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Chemical Kinetics and Thermodynamics

Chemical Kinetics and Thermodynamics. Objectives: explain the collision theory of reactions describe reaction mechanism, rate-determining step, activated complex, and activation energy

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Chemical Kinetics and Thermodynamics

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  1. Chemical Kinetics and Thermodynamics Objectives: explain the collision theory of reactions describe reaction mechanism, rate-determining step, activated complex, and activation energy explain how rate of reaction is affected by the nature, the surface area, and the concentration of reactants; by the temperature; and by the presence of a catalyst interpret potential energy and energy distribution diagrams interpret the significance of changes in enthalpy use Hess’s law to calculate heats of reaction and heats of formation determine values of changes in free energy use them to predict spontaneous reactions Created by C. Ippolito February 2007

  2. Chemical Kinetics • concerned with the rates and mechanisms of chemical reactions • reaction rate • moles of reactant used up over time • moles of product produced over time • reaction mechanism • rearrangement to form products Created by C. Ippolito February 2007

  3. Collision Theory • explains different rates of reaction • collisions may shift electron positions • breaking old bonds and forming new ones Created by C. Ippolito February 2007

  4. Effective Collision • proper angle (orientation) • proper amount of energy • rate of reaction depends on: • more collisions • effectiveness of collisions Created by C. Ippolito February 2007

  5. Factors Affecting Rate • nature of reactants • temperature • concentration of reactants • presence of catalysts Created by C. Ippolito February 2007

  6. Nature of Reactants • slight rearrangements occur rapidly • ionic substances in solution • covalent rearrangements occur more slowly • bonds have to be broken and reformed Created by C. Ippolito February 2007

  7. Temperature • increase in temperature increases rate •  temperature  kinetic energy •  kinetic energy  collisions •  collisions  effective collisions Created by C. Ippolito February 2007

  8. Concentration of Reactants •  concentration  rate of reaction • Homogeneous Reaction • all reactants are in same phase • Heterogeneous Reaction • the reactants are in more than one phase • increase of surface area can increase rate in heterogeneous reactions Created by C. Ippolito February 2007

  9. Catalysts • substance that speeds up a reaction without being changed • usually works by increasing effective collisions • by “positioning” • or orientation Created by C. Ippolito February 2007

  10. Activation Energy • Activated Complex • intermediate particles • collision at proper angle • collision has enough energy • short lived and unstable • changes into product • Activation Energy • minimum energy to form complex Created by C. Ippolito February 2007

  11. Potential Energy Diagram • Y-axis • potential energy • reactants • activated complex • products • X-axis • reaction coordinate • progress of reaction • a.k.a time • see p. 478 in text Created by C. Ippolito February 2007

  12. Activation Energy and Temperature • lower temperature T1 • fewer molecules reach EACT • higher temperature T2 • more molecules reach EACT • see p. 480 in text Created by C. Ippolito February 2007

  13. Activation Energy and Concentration • Higher concentration • more particles reach EAct • Lower concentration • less particles reach EAct • see p. 481 Created by C. Ippolito February 2007

  14. Activation Energy and Catalysts • Catalysts “tunnel” EACT • by “positioning” • more molecules achieve EACT • see p. 482 Created by C. Ippolito February 2007

  15. Heat Content (Enthalpy) • Thermodynamics • study of • changes in energy • influence of temperature • Heat Content (a.k.a Enthalpy) • represented by H • all forms of energy • chemical bond energy • potential energy • kinetic energy Created by C. Ippolito February 2007

  16. Heat of Reaction • H - heat of reaction • measured in kilojoules (kJ) or kilocalories (kcal) see Table I in Regents Reference Tables • H = Hproducts – Hreactants • positive H – energy absorbed - ENDOTHERMIC • negative H – energy released - EXOTHERMIC Created by C. Ippolito February 2007

  17. Heat of Formation • special H (heat of reaction) • 1 mole of a compound made from elements • standard heat of formation symbol  • temperature 25°C (298K) • pressure 1 atmosphere (101.3kPa) • measured • kilojoules/mole • kilocalories/mole • see table p. 485 Created by C. Ippolito February 2007

  18. Stability of Compounds • large negative heat of formation • CO2Hf = -394 kJ/mol • very stable • decomposition would require large input of energy • small negative or positive heat of formation • NO2 Hf = +33.1 kJ/mole • very unstable • decomposition requires little or no energy input to decompose (explosives) Created by C. Ippolito February 2007

  19. Hess’s Law of Constant Heat Summation • sum of two or more reactions • heat of reaction that is the sum of individual heat of reactions (Use Table I) Created by C. Ippolito February 2007

  20. Hess Law Calculations • What is the heat of reaction? CuO(s) + H2(g)  Cu(s) + H2O(g) H = ? kJ Look for heat of formation reactions in data table (p.485 or Table I on Regents) Cu(s) + ½O2(g)  CuO(s) H = -155 kJ • but our reaction is the reverse of one in table CuO(s)  Cu(s) + ½O2(g) H = 155 kJ H2(g) + ½O2(g)  H2O(g) H = -242 kJ • check for “same” coefficient, if true no adjustment needed • combine reactions Created by C. Ippolito February 2007

  21. Hess Law Calculations (cont’d) CuO(s) Cu(s) + ½O2(g) H = 155 kJ H2(g) + ½O2(g)  H2O(g) H = -242 kJ _______________________________________________ CuO(s) + H2(g)  H2O(g) + Cu(s)H = -87 kJ Created by C. Ippolito February 2007

  22. Entropy • measures disorder, randomness, or lack of organization • S – entropy • Si - initial entropy (before change) • Sf –final entropy (after change) • S – change of entropy • S = Sf – Si • S = a positive number • entropy increases, disorder increases, decomposition • S = a negative number • entropy decreases, disorder decreases, synthesis Created by C. Ippolito February 2007

  23. Spontaneous Reactions • occur without external cause • affected by: • H – change in enthalpy • S – change in entropy • Effect of Signs of H and S • H = - (favorable) • reaction will occur • H = + (unfavorable) S = - (unfavorable) • reaction cannot occur • H = - (favorable) S = - (unfavorable) • reaction can only occur if H > S • H = + (unfavorable) S = + (favorable) • reaction can only occur if S > H Created by C. Ippolito February 2007

  24. Gibbs Free Energy Equation • G = H - TS • explains relationship of enthalpy and entropy • G – free energy • H – change in enthalpy (heat) • T – temperature (in kelvin) • S – change in entropy • spontaneous reactions • G - negative Created by C. Ippolito February 2007

  25. Free Energy of Formation • This is the change in free energy when 1 mole of the compound is formed from its constituent elements. Created by C. Ippolito February 2007

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