Unit 2: Moles, Solutions, and Gas Laws

Vocabulary flashcards covering molar calculations, solution concentrations, dilutions, and gas behavior laws based on Unit 2 course material.

Mole ($n$)

• Unit of measurement for amount in chemistry, referring to amount of particles such as atoms, ions, or molecules

Avogadro's Constant ($N_A$)

• The conversion factor between the number of particles and moles, equal to $6.02 \times 10^{23}\,\text{mol}^{-1}$.

Molar Mass ($M_r$) of an Element

• The mass of one mole of an element's atoms, as shown on the periodic table.

Molar Mass of a Compound

• mass of one mole of a compound's particles

• calculated by sum of the molar masses of all atoms

Percent Composition by Mass

• The percentage of a compound's mass contributed by a specific element,

• Calculated as: $\text{Percent Comp. by Mass} = \frac{\text{TOTAL Element Mass}}{\text{Molar Mass of Compound}} \times 100$

Hydrate

• Ionic compound containing unbonded water molecules trapped in the empty spaces of its crystal lattice → evaporate upon heating

Solution

• A homogeneous mixture that is uniformly blended together on the molecular or atomic scale.

Solute

• The substance that is dissolved in a solution

• Usually present in the smaller amount

Solvent

• Substance that dissolves solute in a solution

Aqueous Solution

• A solution in which water is the solvent

• Designated in chemical equations by the symbol $(aq)$.

Concentrated Solution

• A solution containing a large proportion of solute

• Characterized by a high ratio of solute to solvent

Dilute Solution

• A solution containing a small proportion of solute

• Characterized by a low ratio of solute to solvent

Molar Concentration ($C$)

• The amount of solute in moles per volume of solution in cubic decimeters ($C = \frac{n}{V}$)

• Units for concentration are: $\text{mol}\,\text{dm}^{-3}$

Mass Concentration ($\rho$)

• The mass of solute in grams per given volume of solution in cubic decimeters ($\rho = \frac{m}{V}$), commonly measured in $\text{g}\,\text{dm}^{-3}$ or $\text{g}\,\text{cm}^{-3}$

Preparation of a Standard Solution

• The analytical process of dissolving a precisely weighed mass of solute in solvent within a beaker, transferring it, and diluting to the calibration mark in a volumetric flask

Dilution

• The process of decreasing the concentration of a solution by adding more solvent, uses formula $C_1 V_1 = C_2 V_2$

Standard Temperature and Pressure (STP)

• A standard reference condition for comparing gases defined as a pressure of $100\,\text{kPa}$ and a temperature of $273\,\text{K}$ ($0\,^\circ\text{C}$)

Kinetic Molecular Theory (KMT)

1. A model that describes gas behavior at the particle level

2. Particles have negligible volume

3. Exert no intermolecular forces

4. Undergo perfectly elastic collisions

Gas Pressure

• The force per unit area created when moving gas particles collide with the walls of their container.

Relationship Between Volume and Pressure

• A gas law stating that volume and pressure are inversely proportional when temperature and gas amount remain constant ($p_1 V_1 = p_2 V_2$; linear for $P$ vs. $\frac{1}{V}$)

Relationship Between Gas Volume and Temperature

• A gas law stating that gas volume is directly proportional to absolute temperature when pressure and amount of gas remain constant ($\frac{V_1}{T_1} = \frac{V_2}{T_2}$)

Relationship Between Pressure and Temperature

• A gas law stating that gas pressure is directly proportional to absolute temperature when volume and amount of gas remain constant ($\frac{P_1}{T_1} = \frac{P_2}{T_2}$)

Combined Gas Law

• An equation relating initial and final conditions of pressure, volume, and temperature for a fixed amount of gas: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$, where temperature must be in Kelvin

Avogadro's Law

• A principle stating that equal volumes of gases at the same temperature and pressure contain equal numbers of moles

Molar Volume at STP

• The volume occupied by one mole of an ideal gas at standard temperature and pressure ($100\,\text{kPa}$ and $273\,\text{K}$), which equals $22.7\,\text{dm}^3\,\text{mol}^{-1}$.

Ideal Gas Law

• The equation of state of an ideal gas given by $PV = nRT$, where $P$ is pressure ($\text{kPa}$), $V$ is volume ($\text{dm}^3$), $n$ is amount in moles, $R = 8.31\,\text{J}\,\text{K}^{-1}\,\text{mol}^{-1}$, and $T$ is temperature ($\text{K}$).

Real Gas

• A gas that deviates from ideal behavior and the Kinetic Molecular Theory

• Especially, under conditions of very low temperature or very high pressure where particle volume and intermolecular forces become significant.

Empirical Formula

• Simplest ratio of atoms in a compound

Empirical Formula by % Comp

1. % → grams

2. Convert element masses into moles

3. Divide by the smallest no. of moles

4. Answers w/ decimals should be multiplied to get a whole no.

Empirical Formula by Analysis

• Found by analyzing the masses of reactant and products in a reaction

• Used to determine the masses of the elements in the compound

• From this point forward, process for empirical formula finding is the same

Empirical Formula by Combustion

1. Use % comp to find mass of carbon in the mass of CO₂ produced

2. Use % comp to find mass of hydrogen in H₂O produced

3. 3rd element if applicable, find its mass by subtraction → requires initial mass of compound

4. Convert masses to moles

5. Divide all results by the smallest no. of moles to get subscripts

Molecular Formula

• No. of atoms in a covalent molecule

Finding Molecular Formula

1. Know empirical formula + mass of compound

2. Calculate empirical formula → molar mass of emp. formula

3. Divide molecular mass by empirical formula

4. Multiply the result by the subscripts