= Daniell cell
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A <Daniell cell> uses the <zinc>/$\mathrm{Zn}^{2+}$ and $\mathrm{Cu}^{2+}$/<copper> pairs. In the separated-compartment teaching version, a <zinc> <electrode> contacts an aqueous <chemical solution> containing $\mathrm{Zn}^{2+}$, and a <copper> <electrode> contacts one containing $\mathrm{Cu}^{2+}$. A <salt bridge> supplies an ionic connection; a wire and load supply an electronic connection. During discharge,
$$
\begin{aligned}
\text{anode, oxidation:}\quad&\mathrm{Zn}(s)\longrightarrow\mathrm{Zn}^{2+}(aq)+2e^-,\\
\text{cathode, reduction:}\quad&\mathrm{Cu}^{2+}(aq)+2e^-\longrightarrow\mathrm{Cu}(s),\\
\text{net:}\quad&\mathrm{Zn}(s)+\mathrm{Cu}^{2+}(aq)\longrightarrow\mathrm{Zn}^{2+}(aq)+\mathrm{Cu}(s).
\end{aligned}
$$
The <zinc> <electrode> loses material and <copper> is deposited at the <cathode>. Each <half-reaction> conserves <chemical atom> counts and <electric charge>: the first has net charge zero on each side, the second $+2e-2e=0$ on the left and zero on the right. The net reaction has $+2e$ on both sides, with no <Electrons> left in its <chemical equation>. The Zn/Cu teaching cell and these <half-reactions> are described in https://www.chem.fsu.edu/chemlab/chm1lab/redox/background.html[Florida State University, Redox Reactions in Voltaic Cells: Construction of a Potential Series, lab background].
Creating one <mole> of $\mathrm{Zn}^{2+}$ adds $2F$ of positive <electric charge> to the <anode> compartment's ionic inventory; consuming one <mole> of $\mathrm{Cu}^{2+}$ removes $2F$ of positive <electric charge> from the <cathode> compartment. Here $F$ is the <Faraday constant>. In an ideal bookkeeping example using a <potassium> <nitrate> <salt bridge>, two <moles> of <nitrate> <anions> entering the <anode> compartment compensate the first change, and two <moles> of <potassium> <cations> entering the <cathode> compartment compensate the second. Other <ions> in the cell can share the actual transport. The two compensating inventory changes do not mean that $4F$ flowed through the load: the same series <electric current> transports $2F$ through each cross-section of the circuit per <mole> of net reaction. See <salt bridge>.
Under standard-state conditions near $298\,\mathrm K$, tabulated reduction potentials give a cell <electromotive force> of approximately $1.10\,\mathrm V$. This is a reversible reference value, not a promised loaded reading; actual composition matters through <thermodynamic activity>. The tabulated <copper> and <zinc> values and their subtraction are given in https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/electro.php[Purdue University, Electrochemical Reactions: standard-state reduction potentials]. For the net <chemical equation> as written, the molar <Gibbs free energy> change obeys
$$
\Delta_r G=-2F\mathcal E.
$$
Thus $\mathcal E=1.10\,\mathrm V$ corresponds to about $-212\,\mathrm{kJ\,mol^{-1}}$ of net reaction. At fixed <temperature> and <pressure>, $-\Delta_r G$ is the maximum reversible electrical <work> per <mole> for this cell reaction. It is not automatically all the <heat> released by a finite-current cell. See <Chemistry 2e by Paul Flowers and others/17.4 Potential, Free Energy, and Equilibrium> (§17.4: cell potential and free energy).
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