Engineering Chemistry: Module 1 - Chemical Reactions, Electrochemistry, Corrosion, and Batteries

Comprehensive vocabulary flashcards covering chemical reaction classifications, thermodynamics, electrochemical cells, reference electrodes, corrosion mechanisms, protection strategies, and modern battery technologies.

Chemical Reaction

A process in which one or more reactants transform into products through making and breaking bonds between atoms, characterized by changes in electron locations without any modification to atomic nuclei.

Nuclear Fission

The splitting of a heavy atomic nucleus, such as $^{235}\text{U}$, into two or more smaller daughter nuclei upon neutron bombardment, accompanied by the ejection of two or more neutrons and the liberation of large amounts of energy.

Nuclear Fusion

A thermonuclear reaction taking place at approximately $100\text{ million}\,^\circ\text{C}$ in which two light nuclei (such as deuterium) combine to yield a heavier nucleus (such as helium), where a mass deficit of $0.02017\,\text{amu}$ is converted into enormous energy.

Activation Energy ($E_a$)

The minimum threshold energy required to initiate a chemical reaction by reaching the transition state to break bonds in reactants, required by both endothermic and exothermic reactions.

Enthalpy Change ($\Delta H$)

A thermodynamic property representing the overall heat absorbed or released by a reaction at constant pressure, calculated as $\Delta H = H_{\text{products}} - H_{\text{reactants}}$.

Endothermic Reaction

A reaction with a positive enthalpy change ($\Delta H > 0$) where the energy required to break bonds in the reactants is greater than the energy released upon product bond formation, pulling heat from the surroundings and cooling them.

Exothermic Reaction

A reaction with a negative standard enthalpy change ($\Delta H^\circ < 0$) where the energy released in bond formation exceeds the energy needed to break reactant bonds, releasing heat and light into the surroundings.

Electrochemistry

The branch of chemistry dealing with the inter-conversion of chemical energy and electrical energy, studying both electrolytic processes and spontaneous galvanic processes.

Metallic Conductor

A solid conductor (such as copper, silver, aluminum, or graphite) that conducts current via the movement of free valence electrons, undergoing no chemical change and showing decreased conductivity as temperature rises.

Electrolytic Conductor

A substance in aqueous solution or molten state that conducts electricity through mobile ions (cations and anions), experiencing chemical decomposition at the electrodes and increased conductivity at higher temperatures.

Galvanic Cell

An electrochemical cell that spontaneously converts the chemical energy of a redox reaction into electrical energy, utilizing a negative anode (site of oxidation) and a positive cathode (site of reduction).

Daniel Cell

A classic galvanic cell invented by John Frederic Daniell in 1836 consisting of a zinc anode in $\text{ZnSO}_4$ solution and a copper cathode in $\text{CuSO}_4$ solution, producing a standard cell potential of $E^\circ_{\text{cell}} = +1.10\,\text{V}$.

Helmholtz Electrical Double Layer

An equilibrium layer of ions of opposite charge formed at the boundary between a metal electrode and its surrounding electrolytic solution, creating an electrode potential.

Salt Bridge

A U-tube containing an agar-agar or gelatin gel saturated with an inert electrolyte with equal ionic mobilities (such as $\text{KCl}$ or $\text{NH}_4\text{NO}_3$) that connects two half-cells, maintaining electrical neutrality and eliminating the liquid junction potential.

Standard Hydrogen Electrode (SHE)

The universally accepted primary reference electrode consisting of platinized platinum bathed in $\text{H}_2$ gas at $1\,\text{atm}$ and immersed in $1\,\text{mol}\,\text{dm}^{-3}$ $\text{H}^+$ solution at $25\,^\circ\text{C}$, with an assigned potential of $E^\circ = 0.00\,\text{V}$ at all temperatures.

Calomel Electrode

A secondary reference electrode comprising mercury, mercurous chloride ($\text{Hg}_2\text{Cl}_2$), and potassium chloride solution, providing a stable potential of $+0.2422\,\text{V}$ versus SHE in a saturated $\text{KCl}$ solution at $25\,^\circ\text{C}$.

Glass Electrode

An ion-selective membrane electrode sensitive to $\text{H}^+$ ions, constructed with a thin silicate glass bulb doped with metal oxides ($\text{SiO}_2$, $\text{Na}_2\text{O}$, $\text{CaO}$) enclosing $0.1\,\text{M}$ $\text{HCl}$ and an internal $\text{Ag/AgCl}$ wire, whose potential varies linearly with pH.

Electrochemical Series

A sequence of metals and non-metals ordered by their increasing standard reduction potentials ($E^\circ$) relative to the Standard Hydrogen Electrode under standard conditions ($25\,^\circ\text{C}$, $1\,\text{M}$, and $1\,\text{atm}$).

