= Solution
With the electrodes disconnected, the copper has no sustained external route through which to receive electrons from the zinc. It is not appreciably attacked by a clean, dilute nonoxidizing <sulfuric acid> solution. Zinc can in principle reduce acid while dissolving, but amalgamation suppresses much of the local action associated with impurities and promotes a high kinetic barrier for hydrogen evolution on the amalgamated surface. Thus the intended historical observation is \b[little visible action with the circuit open], not a thermodynamic assertion that zinc cannot react with acid.
Closing the external wire supplies an electron path. At the zinc <anode>,
$$
\mathrm{Zn}\longrightarrow\mathrm{Zn}^{2+}+2e^-;
$$
at the copper <cathode>,
$$
2\mathrm H^++2e^-\longrightarrow\mathrm H_2.
$$
The net <redox> is therefore
$$
\boxed{\mathrm{Zn}+2\mathrm H^+\longrightarrow\mathrm{Zn}^{2+}+\mathrm H_2.}
$$
The zinc dissolves, and bubbles of <hydrogen> appear chiefly at the copper. Electrons flow through the wire from zinc to copper; conventional <electric current> flows oppositely there. Ionic transport through the <electrolyte> completes the circuit and prevents indefinite charge accumulation in the solution. The copper primarily provides a conducting surface: there is no stipulated copper-ion solution from which metallic copper must be deposited, so this is not the <Daniell cell> reaction.
The <chemical energy> supplies the electrical <work>. Accumulation of hydrogen and changing interfacial conditions can polarize the cell, reducing the current. A complete description therefore distinguishes the initial closed-circuit response from an indefinitely constant-current idealization. <Electrochemical overpotential> explains why a favorable reaction and its observable rate are different questions.
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