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spontaneous process

A spontaneous process is a physical or chemical change that occurs without the need for external influence or energy input, driven by the system's tendency to minimize its free energy. Such processes naturally proceed in the direction that increases the disorder or entropy of the system.

entropy (S)

A measure of the disorder or randomness in a system, often associated with the number of microscopic configurations that correspond to a macroscopic state.

second law of thermodynamics

The principle states that the total entropy of an isolated system can never decrease over time, and it will tend to increase, leading to irreversible processes.

reversible process

A reversible process is a theoretical thermodynamic process that can be reversed without leaving any net change in the system or surroundings, maintaining equilibrium throughout the process. It occurs infinitely slowly and does not increase the total entropy of the universe.

Gibbs free energy (G)

A thermodynamic potential that measures the maximum reversible work obtainable from a system at constant temperature and pressure. It is defined as the difference between the enthalpy and the product of the temperature and entropy.

standard entropy change for a

reaction ( DeltaS rxn)

The standard entropy change for a reaction is defined as the difference in the standard entropies of the products and the reactants, reflecting the disorder of the system at standard conditions.

standard molar

entropies (S°)

The standard molar entropies (S°) are the entropies of one mole of a substance under standard conditions, typically measured at a temperature of 298.15 K and 1 atm pressure. These values indicate the degree of disorder or randomness associated with the substance.

third law of

thermodynamics

States that the entropy of a perfect crystal approaches zero as the temperature approaches absolute zero, implying that absolute zero is unattainable.

standard change in free

energy (DeltaG rxn)

The standard change in free energy (ΔG rxn) is defined as the difference in free energies of the products and reactants of a chemical reaction under standard conditions, which indicates the spontaneity and equilibrium of the reaction.

free energy of formation

(Delta Gf)

The free energy of formation (ΔG_f) is the change in free energy when one mole of a compound is formed from its elements in their standard states. It reflects the stability and spontaneity of the compound's formation.

free energy change of a

reaction under nonstandard

condition (DeltaG rxn)

The free energy change of a reaction under nonstandard conditions (ΔG rxn) refers to the change in free energy when reactants and products are in concentrations, pressures, or temperatures that differ from standard conditions. It determines the direction and extent of a chemical reaction.

Spontaneous and Nonspontaneous Processes

Both spontaneous and nonspontaneous processes can occur, but only spontaneous processes can take place without outside intervention.

Spontaneous and Nonspontaneous

Processes pt. 2

Thermodynamics is the study of the spontaneity of reactions, not to be confused with kinetics, the study of the rate of reactions.

Entropy and the Second Law of

Thermodynamics

The second law of thermodynamics states that for any spontaneous process, the entropy of the universe increases.

Entropy and the Second Law of

Thermodynamics pt. 2

Entropy (S) is proportional to the number of energetically equivalent

ways in which the components of a system can be arranged

and is a measure of energy dispersal per unit temperature.

The Entropy Changes Associated with Changes

of State

The entropy of a substance increases as it changes state from a solid

to a liquid and when it changes from a liquid to a gas.

The Entropy Changes Associated with Changes

of State pt. 2

The change in entropy for system undergoing a reversible change

in state is DeltaS = qrev/T.

The Entropy Changes Associated with Changes

of State pt. 3

Standard entropies are absolute; an entropy of zero is established

by the third law of thermodynamics as the entropy of a perfect

crystal at absolute zero.

The Entropy Changes Associated with Changes

of State pt. 4

The entropy of a substance at a given temperature depends on

factors that affect the number of energetically equivalent arrangements

of the substance; these include the state, size, and molecular

complexity of the substance.

Free Energy Changes in Chemical Reactions:

Calculating DeltaGrxn

There are three ways to calculate DeltaGrxn n: (1) from DeltaH and DeltaS,

(2) from free energies of formation (only at 25°C), and (3) from

the delta G 's of reactions that sum to the reaction of interest

Free Energy Changes in Chemical Reactions:

Calculating DeltaGrxn pt.2

The magnitude of a negative Delta G rxn represents the theoretical

amount of energy available to do work, while a positive delta G rxn represents

the minimum amount of energy required to make a nonspontaneous

process occurs.

Heat Transfer and Changes in the Entropy of the

Surroundings

For a process to be spontaneous, the total entropy of the universe

(system plus surroundings) must increase.

