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Thermochemistry: Understanding Heat, Enthalpy, and Energy Changes in Chemical Reactions

This section delves into thermochemistry, focusing on the heat exchanged during chemical reactions and physical changes. It covers key concepts such as enthalpy changes and the laws of thermodynamics, including the conservation of energy and entropy. The first law explains the relationship between internal energy, heat, and work. We discuss endothermic and exothermic reactions, highlighting energy transfer, heat capacity, and phase transitions. By examining specific examples, this material illustrates the principles governing energy changes and heat dynamics in various chemical processes.

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Thermochemistry: Understanding Heat, Enthalpy, and Energy Changes in Chemical Reactions

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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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