Unit 9: Thermodynamics and Electrochemistry

19 terms. Say what each one means, then open the card to check.

• Entropy (S)

  A measure of how spread out the energy and matter of a system are; more ways to arrange the particles and their energy means higher entropy.

  Example. A gram of water vapor has far more entropy than a gram of ice, because its molecules move freely through a much larger space.

• Sign of the entropy change

  Entropy usually rises when a solid melts, a liquid boils, a solid dissolves, or a reaction makes more moles of gas, and usually falls in the reverse changes.

  Example. For $2\text{H}_{2}\text{O}_{2}(l) \rightarrow 2\text{H}_{2}\text{O}(l) + \text{O}_{2}(g)$, making a gas from liquids gives a positive $\Delta S^\circ$.

• Standard entropy change

  For a reaction, $\Delta S^\circ$ equals the sum of the standard molar entropies of the products, each times its coefficient, minus the same sum for the reactants.

  Example. If products total 400 J/(mol·K) and reactants total 250 J/(mol·K), then $\Delta S^\circ = +150$ J/(mol·K).

• Gibbs free energy change

  The quantity $\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$; a negative value means the process is thermodynamically favorable under standard conditions.

  Example. With $\Delta H^\circ = -50$ kJ/mol and $\Delta S^\circ = -0.10$ kJ/(mol·K) at 300 K, $\Delta G^\circ = -20$ kJ/mol, so the process is favorable.

• Thermodynamically favorable

  A process with a negative $\Delta G$ that can proceed on its own without outside energy; it may still be too slow to observe.

  Example. Diamond turning into graphite is favorable at room temperature, but a high activation energy keeps it from happening in any human lifetime.

• Temperature and favorability

  When $\Delta H^\circ$ and $\Delta S^\circ$ have the same sign, the temperature decides the sign of $\Delta G^\circ$; the crossover is near $T = \Delta H^\circ/\Delta S^\circ$.

  Example. Ice melting has positive $\Delta H^\circ$ and positive $\Delta S^\circ$, so it is favorable only above 0 °C.

• Free energy and K

  The standard free energy change sets the equilibrium constant: $\Delta G^\circ = -RT\ln K$. A negative value means K is greater than 1.

  Example. A reaction with $\Delta G^\circ = 0$ has K = 1 at that temperature.

• Coupled reactions

  Pairing an unfavorable reaction with a strongly favorable one so that the sum of their $\Delta G^\circ$ values is negative and the combined process can run.

  Example. Cells drive many uphill reactions by linking them to the breakdown of ATP, which releases a large amount of free energy.

• Oxidation and reduction

  Oxidation is the loss of electrons and reduction is the gain of electrons; in a redox reaction one species does each, with the same number of electrons.

  Example. When iron rusts, iron atoms lose electrons (oxidation) and oxygen molecules gain them (reduction).

• Galvanic cell

  A cell in which a thermodynamically favorable redox reaction produces an electric current; its cell potential is positive.

  Example. A flashlight battery is a galvanic cell: its chemical reaction pushes electrons through the bulb.

• Anode and cathode

  Oxidation always happens at the anode and reduction always happens at the cathode, in both galvanic and electrolytic cells.

  Example. In a magnesium-iron cell, the magnesium strip is the anode and slowly loses mass, while iron builds up on the cathode.

• Salt bridge

  A tube of ion solution that joins the two half-cells and lets ions drift between them, keeping each side electrically neutral so current can keep flowing.

  Example. If the salt bridge is lifted out of a working cell, the voltmeter reading drops to zero, because charge builds up in each beaker and the current stops.

• Standard reduction potential

  The voltage of a half-reaction written as a reduction, measured against the standard hydrogen electrode; a more positive value means a stronger pull for electrons.

  Example. Gold(III) ion, at +1.50 V, is reduced far more readily than iron(II) ion, at -0.44 V.

• Standard cell potential

  The cathode potential minus the anode potential, both as reductions. Multiplying a half-reaction by a coefficient does not change its potential.

  Example. A cell with a cathode at +0.54 V and an anode at -0.44 V has $E^\circ_{\text{cell}} = 0.98$ V.

• Cell potential and free energy

  The link $\Delta G^\circ = -nFE^\circ_{\text{cell}}$, where n is the moles of electrons transferred and F is 96,485 C per mole of electrons; a positive cell potential means a negative $\Delta G^\circ$.

  Example. A cell that moves 2 mol of electrons at 1.00 V has $\Delta G^\circ = -193$ kJ per mole of reaction.

• Nonstandard conditions

  When Q is not 1, the cell potential shifts: Q below 1 raises it above the standard value, Q above 1 lowers it, and at equilibrium the potential is zero.

  Example. A cell set up with a very dilute product-ion solution and a concentrated reactant-ion solution reads a little above its standard potential, because Q is below 1.

• Electrolytic cell

  A cell in which an outside power source drives a redox reaction that is not thermodynamically favorable on its own.

  Example. Passing current through molten sodium chloride makes sodium metal at one electrode and chlorine gas at the other.

• Faraday’s constant

  The charge on one mole of electrons, 96,485 C/mol. Charge in coulombs equals current in amperes times time in seconds.

  Example. A current of 0.500 A for 386 s carries 193 C, which is 0.00200 mol of electrons.

• Electroplating

  Using electrolysis to reduce metal ions onto an object made the cathode, so a thin layer of metal coats it.

  Example. A copper ring is made the cathode in a bath of gold ions, and gold is reduced onto its surface as a thin layer.