Solids Liquids and Gases Study Guide

Chapter 8: Solids, Liquids, and Gases - CHEM 121 (Tacoma Community College)

Phase Transitions and Intermolecular Forces

• Phase Change Factors:

  • Temperature and pressure dictate when a substance transitions between solid, liquid, and gas phases.

  • Energy relationship: Increasing Kinetic Energy ($KE$), often measured by temperature ($T$), moves matter from solid to liquid to gas.

  • Attractive relationship: Increasing Intermolecular Forces ($IMF$) pulls matter from gas to liquid to crystalline solid.

• Defining Transitions:

  • Melting Point: The specific temperature that separates the solid phase from the liquid phase.

  • Boiling Point: The specific temperature that separates the liquid phase from the gas phase.

• Phase Transition Data for Selected Substances:

  • Hydrogen ($H_2$): Solid phase below $-259\,^\circ C$; Liquid phase above $-259\,^\circ C$; Gas phase above $-253\,^\circ C$.

  • Water ($H_2O$): Solid phase below $0\,^\circ C$; Liquid phase above $0\,^\circ C$; Gas phase above $100\,^\circ C$.

  • Sodium Chloride ($NaCl$): Solid phase below $801\,^\circ C$; Liquid phase above $801\,^\circ C$; Gas phase above $1413\,^\circ C$.

Intermolecular vs. Intramolecular Forces

• Intramolecular Forces: Forces that act on the inside of a molecule (e.g., covalent or ionic bonds within the molecule). These are very strong compared to intermolecular forces.

• Intermolecular Forces (IMF): Forces that exist between different particles (atoms or molecules) of a substance. The strength of these forces determines the substance's physical properties, such as boiling and melting points.

Hierarchy of Intermolecular Forces (Weakest to Strongest)

1. London Dispersion Forces (LDF):

   • The weakest type of IMF.

   • Present in ALL molecules.

   • Caused by the constant movement of electrons, which leads to a temporary, instantaneous separation of charge ($\delta-$ and $\delta+$). This sets up a temporary dipole that disappears almost immediately but induces similar dipoles in neighboring atoms.

   • Atomic Size/Radius Effect: In larger atoms, valence electrons are farther from the nuclei and less tightly held. This allows them to form temporary dipoles more easily, increasing the strength of the attraction.

   • Data for Halogens ($X_2$):

     • Fluorine ($F_2$): Atomic Radius $72\,pm$; Melting Point $53\,K$; Boiling Point $85\,K$ (Gas at room temperature).

     • Chlorine ($Cl_2$): Atomic Radius $99\,pm$; Melting Point $172\,K$; Boiling Point $238\,K$ (Gas at room temperature).

     • Bromine ($Br_2$): Atomic Radius $114\,pm$; Melting Point $266\,K$; Boiling Point $332\,K$ (Liquid at room temperature).

     • Iodine ($I_2$): Atomic Radius $133\,pm$; Melting Point $387\,K$; Boiling Point $457\,K$ (Solid at room temperature).

   • Real-World Example: Geckos

     • Geckos adhere to surfaces using dispersion forces.

     • Their toes are covered with hundreds of thousands of tiny hairs called setae.

     • Each seta branches into hundreds of tiny, flat, triangular tips called spatulae.

     • The cumulative attraction of millions of spatulae is strong enough to support many times the gecko's weight.

2. Dipole-Dipole Forces:

   • Occur in polar molecules.

   • Polar Covalent Bond: Results from unequal sharing of electrons.

   • Nonpolar Covalent Bond: Results from equal sharing of electrons (e.g., between the same atoms).

   • Dipole: A separation of charge with a partial positive end and a partial negative end in a molecule.

   • Mechanism: The positive end of one polar molecule is attracted to the partially negative end of another.

3. Hydrogen Bonding:

   • An unusually strong type of dipole-dipole attraction.

   • Occurs specifically in molecules where a hydrogen atom is bonded to highly electronegative atoms: Oxygen ($O$), Nitrogen ($N$), or Fluorine ($F$).

   • Biological Application: DNA

     • Deoxyribonucleic acid (DNA) contains genetic information.

     • It consists of two anti-parallel chains of repeating nucleotides in a double helical structure.

     • These two chains are held together via hydrogen bonds.

4. Ionic Bonding:

   • The strongest force between any two particles.

   • Involves the attraction between ions of opposing charges (cations and anions).

   • Compounds with ionic interactions possess very high melting and boiling points due to the strength of these forces.

Boiling Point and Intermolecular Force Comparison

• To enter the gas phase, particles must completely overcome the IMFs holding them together. Higher boiling points indicate stronger IMFs.

• Substance Data Table:

  • $H_2$: Dispersion ($-253\,^\circ C$)

  • $Ne$: Dispersion ($-246\,^\circ C$)

  • $O_2$: Dispersion ($-183\,^\circ C$)

  • $Cl_2$: Dispersion ($-34\,^\circ C$)

  • $HCl$: Dipole-dipole ($-85\,^\circ C$)

  • $HBr$: Dipole-dipole ($-66\,^\circ C$)

  • $H_2S$: Dipole-dipole ($-61\,^\circ C$)

  • $NH_3$: Hydrogen bonding ($-33\,^\circ C$)

  • $HF$: Hydrogen bonding ($20\,^\circ C$)

  • $H_2O$: Hydrogen bonding ($100\,^\circ C$)

Characteristics of Physical Phases

The Unique Properties of Water ($H_2O$)

• Phase Permanence: Strong intermolecular interactions (Hydrogen bonding) allow water to remain liquid at much higher temperatures than other molecules of similar size.

