CHEM 1000 - Chapter 5: Gases Flashcards

A set of 50 vocabulary flashcards covering gas laws, Kinetic Molecular Theory, stoichiometry, effusion, and real gas behavior based on CHEM 1000 lecture slides.

Sodium Chlorate Decomposition

The chemical reaction used in airplane oxygen masks to generate oxygen gas: $2\text{NaClO}_{3(s)} \rightarrow 2\text{NaCl}_{(s)} + 3\text{O}_{2(g)}$.

Sodium Azide Decomposition

The rapid chemical explosion reaction used to inflate vehicle airbags: $2\text{NaN}_{3(s)} \rightarrow 2\text{Na}_{(s)} + 3\text{N}_{2(g)}$.

Macroscopic Gas Properties

Large-scale, observable properties of a bulk gas sample, specifically pressure ($P$), temperature ($T$), and volume ($V$).

Microscopic Gas Properties

Molecular-scale properties describing individual gas particle motion, specifically position ($x_i$), velocity ($v_i$), and mass ($m_i$).

Gas Pressure

The total force per unit area exerted by collisions of gaseous atoms or molecules against a surface.

Pressure Equation

The mathematical formula for pressure: $P = \frac{F}{A} = \frac{(m)(a)}{A}$, where $F$ is force, $A$ is area, $m$ is mass, and $a$ is acceleration.

Pascal (Pa)

The SI unit for pressure, defined as $1\,\text{N}\,\text{m}^{-2}$ or $1\,\text{kg}\,\text{m}^{-1}\,\text{s}^{-2}$.

Atmosphere (atm)

A pressure unit defined as the average pressure at sea level supporting a $760\,\text{mmHg}$ mercury column, equal to $101.3\,\text{kPa}$.

Bar

A unit of pressure defined by IUPAC as exactly $100,000\,\text{Pa}$ ($10^5\,\text{Pa}$) or $1.01325\,\text{bar} = 1\,\text{atm}$.

Torr

A pressure unit defined as exactly equal to $1\,\text{mmHg}$, named after Evangelista Torricelli ($1\,\text{atm} = 760\,\text{Torr}$).

Mercury Barometer

An instrument consisting of an evacuated glass tube submerged in a pool of mercury, used to measure atmospheric pressure.

Barometer Pressure Equation

The formula relating atmospheric pressure to a liquid column height: $P_{\text{atm}} = g \cdot h \cdot d$, where $h$ is height, $d$ is density, and $g$ is acceleration due to gravity.

Manometer

An instrument that measures pressure differences between a gas sample and the atmosphere using a U-tube containing liquid (usually mercury).

Standard Pressure (IUPAC)

The standard reference pressure defined by IUPAC as exactly $100,000\,\text{Pa}$ ($1\,\text{bar}$).

Boyle's Law

The gas law stating that gas volume is inversely proportional to pressure at constant temperature and mole amount ($V \propto \frac{1}{P}$ or $P_1 V_1 = P_2 V_2$).

Boyle's Law Molecular Mechanism

As gas volume decreases, particles hit the container walls more frequently, resulting in higher pressure.

Charles's Law

The gas law stating that gas volume is directly proportional to absolute temperature at constant pressure and mole amount ($V \propto T$ or $\frac{V_1}{T_1} = \frac{V_2}{T_2}$).

Absolute Zero

The theoretical lowest temperature where ideal gas volume extrapolates to zero, defined as $-273.15\,^\circ\text{C}$ or $0.00\,\text{K}$.

Combined Gas Law

The equation combining Boyle's and Charles's laws for a fixed amount of gas: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$.

Avogadro's Law

The gas law stating that gas volume is directly proportional to the number of moles of gas at constant pressure and temperature ($V \propto n$ or $\frac{V_1}{n_1} = \frac{V_2}{n_2}$).

Ideal Gas Law

The state equation combining Boyle's, Charles's, and Avogadro's laws: $PV = nRT$.

Ideal Gas Constant ($R$ in SI Units)

The universal gas constant expressed in SI units: $8.314\,\text{J}\,\text{mol}^{-1}\,\text{K}^{-1}$ (or $\text{Pa}\,\text{m}^3\,\text{K}^{-1}\,\text{mol}^{-1}$).

Ideal Gas Constant ($R$ in L atm Units)

The universal gas constant expressed using atmospheres: $0.08206\,\text{L}\,\text{atm}\,\text{mol}^{-1}\,\text{K}^{-1}$.

