AP®︎ Chemistry: Unit 9 Flashcards

A measure of matter and energy dispersal, or the number of possible arrangements.

Positive when entropy increases, negative when entropy decreases, and zero when disorder is unchanged.

Solids have lowest entropy, liquids intermediate, and gases highest because particle freedom increases.

Melting, vaporization, and sublimation increase entropy; freezing, condensation, and deposition decrease it.

Entropy increases when particles spread out more freely through space or occupy larger volumes.

Increasing a gas's volume raises entropy because particles can occupy a larger space.

Entropy usually increases when products contain more total moles of gas than reactants.

In an isolated system, entropy tends to increase or remain constant for spontaneous processes.

Higher temperature increases entropy because kinetic energy is spread over a wider range.

The entropy of 1 mole of a substance at standard conditions, measured absolutely.

A property whose change depends only on initial and final states, not pathway.

Positive means greater disorder; negative means greater order or fewer accessible microstates.

Entropy usually increases with more gas particles, phase changes toward gas, or more dispersed matter.

Calculate by summing coefficient-weighted product entropies and subtracting coefficient-weighted reactant entropies.

The free energy change for a process when all substances are in standard states.

The free energy change when one mole of a compound forms from elements in standard states.

Add product formation values, subtract reactant formation values, and include stoichiometric coefficients.

Any element in its standard state has a formation free energy of zero.

Exergonic processes have ΔG° < 0; endergonic processes have ΔG° > 0.

Exothermic and entropy-decreasing, so it is favored only at low temperatures.

Endothermic but entropy-increasing, so it can be favored at sufficiently high temperatures.

Free energy equals ΔH° − TΔS°, so ΔH° and ΔS° signs predict favorability.

A situation where reaction rate, often limited by high activation energy, prevents noticeable product formation.

A negative ΔG means a process is favorable, but it may still occur extremely slowly.

Reaction behavior determined mainly by the free-energy difference between reactants and products.

A system may have a favorable forward process yet remain far from equilibrium if rate is negligible.

A catalyst provides a lower-energy pathway, increasing rate without changing overall thermodynamic favorability.

A spontaneous carbon allotrope change with negative ΔG that occurs extremely slowly because of a large barrier.

A usually slow favorable reaction that becomes measurable when iodide lowers the activation energy.

A large activation-energy barrier can keep a thermodynamically favored reaction extremely slow.

Kinetic: forward and reverse rates are equal; thermodynamic: free energy is at a minimum.

ΔG is free energy change under current conditions; ΔG° is free energy change under standard conditions.

Use ΔG = ΔG° + RT ln Q to relate nonstandard free energy change to reaction conditions.

ΔG° = -RT ln K and K = e^(-ΔG°/RT), linking standard free energy to equilibrium position.

Equilibrium has minimum free energy, while K > 1 favors products and K < 1 favors reactants.

Negative ΔG° gives K > 1, positive ΔG° gives K < 1, and ΔG° near zero gives K near 1.

A process is thermodynamically favored when ΔG° < 0 under standard conditions, meaning products are favored at equilibrium; this was historically called spontaneous.

The overall Gibbs free energy change when a solute dissolves in a solvent.

Lattice enthalpy is positive to separate ions; hydration enthalpy is negative when water surrounds ions.

Negative ΔH releases heat and warms solution; positive ΔH absorbs heat and cools solution.

Breaking a solid raises disorder, while forming organized solvent shells often lowers disorder.

Negative ΔH and positive ΔS always favor dissolving; positive ΔH and negative ΔS never do.

Higher temperature favors dissolution when ΔH is positive and disfavors it when ΔH is negative.

Small, highly charged ions increase both lattice strength and hydration strength, creating competing effects.

Large opposing enthalpy and entropy contributions often cancel, making overall free energy hard to predict.

Lattice breaking, solvent reorganization, and solute-solvent interactions together determine overall dissolution free energy.

The overall heat and disorder changes together determine whether dissolving is thermodynamically favored.

Dissolving is thermodynamically favored when ΔG = ΔH - TΔS is negative.

A process with positive ΔG that does not occur on its own.

Added energy, such as electricity or light, used to drive a nonspontaneous process.

A spontaneous and nonspontaneous reaction linked by shared intermediates to give an overall negative ΔG.

A species produced in one step and consumed in another, canceling in the overall reaction.

Add the ΔG values of the individual reactions; the total must be negative.

When a reaction is multiplied, multiply its ΔG by the same factor.

Using electrical energy to reverse the spontaneous cell reaction and store chemical energy.