Heat Transfer and Changes in the Entropy of the

Surroundings pt.2

The entropy of the surroundings increases when the change in

enthalpy of the system (deltaHsys) is negative (i.e., for exothermic reactions)

Heat Transfer and Changes in the Entropy of the

Surroundings pt.3

The change in entropy of the surroundings for a given 6Hsys depends

inversely on temperature—the greater the temperature, the

smaller the magnitude of deltaS surr

Gibbs Free Energy

Gibbs free energy, G, is a thermodynamic function that is proportional

to the negative of the change in the entropy of the universe.

A negative deltaG represents a spontaneous reaction, and a positive deltaG represents a nonspontaneous reaction.

Gibbs Free Energy pt.2

We can calculate the value of deltaG for a reaction from the values of

deltaH and deltaS for the system using the equation deltaG = deltaH - TdeltaS.

Entropy Changes in Chemical Reactions: Calculating deltaSrxn

We calculate the standard change in entropy for a reaction similarly

to the way we calculate the standard change in enthalpy for a reaction: by subtracting the sum of the standard entropies of the

reactants multiplied by their stoichiometric coefficients from the

sum of the standard entropies of the products multiplied by their

stoichiometric coefficients.

Free Energy Changes for Nonstandard States:

The Relationship between Delta*Grxn and DeltaGrxn

The value of delta*Grxn applies only to standard conditions, and most

Real conditions are not standard.

Free Energy Changes for Nonstandard States:

The Relationship between Delta*Grxn and DeltaGrxn pt.2

Under nonstandard conditions, we can calculate t!.Grxn from the

equation DeltaGxn = Delta*Gxn + RT In Q.

Free Energy and Equilibrium: Relating DeltaG*rxn to

the Equilibrium Constant

Under standard conditions, the free energy change for a reaction

is directly proportional to the negative of the natural log of the

equilibrium constant, K; the more negative the free energy change

(i.e., the more spontaneous the reaction), the larger the equilibrium

constant.

Free Energy and Equilibrium: Relating DeltaG*rxn to

the Equilibrium Constant pt.2

We can use the temperature dependence of DeltaG*rxn, as given by

DeltaG*=DeltaH*-TDeltaS*, to derive an equation for the temperature

dependence of the equilibrium constant.

electrical current

The flow of electric charge in a circuit, typically measured in amperes. It represents the rate at which charge is transferred through a conductor.

electrochemical cell

A device that converts chemical energy into electrical energy through redox reactions, typically consisting of two electrodes submerged in an electrolyte.

voltaic (galvanic) cell

A type of electrochemical cell that generates electrical energy from spontaneous chemical reactions, converting chemical energy directly into electrical energy.

half-cell

A portion of an electrochemical cell, consisting of a conductive electrode and an electrolyte, that participates in electrochemical reactions.

electrode

A conductive component in an electrochemical cell that facilitates the transfer of electrons during redox reactions.

ampere (A)

The unit of electric current, defined as one coulomb per second, used to measure the flow of electric charge.

potential difference

The difference in electric potential energy per unit charge between two points in an electric circuit, commonly measured in volts.

volt (V)

The unit of electric potential, defined as the potential difference that will drive one ampere of current against one ohm of resistance. It is commonly used in electrical systems to measure the energy per unit charge.

electromotive force (emf)

The energy provided per unit charge by a source of electrical energy, such as a battery or generator, measured in volts. It represents the potential energy available to drive electric current in a circuit.

cell potential (cell emf)

(Ecell)

The measure of the potential difference between the electrodes of a galvanic or electrolytic cell, indicating the cell's ability to drive an electric current.

standard cell potential

(standard emf) ( E*cell)

The measure of cell potential under standard conditions, typically at 25°C and 1 M concentration, reflecting the maximum voltage obtainable from a galvanic cell.

anode

The electrode where oxidation occurs in an electrochemical cell, responsible for supplying electrons to the external circuit.

cathode

The electrode where reduction occurs in an electrochemical cell, receiving electrons from the external circuit.

salt bridge

A device used in electrochemical cells to maintain electrical neutrality by allowing ions to flow between the two half-cells, preventing the buildup of charge that would impede the reaction.

standard electrode potential

The measure of the tendency of an electrode to gain or lose electrons, expressed in volts, under standard conditions.

standard hydrogen electrode

(SHE)

is a reference electrode used to measure the standard electrode potential of other electrodes, defined as having a potential of 0.00 volts at all temperatures.

Faraday's constant (F)

The charge per mole of electrons, approximately equal to $96485 ext{ C mol}^{-1}$, used in electrochemical calculations.

