Intermolecular Forces and the Uniqueness of Water — Comprehensive Study Notes Flashcards

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Intramolecular Forces

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Forces operating within a molecule that arise from the combination of attractive and repulsive electrical interactions between pairs of electrons and nuclei. These forces determine molecular structure and stability rather than interactions between separate molecules. Intramolecular forces include covalent and ionic bonding.

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Intermolecular Forces

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Forces that act between molecules or between molecules and ions and strongly influence physical properties such as boiling point, melting point, and viscosity. Examples include dispersion forces, dipole–dipole interactions, hydrogen bonding, and ion–dipole forces. Their relative strengths determine macroscopic behavior of substances.

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States of Matter

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The distinct physical forms (solid, liquid, gas) determined by the balance of kinetic energy and intermolecular forces. Stronger intermolecular attractions favor condensed phases (liquid or solid), while weaker attractions favor the gas phase. Phase changes occur when temperature or pressure shifts this balance.

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Boiling Point Trends

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Boiling points generally increase as the strength of intermolecular forces increases because more energy is required to separate attracted particles. For homologous series or noble gases, larger atoms/molecules (higher atomic number) have higher boiling points due to stronger dispersion forces. Molecular polarity and hydrogen bonding also raise boiling points.

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Dispersion Forces

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Also called London dispersion forces, these are instantaneous attractive interactions between nonpolar molecules caused by momentary induced dipoles. Dispersion is universal (present in all molecules) and becomes stronger for larger, more polarizable species. It often dominates in nonpolar substances.

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Temporary Dipole

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A momentary separation of charge in an atom or molecule produced by an uneven, instantaneous distribution of electrons. Temporary dipoles induce dipoles in neighboring particles, producing dispersion (London) forces. These fluctuations underpin attraction in nonpolar substances.

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Polarizability

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A measure of how easily the electron cloud of an atom, ion, or molecule can be distorted to form an induced dipole. Polarizability increases with larger, more diffuse electron clouds and contributes to stronger dispersion forces. Higher polarizability typically raises boiling points.

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Factors Affecting Dispersion

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Primary factors are size (molar mass/polarizability) and molecular shape (surface area). Larger, more polarizable species and more extended (less compact) shapes increase dispersion interactions and therefore raise boiling points. Linear molecules usually experience stronger dispersion than branched isomers of similar mass.

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Ion–Dipole Interaction

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An attractive force between an ion and a polar molecule that has a permanent dipole moment. This interaction is especially important in solutions, for example when ionic salts dissolve in polar solvents like $H_2O$. Ion–dipole forces are among the strongest intermolecular attractions.

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Sphere of Hydration

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The cluster of $H_2O$ molecules that surround an ion in aqueous solution, stabilizing the ion via ion–dipole interactions. When the solvent is not water, the same concept is called the sphere of solvation. Hydration lowers the effective ion–ion attraction and facilitates solubility.

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Dipole–Dipole Interaction

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An attractive force between polar molecules arising from electrostatic interactions between permanent molecular dipoles. Dipole–dipole forces affect physical properties like boiling point and solubility and are stronger when dipole moments are larger and molecular orientations favor attraction.

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Hydrogen Bonding

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A particularly strong type of dipole–dipole interaction that occurs when a hydrogen atom bonded to a small, highly electronegative atom (F, O, or N) interacts with a lone pair on F, O, or N in another molecule. Hydrogen bonding significantly raises boiling points, increases viscosity, and imparts unique properties to substances such as $H_2O$. It is directional and stronger than typical dipole–dipole forces.

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Dipole–Induced Dipole

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An interaction in which a permanent dipole of a polar molecule induces a temporary dipole in a nearby nonpolar molecule by distorting its electron cloud. This results in an attractive force that is weaker than permanent dipole–dipole interactions but stronger than pure dispersion between very small species.

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Binary Hydride Boiling Points

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The anomalously high boiling points of hydrides like $H_2O$, $HF$, and $NH_3$ are due to hydrogen bonding between molecules. Even though these molecules have low molar masses, hydrogen bonds increase intermolecular attraction and therefore raise boiling points relative to non-hydrogen-bonding analogs.

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Hydrogen Bonding in DNA

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Hydrogen bonds hold complementary DNA strands together: adenine (A) hydrogen-bonds to thymine (T) via two hydrogen bonds, while guanine (G) hydrogen-bonds to cytosine (C) via three hydrogen bonds. These specific pairings stabilize the double helix and contribute to the fidelity of base pairing.

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Relative Strengths

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Intermolecular forces can be ranked roughly as: ion–dipole (highest), hydrogen bonding (medium-high), dipole–dipole (medium), dipole–induced dipole (medium-low), and dispersion (lowest). The ranking explains many physical trends such as solubility, boiling point, and phase at room temperature.

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Boiling Point Ranking

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To rank boiling points, identify the dominant intermolecular forces, then order by their relative strengths and consider molecular size/shape. For example, among $CH_3OH$, $CH_3CH_2CH_2CH_3$, and $CH_3CH_2OCH_3$, $CH_3CH_2CH_2CH_3$ is lowest (dispersion only), $CH_3CH_2OCH_3$ is middle (dipole + dispersion), and $CH_3OH$ is highest (hydrogen bonding).

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Vapor Pressure

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The pressure exerted by a substance’s molecules in the gas phase when in equilibrium with its liquid at a given temperature. Vapor pressure rises with temperature and decreases with stronger intermolecular forces, and it is also influenced by surface area. High vapor pressure corresponds to volatile substances.

