Chapter 03_liquid,solid N Phase Changes

Chapter 3: Liquids, Solids & Phase Changes

Chapter Objectives: • To learn the differences between the solid, liquid, and gas state, and how the polarity of molecules influences those states. • To learn the different types of intermolecular forces between different molecules.

Kinetic Molecular Theory of Liquids and Solids • • •

Liquids and solids have significant interactions. Liquids and solids have well-defined volume. Liquid molecules “flow,” while solids are held “rigid.”

A molecular comparison of gases, liquids, and solids. (b) In gases, the particles feel little attraction for one another and are free to move about randomly. (b) In liquids, the particles are held close together by attractive forces but are free to move over one another. (c) In solids, the particles are rigidly held in an ordered arrangement.

Polar Bonds

• Individual covalent bonds are polar if the atoms being connected are of different electronegativities. This is described as a bond dipole. Example: CH3Cl The C—H bonds are nonpolar since C and H have about the same electronegativity. Since Cl is more electronegative than C, the C—Cl bond is polarized so that the Cl atom is slightly electron-rich (partial negative charge,δ-) and the C atom is slightly electron-poor (partial positive charge, δ+). This bond is a polar covalent bond (or just polar bond).

Polar Bonds •

Dipole Moment (µ): the sum of all the bond dipoles within a molecule, (net polarity) can be illustrated with an electrostatic potential map.

Since CH3Cl has a tetrahedral shape, with one polar bond and three nonpolar bonds, there is an overall dipole moment pointing towards the Cl end of the molecule. Depending on the number and orientation of the bond dipoles, the molecule may possess an overall molecular dipole.

Ammonia and water

Structures and dipole moments for ammonia and water are shown

Carbon dioxide and tetrachloromethane

Structures and dipole moments are shown for carbon dioxide and tetrachloromethane. Although each molecule has bond dipoles, they do not have molecular dipoles

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Intramolecular & Intermolecular Forces Intramolecular forces operate within each molecule, influencing the chemical properties of the substance (i.e., covalent bonds). These are the forces that hold the atoms in a molecule together. They are very strong forces which result from large charges (on protons and electrons)interacting over very short distances.

• Intermolecular forces (van der Waals forces) operate between separate molecules, influencing the physical properties of the substance. These are the forces that hold liquids and solids together, and influence their melting and boiling points. They are weaker forces, because they result from smaller charges, or partial charges, interacting over much larger distances.

Intramolecular & Intermolecular Forces To break an O—H bond in water, the water must be heated to thousands of degrees C; to completely overcome the intermolecular forces, all you have to do is boil it — 100ºC.

Intramolecular & Intermolecular Forces

(a) In an individual N2 molecule, atoms are held together by strong intramolecular force(covalent bond). Different N2 molecules are weakly attracted to one another at low temperature by intermolecular forces, causing nitrogen to become liquid. (b) At a higher temperature, intermolecular forces are no longer able to keep molecules close together, so nitrogen becomes a gas.

Intermolecular Forces • An attractive interaction between molecules – Determine bulk properties of matter.

• Much weaker than intramolecular forces • Several types of intermolecular (IM) forces: Ion–dipole Dipole–dipole London dispersion forces Hydrogen “bonds.”

London Dispersion Forces  Attraction is due to instantaneous, temporary dipoles formed due to electron motions  The larger the molar mass of a molecule, generally the greater the LDF’s.  Dispersion forces exist between all molecules, but they are the only forces that exist between nonpolar molecules.

London Dispersion Forces

Using molecular bromine as an example, each molecule has zero polarity. However, due to the motion of electrons at any given instant a temporary dipole would arise that would then induce a dipole in an adjacent molecule. This type of intermolecular attraction is called London dispersion forces.

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Dipole-Dipole Forces

Dipole-Dipole forces are the attractions between the opposite partial charges in the permanent dipoles of polar molecules.  exit between all polar molecules, (HCl – HCl)

(a) Polar molecules attract one another and approach closely when oriented with unlike charges together, but (b) they repel one another and push apart when oriented with like charges together.

