Consider the standard galvanic cell based on the following half-reactions: $$ \begin{aligned} \mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu} \\ \mathrm{Ag}^{+}+\mathrm{e}^{-} \longrightarrow \mathrm{Ag} \end{aligned} $$

When delving into the world of electrochemistry, the concept of standard reduction potentials is critical. These potentials, often denoted as E°, are used to measure the tendency of a species to gain electrons, or reduce, under standard conditions, which are 25°C, 1 M concentration for each ion participating in the reaction, and a pressure of 1 atm for gases.

Standard reduction potentials are listed in tables for half-reactions, which show a reagent gaining electrons. A higher E° value indicates a greater likelihood of the species being reduced. In turn, these potentials are instrumental in predicting the direction of electron flow in a galvanic cell. To understand and predict galvanic cell outcomes, students should familiarize themselves with this concept and how to use standard reduction potential tables.

The heart of a galvanic cell lies in the anode and cathode reactions. In such cells, we have two separate half-reactions – oxidation at the anode and reduction at the cathode. The distinction between the two is straightforward: oxidation corresponds to the loss of electrons, while reduction is the gain of electrons.

Determining which reaction occurs at which electrode requires a look at the standard reduction potentials. The substance with the lower reduction potential will undergo oxidation at the anode; conversely, the substance with the higher potential will undergo reduction at the cathode. This concept is paramount when constructing a galvanic cell or trying to understand its operation. It helps predict the flow of electrons and the direction of chemical reactions.

To harness the power of a galvanic cell, we must know how to perform a cell potential calculation. The cell potential, E°(cell), signifies the electromotive force of the cell — essentially, it measures how much voltage the cell can produce. It is calculated by subtracting the anode's standard reduction potential from the cathode's:
\( E_{cell}^{\boldsymbol{\theta}} = E_{cathode}^{\boldsymbol{\theta}} - E_{anode}^{\boldsymbol{\theta}} \)

For a positive cell potential, the reaction is spontaneous, indicating that electrons will spontaneously flow from anode to cathode. Remember, a negative cell potential implies a non-spontaneous reaction under standard conditions. While this formula seems straightforward, it's important for learners to practice this calculation regularly using various half-reactions to become adept at predicting cell behavior.