Chapter 17

Reaction Energy and Reaction Kinetics

 

Objectives

Define temperature and state the units in which it is measured.

Define heat and state its units.

Perform specific heat calculations.

Explain heat of reaction, heat of formation, heat of combustion, and enthalpy change.

Solve problems involving heats of reaction, heats of formation, and heats of combustion.

  1. Section 17-1 Thermochemistry
    1. Defined- is the study of the transfers of energy as heat that accompany chemical reactions and physical changes.
    2. Heat and Temperature
      1. The energy absorbed or released as heat in a chemical or physical change is measured in a calorimeter.
        1. An enclosed chamber
        2. Measures heat gained and heat loss

         

      2. Temperature is a measure of the average kinetic energy of the particles in a sample of matter
        1. SI unit Joule (J)
        2. N x m = (Kg x m2)/s2
      3. Heat can be thought of as the energy transferred between samples of matter because of a difference in their temperatures.
    3. Heat capacity and Specific Heat
      1. What affects the amount of heat exchanged?
        1. Nature of material
        2. Mass of material
        3. Size of temperature change
      2. Specific Heat is the amount of energy required to raise the temperature of one gram of substance by one-Celsius degree or one Kelvin.
        1. Values- joules per gram per Celsius
        2. J/ (g oC)
        3. Measured under constant pressure
        4. symbol cp
        5. Table 17-1

          6.  

           

           

          Specific Heats of Some Common Substances at 298.15K

             
             
             
             
             
             
             
             
             
             
             
             
             
             
             
             

        6. Equation

          7. cp=q/(m D T) or q= m cp D T

        7. Practice

       

       

       

       

    4. Heats of reaction
      1. Defined- is the quantity of energy released or absorbed as heat during a chemical reaction.
      2. Expressed as a
        1. Thermochemical equation- an equation that includes the quantity of energy released or absorbed as heat during the reaction as written.
        2. Coefficient- interpreted as number or moles
        3. 2H2(g) + O2(g) -> 2H2O(g) + 483.6 KJ
        4. Physical state must be included
        5. The energy absorbed or lost is defined as ENTHALPY, D H
          1. Enthalpy change- is the amount of energy absorbed or lost by a system as heat during a process at constant pressure.
          2. D H = H products –H reactants
          3. D H is positive then endothermic
          4. D H is negative then exothermic
          5. See figure 17-2 & 17-3

           

        6. Important facts
          1. Coefficients represent moles
          2. Physical state involved is important
          3. Energy represented is proportional to the number or moles
          4. Value is not usually influenced by changing temperature
    5. Heat or formation- is the energy released or absorbed as heat when one mole of a compound is formed by combination of its elements.
    6. Stability and heat of formation
      1. Large amount of energy released
        1. Negative heat of formation
        2. Very stable
        3. Once they start they continue
        4. Ex. Carbon dioxide is –393.5 so it is more stable than its elements.
      2. Small amounts of energy released
        1. Slightly negative
        2. Relatively unstable
        3. Will spontaneously decompose into its elements

         

         

      3. Large amounts of energy absorbed
        1. High positive heat of formation
        2. Very unstable
        3. May react or decompose violently
        4. Used as detonators for explosives
    7. Heat of combustion-is the energy released as heat by the complete combustion of one mole of a substance.
    8. Calculation Heats of Reaction
      1. 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.
      2. Heat loss = Heat gained

     

     

     

     

     

     

     

     

     

     

     

     

     

     

    Objectives

    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.

  2. Section 17-2 Driving Force of Reactions
    1. Enthalpy and reaction tendency
      1. Tendency to move toward lower energy
    2. Entropy and reaction tendency
      1. Tendency to move toward more disorder without enthalpy change
      2. Entropy, S, can be defined as a measure of the degree of randomness of the particles, such as molecules, in a system.
    3. Free energy
      1. G, is the combined enthalpy-entropy of a system

        2.  

      2. Free energy change-is defined as the difference between the change in enthalpy and the product of the Kelvin temperature and the entropy change which is defined as TD S.
      3. Equation

        4. D G0 = D H0 - TD 0

      4. Value of G
        1. Positive- reaction will not occur
        2. Negative- reaction will occur
      5. Table 17-2 Relating Enthalpy, Entropy, and Free energy changes to Reaction Occurrence

    D H

    D S

    D G
         
         
         
         

     

     

     

     

     

     

     

     

     

    Objectives

    Explain the concept of reaction mechanism.

    Use the collision theory to interpret chemical reactions.

    Define activated complex.

    Relate activation energy to heat of reaction.

  3. Section 17-3 The Reaction Process
    1. Reaction Mechanisms
      1. Defined as the step-by-step sequence of reactions by which the overall chemical change occurs.
        1. Intermediates- species that appear in some steps but not in the net equation.
        2. Homogeneous reaction-a reaction whose reactants and products exist in a single phase
    2. Collision Theory
      1. Defined-a set of assumptions regarding collisions and reactions.
      2. Collision theory shows two things
        1. Collision energy
        2. Colliding molecules orientation

       

       

    3. Activation Energy
      1. Defined as the minimum energy required to transform the reactants into an activated complex.
      2. Activated complex- a transitional structure that results from an effective collision and that persists while old bonds are breaking and new bonds are forming.
      3. See figure 17-11

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

    Objectives

    Define chemical kinetics, and explain the two conditions necessary for chemical reactions to occur.

    Discuss the five factors that influence reaction rate.

    Define catalyst, and discuss two different types.

    Explain and write rate laws for chemical reactions.

  4. Section 17-4 Reaction Rate
    1. Reaction rate- the change in the concentration of reactants per unit time as a reaction proceeds.
    2. Chemical Kinetics- the area of chemistry that is concerned with reaction rates and reaction mechanisms.
    3. Rate-influencing factors
      1. Nature of reactants
      2. Surface area
      3. Temperature
      4. Concentration
      5. Presence of catalysts
        1. Catalyst- is a substance that changes the rate of a chemical reaction without itself being permanently consumed.
        2. Catalysis- is the action of a catalyst.

       

       

    4. Rate Laws for Reactions
      1. Rate law is an equation that relates reaction rate and concentrations of reactants.