Nernst equation

A mathematical relationship that relates the reduction potential of a half-cell to the standard electrode potential and the concentration of ions in solution, often expressed as $E = E^\circ - \frac{RT}{nF} \ln(Q)$.

dry-cell battery

A type of electrochemical cell that consists of a paste electrolyte and is commonly used in portable devices. It has a longer shelf life compared to wet batteries and is non-leakable.

alkaline battery

A type of dry battery that uses an alkaline electrolyte, typically potassium hydroxide, resulting in a higher energy density and longer shelf life than standard batteries.

lead- acid storage

battery

A type of rechargeable battery that uses lead dioxide and sponge lead as electrodes, with sulfuric acid as the electrolyte. It is commonly used in automobiles due to its ability to provide high surge currents.

nickel-cadmium (NiCad)

battery

A type of rechargeable battery that uses nickel oxide hydroxide and metallic cadmium as electrodes. It is known for its ability to perform well in low temperatures and can be recharged many times.

nickel-metal hydride (NiMH)

battery

A type of rechargeable battery that utilizes nickel metal hydride as one of the electrodes. It offers a higher capacity than nickel-cadmium batteries and is commonly used in hybrid vehicles and portable electronics.

lithium ion battery

A type of rechargeable battery that uses lithium ions as the primary charge carriers. It is known for its high energy density, lightweight, and is widely used in mobile devices and electric vehicles.

fuel cell

A device that converts the chemical energy of a fuel, typically hydrogen, directly into electricity through a chemical reaction with oxygen, producing water as a byproduct. Fuel cells are used in various applications, including vehicles and stationary power generation.

electrolysis

A chemical process that uses electricity to drive a non-spontaneous reaction, often used to split water into hydrogen and oxygen.

corrosion

The process of deterioration of materials, typically metals, due to chemical reactions with their environment, often involving oxidation. Corrosion can cause significant structural damage over time.

Pulling the Plug on the Power Grid

■ Oxidation-reduction reactions are reactions in which electrons

are transferred from one reactant to another.

■ In the most common form of fuel cell, an electrical current is created

as hydrogen is oxidized and oxygen is reduced, water is the

The only product.

Balancing Oxidation-Reduction Equations

■ Oxidation is the loss of electrons and corresponds to an increase

in oxidation state; reduction is the gain of electrons and corresponds

to a decrease in oxidation state.

■ We can balance redox reactions using the half-reaction method.

in which the oxidation and reduction reactions are balanced separately

and then added. This method differs slightly for redox reactions.

in acidic and in basic solutions.

Voltaic (or Galvanic) Cells: Generating Electricity

from Spontaneous Chemical Reactions

■ A voltaic electrochemical cell separates the reactants of a spontaneous redox reaction into two half-cells that are connected by a wire and a means to exchange ions so that electricity is generated.

■ In an electrochemical cell, the electrode where oxidation occurs

is the anode and the electrode where reduction occurs is the cathode;

Electrons flow from the anode to the cathode.

■ The rate of electrons flowing through a wire is measured in amperes

(A), and the cell potential is measured in volts (V).

■ A salt bridge allows ions to flow between the half-cell solutions

and prevent the buildup of charge.

■ The cell diagram or line notation symbolizes electrochemical cells

concisely by separating the components of the reaction using

lines or commas.

Standard Electrode Potentials

■ The electrode potentials of half-cells are measured in relation to

that of a standard hydrogen electrode, which is assigned an electrode

potential of zero under standard conditions (solution concentrations

of 1 M, gas pressures of 1 atm, and a temperature of 25 °C).

■ A species with a highly positive E₀ has a strong tendency to attract

electrons and undergo reduction (and is therefore an excellent

oxidizing agent).

■ A species with a highly negative E₀ has a strong tendency to repel

electrons and undergo oxidation (and is therefore an excellent

reducing agent).

Cell Potential, Free Energy, and the Equilibrium

Constant

In a spontaneous reaction, Ecell* is positive, the change in free energy

(DeltaG*) is negative, and the equilibrium constant (K) is greater than 1.

■ In a nonspontaneous reaction, Ecell* is negative, DeltaG* is positive,

and K is less than 1.

■ Because Ecell, DeltaG*, and K all relate to spontaneity, we can derive

equations relating all three quantities.

Cell Potential and Concentration

The standard cell potential (E*cell) is related to the cell potential

(Ecell) by the Nernst equation, E = E*cell- (0.0592 V/n) log Q.

Cell Potential and Concentration

As shown by the Nernst equation, Ecell is related to the reaction

quotient (Q); Ecell equals zero when Q equals K.

Cell Potential and Concentration pt.2

In a concentration cell, the reactions at both electrodes are identical

and electrons flow because of a difference in concentration.

Nerve cells are a biological example of concentration cells.

