Liquids, Solids, and Intermolecular Forces (Sec 3)

Flashcards created using AI and the note slideshow, but (currently) being looked over by me.

What is the primary difference between intramolecular forces and intermolecular forces (IMFs)?

Intramolecular forces are true chemical bonds that hold atoms together within a single molecule and are stronger, whereas intermolecular forces are weaker attractions that keep separate molecules together.

How are intermolecular forces typically represented in chemical diagrams, and what physical properties do they influence?

They are represented by dashed lines; their strength and quantity determine physical properties such as melting point, boiling point, and vapor pressure.

What are London dispersion forces (LDFs), and what mechanism causes them?

LDFs are the weakest intermolecular forces, caused by instantaneous, momentary dipoles that arise from the random movement of electrons and collisions between molecules.

In which types of molecules are London dispersion forces found?

London dispersion forces exist in all molecules, including nonpolar atoms and molecules.

What is polarizability, and how does it correlate with the strength of London dispersion forces?

Polarizability is the ease with which a molecule's electron cloud can be distorted; molecules with more electrons have greater polarizability, resulting in stronger London dispersion forces.

On the AP Chemistry exam, what factor must be cited to explain why heavier halogens have higher boiling points ($F_2 < Cl_2 < Br_2 < I_2$)?

The explanation must focus on the number of electrons (greater electron cloud polarizability leading to stronger LDFs), not molecular mass.

What are dipole-dipole forces, and how do polar molecules orient themselves?

They are attractions between polar molecules with permanent dipoles, where molecules arrange themselves so that opposite partial charges align ($\delta^+$ to $\delta^-$).

How does molecular distance affect dipole-dipole attractions in the gas phase?

Attractive dipole-dipole forces dissipate significantly with increased distances, rendering them minimal in gases.

What is an ion-dipole interaction, and how does its strength compare to dipole-dipole forces?

An attraction between an ion and a polar molecule; it is stronger than a standard dipole-dipole interaction.

Why is the triiodide ion ($I_3^-$) significantly more soluble in water than elemental iodine ($I_2$)?

$I_3^-$ forms strong ion-dipole interactions with polar water molecules, whereas nonpolar $I_2$ cannot.

What specific chemical criteria are necessary for a molecule to form hydrogen bonds?

A hydrogen atom must be directly covalently bonded to a highly electronegative atom: nitrogen ($N$), oxygen ($O$), or fluorine ($F$).

Why do the period 2 hydrides $H_2O$, $HF$, and $NH_3$ display anomalously high boiling points compared to heavier group hydrides?

They form strong hydrogen bonds and have a small molecular size that permits closer packing and stronger dipole attractions.

What intermolecular forces are active in a pure sample of a substance exhibiting hydrogen bonding?

All three types of intermolecular forces: London dispersion forces, dipole-dipole forces, and hydrogen bonding.

What is surface tension, and how is it related to intermolecular forces?

Surface tension is the resistance of a liquid to increase its surface area (causing liquid beading); stronger IMFs produce greater surface tension.

What is capillary action, and what two distinct forces govern it?

Capillary action is the spontaneous rising of a liquid in a narrow tube; it is governed by cohesive forces (between liquid molecules) and adhesive forces (between liquid molecules and the container walls).

Why does water exhibit a concave meniscus in glass while liquid mercury ($Hg$) exhibits a convex meniscus?

In water, adhesive forces with glass exceed cohesive forces; in mercury, cohesive forces between $Hg$ atoms exceed adhesive forces with glass.

How does tube diameter influence the height of liquid rise in capillary action?

A narrower tube provides more relative surface area of glass, causing the liquid to rise higher.

What is viscosity, and what conditions cause it to increase?

Viscosity is a liquid's resistance to flow; it increases with stronger intermolecular forces, larger molecular size, and lower temperatures.

What is volatility, and how does it depend on intermolecular forces?

Volatility is the ability of a liquid to evaporate readily; liquids with higher IMFs have lower volatility (are less volatile).

What thermodynamic energy steps occur during the dissolution of a solute in a solvent?

Energy is required to overcome solute-solute IMFs and solvent-solvent IMFs ($\Delta H > 0$), while energy is released when solute-solvent attractions form ($\Delta H < 0$).