Dipole-Dipole Forces • In general, for molecules of the same molecular weight, a polar molecule (dipole-dipole + London) will have a higher boiling point than a nonpolar molecule (London only):

Hydrogen Bonding

Attractive interaction between a hydrogen atom bonded to a very electronegative atom (O, N, or F) and an unshared electron pair on another electronegative atom.

Hydrogen bonds are also found between molecules of water and molecules of ammonia.

Ion-Dipole Forces • Ion-Dipole forces are the result of electrical interactions between an ion and the partial charges on a polar molecule.

Polar molecules orient toward ions so that (a) the positive end of the dipole is near an anion and (b) the negative end of the dipole is near a cation.

Ion-Dipole Forces • These forces are responsible for the ability of polar solvents like water to dissolve ionic compounds.

Ion- and Dipole-Induced Dipole Forces •

Ion - induced dipole forces are the attractive forces that exist between ions and nonpolar molecules.

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Being next to an ion “induces” a dipole in a nonpolar molecule, attracting it towards the ion.  These forces are responsible for the attraction between Fe2+ and O2 molecules in the bloodstream, and contributes to the solvation of ions in water.

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Dipole - induced dipole forces are the attractive forces that exist between polar molecules and nonpolar molecules. Being next to a polar molecule “induces” a dipole in a nonpolar molecule.  These forces are responsible for the solvation of gases (nonpolar) in water (polar).

Summary: Intermolecular Forces Intermolecula Formed by the Examples r Forces Ion-dipole

attraction … an between ion and a polar molecule

Strength

Hydrogen bond Dipole - Dipole

molecules which have H on N, O, or F atoms two polar molecules

Ion - Induced dipole

an ion and a nonpolar molecule

Dipole Induced dipole London

a polar molecule and a nonpolar molecule two nonpolar molecules

(dispersion) forces

Na+ and H2O H2O and H2O; H2 O and CHand OH CH3Br 3CH2ICl; CH3Br and H2O Fe2+ and O2 HCl and Cl2 CH4 and CH4; F2 and F2; CH4 and F2

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Intermolecular Force Surface TensionEffects  The resistance of a liquid to spread out and increase its surface area.  Surface tension results from intermolecular force differences between molecules in the interior of a liquid and those on the surface. Molecules at the surface of a liquid feel attractive forces only one side and are thus pulled in toward the liquid, while molecules in interior are surrounded and are pulled equally in all direction.

Intermolecular Force Effects Surface tension, which causes these drops of liquid mercury to form beads

Atoms on the surface are less stable because they have fewer neighbors and feel fewer attractive forces than atoms in the interior, so the liquid acts to minimize their number by minimizing the area of the surface.

Intermolecular Force Effects

More surface tension examples…

Intermolecular Force Effects

2. Viscosity

 The measure of a liquid’s resistance to flow and is related to the ease with which molecules move around, and thus to the intermolecular forces.  Substances composed of small, nonpolar molecules (such as gasoline and benzene) have low viscosities.  Polar molecules (such as glycerol) and molecules composed of long chains of atoms (such as oil and grease) have higher viscosities.  The viscosity of a liquid decreases at higher temperatures.

Intermolecular Force Effects

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Phase Changes

Enthalpy  the heat flow associated with making or breaking intermolecular attractions that hold liquids and solids together Entropy  change in a molecular randomness between various phases

Changes from a less random phase to a more random one have positive values of ΔH and Δ S. Changes from a more random phase to a less random one have negative values of Δ H and Δ S.

Phase Changes • Sublimation: The process in which molecules go directly from the solid into the vapor phase. • Deposition: The process in which molecules go directly from the vapor into the solid phase. • Molar heat of sublimation (∆Hsub): The energy (kJ) required to sublime one mole of solid.

∆ Hsub = ∆ H fus + ∆ Hvap

Phase Changes • Molar Heat of Fusion (∆Hfus): The energy required to melt one mole of solid (in kJ). • Molar Heat of Vaporization (∆Hvap): The energy (in kJ) required to vaporize one mole of liquid.

