Unit 2: Energy, Thermodynamics, and Electrochemistry

Comprehensive vocabulary flashcards covering stoichiometry, thermochemistry, calorimetry, bond energies, Hess's law, entropy, Gibbs energy, redox, galvanic cells, expansion work, and historical chemist contributions from Unit 2.

Kinetic Energy ($\text{KE}$) Changes

Energy changes in a substance that alter molecular motion and manifest directly as temperature changes, modeled using $q = mc\theta$ or $q=mc\Delta T$ . Change of thermal kinetic energy is measured in q.

Potential Energy ($\text{PE}$) Changes

Energy changes associated with altered particle positions or attractions (intermolecular forces during phase changes or chemical bonds during reactions), modeled using $\Delta H$.

Law of Conservation of Mass

Energy is neither created nor destroyed (Hence the need for Balanced Equations)

Limiting Reactant

The reactant in a chemical reaction that yields the smallest amount of product and is completely consumed first, dictating the maximum product that can form.

Excess Reactant

A reactant that is not completely consumed once the limiting reactant has been entirely converted into products.

Heat Capacity ($C$)

The quantity of heat energy required to raise the temperature of an entire sample by $1\,^{\circ}\text{C}$ ($\text{J}/^{\circ}\text{C}$), calculated as $C = mc$.

Specific Heat ($c$)

The amount of heat energy required to raise the temperature of $1\,\text{g}$ of a substance by $1\,^{\circ}\text{C}$ ($\text{J}/(\text{g}\cdot^{\circ}\text{C})$); for liquid water, $c = 4.18\,\text{J}/(\text{g}\cdot^{\circ}\text{C})$.

Molar Heat Capacity ($c_m$)

The quantity of heat energy required to raise the temperature of $1\,\text{mol}$ of a substance by $1\,^{\circ}\text{C}$ ($\text{J}/(\text{mol}\cdot^{\circ}\text{C})$).

Degrees of Freedom

A degree of freedom in chemistry and physics is an independent physical parameter or way in which a molecule or particle can move, rotate, or vibrate, and thus store thermal energy

Calorimetry

The experimental measurement of heat transferred between an event/system and its surroundings (typically water), governed by the relation $q_{\text{water}} = -q_{\text{event}}$.

Exothermic Process

A thermodynamic change in which heat exits the system into the surroundings ($q < 0$, $\Delta H < 0$), shifting particles toward stronger attractions, lower potential energy, and feeling hot.

Endothermic Process

A thermodynamic change in which heat enters the system from the surroundings ($q > 0$, $\Delta H > 0$), moving particles toward weaker attractions, higher potential energy, and feeling cold.

Enthalpy of Fusion ($\Delta H_{\text{fus}}$)

DELTA H - heat absorbed or released(enthalpy) The energy input required to melt $1\,\text{mol}$ of a solid substance at constant temperature, equal in magnitude and opposite in sign to the enthalpy of freezing ($\Delta H_{\text{fus}} = -\Delta H_{\text{freeze}}$).

Enthalpy of Vaporization ($\Delta H_{\text{vap}}$)

The energy input required to boil or vaporize $1\,\text{mol}$ of a liquid at constant temperature; it is much larger than $\Delta H_{\text{fus}}$ because separating particles from liquid to gas requires overcoming all intermolecular attractions.

Electrostatic Potential Energy

The potential energy between two charged particles given by, Coloumbs Law, $\text{PE} = \frac{k\cdot q_1\cdot q_2}{r}$, where like charges repel ($\text{PE} > 0$) and opposite charges attract ($\text{PE} < 0$), becoming more negative as distance $r$ decreases.

Bond Length

The distance $r$ at the minimum of a potential energy curve where attractive forces between nuclei and electrons are maximized before electron-electron and nuclear repulsions dominate.

Bond Energy

The positive amount of energy ($\text{kJ}/\text{mol}$) required to break a covalent bond; bond breaking is always endothermic, and energy is released only when stronger bonds form.

Activation Energy ($E_a$)

The initial energy barrier that reacting molecules must overcome to initiate a reaction; it determines the reaction rate but does not affect the overall values of $\Delta H$ or $\Delta G$.

