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Liquids and Solids

by: Caitrín Hall

Liquids and Solids CHEM 1127Q 001

Caitrín Hall
GPA 3.9

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About this Document

These notes cover chapter 10 of the textbook. It will be included on the final exam.
General Chemistry
Fatma Selampinar (TC), Joseph Depasquale (PI)
Class Notes
Chemistry, notes, outline
25 ?




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This 5 page Class Notes was uploaded by Caitrín Hall on Saturday April 30, 2016. The Class Notes belongs to CHEM 1127Q 001 at University of Connecticut taught by Fatma Selampinar (TC), Joseph Depasquale (PI) in Spring 2016. Since its upload, it has received 12 views. For similar materials see General Chemistry in Chemistry at University of Connecticut.

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Date Created: 04/30/16
Chapter 10Liquids and Solids  Unlike with gases, the properties of liquids and solids depend on chemical identity 10.1 Intermolecular Forces   Intermolecular interaction refers to attractive forces between the particles of a substance, regardless of whether these particles are molecules, atoms, or ions  Intermolecular forces (IMFs) – the various forces of attraction that may exist between the atoms and molecules of a substance due to electrostatic phenomena; serve to hold particles close together  KE provides the energy required to overcome attractive forces  Phase changes occur when conditions of temperature or pressure favor the associated changes in IMF  Increased pressure brings molecules closer, while increased temp increases KE  If temp becomes sufficiently low or pressure becomes sufficiently high, the molecules don’t have enough KE to overcome the IMF  solid forms Forces between Molecules  Intermolecular forces occur between particles, while intramolecular forces are those within the molecule that keep it together (ex: bonds between atoms)  IMF determine many physical properties  IMFs between small molecules are often weak compared to the intramolecular forces that bond atoms together  Van der Waals forces – all attractive forces between neutral atoms and molecules Various types of IMFs:  Dispersion forces, dipole-dipole forces, and hydrogen bonding Dispersion Forces  The London dispersion force is present in all condensed phases  Dispersion force – attraction between two rapidly fluctuating, temporary dipoles;  significant only when particles are very close together   Because electrons are in constant motion, an atom/molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically  An induced dipole results when an instantaneous dipole distorts the electrons of a neighboring atom/molecule o Both result in weak, electrostatic dispersion forces  Polarizability – the measure of how easy or difficult it is for another electrostatic charge to distort a molecule’s charge distribution (its electron cloud) o If a charge cloud is easily distorted, it is very polarizable and will have large dispersion forces o In larger atoms, valence electrons are farther from the nuclei  less tightly held  more easily form temporary dipoles that produce the attraction  Shapes of molecules also affect magnitudes of dispersion forces o Greater surface area available for contact between molecules  stronger dispersion forces  higher boiling point Dipole-Dipole Attractions  Dipole-dipole attraction – the electrostatic force between the partially positive end of one polar molecule and the partially negative end of another  Present in polar molecules only  Stronger than dispersion forces  Two different substances with the same MM can have different boiling points if one substance is polar and the other isn’t o Presence of dipole-dipole attraction  higher boiling point Hydrogen Bonding  Strongest van der Waals force, but much weaker than covalent bonds  Hydrogen bonding – strong, type of dipole-dipole attraction that occurs when a molecule contains an H atom bonded to F, O, or N (most electronegative atoms)  Intermolecular attractive force  The large difference in electronegativity between hydrogen and F, O, or N combined with the very small size of an H atom and the relatively small sizes on F, O, or N atoms leads to highly concentrated partial charges  Hydrogen bonds are denoted by dots connecting atoms  The effect of increasingly stronger dispersion forces down a group dominates that of increasingly weaker dipole-dipole attractions  boiling points increase steadily  Hydrogen bonding molecules exhibit anomalously high boiling points  Hydrogen bonding in DNA – A and T share two H bonds while C and G share three 10.3 Phase Transitions Vaporization and Condensation  Gas  liquid = condensation  Liquid  gas = vaporization  In a closed container, gas molecules collide with the surface of the condensed phase; some collisions result in molecules re-entering the condensed phase  When rate of condensation = rate of vaporization, neither the amount of liquid nor the amount of vapor in the container changes  Dynamic equilibrium – the status of a system in which reciprocal processes occur at equal rates  Vapor pressure – the pressure exerted by the vapor in EQ with a liquid in a closed container at a given temperature o Does NOT depend on surface area of contact with container o Does depend on IMF o Strong IMF impede vaporization and favor the recapture of gas- phase molecules  low vapor pressure o Weak IMF prevent less of a barrier to vaporization and a reduced likelihood of gas recapture  high vapor pressure o As temp increases, VP of a liquid also increases due to increased average KE o Escape of more molecules per unit of time and greater average speed of molecules that escape both contribute to higher VP Boiling Points  Boiling point – the temperature at which a liquid’s EQ VP = the pressure exerted on the liquid by its gaseous surroundings  Normal boiling point – boiling point of a liquid when surrounding pressure = 1 atm  Temp remains constant throughout boiling process  Quantitative relation between a substance’s VP and its temperature: −ΔH vap /RT P = Ae Enthalpy of Vaporization  Vaporization is endothermic  Energy change associated with vaporization is enthalpy of vaporization, ΔHvap Melting and Freezing  Melting – energy becomes large enough to overcome molecules in their fixed positions, and the solid transitions to liquid state o Temp of solid stops rising despite continual input of heat and remains constant until all of the solid is melted  Freezing – the reciprocal process of melting  The temp at which the solid and liquid phases of a given substance are in EQ is called the melting point of the solid or the freezing point of the liquid Sublimation and Deposition  Sublimation – solids transition directly into the gaseous state; ex: dry ice  Deposition – the reverse of sublimation; ex: formation of frost


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