General Chemistry Chapters 4-7 Vocabulary Review

Comprehensive vocabulary flashcards covering Redox Reactions (Chapter 4), Gas Laws and Kinetic-Molecular Theory (Chapter 5), Thermochemistry and Thermodynamics (Chapter 6), and Classical Waves and Quantum Theory (Chapter 7).

Oxidation-Reduction (Redox) Reaction

A chemical reaction involving the transfer of electrons.

Half-Reaction

An equation showing either the oxidation process or the reduction process separately.

Oxidation

The loss of electrons.

Reduction

The gain of electrons.

Oxidation Number (Oxidation State)

A bookkeeping device used to track electron transfer by assigning hypothetical charges to atoms.

Activity Series

A ranking of elements according to how readily they lose electrons (are oxidized), used to predict whether redox reactions will occur.

Gas

A substance that normally exists in the gaseous state at ordinary temperatures and pressures.

Vapor

The gaseous form of a substance that is usually a liquid or solid under ordinary conditions.

Pressure ($P$)

The force exerted per unit area by collisions of gas molecules with a surface: $P = \frac{F}{A}$.

Pascal ($\text{Pa}$)

The SI unit of pressure; one newton per square meter.

Atmosphere ($\text{atm}$)

A common pressure unit equal to the pressure exerted by a $760\,\text{mm}$ column of mercury at sea level and $0\,^\circ\text{C}$. Equivalent to $1\,\text{atm} = 760\,\text{mm\,Hg} = 760\,\text{torr} = 101{,}325\,\text{Pa}$.

Bar

A pressure unit equal to $100{,}000\,\text{Pa}$.

Barometer

An instrument used to measure atmospheric pressure.

Manometer

A device used to measure the pressure of a gas.

Boyle’s Law

At constant temperature, the pressure of a gas is inversely proportional to its volume ($P_1V_1 = P_2V_2$).

Charles’s Law

At constant pressure, the volume of a gas is directly proportional to its absolute temperature ($\frac{V_1}{T_1} = \frac{V_2}{T_2}$).

Absolute Zero

The lowest possible temperature, equal to $-273.15\,^\circ\text{C}$ ($0\,\text{K}$), at which an ideal gas would have zero volume.

Law of Combining Volumes

The observation that gases involved in chemical reactions combine in simple whole-number volume ratios when measured at the same temperature and pressure.

Avogadro’s Law

Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

Ideal Gas Law

A mathematical relationship among pressure, volume, temperature, and moles of gas: $PV = nRT$.

Universal Gas Constant ($R$)

The proportionality constant in the ideal gas law, equal to $0.082057\,\text{L}\cdot\text{atm}/(\text{mol}\cdot\text{K})$ or $8.314\,\text{J}/(\text{mol}\cdot\text{K})$.

Standard Temperature and Pressure (STP)

Conditions of $0\,^\circ\text{C}$ and $1\,\text{atm}$ pressure.

Ideal Gas

A hypothetical gas whose particles have negligible volume and exert no intermolecular forces except during collisions.

Gas Density ($d$)

Mass per unit volume of a gas, calculated by $d = \frac{P \cdot \text{MM}}{RT}$.

Molar Mass from Gas Density

The calculation of gas molar mass defined by $\text{MM} = \frac{dRT}{P}$.

Dalton’s Law of Partial Pressures

The total pressure of a mixture of gases equals the sum of the partial pressures of the individual gases: $P_t = P_1 + P_2 + P_3 + \dots$.

Partial Pressure

The pressure that a gas would exert if it occupied the container alone at the same temperature and volume.

Mole Fraction ($X$)

The ratio of the moles of one component to the total moles of all components, related to partial pressure by $P_i = X_i P_t$.

Kinetic-Molecular Theory

A model describing gas behavior based on five main assumptions: continuous random motion of particles, negligible particle volume compared with container volume, negligible intermolecular forces, perfectly elastic collisions, and temperature proportional to average kinetic energy.

Elastic Collision

A collision in which total kinetic energy is conserved.

Kinetic Energy ($\text{KE}$)

Energy associated with motion: $\text{KE} = \frac{1}{2}mu^2$.

Maxwell–Boltzmann Distribution

The statistical distribution describing the range of molecular speeds in a gas.

Root-Mean-Square Speed ($u_{\text{RMS}}$)

The square root of the average of the squares of molecular speeds: $u_{\text{RMS}} = \sqrt{\frac{3RT}{\text{MM}}}$.

Effusion

The escape of a gas through a tiny hole into a vacuum.

Diffusion

The spreading of a gas throughout a space, often into another gas.

Graham’s Law of Effusion

The rate of effusion is inversely proportional to the square root of molar mass: $\frac{\text{rate}_A}{\text{rate}_B} = \sqrt{\frac{\text{MM}_B}{\text{MM}_A}}$.

Real Gas

An actual gas whose behavior deviates from ideal gas behavior because molecules have volume and intermolecular forces.

Deviation from Ideal Behavior

Differences between measured gas behavior and predictions of the ideal gas law, especially at high pressures and low temperatures.

van der Waals Equation

A modified gas equation that accounts for intermolecular attractions and molecular volume: $\left(P + \frac{n^2a}{V^2}\right)(V - nb) = nRT$.

van der Waals Constant $a$

A correction factor in the van der Waals equation accounting for intermolecular attractions.

van der Waals Constant $b$

A correction factor in the van der Waals equation accounting for the finite volume occupied by gas.

Thermochemistry

The study of energy changes that accompany chemical reactions.

Energy

The capacity to do work.

Work ($w$)

Work is done when an object is moved by a force: $w = F \cdot d$.

