1 / 46

Chapter 16 Reaction Energy

Chapter 16 Reaction Energy. Chapter 16. Section 1 – Thermochemistry Section 2 – Driving Force of Reactions. Thermochemistry. Define temperature and state the units in which it is measured. Define heat and state its units. Perform specific-heat calculations.

inara
Télécharger la présentation

Chapter 16 Reaction Energy

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 16Reaction Energy

  2. Chapter 16 • Section 1 – Thermochemistry • Section 2 – Driving Force of Reactions

  3. Thermochemistry • Define temperature and state the units in which it is measured. • Define heat and state its units. • Perform specific-heat calculations. • Explain enthalpy change, enthalpy of reaction, enthalpy of formation, and enthalpy of combustion. • Solve problems involving enthalpies of reaction, enthalpies of formation, and enthalpies of combustion.

  4. Thermochemistry • Virtually every chemical reaction is accompanied by a change in energy. • What two ways can energy be involved? • Thermochemistry: study of the transfer of energy as heat that accompanies chemical reactions and physical changes.

  5. Calorimeter • Calorimeter: used to measure heat absorbed or released. • Energy released from the reaction is measuredfrom the temp change ofthe water

  6. Temperature • measure of the average kinetic energy of the particles in a sample of matter. • The greater the kinetic energy of the particles in a sample, the hotter it feels. • Remember: How do we convert Celsius to Kelvin?

  7. Heat • Joule: SI unit of heat and energy • Heat: the energy transferred between samples of matter because of a difference in their temperatures. • Energy transferred as heat moves spontaneously  from matter at a higher temp  lower temp

  8. Specific Heat • Energy transferred depends on: • The material • The mass • Specific Heat: specific heat of a substance is the amount of energy required to raise the temperature of one gram by one Celsius degree (1°C) or one kelvin (1 K). • Units: J/(g•°C) or J/(g•K) • The temperature difference as measured in either Celsius degrees or kelvins is the same.

  9. Molar Heat Capacities of Elements and Compounds

  10. Specific Heat Calculation Heat released Specific heat: or absorbed: cp = specific heat at a given pressure q = energy lost or gained m = mass of the sample ∆T = difference between the initial and final temperatures.

  11. Example Problem 1 • A 4.0 g sample of glass was heated from 274 K to 314 K, and was found to have absorbed 32 J of energy as heat.a. What is the specific heat of this type of glass? b. How much energy will the same glass sample gain when it is heated from 314 K to 344 K?

  12. Example Problem 1 Solution Given:m = 4.0 g ∆T = 314 – 274 = 40. K q = 32 J Unknown: a. cp in J/(g•K) b. q for ∆T of 314 K → 344 K Solution: a.

  13. Example Problem 1 Solution b.

  14. Enthalpy of Reaction • Enthalpy(∆H): “change in enthalpy”energy absorbed as heat during a chemical reaction at constant pressure

  15. Enthalpy change • The difference between the enthalpies of products and reactants. ∆H = Hproducts – Hreactants

  16. Thermochemical Equation • Equation that includes the quantity of energy released or absorbed as heat 2H2(g) + O2(g) → 2H2O(g) + 483.6 kJ • States of matter are ALWAYS shown because it can influence the energy of released or absorbed. • Coefficients are always viewed as number of moles in the reaction. • If 2x as many reactants were provide, 2x as many moles of water would be produced and 2x as many kJ of energy would be released.

  17. Exothermic • chemical reaction that releases energy the energy of the products is less than the energy of the reactants • example: 2H2(g) + O2(g) → 2H2O(g) + 483.6 kJ

  18. Endothermic • chemical reaction that absorbs energythe products have a larger enthalpy than the reactants 2H2O(g) + 483.6 kJ → 2H2(g) + O2(g)

  19. Enthalpy of a Reaction • Enthalpy is often not included in the reaction expression, instead as a ∆H value • Example: 2H2(g) + O2(g) → 2H2O(g) ∆H = –483.6 kJ • Negative ∆H is exothermic – releases heat • Positive ∆H is endothermic – absorbs heat

  20. Exothermic Reaction Energy released, ∆H is negative.

  21. Endothermic Energy is absorbed, ∆H is positive.

  22. Enthalpy of Formation • molar enthalpy of formation: enthalpy change that occurs when one mole of a compound is formed from its elements in their standard state at 25°C and 1 atm. • Enthalpies of formation are given for a standard temperature and pressure so that comparisons between compounds are meaningful. standard states change in enthalpy heat of FORMATION

  23. Where do I find ΔHf? • In your textbook! • The values are given as the enthalpy of formation for one mole of the compound from its elements in their standard states.

  24. Stability • Compounds with a large negative enthalpy of formation are very stable. • Compounds with positive values of enthalpies of formation are typically unstable.