Catalyst

A substance that accelerates a chemical reaction by providing an alternative reaction pathway that lowers the activation energy ($E_a$) without altering $\Delta H$ or $\Delta G$.

Hess's Law

The thermodynamic principle stating that if a reaction can be expressed as the sum of multiple steps, the overall enthalpy change $\Delta H$ equals the sum of the enthalpy changes of each individual step.

Standard Enthalpy of Formation ($\Delta H^\circ_f$)

The heat energy released or consumed when $1\,\text{mol}$ of a chemical compound is formed from its constituent pure elements in their most stable states at standard conditions ($1\,\text{atm}$, $298\,\text{K}$).

Standard State of an Element

The most stable physical state of a pure element at $1\,\text{atm}$ and $298\,\text{K}$ (e.g., $\text{C}(s,\,\text{graphite})$, $\text{O}_2(g)$, $\text{N}_2(g)$, $\text{He}(g)$, $\text{Br}_2(l)$, $\text{Fe}(s)$), which has $\Delta H^\circ_f = 0\,\text{kJ}/\text{mol}$.

Entropy ($S$)

A thermodynamic measure of the dispersion of matter and energy, quantified by the number of microscopic arrangements or microstates available to a system.

First Law of Thermodynamics

The physical law stating that energy cannot be created or destroyed, only transferred or transformed, represented by $\Delta E = q + w$.

Second Law of Thermodynamics

The fundamental law stating that spontaneous changes occur such that the total entropy of the universe increases: $\Delta S_{\text{sys}} + \Delta S_{\text{surr}} = \Delta S_{\text{tot}} > 0$.

Third Law of Thermodynamics

The thermodynamic law stating that the entropy of a pure, perfectly crystalline substance is zero at absolute zero ($0\,\text{K}$).

Entropy of the Surroundings ($\Delta S_{\text{surr}}$)

The entropy change of the environment around a system, calculated as $\Delta S_{\text{surr}} = -\frac{\Delta H_{\text{sys}}}{T}$, which is always opposite in sign to $\Delta H_{\text{sys}}$.

Gibbs Free Energy ($\Delta G$)

A thermodynamic potential defined as $\Delta G = \Delta H - T\Delta S$ that combines system enthalpy and total entropy; a negative value ($\Delta G < 0$) indicates a spontaneous, product-favored process.

Exergonic Reaction

A spontaneous, product-favored chemical reaction characterized by a negative change in Gibbs free energy ($\Delta G < 0$).

Endergonic Reaction

A nonspontaneous, reactant-favored chemical reaction characterized by a positive change in Gibbs free energy ($\Delta G > 0$).

Crossover Temperature

The temperature at which a reaction flips between spontaneous and nonspontaneous, defined by $\Delta G = 0$ and computed as $T = \frac{\Delta H}{\Delta S}$.

Oxidation

The loss of electrons by a chemical species during a redox reaction, occurring at the anode of an electrochemical cell.

Reduction

The gain of electrons by a chemical species during a redox reaction, occurring at the cathode of an electrochemical cell.

Standard Reduction Potential ($E^\circ$)

A voltage metric describing the intrinsic tendency of a chemical species to gain electrons under standard conditions ($1\,\text{M}$, $1\,\text{atm}$, $298\,\text{K}$); larger positive values designate stronger oxidizing agents.

Cell Potential ($E^\circ_{\text{cell}}$)

The electromotive force of an electrochemical cell calculated as $E^\circ_{\text{cell}} = E^\circ_{\text{red}} - E^\circ_{\text{ox}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$, where $E^\circ_{\text{cell}} > 0$ indicates a spontaneous reaction.

Faraday Constant ($F$)

The electrical charge per mole of electrons, given in these notes as $96.5\,\text{kJ}/(\text{V}\cdot\text{mol }e^-)$ in the relationship $\Delta G^\circ = -nFE^\circ$.

Galvanic Cell

An electrochemical cell that physically separates spontaneous oxidation and reduction half-reactions so that electrons travel through an external wire to do electrical work.

