chemistry

How much energy is evolved when 2.65 mg of Cl(g) atoms adds electrons to give Cl1-(g) ions.

No clue how to go about solving this problem as it doesn't really relate to any of my examples in my text book or notes. All I can find that related to this problem is electron affinity and E=(K*Q1*Q2)/r where K is a constant. The EA for Cl is -349 kJ/mol. But I just cant find anything else out about this problem. Any help is great as I'm out of ideas.

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  1. What about delta Hfsup>o?
    For the reaction Cl + e ==> Cl^-, you can look up delta Hof of Cl and for Cl^- and delta H rxn = delta H products - delta H reactants, then convert from mol to 2.65 mg. My text gives +121.7 kJ/mol for deltHof Cl and -226 for Cl^-. Of course that's just delta H; but you could do the same for delta Go.

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  2. @DrBob222 - I understand up to the part about the conversion of mol to 2.65 mg. I understand that the enthalpy change is 121.3 kJ/mol for Cl(g) and that would be the reactant part of the equation. But whats the deal with working with 2.65mg and the delta Go?

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  3. The ionization (is always concerned with the formation of positive ions) of Chlorine is given by the following equation;
    Cl --> Cl+ + e- IE = 1251.19 KJ/mol

    While the electron affinity (is always concerned with the formation of negative ions) of Chlorine is given by the following equation;
    Cl + e- --> Cl- EA = -349 kJ/mol

    # moles of Cl = 0.00265 g/35.5 g/mol
    = 0.0000746479 mol
    Therefore, the energy involved in the reaction is E = EA x nCl
    = -0.026 kJ
    Exothermic reaction because the energy is released when the atom gains electrons

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