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Chapter 12

Chapter 12. Section 1 Chemical Reactions that involve heat Section 2 Heat and Enthalpy Changes. Thermochemistry. Study of heat released or absorbed during a chemical reaction or a physical change. 3 Laws of Thermodynamics. First law Conservation of energy ∆U = Q + W

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Chapter 12

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  1. Chapter 12 Section 1 Chemical Reactions that involve heat Section 2 Heat and Enthalpy Changes

  2. Thermochemistry • Study of heat released or absorbed during a chemical reaction or a physical change

  3. 3 Laws of Thermodynamics • First law • Conservation of energy • ∆U = Q + W ∆U = Change in thermal energy (= internal energy) in a system +Q = Heat absorbed by the system ‒Q = Heat given off by the system +W = work done on the system ‒W = work done by the system

  4. Second law • Entropy, S, (randomness, disorder, or chaos) in Universe is always increasing (= spontaneity) • Heat added increases entropy (Ex) Which will add more disorder? a) 10 J heat added to a system at 100 ˚C b) 10 J heat added to a system at -10 ˚C (3) (T must be in K)

  5. Third law: The entropy of a “perfect” crystal is zero.

  6. Terms to Know • Energy = the ability to do work • Chemical potential energy: energy stored in chemical bonds of a substance • Breaking a chemical bond requires energy • Forming a chemical bond releases energy • Energy change = heat (q) or work (w) • heat = the transfer of energy from high to low temperature

  7. System: a part of universe that is under the observation, usually the chemical reaction • Surrounding: the rest of universe • Endothermic • energy going into the system • energy shown as a reactant (Ex) Ice + heat → water • positive “q” value • You, as part of surrounding, will feel cold • boiling, melting, evaporation

  8. Exothermic • energy coming out of the system • energy shown as a product (Ex) water → ice + heat • negative “q” value • You, as part of surrounding, will feel hot • freezing

  9. Examples • Endothermic reaction (Ex) CO2 + 2H2O  CH4 + 2O2ΔH = 890.4 kJ (Ex) CO2 + 2H2O + 890.4 kJ  CH4 + 2O2 • Exothermic reaction (Ex) CH4 + 2O2 CO2 + 2H2O ΔH = ‒890.4 kJ (Ex) CH4 + 2O2 CO2 + 2H2O + 890.4 kJ

  10. Units of Energy • calorie (=cal) – Do not capitalize c! 1 cal: the amount of energy needed to raise the temperature of 1 g water 1 °C (Ex) A candy bar has about 200 calories. 2. Calorie (=Cal) 1 Cal = 1 kcal = 1000 cal 3. joule (=J) – the SI unit of energy 1 cal = 4.184 J (Two conversion factors are: )

  11. Example Express 60.1 cal to J.

  12. Enthalpy Change, ΔH • heat released or absorbed when reactions take place under a constant pressure • From now on, we will assume that energy is only converted to heat • No energy is converted to work • ΔH= Hproducts – Hreactants • H = enthalpy = heat content in a substance • positive ΔH = endothermic • negative ΔH= exothermic

  13. Examples

  14. Exothermic and Endothermic

  15. Standard Enthalpy Change, ΔHo • The heat released or absorbed by a chemical reaction at 25oC and 1 atm

  16. Enthalpy Change of Reactions • The coefficients reflects the moles (Ex) 2H2O2(l)  2H2O(l) + O2(g) ΔHo=-190 kJ 2 mol of H2O2(l) release 190 kJ of heat (Ex) 1H2O2(l)  1H2O(l) + ½ O2(g) ΔHo=-95 kJ 1 mol of H2O2(l) releases 95 kJ of heat • If the reaction is reversed, the sign of enthalpy changes to the opposite (Ex) 2H2O(l) + O2(g)  2H2O2(l) ΔHo=190 kJ

  17. Examples (1) How much heat is released if 1.0 g H2O2 decomposes? 2H2O2(l)  2H2O(l) + O2(g) ΔHo=-190 kJ

  18. (2) How much heat is transferred when 9.22 g of glucose (C6H12O6) in your body reacts with O2? C6H12O6+ 6O2 6CO2 + 6H2O ΔHo=-2803 kJ

  19. Changing States of Matter • States of Matter • Phase Transitions

  20. Heat & Temperature

  21. Specific Heat (Capacity), C • The amount of heat energy needed to raise 1 g of a substance by 1 K or 1 ˚C • Intrinsic property • can be used to identify the substance (Ex) For water, C = 4.184 J/g∙K For aluminum, C = 0.90 J/g∙K • Unit: J/g∙K, J/g∙˚C, cal/g∙K, cal/g∙˚C, J/mol ∙˚C • Amount of heat required or released: q = m∙C∙ΔT • q = amount of heat • m = mass • C = specific heat • ΔT= change in temperature =Tfinal– Tinitial

  22. Example • How many joules of energy must be absorbed to raise the temperature of 20 grams of water from 25°C to 30°C? The specific heat of water is 4.184 J/g∙K. Is this an exothermic or endothermic reaction? (Answer) 420 J; endothermic

  23. 2. An 50 gram sample of an unknown metal warms from 18°C to 58°C after absorbing 800 joules. What is the specific heat of the metal? (Answer) C=0.40 J/g∙°C

  24. Heat Energy Diagram endothermic exothermic

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