Corrosion

The spontaneous deterioration of metals or alloys caused by chemical or electrochemical reactions with the surrounding environment, converting refined metals back into lower Gibbs free energy compounds ($\Delta G < 0$) such as oxides, hydroxides, or sulfides.

Dry Corrosion

The direct chemical attack of atmospheric gases such as $\text{O}_2$, $\text{H}_2\text{S}$, and $\text{SO}_2$ on dry metal surfaces without moisture, where corrosion products accumulate directly at the reaction site.

Pilling-Bedworth Rule

A quantitative criterion predicting oxide film protectiveness using the ratio of the volume of metal oxide formed to the volume of metal consumed; films with a ratio between $1$ and $2$ (such as $\text{Al}_2\text{O}_3$ at $1.28$) are protective, while those below $1$ are porous and non-protective.

Hydrogen Embrittlement

The mechanical degradation of metals exposed to $\text{H}_2\text{S}$ where generated atomic hydrogen ($[\text{H}]$) diffuses into grain boundaries and voids, recombining into high-pressure $\text{H}_2$ gas that causes internal blistering, cracking, and brittle failure.

Decarburisation

The high-temperature (above $700\,^\circ\text{C}$) depletion of carbon from steel caused by atomic hydrogen reacting with carbon to form volatile methane gas ($\text{CH}_4$), resulting in a loss of hardness and tensile strength.

Wet Corrosion

Electrochemical degradation occurring when a metal contacts an aqueous electrolyte, dividing the metal into distinct anodic sites (where metal dissolves) and cathodic sites (where oxygen reduction or hydrogen evolution occurs).

Galvanic Corrosion

An accelerated form of electrochemical corrosion occurring when two dissimilar metals are in direct electrical contact within an electrolyte, forcing the metal higher in the EMF series to act as an anode and corrode preferentially.

Differential Aeration Corrosion

Concentration cell corrosion that develops when a single piece of metal is exposed to unequal oxygen concentrations; poorly aerated areas act as anodes and corrode, while well-aerated areas act as protected cathodes.

Pitting Corrosion

An intense localized form of differential aeration corrosion where a small oxygen-deficient area under a droplet or cavity behaves as a small anode that dissolves rapidly under the high current density of the surrounding large cathode.

Water Line Corrosion

A specific manifestation of differential aeration corrosion on partially submerged structures (such as ships and storage tanks), where intense corrosion occurs along a band just below the waterline due to lower dissolved oxygen relative to the air-exposed upper section.

Pourbaix Diagram

A potential-pH ($E\text{-pH}$) stability diagram that thermodynamically maps the distinct domains of corrosion, passivity (protective oxide formation), and immunity for a metal-water system at a given temperature.

Sacrificial Anode Protection

A cathodic protection method where a more active metal (such as $\text{Mg}$, $\text{Zn}$, or $\text{Al}$) is electrically coupled to a structure, preferentially oxidizing and sacrificing itself while supplying electrons to make the structure a protected cathode.

Impressed Current Cathodic Protection (ICCP)

An engineered cathodic protection system using an external DC power source whose negative terminal is connected to the protected structure and positive terminal is wired to an inert anode (e.g., graphite, platinised titanium) buried in backfill.

Primary Battery

A non-rechargeable electrochemical cell in which the chemical reactions are irreversible, generating electricity until internal active components are depleted, such as the Leclanché dry cell.

Lead-Acid Storage Battery

A secondary storage battery operating at $2\,\text{V}$ per cell consisting of a spongy lead ($\text{Pb}$) anode, a lead dioxide ($\text{PbO}_2$) cathode, and a $20\%$ sulfuric acid ($\text{H}_2\text{SO}_4$) electrolyte, whose state of charge is monitored by electrolyte specific gravity.

Nickel-Cadmium Battery

A rechargeable secondary battery providing a cell voltage of $1.4\,\text{V}$ composed of a cadmium anode, a metal grid containing $\text{NiO}_2$ as cathode, and an alkaline $\text{KOH}$ electrolyte.

Hydrogen-Oxygen Fuel Cell

An electrochemical cell that continuously converts the chemical energy of supplied hydrogen fuel and oxygen directly into electricity using platinum catalysts and a Proton Exchange Membrane, yielding only water as a by-product.

Solid-State Battery (SSB)

An advanced rechargeable battery architecture replacing flammable organic liquid electrolytes with a solid electrolyte (such as ceramic, polymer, glass, or composite), preventing leakage and thermal runaway while increasing energy density.

Sodium-Ion Battery (SIB)

A rechargeable energy storage device that uses abundant sodium ions ($\text{Na}^+$) as charge carriers between a hard carbon anode and a Prussian Blue analogue or layered oxide cathode in an organic electrolyte.

Thermal Runaway

A hazardous chain reaction occurring in batteries where a rise in internal temperature initiates exothermic reactions, rapidly increasing heat output and leading to swelling, toxic leakage, fire, or explosion.