Light energy drives the overall conversion of carbon dioxide and water into glucose and oxygen.

ATP hydrolysis releases free energy that can drive otherwise unfavorable biological reactions.

The study of redox reactions that convert chemical energy and electrical energy into each other.

A device in which separated half-reactions transfer electrons through a wire and ions through solution.

Separate oxidation and reduction equations that show electron loss and gain in a redox process.

Tabulated voltages for reduction half-reactions under standard conditions, used to calculate overall cell voltage.

Electrons move through the wire from the anode to the cathode.

Solutions containing the ions involved in each electrode reaction, allowing charge transfer and reaction progress.

An external device connected to the circuit to measure the cell's potential difference or current.

Gas forms at an electrode when a half-reaction produces a gaseous product, often seen as bubbles.

Overall voltage equals cathode reduction potential minus anode reduction potential; reversing changes sign only.

A tube of inert ions that maintains neutrality by sending anions to the anode and cations to the cathode.

Conductive surfaces where redox occurs; some are inert and only provide electron-transfer surfaces.

The flow of positive charge through a circuit, opposite the direction of electron flow.

Oxidation occurs at the anode and reduction occurs at the cathode in all electrochemical cells; electrons flow from anode to cathode.

Galvanic cells are spontaneous and generate electrical energy, whereas electrolytic cells are nonspontaneous and require external electrical energy.

The voltage that drives electrons from anode to cathode in an electrochemical cell.

Use E degrees cell = E degrees cathode minus E degrees anode, with both values as reductions.

If a reduction half-reaction is reversed to oxidation, its listed potential changes sign.

The voltage for a half-reaction written as a reduction under standard conditions.

The reference half-cell assigned 0.00 V, used to define all reduction potentials.

The reducing agent is oxidized at the anode, and the oxidizing agent is reduced at the cathode.

A positive standard cell potential means a spontaneous reaction, while a negative one requires external energy.

Standard free energy change equals −nFE°, so positive cell potential gives negative ΔG°.

Standard cell potential, E°cell, is the voltage of an electrochemical cell under standard conditions: 298 K, 1 atm, 1 M dissolved species, and Q = 1.

Any cell conditions where concentrations, pressures, temperature, or Q differ from standard values.

E°cell is the voltage at standard conditions; Ecell is the actual voltage under current conditions.

Q is products over reactants using activities or concentrations, omitting pure solids and liquids.

Q = 1 at standard conditions; Q < 1 means reactant-favored; Q > 1 means product-favored.

Voltage measures the tendency for a redox reaction to proceed until equilibrium is reached.

The farther a system is from equilibrium, the larger the absolute value of its voltage.

If Q < 1, Ecell > E°cell; if Q > 1, Ecell < E°cell; if Q = 1, Ecell = E°cell.

Changes moving Q toward K lower |Ecell|, while changes moving Q away from K raise |Ecell|.

Changing ion concentrations changes Q, which changes the cell voltage under nonstandard conditions.

Their activities are treated as 1, so they are omitted from the reaction quotient.

Changing the amount or surface area of a pure solid electrode does not change Q or voltage.

Electrons flow from the lower-ion-concentration half-cell toward the higher-ion-concentration half-cell to equalize concentrations.

Substituting equilibrium conditions gives Ecell = 0 and connects standard voltage to the equilibrium constant.

A functioning electrochemical cell is not at equilibrium, so Le Châtelier’s principle does not apply.

At equilibrium, Q equals K and the cell potential is zero.

It relates cell potential to Q, letting you predict voltage changes qualitatively under nonstandard conditions.

Standard conditions are 298 K, 1 atm for gases, 1 M for dissolved species, and for electrochemical cells correspond to Q = 1.

Galvanic cells are spontaneous and produce electricity; electrolytic cells are nonspontaneous and require electricity.

The applied voltage must exceed the magnitude of the negative cell potential to force the reaction.

The amount of substance produced or consumed is proportional to the total charge passed.

Current equals charge divided by time, with I = q/t.

A unit of current equal to one coulomb of charge passing per second.

Use balanced half-reactions to relate moles of electrons to moles of substance produced or consumed.

Use unit conversions among amps, seconds, coulombs, electrons, moles, and grams.

An electrochemical cell that uses external electrical energy to drive a nonspontaneous redox reaction.

The Faraday constant is 96,485 coulombs per mole of electrons and is used to relate charge, moles of electrons, and free energy in equations such as ΔG° = -nFE°.

During electrolysis or other electrochemical processes involving solid electrodes, the anode loses mass as oxidation occurs and the cathode gains mass as reduction deposits material.