Heat ($q$)

The transfer of thermal energy between two objects that are at different temperatures.

Thermal Energy

The kinetic energy associated with the random motion of atoms and molecules.

Potential Energy

Energy available due to the position of an object in a force field.

Chemical Energy

The form of energy stored within the structural units of chemical substances.

Joule ($\text{J}$)

The SI unit of energy: $1\,\text{J} = 1\,\text{kg}\cdot\text{m}^2/\text{s}^2$.

System

The part of the universe that is of interest to us.

Surroundings

The remainder of the universe outside the system.

Open System

A system that can exchange both matter and energy with its surroundings.

Closed System

A system that can exchange energy, but not matter, with its surroundings.

Isolated System

A system that can exchange neither matter nor energy with its surroundings.

Thermodynamics

The study of all types of energy changes.

Law of Conservation of Energy

Energy can neither be created nor destroyed; the total quantity of energy in the universe remains constant.

First Law of Thermodynamics

The total energy of the universe is constant: $\Delta E_{\text{univ}} = \Delta E_{\text{sys}} + \Delta E_{\text{surr}} = 0$.

Internal Energy ($E$)

The total energy possessed by a system, including thermal, chemical, and potential energies.

State of a System

The values of all pertinent macroscopic properties of the system.

State Function

A property that depends only on the current state of a system and not on the path used to reach that state.

Exothermic Process

A process in which the system releases heat to the surroundings.

Endothermic Process

A process in which the system absorbs heat from the surroundings.

Constant-Pressure Process

A process occurring at constant external pressure.

Enthalpy ($H$)

A thermodynamic quantity defined as: $H = E + PV$.

Change in Enthalpy ($\Delta H$)

The heat absorbed or released during a constant-pressure process: $\Delta H = H_{\text{final}} - H_{\text{initial}}$.

Heat of Reaction ($\Delta H_{\text{rxn}}$)

The enthalpy change associated with a chemical reaction: $\Delta H_{\text{rxn}} = H_{\text{products}} - H_{\text{reactants}}$.

Thermochemical Equation

A balanced chemical equation that includes the associated enthalpy change.

Molar Enthalpy Change

The enthalpy change corresponding to one mole of reaction.

Calorimetry

The measurement of heat absorbed or released during a process.

Heat Capacity ($C$)

The amount of heat required to raise the temperature of a sample by one kelvin.

Molar Heat Capacity

The amount of heat required to raise the temperature of one mole of a substance by one kelvin.

Specific Heat Capacity ($c$)

The amount of heat required to raise the temperature of one gram of a substance by one kelvin.

Temperature Change

The difference between the final and initial temperature: $\Delta T = T_{\text{final}} - T_{\text{initial}}$.

Heat-Temperature Relationship

The mathematical formula relating heat to heat capacity and temperature change: $q = C\Delta T$.

Bomb Calorimeter

A rigid, constant-volume calorimeter used primarily for combustion reactions.

Constant-Volume Calorimetry

Calorimetry performed under constant-volume conditions, where the measured heat corresponds to $\Delta E$.

Constant-Pressure Calorimetry

Calorimetry performed under constant-pressure conditions, where the measured heat corresponds to $\Delta H$.

Hess's Law of Heat Summation

The enthalpy change of an overall process equals the sum of the enthalpy changes of its individual steps.

Formation Reaction

A reaction in which a compound is formed from its constituent elements in their standard states.

Enthalpy of Formation ($\Delta H_f$)

The enthalpy change associated with a formation reaction.

Standard Conditions (Thermodynamic Standard Conditions)

Conditions where gases are at $1\,\text{atm}$ pressure, aqueous species are $1\,\text{M}$, and substances are in their standard states.

Standard State

The most stable form of a substance at $1\,\text{atm}$ pressure and the specified temperature.

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

The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions.

Standard Heat of Reaction ($\Delta H^\circ$)

The enthalpy change for a reaction calculated from standard heats of formation: $\Delta H^\circ = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants})$.

Wave

A disturbance that propagates from point to point in space without any net displacement of the medium (if a medium exists).

Frequency ($\nu$)

The number of complete cycles that pass through a particular point in space during one second. Units: $\text{s}^{-1}$ or hertz ($\text{Hz}$).

Wavelength ($\lambda$)

The distance between identical points on successive cycles of a wave.

Amplitude ($A$)

The vertical distance from the midline of a wave to its peak or trough.

Hertz ($\text{Hz}$)

A unit of frequency equal to one cycle per second: $1\,\text{Hz} = 1\,\text{s}^{-1}$.

Electromagnetic Radiation

A form of energy consisting of oscillating electric and magnetic fields that can travel through a vacuum.

Speed of Light ($c$)

The speed at which electromagnetic radiation travels in a vacuum: $c = 2.99792458 \times 10^8\,\text{m/s}$.

Visible Region

The portion of the electromagnetic spectrum visible to humans, approximately $400\text{--}750\,\text{nm}$.

Refraction

The change in direction of a wave resulting from a change in its speed when passing from one medium to another.

Diffraction

The bending and spreading of a wave as it encounters an obstacle or opening.

Interference

The interaction of two waves that meet and combine.

Constructive Interference

Occurs when waves combine in phase, producing a larger amplitude.

Destructive Interference

Occurs when waves combine out of phase and partially or completely cancel each other.

Blackbody Radiation

Electromagnetic radiation emitted by a heated object whose intensity depends on frequency.

Quantum (plural: Quanta)

A discrete packet of energy.

Planck’s Constant ($h$)

The proportionality constant relating photon energy to frequency: $h = 6.62607015 \times 10^{-34}\,\text{J}\cdot\text{s}$.