Ideal Gas Constant ($R$ in L bar Units)

The universal gas constant expressed using bar: $0.08314\,\text{L}\,\text{bar}\,\text{mol}^{-1}\,\text{K}^{-1}$.

Gas Density Equation

Formula derived from the ideal gas law calculating gas mass per unit volume: $d = \frac{P \cdot M}{R \cdot T}$, where $M$ is molar mass.

Molar Mass Gas Formula

Rearranged form of the ideal gas law used to calculate molar mass ($M$) from sample mass ($m$): $M = \frac{m \cdot R \cdot T}{P \cdot V}$.

Standard Temperature and Pressure (STP)

Standard reference conditions defined as $273.15\,\text{K}$ ($0\,^\circ\text{C}$) and $1.00\,\text{bar}$.

Molar Volume of an Ideal Gas at STP

The volume occupied by one mole of an ideal gas at STP ($1.00\,\text{bar}$, $273.15\,\text{K}$), equal to $22.71\,\text{L}$.

Dalton's Law of Partial Pressures

Statement that total pressure of a non-reacting gas mixture equals the sum of the partial pressures of each component: $P_{\text{total}} = P_1 + P_2 + \dots + P_n$.

Mole Fraction ($X_i$)

The dimensionless ratio of moles of component $i$ to total moles in a mixture: $X_i = \frac{n_i}{n_{\text{total}}}$.

Partial Pressure Equation

Formula calculating component pressure $P_i$ from its mole fraction $X_i$ and total pressure $P_{\text{total}}$: $P_i = X_i \cdot P_{\text{total}}$.

Vapor Pressure of Water

The temperature-dependent pressure exerted by water vapor in dynamic equilibrium with liquid water when collecting gases over water.

Dry Gas Pressure Over Water

Application of Dalton's Law to calculate dry target gas pressure collected over water: $P_{\text{dry gas}} = P_{\text{total}} - P_{\text{H}_2\text{O}}$.

KMT Postulate 1 (Particle Motion)

Gas consists of tiny particles (atoms or molecules) moving randomly that do not interact with one another.

KMT Postulate 2 (Particle Volume)

The physical size of gas particles is extremely small compared to the total volume of the container.

KMT Postulate 3 (Kinetic Energy)

The average kinetic energy of gas particles is directly proportional to the temperature in Kelvin ($\text{KE}_{\text{avg}} \propto T$).

KMT Postulate 4 (Collisions)

Collisions between gas particles or against container walls are completely elastic.

Elastic Collision

A collision in which no net kinetic energy is converted into heat or other internal forms of energy.

Molar Translational Kinetic Energy

The total average kinetic energy in one mole of ideal gas molecules: $\text{KE}_{\text{avg}} = \frac{3}{2}RT = \frac{1}{2}N_A m u^2$.

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

The square root of the average of the squared molecular speeds: $u_{\text{rms}} = \sqrt{\frac{3RT}{M}}$, where $R = 8.314\,\text{J}\,\text{mol}^{-1}\,\text{K}^{-1}$ and $M$ is in $\text{kg}\,\text{mol}^{-1}$.

Most Probable Speed ($u_m$)

The speed corresponding to the maximum peak of the Maxwell-Boltzmann molecular speed distribution.

Average Speed ($u_{\text{av}}$)

The arithmetic mean speed of all gas molecules present in a sample.

Mean Free Path

The average distance a gas particle travels between successive collisions with other particles.

Diffusion

The process by which gas molecules spread out throughout a space in response to a concentration gradient.

Effusion

The process by which gas molecules escape from a container through a small hole into a vacuum or region of lower pressure.

Graham's Law of Effusion

Law stating that the effusion rate of a gas is inversely proportional to the square root of its molar mass: $\frac{\text{rate}_A}{\text{rate}_B} = \sqrt{\frac{M_B}{M_A}}$.

van der Waals Equation

Equation modifying the ideal gas law for real gas deviations: $\left(P + a\left(\frac{n}{V}\right)^2\right)(V - nb) = nRT$.

van der Waals Constant $a$

An empirical parameter in the van der Waals equation correcting for attractive intermolecular forces between gas molecules.

van der Waals Constant $b$

An empirical parameter in the van der Waals equation correcting for the finite physical volume occupied by gas molecules.

Positive Deviation from Ideality

Behavior observed at high pressure where $PV/RT > 1$ because molecular particle volume ($b$) makes actual volume higher than predicted.