1 / 30

Chapter 17: Energy and Kinetics

Pages 510-547. Chapter 17: Energy and Kinetics. Thermochemistry : Causes of change in systems Kinetics : Rate of reaction progress (speed). Heat , Energy , and Temperature changes. Standard unit of heat is the Joule , J Standard unit of temperature is Kelvin , K.

ofira
Download Presentation

Chapter 17: Energy and Kinetics

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. Pages 510-547 Chapter 17: Energy and Kinetics Thermochemistry: Causes of change in systems Kinetics: Rate of reaction progress (speed) Heat,Energy, and Temperature changes Standard unit of heat is the Joule, J Standard unit of temperature is Kelvin, K

  2. Heat vs Temperature • Heat • measure of energy change in a system. • Temperature • measure of the average kinetic energy (movement) of the particles in a system. • Exothermic • System loses energy to surroundings • Endothermic • System gains energy from surroundings K = C + 273.15

  3. Specific Heat • Specific Heat • measure of how a substance reacts to heat energy changes. • Think thermal inertia • is the heat energy required to raise one gram of a pure substance one degree Celsius. • is a property of matter; different species have different Specific Heats. • The symbol we use is cp. • The “p” stands for constant pressure while heat is added or lost.

  4. Specific Heat Capacity • Metals have very low cp, • which is why metals often feel cold to the touch. • Water has a very highcp, • 4.184 J/g·0C • Substances with lower cp will rise in temperature faster and require less energy to do so than do substances with high cp.

  5. Specific Heat Capacity Cp (H2O) = 4.184 • Cpunits are J/g0C) • Change inheat(joules, J) • = • Change intemperature(degree, 0C) • x • Mass (mass, g) • x • Specific Heat Capacity(4.184 for water) 1 calorie = 4.184 Joules

  6. Specific Heat Example Exercise • Determine the specific heat of 34 grams of an unknown material if 485 J of heat are absorbed to change the temperature by 20.0 oC. • If 950 J of heat are added to 5.4 mL of water at 280 K, what will be the resulting temperature of the water? (hint: mL  g)

  7. The Calorimeter The Calorimeter (shown) Heat energy is transferred from a reaction inside the calorimeter to the water in the calorimeter. The temperature change of the water is observed. Text page 519 When two objects are in contact, they eventually obtain Thermal Equilibrium; their temperatures become equal.

  8. Enthalpy,ΔH • Enthalpy • heat energytransferred for a specific change to take place. • We specify enthalpy with ΔH. “Δ” means “change in”. • Exothermic reaction • negative enthalpy (-ΔH ) • Endothermic reaction • Positive enthalpy (+ΔH). • Elements in their standard (elemental) state have a ΔH of zero. • O2, Fe, Cu, N2, He, etc are have Hf = 0 kj/mol The universe favors LOW energy states - if the products have lower energy reaction is favored.

  9. Enthalpy,ΔH State change to/from? Sign of ΔH? • Some common changes involving ΔH: • ΔHfus=heat of fusion • ΔHvap=heat of vaporization • ΔHcond=heat of condensation • ΔHsub=heat of sublimation • ΔHrxn=heat of reaction • ΔHf=heat of formation • ΔHsol=heat of solution • ΔHcomb=heat of combustion

  10. Phase Changes • Solid + heat = temp ▲ • Solid + heat = phase change • Liquid + heat = ? • And then…? • Gas + heat = ? temperature temperature Heat added Heat added

  11. Reaction Enthalpy • If ΔH is negative, the reaction is exothermic. • C6H12O6 + 6O26CO2 + 6H2O + 2870kJ • ΔHrxn = -2870 kJ/mol • If ΔH is positive, the reaction is endothermic. • 2H2O + 571.6kJ 2H2 + O2 • ΔHrxn = +571.6 kJ/mol energy energy

  12. Spontaneity • Spontaneous • A reaction that will proceed on its own once started. • Sometimes, all the reaction needs to get going is the kinetic energy of nearby colliding atoms. • Kinetic Molecular Theory: • All Matter is made of particles in constant motion • Some collisions are more energetic than others. Why? • Spontaneous combustion • occurs when the kinetic energy of colliding oxygen molecules striking a fuel have enough energy on their own to start the combustion reaction.

  13. DHrxn Exothermic Reaction: products have lower energy than do the reactants. What if endothermic?

  14. In the diagram, the hump is called a activation energy barrier - the amount of energy required for the reaction to begin. We can reduce the activation energy with a catalyst. Activated complex All reactions require some sort of activation energy , Ea.