Batteries: Using Chemistry to Generate

Electricity

Batteries are packaged voltaic cells

Batteries: Using Chemistry to Generate

Electricity pt.2

Dry-cell batteries, including alkaline batteries, do not contain

large amounts of water.

Batteries: Using Chemistry to Generate

Electricity pt.3

The reactions in rechargeable batteries, such as lead- acid storage,

nickel-cadmium, nickel-metal hydride, and lithium ion batteries,

can be reversed.

Batteries: Using Chemistry to Generate

Electricity pt.4

Fuel cells are similar to batteries except that fuel cell reactants

must be continually replenished from an external source

Electrolysis: Driving Nonspontaneous Chemical

Reactions with Electricity

An electrolytic electrochemical cell differs from a voltaic cell

in that (1) an electrical charge is used to drive the reaction, and

(2) although the anode is still the site of oxidation and the cathode the site of reduction, they are represented with signs opposite

those of a voltaic cell (anode+, cathode-).

Electrolysis: Driving Nonspontaneous Chemical

Reactions with Electricity pt.2

In electrolysis reactions, the anion is oxidized; if there is more

than one anion, the anion with the more negative E* is oxidized.

Electrolysis: Driving Nonspontaneous Chemical

Reactions with Electricity pt.3

We can use stoichiometry to calculate the quantity of reactants

consumed or products produced in an electrolytic cell.

Corrosion: Undesirable Redox Reactions

Corrosion is the undesired oxidation of metal by environmental

oxidizing agents.

Corrosion: Undesirable Redox Reactions pt.2

When some metals, such as aluminum, oxidize, they form a stable

compound that prevents further oxidation. Iron, in contrast,

does not form a structurally stable compound when oxidized, and

therefore, rust flakes off and exposes more iron to corrosion.

Corrosion: Undesirable Redox Reactions pt.3

Iron corrosion can be prevented by protecting iron from water,

minimizing the presence of electrolytes and acids, or using a sacrificial

electrode.

crystal field theory

is a model that describes the electronic structure of transition metal complexes, focusing on the effect of the surrounding ligands on the d-orbitals.

lanthanide contraction

refers to the phenomenon where the ionic radii of lanthanide elements decrease with increasing atomic number, due to the insufficient shielding effect of the 4f electrons.

complex ion

is a charged species consisting of a central metal atom bonded to one or more ligands, which can be neutral or anions.

ligand

is an ion or molecule that donates a pair of electrons to a central metal atom in a complex, helping to stabilize its structure and influence its properties.

coordination

compound

that forms when a central metal atom is bonded to one or more ligands through coordinate covalent bonds, resulting in a stable complex.

primary valence

refers to the number of bonds formed by a central metal ion with ligands in a complex ion, indicating its oxidation state.

secondary valence

is the number of ligand donor atoms that can bind to a central metal atom within a coordination complex, reflecting its ability to accommodate different ligands.

coordination number

is the total number of ligands surrounding a central metal atom in a coordination complex.

coordinate covalent bond

is a type of bond formed when one atom donates a pair of electrons to another atom, typically involving metal complexes and ligands.

monodentate

ligands that can form only one bond with a central metal atom in a coordination complex.

bidentate

ligands that can form two bonds with a central metal atom in a coordination complex.

polydentate

ligands that can form multiple bonds with a central metal atom in a coordination complex.

chelate

A type of ligand that binds to a central metal atom through multiple donor atoms, forming a stable cyclic structure.

chelating agent

a substance that can form multiple bonds to a central metal atom, stabilizing it in solution.

structural isomers

molecules with the same molecular formula but different structural arrangements.

stereoisomers

molecules that have the same molecular formula and structural arrangement but differ in the spatial orientations of their atoms.

coordination isomers

isomers that differ in the arrangement of ligands around a central metal atom, resulting in distinct coordination structures.

linkage isomers

isomers that differ in the way ligands are attached to the central metal atom, resulting in different coordination complexes.

geometric isomers

isomers that differ in the spatial arrangement of groups around a rigid structure, such as a double bond or a ring.

optical isomers

isomers that are non-superimposable mirror images of each other, typically involving chiral molecules.

strong-field complex

A type of coordination complex that causes a large splitting of d-orbitals, leading to increased pairing of electrons and typically low spin states.

weak-field complex

a type of coordination complex that typically has a low crystal field splitting energy, allowing for spin-free configurations.

low-spin complex

A coordination complex characterized by a larger crystal field splitting energy, which encourages electron pairing and results in fewer unpaired electrons.

high-spin complex

A coordination complex with a smaller crystal field splitting energy, leading to a larger number of unpaired electrons and typically high spin states.