Heating curve for water

A heating curve for H2O, showing the temperature changes and phase transitions that occur when heat is added. The plateau at 0°C represents the melting of solid ice, and the plateau at 100°C represents the boiling of liquid water. Plateau regions in a heating curve indicate a change in phase of the substance where the temperature remains constant

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Phase Changes Vapor Pressure: The pressure exerted by gaseous molecules above a liquid.

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Liquids after sitting for a length of time in (a) an open container and (b) a closed container. The liquid in the open container has evaporated, but the liquid in the closed container has brought about a rise in pressure. • Evaporation of a liquid results in more gas phase molecules which exert a pressure in a closed container

Phase changes depend on temperature.

The distribution of molecular kinetic energies in a liquid at two temperatures. Only the faster-moving molecules have sufficient kinetic energy to escape from the liquid and enter the vapor. The higher the temperature, the larger the number of molecules with enough energy to escape.

Kinds of Solids • Solids are divided into two categories:  Crystalline: - Possesses rigid and long-range order - Flat faces - Distinct angles - eg. NaCl  Amorphous - Lacks well-defined arrangement (particles are randomly arranged) - have no long ranged structure - eg. rubber • Structure of a crystalline solid is based on the unit cell, a basic repeating structural unit.

Crystalline Solids

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Crystal structures of (a) ice, a molecular solid, and (b) quartz, a covalent network solid. Ice consists of individual H2O molecules held together in a regular manner by hydrogen bonds. Quartz (SiO2) is essentially one very large molecule whose Si and O atoms are linked by covalent bonds. Each silicon atom has tetrahedral geometry and is bonded to four oxygens; each oxygen has approximately linear geometry and is bonded to two silicons. The shorthand representation on the right shows how SiO4 tetrahedra join at their corners to share oxygen atoms.

Cubic Packing

Simple Cube and Body-Centered Cube:

(a) Simple cubic packing of spheres  all the layers are identical and all atoms are lined up in stacks and rows. (b) Body-centered cubic packing of spheres  spheres in layer a are separated slightly and the spheres in layer b are offset so that they fit into the depressions between atoms in layer a.

Hexagonal and Cubic ClosestPacking a)

Hexagonal closest-packing. Two alternating hexagonal layers (a and b) offset from each other so that the spheres in one layer sit in the small triangular depressions of neighboring layers.

i)

Cubic closest-packing of spheres Three alternating hexagonal layers, (a, b, and c) offset from one another so that the spheres in one layer sit in small triangular depressions of neighboring layers.

Cubic Unit Cells

Simple Cube and Body-Centered Cube:

Geometries of (a) primitive-cubic and (b) body-centered cubic unit cells in both a skeletal view (top) and a space-filling view (bottom). Part (c) shows how eight primitive-cubic unit cells stack together to share a common corner where they meet.

Cubic Unit Cells Face-Centered Cube:

(a) Geometry of a facecentered cubic unit cell, (b) a view showing how this unit cell is found in cubic closest-packing. The faces are tilted at 54.7° angles to the three repeating atomic layers

Phase Diagrams • A Phase Diagram is a graphical display of the temperatures and pressures at which two phases of a substance are in equilibrium. – Triple Point: The only condition under which all three phases can be in equilibrium with one another. – Critical Temperature (Tc): The temperature above which the gas phase cannot be made to liquefy at any pressure. – Critical Pressure (Pc): The minimum pressure required to liquefy a gas at its critical temp.

Phase diagram for water

A phase diagram for H2O, showing a negative slope for the solid/liquid boundary.

Phase diagram for carbon dioxide

A phase diagram for CO2, showing a positive slope for the solid/liquid boundary. The pressure and temperature axes are not to scale.

Formation of supercritical fluid Here we can see the separate phases of carbon dioxide. The meniscus is easily observed.

With an increase in temperature the meniscus begins to diminish. Increasing the temperature further causes the gas and liquid densities to become more similar. The meniscus is less easily observed but still evident. Once the critical temperature and pressure have been reached the two distinct phases of liquid and gas are no longer visible. The meniscus can no longer be seen. One homogenous phase called the "supercritical fluid" phase occurs which shows properties of both liquids and gases.