C−Br Bond

  • 34. 34. Assume that the following reaction is a single step reaction in which a C−Br bond is broken

    as the C−I bond is formed. The heat of reaction is +38 kJ/mol.

    a. With reference to collision theory, describe the general process that takes place as this

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    reaction moves from reactants to products.

    b. List the three requirements that must be met before a reaction between I−and CH3Br is

    likely to take place.

    c. Explain why an I− ion and a CH3Br molecule must collide before a reaction can take

    place.

    d. Explain why, in the process of this reaction, it is usually necessary for the new C−I

    bonds to form at the same time as the C−Br bonds are broken. f. Explain why a collision between an I− ion and a CH3Br molecule must have a certain

    minimum energy (activation energy) in order to proceed to products.

    g. The activation energy for this reaction is 76 kJ/mol. Draw an energy diagram for this reaction, showing the relative energies of the reactants, the activated complex, and the products. Using arrows show the activation energy and heat of reaction.

    h. Is this reaction exothermic or endothermic? i. Explain why an I− ion and a CH3Br molecule must collide with the correct orientation

    if a reaction between them is going to be likely to take place.

    40. If both systems described by the energy diagrams in the previous problem are at the same

    temperature, if the concentrations of initial reactants are equivalent for each reaction, and if the

    orientation requirements for each reaction are about the same, which of these reactions would

    you expect to have the greatest forward reaction rate? Why?

    44. Reversible chemical reactions lead to dynamic equilibrium states. What is dynamic about

    these states? Why are they called equilibrium states?

    55. Predict whether each of the following reactions favor reactants, products, or neither at the

    temperature for which the equilibrium constant is given.

    a. 2COF2(g) CO2(g) + CF4(g) KP= 2 at 1000 °C

    b. ⅛S8(s) + O2(g) SO2(g) KP= 4.2 × 1052 at 25 °C

    c. C2H6(g) C2H4(g) + H2(g) KP= 1.2 × 10−18 at 25 °C

    63. Ethylene, C2H4, used to make polyethylene plastics, can be made from ethane, C2H6, one of

    the components of natural gas. The heat of reaction for the decomposition of ethane gas into

    ethylene gas and hydrogen gas is 136.94 kJ per mole of C2H4 formed, so it is endothermic. The

    reaction is run at high temperature, in part because at 800-900 °C, the equilibrium constant for

    the reaction is much higher, indicating that a higher percentage of products forms at this

    temperature. Explain why increased temperature drives this reversible chemical reaction in the

    endergonic direction and why this leads to an increase in the equilibrium constant for the

    reaction. C2H6(g) C2H4(g) + H2(g)

    67. Acetic acid, which is used to make many important compounds, is produced from methanol

    and carbon monoxide (which are themselves both derived from methane in natural gas) by a

    process called the Monsanto process. The endothermic reaction is run over a rhodium and iodine

    catalyst at 175 °C and 1 atm of pressure. Predict whether each of the following changes in the

    equilibrium system will shift the system to more products, to more reactants, or neither. Explain

    each answer in two ways, (1) by applying Le Chatelier’s principle and (2) by describing the

    effect of the change on the forward and reverse reaction rates. CH3OH(g) + CO(g) + 207.9 kJ CH3CO2H(g)Rh/I2 175 °C1 atm

    a. The concentration of CO is increased by the addition of more CO.

    b. The concentration of CH3OH is decreased.

    c. The concentration of CH3CO2H(g) is decreased by removing the acetic acid as it

    forms.

    d. The temperature is decreased from 300 °C to 175 °C.

    e. The Rh/I2 catalyst is added to the equilibrium system.

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