Why shouldn't a student simply write 'like dissolves like' on the AP Chemistry exam?

AP credit requires explaining solubility through the balance of energy required to overcome solute-solute and solvent-solvent IMFs versus the energy released when solute-solvent attractions form.

Why do polar or ionic substances fail to dissolve in nonpolar solvents?

The energy released ($\Delta H$) by weak solute-solvent interactions is too small to overcome the strong solute-solute attractions present in the polar or ionic compound.

How does molecular geometry (linear vs. spherical/branched) influence the magnitude of dispersion forces?

Linear or flat molecules have greater surface area contact, leading to stronger IMFs, whereas spherical or branched molecules have reduced surface contact and weaker IMFs.

What distinguishes crystalline solids from amorphous solids?

Crystalline solids possess a highly regular, repeating arrangement of components (e.g., ionic crystals), whereas amorphous solids exhibit considerable disorder in their structure (e.g., glass).

What are molecular solids, and what melting point characteristics do they typically have?

Solids containing discrete molecules held at lattice points by intermolecular forces; they typically have low melting points (e.g., ice, dry ice, $P_4$, $S_8$).

How does closest packing in ionic solids minimize electrostatic potential energy?

Large anions pack tightly while small cations nestle into interstitial holes, maximizing cation-anion ($+/-$) attractions and minimizing like-charge repulsions.

What is the electron sea model of metallic bonding?

A model representing a regular lattice of metal cations surrounded by a delocalized, freely moving 'sea' of valence electrons that conduct electricity and heat.

What type of alloy is depicted in this atomic diagram containing copper ($\text{Cu}$) and zinc ($\text{Zn}$)?

A substitutional alloy, where host metal atoms are replaced by solute atoms of similar atomic radius.

What type of alloy is represented in this diagram showing iron ($\text{Fe}$) and carbon ($\text{C}$)?

An interstitial alloy, where small nonmetal atoms occupy the interstitial holes in a closely packed metal lattice, making it more rigid and dense.

What is enthalpy of vaporization ($\Delta H_{vap}$), and why is vaporization inherently endothermic?

It is the energy required to vaporize one mole of liquid at $1\,atm$; it is always endothermic because energy must be absorbed to overcome intermolecular forces.

What is vapor pressure, and how does it vary with temperature and molar mass?

The pressure of vapor at dynamic equilibrium with its liquid in a closed container; it increases as temperature rises and decreases as molar mass (and LDF strength) increases.

What is defined as the normal boiling point of a liquid?

The temperature at which the vapor pressure of the liquid equals exactly $1\,atm$ ($760\,mmHg$).

Why is $\Delta H_{fus}$ always significantly less than $\Delta H_{vap}$ for any given substance?

Melting only partially disrupts some of the intermolecular attractions holding particles together, whereas vaporization requires overcoming all intermolecular forces.

On this heating curve for water, why does temperature remain constant across the horizontal plateaus at $0\,^\circ\text{C}$ and $100\,^\circ\text{C}$?

Energy supplied during phase changes alters potential energy by breaking intermolecular forces ($\Delta H_{fus}$ and $\Delta H_{vap}$), keeping average kinetic energy constant.

What are supercooling and superheating?

Supercooling is when rapid cooling prevents molecules from organizing, keeping a liquid below its freezing point; superheating is when rapid heating prevents bubble nucleation, keeping a liquid above its boiling point.

What do the lines (phase boundaries) on a pressure-temperature phase diagram represent?

Sets of temperature and pressure conditions where two adjacent phases coexist in dynamic equilibrium.

What is the critical point on a phase diagram?

The point defined by critical temperature ($T_c$) and critical pressure ($P_c$) beyond which liquid and gas become indistinguishable and form a supercritical fluid.

What physical property of water is indicated by the negative slope (left tilt) of the solid-liquid line in its phase diagram?

Ice (solid water) is less dense than liquid water, meaning increasing pressure favors the liquid state.

According to this phase diagram for carbon dioxide ($CO_2$), what occurs when solid $CO_2$ is warmed at $1.0\,atm$?

It sublimes directly from solid to gas at $-78\,^\circ\text{C}$ without passing through a liquid phase, because its triple point lies above atmospheric pressure at $5.1\,atm$.