The reduction potential is a measure of the tendency of the oxidizing agent to be reduced. Its value is zero for H+ + e− →  H2 by definition, positive for oxidizing agents stronger than H+ (e.g., +2.866 V for F2) and negative for oxidizing agents that are weaker than H+ (e.g., −0.763 V for Zn2+). For a redox reaction that takes place in a cell, the potential difference is:  However, the potential of the reaction at the anode was sometimes expressed as an "oxidation potential": The oxidation potential is a measure of the tendency of the reducing agent to be oxidized, but does not represent the physical potential at an electrode. With this notation, the cell voltage equation is written with a plus sign  In the reaction between hydrogen and fluorine, hydrogen is being oxidized and fluorine is being reduced:  This reaction is spontaneous and releases 542 kJ per 2 g of hydrogen because the H-F bond is much stronger than the weak, high-energy F-F bond. We can write this overall reaction as two half-reactions:  the oxidation reaction:  and the reduction reaction:  Analyzing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above). Elements, even in molecular form, always have an oxidation state of zero. In the first half-reaction, hydrogen is oxidized from an oxidation state of zero to an oxidation state of +1. In the second half-reaction, fluorine is reduced from an oxidation state of zero to an oxidation state of −1. When adding the reactions together the electrons are canceled:  And the ions combine to form hydrogen fluoride:  The overall reaction is:  In this type of reaction, a metal atom in a compound (or in a solution) is replaced by an atom of another metal. For example, copper is deposited when zinc metal is placed in a copper(II) sulfate solution:  Zn(s)+ CuSO4(aq) → ZnSO4(aq) + Cu(s)  In the above reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc because relative to zinc, copper metal is lower in energy due to bonding via its partially filled d-orbitals.