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Factors Affecting Vapor Pressure

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Key factors are temperature (higher temperature increases vapor pressure), surface area (more molecules at the surface can evaporate), and intermolecular forces (stronger forces lower vapor pressure). These factors determine how readily molecules escape the liquid into the gas phase.

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Clausius–Clapeyron Equation

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A thermodynamic relation that quantifies how vapor pressure changes with temperature. In logarithmic form it is written as: $$\ln\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_{vap}}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)$$ where $\Delta H_{vap}$ is the enthalpy of vaporization, $R$ is the gas constant, and $T$ values are in kelvin. This equation is used to calculate vapor pressures at different temperatures.

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Vapor Pressure Calculation

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To calculate vapor pressure at a new temperature use the Clausius–Clapeyron equation, converting $\Delta H_{vap}$ to J/mol and temperatures to kelvin. For example, using $T_1=373\,\text{K}$ (normal boiling) and $T_2=308\,\text{K}$ with $\Delta H_{vap}=40.7\,\text{kJ/mol}$ yields the vapor pressure at 35°C. Careful unit conversion and significant-figure handling are essential.

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Phase Diagram

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A graphical representation of the stability regions of solid, liquid, and gas as functions of temperature and pressure, with equilibrium lines showing phase boundaries. It identifies key features such as the triple point and critical point, and indicates a supercritical region where liquid and gas phases are indistinguishable.

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Triple Point

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The single temperature and pressure at which all three phases (solid, liquid, gas) of a substance coexist in equilibrium. The triple point is a unique thermodynamic state useful for defining temperature scales and characterizing substances.

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Critical Point

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The temperature and pressure at which the densities of liquid and gas phases become equal and the distinction between the two phases vanishes. Above this point, the substance becomes a supercritical fluid with properties intermediate between a liquid and a gas.

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Surface Tension

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The energy required to increase the surface area of a liquid, reflecting the net inward cohesive forces experienced by molecules at the surface. Stronger intermolecular attractions increase surface tension, which explains phenomena like droplets and the ability of small insects to walk on water.

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Viscosity

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A measure of a liquid’s resistance to flow that depends on intermolecular forces, molar mass, molecular shape, and temperature. Stronger interactions and larger or more entangled molecules increase viscosity, while raising temperature typically lowers viscosity.

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Cohesive vs Adhesive

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Cohesive forces are attractions between like molecules, while adhesive forces are attractions between unlike substances (e.g., liquid and glass). The balance between them determines meniscus shape: concave when adhesive ≥ cohesive (e.g., $H_2O$ on glass) and convex when cohesive > adhesive (e.g., mercury on glass).

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Capillary Action

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The rise or fall of a liquid in a narrow tube due to adhesive forces between the liquid and tube walls and cohesive forces within the liquid. Capillary action is driven by surface tension and explains how plants move water through xylem.

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Density of Ice

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Hydrogen bonding arranges water molecules into an open, cage-like lattice in the solid state, making ice less dense than liquid $H_2O$. As a result, ice floats and bodies of water freeze from the top down, which has important ecological consequences for aquatic life.

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Solubility Principle

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Solubility depends on the relative strengths of solute–solvent interactions versus solute–solute and solvent–solvent interactions; a useful heuristic is "like dissolves like." Polar solvents favor dissolution of polar or ionic solutes, while nonpolar solvents favor nonpolar solutes. Solubility also depends on temperature and pressure for gases.

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Predicting Solubility Example

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To predict solubility in $H_2O$, evaluate molecular polarity: polar species such as $NH_3$ and $HF$ form favorable interactions with water and are highly soluble, whereas nonpolar species like $CCl_4$ and $O_2$ have limited solubility. Molecular geometry and hydrogen-bonding capability are key determinants.

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Hydrophobic vs Hydrophilic

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Hydrophobic describes species that repel water and have diminished water solubility due to nonpolar character and strong dispersion among themselves. Hydrophilic describes species that attract water and readily dissolve because they form favorable dipole–dipole or hydrogen-bonding interactions with the solvent.

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Solubility of Gases

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Gas solubility in water decreases with increasing temperature because higher kinetic energy allows gas molecules to escape the solvent. Solubility increases with the partial pressure of the gas above the liquid due to more frequent collisions and greater opportunity for solute–solvent interactions.

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Henry’s Law

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A quantitative relation between gas solubility and partial pressure above a solvent given by $C_{gas}=k_H P_{gas}$, where $k_H$ is the Henry’s law constant. This linear relation allows prediction of dissolved gas concentration from the gas partial pressure and is temperature dependent.

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Henry’s Law Example

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For $CO_2$ in a soft drink bottled at $P_{CO_2}=5.0\,\text{atm}$ with $k_H=3.1\times10^{-2}\,\text{mol⋅L}^{-1}\text{⋅atm}^{-1}$, the dissolved concentration is calculated as $$C = k_H P = (3.1\times10^{-2})(5.0) = 0.155\,\text{mol L}^{-1}. $$ This shows that nonpolar $CO_2$ has modest solubility in polar $H_2O$ due to relatively weak dipole–induced interactions.

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Combination of Forces

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Real systems often experience multiple intermolecular forces simultaneously; the net behavior reflects the balance of these interactions. For example, extensive dispersion interactions can limit octanol solubility in water, while polar regions and hydrogen-bonding groups promote aqueous solubility. Understanding combinations explains microscopic and macroscopic properties.