES THE ENERGY TERM. • 136 kcal + 2H2O(l)  2H2(g) + O2(g) • C(s) + O2(g)  CO2(g) + 94.0 kcal • How much heat is absorbed when 30.0 grams of water decomposes? • What amount of heat is released when 4.4 grams of CO2 form?

  31. 5. 3H2(g) + N2(g)  2NH3 H = -11.0kcal/mole NH3 a. What is the thermochemical reaction? b. What amount of heat is released if 1.6 grams of H2 reacts with nitrogen? 6. N2(g) + 2O2(g)  2NO2(g) H = 8.1kcal/mole NO2 a. What is the thermochemical reaction? b. What amount of heat is needed when 3.50 liters of O2 react with nitrogen?

  32. Changes of State • Molar enthalpy (heat) of vaporizationrefers to the heat required to vaporize one mole of a liquid substance. • Molar enthalpy (heat) of fusionis the amount of heat required to melt one mole of a solid substance.

  33. Molar enthalpy (heat) of vaporizationrefers to the heat required to vaporize one mole of a liquid substance. • Molar enthalpy (heat) of fusionis the amount of heat required to melt one mole of a solid substance.

  34. Phase Change Diagram

  35. Phase Changes are physical changes. Na+Cl-(s) Na+Cl-(l) - MELTING H2O(l)  H2O(g) -EVAPORATION CO2(s)  CO2(g) -SUBLIMATION Ne(g)  Ne(l) - CONDENSATION H2O(g)  H2O(s) -DEPOSITION Cu(l)  Cu(s) -FREEZING

  36. How much heat does it take to completely vaporize 100. grams of room temperature water? • The total heat required is the heat to raise the temperature of the water to its boiling point PLUS the heat to vaporize it at its boiling point.

  37. Q = m c ΔT • H = n (ΔHpc ) • n is in moles • ΔH is in joules/mole or kcal/mol

  38. How much heat does it take to completely vaporize 100. grams of room temperature(25 ºC) water? • Q = 100. g * 4.18 J/g ºC * (100-25) ºC • ΔHvap=(100. g/18.0 g/mol)*40.7 kJ/mol • Total Heat Needed = Q + ΔH = 31,400 J + 224,000 J = 255,000 J

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