• Density Anomalies: Unlike most substances, the solid form (ice) is less dense than the liquid form, causing ice to float.

• Solvent Capabilities: Water is an excellent solvent, dissolving many substances and facilitating chemical reactions in solution.

Kinetic Molecular Theory (KMT) of Gases

1. Gases consist of tiny particles separated by large distances (mostly empty space).

2. Gas particles are in constant motion, colliding with each other and the walls of their container.

3. Particle velocity is directly related to the temperature of the gas.

4. Particles experience no attractive or repulsive forces toward each other.

• Ideal Gas: A gas that perfectly adheres to all these statements.

Gas Pressure

• Mechanism: Pressure is generated by the force of gas particles colliding with the container walls.

• Formula: $\text{Pressure} = \frac{\text{Force}}{\text{Area}}$

• Pressure Units and Conversions (Equal to 1 atm):

  • Atmosphere ($atm$): $1\,atm$

  • Millimeters of Mercury ($mmHg$): $760\,mmHg$

  • Torr ($Torr$): $760\,Torr$

  • Pascal ($Pa$): $101326\,Pa$

  • Bar ($bar$): $1.01326\,bar$

  • Pounds per square inch ($psi$): $14.7\,psi$

The Empirical Gas Laws

• Variables: Pressure ($P$), Volume ($V$), Temperature in Kelvins ($T$), and amount of material in moles ($n$).

• Boyle’s Law:

  • Relationship: Pressure and volume are inversely proportional ($P \propto \frac{1}{V}$) at constant $T$ and $n$.

  • Equation: $P_1V_1 = P_2V_2$

  • Application: Lungs. Inhaling requires the lungs to change volume (typically $0.5\text{--}1.0\,L$), which alters internal pressure. Only $1\text{--}2\,torr$ of extra pressure is needed for exhalation.

• Charles’s Law:

  • Relationship: Volume and temperature are directly proportional ($V \propto T$) at constant $P$ and $n$.

  • Equation: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

  • Critical Rule: Temperature must always be in Kelvin ($K$).

• Combined Gas Law:

  • Relates pressure, volume, and temperature simultaneously for a constant amount of gas ($n$).

  • Equation: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Ideal Gas Law:

  • Relates all properties under any set condition (not just changes in conditions).

  • Equation: $PV = nRT$

  • Ideal Gas Constant ($R$):

    • $R = 0.0821\,\frac{L \cdot atm}{mol \cdot K}$

    • $R = 62.4\,\frac{L \cdot mmHg}{mol \cdot K}$

• Standard Temperature and Pressure (STP):

  • Defined as $273\,K$ ($0\,^\circ C$) and $1.00\,atm$.

  • At STP, 1 mole of any gas occupies a volume of $22.4\,L$.

  • Conversion factor: $1\,mol = 22.4\,L$ (at STP).

Dalton’s Law of Partial Pressures

• Proposed by John Dalton (1766-1844).

• Because gas particles are mostly empty space, different gases in a mixture behave independently.

• Partial Pressure: The contribution each individual gas makes to the total pressure.

• Equation: $P_{Total} = P_1 + P_2 + P_3 + \dots$

• Composition of Air (Total $P = 760\,mmHg$):

  • Nitrogen ($N_2$): $P_{N2} = 594\,mmHg$ ($78\%$)

  • Oxygen ($O_2$): $P_{O2} = 160\,mmHg$ ($21\%$)

  • Carbon Dioxide ($CO_2$): $P_{CO2} = 0.25\,mmHg$ ($0.033\%$)

  • Water Vapor ($H_2O$): $P_{H2O} = 5.7\,mmHg$ ($0.75\%$)

  • Other trace gases: $P_{Other} = 0.05\,mmHg$ ($0.22\%$)

Henry’s Law

• Formulated by William Henry in 1803.

• Definition: At constant temperature, the amount of gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid.

• Equation: $C = kP_{gas}$

• Implications: Higher pressure above a liquid leads to higher gas concentration within the liquid.

• Application: Scuba Diving

  • Water pressure increases by $1\,atm$ for every $10.3\,m$ of depth.

  • Deep-sea air mixtures contain less $O_2$ to prevent toxicity. Helium is used as a diluting agent because it has a very low tendency to dissolve in the blood, even at high pressures.

Practice Problems and Examples

• IMF Identification:

  • $N_2$: Dispersion forces (nonpolar).

  • $HCl$: Dipole-dipole forces (polar).

  • $CCl_4$: Dispersion forces (nonpolar symmetric molecule).

  • $HF$: Hydrogen bonding.

  • $H_2S$: Dipole-dipole forces.

• Pressure Unit Conversion Example:

  • Convert $4.5\,atm$ to other units.

  • To Torr: $4.5\,atm \times 760\,Torr/atm = 3420\,Torr$.

• Boyle's Law Example:

  • Initial: $V_1 = 1.5\,L$, $P_1 = 0.87\,atm$. Final: $V_2 = 1.0\,L$.

  • Find $P_2$: $(0.87 \times 1.5) / 1.0 = 1.305\,atm$.

• Charles's Law Example:

  • Initial: $T_1 = 22.0\,^\circ C$ ($295.15\,K$), $V_1 = 1.0\,L$. Final: $T_2 = 30.0\,^\circ C$ ($303.15\,K$).

  • Find $V_2$: $(1.0 \times 303.15) / 295.15 \approx 1.03\,L$.