Chemistry Exam 3 Study Guide: Gases, Thermochemistry, and Significant Figures

Comprehensive vocabulary flashcards covering gas laws, kinetic molecular theory, thermochemistry, enthalpy rules, and significant figure operations.

Gas Pressure

The force exerted by gas particles per unit surface area, caused by constant molecular collisions with container walls.

Boyle's Law

A gas law stating that at constant temperature and amount of gas ($T$ and $n$), volume is inversely proportional to pressure ($V \propto \frac{1}{P}$, $P_1V_1 = P_2V_2$).

Charles's Law

A gas law stating that at constant pressure and amount of gas ($P$ and $n$), volume is directly proportional to absolute temperature in Kelvin ($V \propto T$, $\frac{V_1}{T_1} = \frac{V_2}{T_2}$).

Charles's Law Graph

A direct, linear plot of gas volume versus temperature in Kelvin ($\text{K}$) showing that volume extrapolates to zero at absolute zero ($0\,\text{K}$).

Avogadro's Law

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

Ideal Gas Law

The equation of state for an ideal gas relating pressure, volume, moles, and temperature: $PV = nRT$.

Standard Temperature and Pressure (STP)

Defined standard conditions of $T = 0\,^\circ\text{C} = 273\,\text{K}$ and $P = 1\,\text{atm}$, under which $1\,\text{mol}$ of an ideal gas occupies approximately $22.4\,\text{L}$.

Dalton's Law of Partial Pressures

A law stating that the total pressure of a gas mixture is equal to the sum of the partial pressures of its individual components: $P_T = P_A + P_B + P_C + \dots$.

Mole Fraction ($X$)

The ratio of moles of a single component to total moles in a mixture ($X_B = \frac{n_B}{n_{\text{total}}}$), where the partial pressure is $P_B = X_B P_T$.

Diffusion

The process by which gas molecules disperse randomly throughout a container or space from higher concentration to lower concentration until evenly spread out.

Effusion

The escape of gas molecules through a tiny hole (such as a pinhole in a balloon) one particle at a time, occurring faster for lighter gases.

Kinetic Molecular Theory (KMT) Postulates

Postulates for ideal gases stating particles are in continuous random motion, occupy negligible volume ($V \approx 0$), experience no intermolecular forces, have perfectly elastic collisions, and have average kinetic energy directly proportional to absolute temperature ($KE_{\text{avg}} \propto T$).

Real Gases vs. Ideal Gases

Real gas molecules possess actual volume and experience intermolecular attractive forces, which slightly decrease both the effective volume and collision force (pressure) compared to ideal gases.

Conditions for Ideal Gas Behavior

Low pressure (which increases volume and space between particles, preventing interactions) and high temperature (particles move fast enough to overcome intermolecular attractions).

Potential Energy

Energy an object possesses due to its relative position, composition, or condition.

Kinetic Energy

Energy an object possesses due to its motion.

Law of Conservation of Energy

A fundamental principle stating that during a chemical or physical change, energy can neither be created nor destroyed; rather, it changes form.

Thermal Energy

A form of kinetic energy associated with the random motion of atoms and molecules within matter.

Temperature

A quantitative measurement of "hot" or "cold" directly reflecting the average thermal energy and speed of moving particles.

Heat ($q$)

The transfer of thermal energy between two bodies of matter at different temperatures.

Thermal Equilibrium

The state reached when heat flow between two contacting substances ceases because both substances have reached the same temperature.

System

The specific portion of the universe chosen for study, such as the reactants and products undergoing a chemical or physical change.

Surroundings

The rest of the universe outside the system, including the solvent, container, beaker, and environment.

Endothermic Reaction

A process where the system absorbs heat from its surroundings ($q > 0$, $\Delta H > 0$), causing the surroundings to get cooler and involving bond breaking.

Exothermic Reaction

A process where the system releases heat to its surroundings ($q < 0$, $\Delta H < 0$), causing the surroundings to get warmer and involving bond making.

Heat Capacity ($C$)

An extensive property representing the quantity of heat required to raise the temperature of an entire substance or object by $1\,^\circ\text{C}$ or $1\,\text{K}$ ($C = \frac{q}{\Delta T}$).

Specific Heat Capacity ($c$)

An intensive property representing the quantity of heat required to raise the temperature of $1\,\text{g}$ of a substance by $1\,^\circ\text{C}$ or $1\,\text{K}$ ($c = \frac{q}{m\Delta T}$).

Hess's Law

A principle stating that if a reaction is carried out in a series of steps, the overall enthalpy change ($\Delta H$) equals the sum of the enthalpy changes of the individual steps.

Leading Zeros

Zeros that precede all nonzero digits and indicate the position of the decimal point; they are never significant (e.g., $0.0026$ has $2$ significant figures).

Trailing Zeros

Zeros at the end of a number that are significant only if the number contains an explicit decimal point (e.g., $300$ has $1$ sig fig, whereas $300.$ has $3$ sig figs and $300.00$ has $5$ sig figs).

Multiplication and Division Sig Fig Rule

A mathematical rule requiring the final calculated answer to be rounded to the same number of significant figures as the measurement with the fewest significant figures.

Addition and Subtraction Sig Fig Rule

A mathematical rule requiring the final calculated answer to be rounded to the same number of decimal places as the measurement with the fewest decimal places.