  15. Hess’s Law: • If two reactions begin with the same reactants in the same condition and end with the same products in the same condition, they must have the same enthalpy change. • It doesn’t matter if you perform a reaction in several steps or produce your final product in one step, the enthalpy change will be the same. • Consider the reaction A +B D : • A + B + 100 kJ CthenC + 50 kJ  D • Must be the same as A + B + 150 kJ  D

  16. Hess’s LawEnthalpy of Reaction ΔHrxn = ∑Hproducts – ∑Hreactants • = ΣHf,all the products – ΣHf,all the reactants “ Sum of ” Enthalpy of Formation

  17. Hess’s Law Example Exercise ΔHrxn = ∑Hproducts – ∑Hreactants • Calculate the heat of reaction when 350 grams of methane, CH4 are burned in excess oxygen. Hf book values for each species are: • CH4(g) = -74.8 kJ/mol • O2(g) = ? • H2O(g) = -285.83 kJ/mol • CO2(g) = -393.5 kJ/mol

  18. Entropy, ΔS • Entropy • is a measure of relative disorder. Thermodynamics tells us that the universe tends towards disorder or entropy. • Temperature affects entropy (why?) • Entropy calculations are very similar to enthalpy calculations: ΔSrxn= ΣSproducts – ΣSreactants Entropy has the unit J/K*mol

  19. Entropy, ΔS • The universe tends towards entropy • entropy plays a part in predicting whether or not a reaction will be spontaneous. • Solids have very low entropy • Gases have very high entropy • Solutions also have high entropy

  20. Qualitative Entropy Values • We can make generalizations about a reaction’s entropy; • 2KClO3(s) 2KCl(s) + 3O2(g) • 2 solids  2 solids + 3 gases • Entropy appears to increase in this reaction.

  21. Quantitative Entropy Values • 2KClO3(s) 2KCl(s) + 3O2(g) • S of KClO3(s)= 143.7 J/mol*K • S of KCl(s)= 82.6 J/mol*K • S of O2(g)= 205.1 J/mol*K • Using ΔSrxn= Sproducts – Sreactants, the reaction has a total entropy change of +493.1 J/mol*K

  22. Entropy Values • A positiveΔS = increase in entropy • A negativeΔS = decrease in entropy Do not confuse entropy and enthalpy! • Tending toward spontaneity: • Negative Enthalpy(-ΔH) • Positive Entropy(+ΔS)

  23. Free Energy, ΔG • Free energy, ΔG: • allows us to assign a value to an entire reaction to predict whether a reaction is spontaneous, product favored. • or nonspontaneous, reactant-favored. • Named for American Chemist, J. Willard Gibbs ΔG = ΔH -TΔS Free Energy kJ/mol EnthalpykJ/mol Entropy J/mol·K temperature inKelvin

  24. Gibbs Free Energy, ΔGrxn Negative Gibbs Energy (-ΔGrxn) Spontaneous, Product favored Positive Gibbs Energy (+ΔGrxn) Nonspontaneous, Reactant favored A ΔG of zero means that neither the products nor reactants are favored-the reaction is in equilibrium.

  25. DG = DH - TDS

  26. Reaction Rates • Reaction rates • how fast a reaction proceeds. • Some factors will affect reaction rate: • Temperature of reactants: higher = faster • Concentration of reactants: greater = faster • Surface area of reactants: greater = faster • (powders react faster than chunks) • Pressure of gaseous reactants: greater = faster • Catalyst presence: catalysts make rxns faster • reduce activation energy! • are not used up (not reactants)

  27. Rate Laws • For any reaction: • The rate is based on the [reactants]: • [X] : “1st order” : 2x[A] , 2x rate • [X]2 : “2nd order” : 2x[A] , 4x rate • [X]3 : “3rd order” : 2x[A] , 8x rate End of C17, conclusion follows

  28. In conclusion… Recall that K = C + 273.15 • Specific Heat Capacity, cp(J/gK) • the amount of heat energy required to raise 1 gram, 1 degree • Enthalpy, ΔH (kJ/mol) • the heat energy transferred in a reaction • Entropy, ΔS(J/mol-K) • the change in disorder of the species in a reaction • Gibbs Free Energy, ΔG(kJ/mol) • measure of spontaneity; how product favored or reactant favored a reaction is

  29. CCSD Syllabus Objectives • 16.1: Thermodynamics, definition • 16.2: Exothermic/Endothermic • 16.3: Changes in Enthalpy • 16.4: Thermochemical Calculations • 16.5: Energy Diagrams • 16.6: Enthalpy-Entropy-Free Energy • 17.1: Kinetics Definition • 17.2: Factors that Affect Reaction Rate

  30. Aligned Labs and Demos • Lab: Flaming Cheeto Calorimetry Lab • Lab: Metals Calorimetry Lab • Lab: NaOH-HCl Enthalpy of Reaction Lab • Lab: Ba(OH)2-8H2O w/ NH4NO3 and H2O2 with a catalyst Free Energy Lab • Lab: KI-H2O2 Kinetics Lab • Demo: Carbon Snake with powdered vs granular sugar (Kinetics)

More Related