What makes a change favorable
Entropy describes how energy and matter can be distributed among accessible arrangements. Gibbs free energy combines enthalpy and entropy at constant temperature and pressure: a negative change indicates thermodynamic favorability in that setting, not a fast reaction. Temperature can change which term dominates. At equilibrium, the reaction’s free-energy change is zero under the current conditions.
An exothermic process is not automatically spontaneous. Free-energy change weighs enthalpy against the temperature-scaled entropy change for a specified process at constant temperature and pressure. A reaction that absorbs heat can still be favorable if dispersal increases sufficiently; one that releases heat can be unfavorable under some conditions. Always ask for the direction and current composition, not only the sign of one energy term.
Entropy is often described as disorder, but “number of accessible microscopic arrangements” is more useful. Dissolving, mixing, phase changes, and changes in gas amount can alter that count. The surroundings matter too: a local decrease in entropy can occur while the total entropy of system plus surroundings increases. Thermodynamics identifies an equilibrium tendency, not a microscopic movie of how fast molecules rearrange.
Building a cell
Oxidation loses electrons at the anode; reduction gains electrons at the cathode. In a galvanic cell, a favorable redox reaction drives electron flow through an external circuit while ions move to maintain charge balance. In an electrolytic cell, an external source drives an otherwise unfavorable process. The sign of the electrodes can differ between these cell types, but anode is always oxidation.
At the anode, oxidation produces electrons; at the cathode, reduction consumes them. Those definitions stay fixed even when the signs of the electrodes change between a spontaneous galvanic cell and a driven electrolytic cell. A positive standard cell potential for a reaction as written corresponds to favorable standard-state electrical work. Reverse the reaction and the cell potential changes sign.
A functioning cell needs more than a wire: ions must move internally to avoid charge separation that would halt electron flow. A salt bridge, porous separator, or ion-conducting membrane serves this role. The chemical reaction does not send one electron all the way through the electrolyte as if it were another wire. In a battery, capacity and usable voltage also depend on materials, concentration, internal resistance, and side reactions.
Voltage and conditions
Cell potential measures electrical work available per unit charge under specified conditions. Standard potential uses standard-state activities; changing concentrations changes the potential. The Nernst relationship links potential to reaction quotient, and free energy connects voltage to chemical driving force.
Standard reduction potentials are tabulated as reductions. To calculate a standard cell potential, use the cathode reduction value minus the anode reduction value. Do not multiply a voltage by stoichiometric coefficients when balancing electrons: potential is energy per unit charge, not total energy. Free energy, by contrast, scales with the amount of reaction and the number of electrons transferred.
Changing ion concentrations alters the reaction quotient and therefore the cell potential. The Nernst relationship formalizes that dependence. As a galvanic cell runs, reactants are used and products accumulate, so its voltage can change even without a change in temperature. At equilibrium, the reaction's net driving force and ideal cell potential for that composition reach zero.
Free energy and equilibrium
The standard free-energy change relates to the equilibrium constant: a product-favored equilibrium corresponds to a negative standard change. Under nonstandard conditions, the reaction quotient changes the current driving force. This is why a reaction mixture can move back toward reactants even when the standard-state equation appears favorable in the forward direction.
The standard free-energy change and K describe the same reaction under a common reference temperature. If the standard change is strongly negative, K is large for the forward reaction; if positive, K is small. But the actual free-energy change also depends on Q, the current product-to-reactant activity ratio. A mixture rich in products may move backward even when the standard comparison favors forward change.
This distinction is the chemical version of a hill whose slope depends on where you stand. Standard conditions specify one reference point; they do not dictate the direction from every possible composition. When Q equals K, the reaction is at equilibrium and the current free-energy change is zero. This connects the equilibrium chapter to electrochemical voltage without treating either equation as an isolated fact.
Tracking charge through a device
Electrons move through an external wire from the anode toward the cathode in a galvanic cell. A salt bridge or membrane lets ions move to prevent charge buildup in the solutions; electrons do not cross through the salt bridge. During electrolysis, electrical energy forces the net reaction in the opposite direction. Faraday’s constant connects moles of electrons to total charge, allowing a current and time to predict material deposited or consumed.
Oxidation at the anode releases electrons to the external circuit in a galvanic cell. The cathode receives them for reduction. Positive ions may travel toward a cathode compartment and negative ions toward an anode compartment through a salt bridge, depending on the changing charge balance. The direction of ionic motion cannot be replaced by the slogan “electrons move through the bridge.”
In electrolysis, the power supply drives electrons and ions in ways that sustain an otherwise unfavorable net reaction. The amount of product depends on charge passed, the number of electrons required per product unit, and efficiency if competing reactions occur. Current multiplied by time gives charge, and Faraday's constant converts that charge to moles of electrons. This is why plating thickness can be related to operating conditions, at least under controlled assumptions.