  25. Enthalpy of Combustion • ∆Hc : enthalpy change that occurs during the complete combustion of one mole of a substance **Enthalpy of combustion is defined in terms of one mole of reactant **Enthalpy of formation is defined in terms of one mole of product

  26. Combustion/Bomb Calorimeterused to find enthalpy of combustion

  27. Enthalpy of Reaction • Hess’s Law: the overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process. • This means that the energy difference between reactants and products is independent of the route taken to get from one to the other.

  28. Example 1 • If you know the reaction enthalpies of individual steps in an overall reaction, you can calculate the overall enthalpy without having to measure it experimentally. What is the heat of formation of methane? C(s) + 2H2(g) → CH4(g)

  29. Calculating using Hess’s Law • The component reactions in this case are the combustion reactions of carbon, hydrogen, and methane: Goal Reaction: C(s) + 2H2(g) → CH4(g) C(s) + O2(g) → CO2(g) H2(g) + ½O2(g) → H2O(l) 2 [ H2(g) + ½O2(g) → H2O(l) ] CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) CO2(g) + 2H2O(l) → CH4(g) + 2O2(g) ∆H0 = +890.8 kJ This reaction must x2! The ΔH is x2! This reaction must be reversed! The ΔH is changed to positive!

  30. ADD IT UP!! C(s) + O2(g) → CO2(g) 2H2(g) + O2(g) → 2H2O(l) CO2(g) + 2H2O(l) → CH4(g) + 2O2(g) C(s) + 2H2(g) → CH4(g)

  31. Enthalpy of Formation ∆H0= Σ[(ΔH0fof products) × (mol of products)]– Σ[(ΔH0f of reactants) × (mol of reactants)]

  32. 16.2 Driving Force of Reactions • Explain the relationship between enthalpy change and the tendency of a reaction to occur. • Explain the relationship between entropy change and the tendency of a reaction to occur. • Discuss the concept of free energy, and explain how the value of this quantity is calculated and interpreted. • Describe the use of free energy change to determine the tendency of a reaction to occur.

  33. Spontaneous Reaction • Majority of spontaneous reactions are exothermic • Reactions proceed in a direction that leads to a lower energy state • Some spontaneous reactions are endothermic • Example: melting (particles have more energy in the liquid state)

  34. Entropy • Entropy: • Increases when a system goes from one state to another without an enthalpy change 2NH4NO3(s)  2N2(g) + 4H2O(l) + O2(g) • The arrangement of particles on the right-hand side of the equation is more random than the arrangement on the left side and hence is less ordered.

  35. Entropy • Entropy (S) : the degree of randomness or disorder of the particles in a system • There is a tendency of nature to go from order to disorder • Think about your bedroom, does it always stay neat and organized? Or does it get messier and messier? • Increase in disorder  Increase in entropy • The entropy of the universe is ALWAYS increasing

  36. Standard Entropy Changes for Some Reactions

  37. Entropy in States of Matter • As you go from a solid  liquid  gas What happens to the entropy? • As you go from a gas  liquid  solid What happens to the entropy?

  38. Entropy

  39. Change in Entropy • Unit: kJ/(molK) • Change in Entropy ΔS • Difference between the entropy of the products and the reactants • Increase in Entropy  Positive Value • Decrease in Entropy  Negative Value

  40. Gibb’s Free Energy • Spontaneous Reactions: • Least enthalpy • Greatest entropy • Gibb’s Free Energy (G): • Combined enthalpy-entropy function

  41. Change in Free Energy • Only the change in free energy can be measured. • At a constant pressure and temperature, the free-energy change, ∆G, of a system is defined as the difference between the change in enthalpy, ∆H, and the product of the Kelvin temperature and the entropy change, which is defined as T∆S: ∆G0= ∆H0– T∆S0

  42. Relating Enthalpy and Entropy to Spontaneity

  43. Gibb’s Free Energy ∆G0 = ∆H0 – T∆S0 • The expression for free energy change is for substances in their standard states. Unit: kJ/mol • If ∆G < 0  the reaction is spontaneous.

  44. Ice - Enthalpy vs Entropy

  45. Example Problem For the reaction NH4Cl(s) → NH3(g) + HCl(g) ∆H0 = 176 kJ/mol and ∆S0 = 0.285 kJ/(mol•K) at 298.15 K. Calculate ∆G0, and tell whether this reaction is spontaneous in the forward direction at 298.15 K.

  46. Example Problem Solution Given:∆H0 = 176 kJ/mol at 298.15 K ∆S0 = 0.285 kJ/(mol•K) at 298.15 K Unknown:∆G0 at 298.15 K Solution:The value of ∆G0 can be calculated according to the following equation: ∆G0= ∆H0– T∆S0 ∆G0= 176 kJ/mol – 298 K [0.285kJ/(mol•K)] ∆G0= 176 kJ/mol – 84.9 kJ/mol ∆G0= 91 kJ/mol

More Related