Anode

The electrode in an electrochemical cell where oxidation occurs (lower $E^\circ$), serving as the source from which electrons depart.

Cathode

The electrode in an electrochemical cell where reduction occurs (higher $E^\circ$), serving as the destination at which electrons arrive.

Salt Bridge

A component in a galvanic cell containing an inert ionic solution (such as $\text{K}_2\text{SO}_4$ or $\text{KNO}_3$) that allows cations to flow toward the cathode and anions toward the anode to prevent charge accumulation.

Concentration Cell

A nonstandard galvanic cell constructed with identical chemical species and electrodes at differing solute concentrations, driven toward equilibrium ($E^\circ_{\text{cell}} = 0$) purely by entropy.

Battery Capacity

The total amount of work a battery can deliver ($\Delta G$), which scales with the quantity of reactant present, whereas voltage ($E$) depends solely on the underlying chemistry.

Expansion Work ($w$)

Mechanical work performed by an expanding or compressing gas against an external pressure, calculated as $w = -P_{\text{ext}}\Delta V$, where $1\,\text{L}\cdot\text{atm} = 0.101\,\text{kJ}$.

Thermal Efficiency

The fraction of heat input successfully transformed into work, given by $\text{efficiency} = \frac{w}{q} \times 100\%$, which is inherently limited to approximately $\frac{1}{3}$ ($\approx 29\text{--}34\%$) in heat engines because thermal motion is isotropic across three dimensions.

Fuel Cell

An electrochemical cell that continuously converts the chemical energy of a fuel (such as $\text{H}_2$) and oxidant ($\text{O}_2$) directly into directional electrical work, producing water as a nontoxic byproduct without needing recharging.

Alice Ball

The chemist who developed the 'Ball Method', converting sticky chaulmoogric acid into a water-soluble, injectable ethyl ester for the treatment of leprosy.

Why do stable molecules exist?

Covalent Bonds hold the atoms together.

Bond Energy

Also referred to as bond dissociation energy, it is the amount of energy required to break a bond between two atoms in a molecule. Also defined as the standard enthalpy change for an endothermic reaction.

Why are Ionic Compounds stable?

Electrostatic Attraction between negative and positive ions.

Lattice Energy of a Compound

Lattice Energy of an Ionic Compound(Signifying Non-metal and Metal). The energy required to separate one mole of solid into its component gaseous ion.

1st Law of Thermodynamics

Energy cannot be created or destroyed. It only changes form.

2nd Law of Thermodynamics

Changes occur so that total entropy increases.

∆𝑆𝑠𝑦𝑠 + ∆𝑆𝑠𝑢𝑟𝑟 = ∆𝑆𝑡𝑜𝑡 > 0

the entropy (disorder or randomness) of an isolated system always increases or stays the same over time, and never decreases

3rd Law of Thermodynamics

The zero of entropy is defined as that of a perfect

crystal at 0 K.

Covalent Bond

Non-metal compound structured by the sharing of electrons.

Ionic Bond

Occurring between a non-metal and metal, the gift of one ion losing an electron and the other gifting and electron.

Gibbs Energy

Thermodynamic Potential for Spotaneous Processes and Chemical Reactions.

Enthalpy > 0

Indicates Energy is Increasing

Enthalpy < 0

Indicating Energy is Decreasing

Entropy > 0

Indicates Spread of Arrangements are Increasing

Entropy < 0

Indicates a Decreasing in Spread of Arrangements

Spontaneous

ΔG = ΔH − TΔS < 0, when ΔG < 0.

E dot Red

When a species gains an electron.

E dot Ox

When a species loses an electron

E dot Cell

Standard Cell Potential, Measure of Voltage

+q

energy enters the system (endothermic)

-q

energy exits the system (exothermic)

positive enthalpy

energy increases, endothermic process

negative enthalpy

energy exists the system (exothermic)

positive gibbs

non-spontaneous (endergonic)

negative gibbs

spontaneous(exergonic)

e dot cell positive

indicates a spontaneous reaction in an electrochemical cell, signifying that the cell can generate electrical energy.

e dot cell negative

indicates a non-spontaneous reaction in an electrochemical cell, meaning it cannot generate electrical energy.