Chemistry: The Central Science Chapter 11: Intermolecular Forces, Liquids, and Solids Intramolecular forces within molecules that give rise to covalent bonding influence molecular shape, bond energies, and many aspects of chemical behavior The physical properties of molecular liquids and solids are largely due to the intermolecular forces between the molecules 11.1: A Molecular Comparison of Gases, Liquids, and Solids Gases consist of collection of widely separated particles in constant, chaotic motion o The average energy of the attractions between the particles is much smaller than their average kinetic energy Lack of strong attractive forced allows a gas to expand to fill its container Liquids have stronger intermolecular attractive forces so they are much denser and far less compressible than gas o The intermolecular attractive forces are no strong enough to keep particles from moving past each other thus liquid flows Solids have stronger intermolecular forces than liquid that the molecules are virtually locked in a place o Often referred to solids and liquids as condensed phases because the particles are fairly close together compared to gas o Possessed highly ordered three-dimensional structures are said to be crystalline o Rigid because the particles are not free to move around Units that form the solid, whether ions or molecules, possess thermal energy and vibrate in place The state of a substance depends largely on the balance between the kinetic energies of the particles and the interparticle energies of attraction o Heating give the particles higher average kinetic energy, allowing them to be able to overcome the intermolecular attractive forces o Cooling take away the particles average kinetic energy, causing the intermolecular energy to overcome the average kinetic energy o Increasing pressure forces the particles to be closer together, which in turn increases the strength of the intermolecular forces of attraction 11.2: Intermolecular Forces
The strength of intermolecular forces of different substances vary over a wide range, but they are generally much weaker than ionic or covalent bonds o Less energy is required to overcome the forces Properties such as boiling points and melting points reflect the strengths of the intermolecular forces o The stronger the intermolecular forces, the higher the boiling and melting points are Intermolecular attractions o Three types of intermolecular attractions exist between neutral molecules Dipole-dipole attractions, London dispersion forces, and hydrogen bonding Dipole-dipole attractions and London dispersion forces are collectively called van der Waals forces o Another kind of attractive force, the ion-dipole force, is important in solutions o All of these intermolecular interactions are electrostatic in nature, involving attractions between positive and negative species o Even at their strongest, these interactions are much weaker than covalent or ionic bonds Ion-Dipole Forces o Exists between an ion and the partial charge on the end of a polar molecule o Positive ions are attracted to the negative end of a dipole, whereas negative ions are attracted to the positive end o Ion-dipole forces are especially important for solutions of ionic substances in polar liquids, such as NaCl in water Dipole-Dipole Forces o Neutral polar molecules attract each other when the positive end of one molecule is near the negative end of another o Are effective only when polar molecules are very close together o Generally weaker than ion-dipole forces o The molecules orientation change according to the repulsive and attractive force between molecules o For molecules of approximately equal mass and size, the strengths of intermolecular attractions increase with increasing polarity o For dipole-dipole forces to operate, the molecules must be able to get close together in the correct orientation Those with smaller molecular volume generally experience higher dipole-dipole attractive forces London Dispersion Forces
o Attraction results of the movement of electron to one side, causing the molecule or atom to have a temporary induce dipole Temporary dipole on one atom can induce a similar temporary dipole on adjacent atom o The dispersion force is significant only when molecules are very close together o The strength of the dispersion force depends on the ease with which charge distribution in a molecule can be distorted, a.k.a. the molecule s polarizability More polarizable molecules have larger dispersion forces In general, larger molecules tend to have greater polarizabilities because they have a greater number of electrons, and their electrons are farther from the nuclei o The dispersion force tend to increase in strength with increasing molecular weight (size and weight are generally parallel to each other) o Shapes of molecules also influence the magnitudes of dispersion forces The more contact, the greater the dispersion force o Dispersion forces operate between all molecules Polar molecules experience dipole-dipole interactions and dispersion forces at the same time Dispersion forces between polar molecules commonly contribute more to intermolecular attractions than do dipole-dipole Some generalizations should be considered when comparing the relative strengths of intermolecular attractions When the molecules of two substances have comparable molecular weights and shapes, dispersion forces are approximately equal in the two substances o In this case, differences in the magnitudes of the attractive forces are due to differences in the strengths of dipole-dipole attractions, with the more polar molecules having the stronger attractions When the molecules of two substances differ widely in molecular weights, dispersion forces tend to be decisive in determining which substance has the stronger intermolecular attractions o In this case, differences in the magnitudes of attractions can usually be associated with difference in molecular weights, with the more massive substance having stronger attractions
Hydrogen Bonding o Are generally stronger than dipole-dipole or dispersion force o Typically much weaker than ordinary chemical bonds o Is a special type of intermolecular attraction between the hydrogen atom in a polar bond (particularly and H F, H O, or H N bond) and nonbonding electron pair of F, O, or N in another molecule o Because the electron-poor hydrogen is so small, it can approach an electronegative atom very closely and thus interact strongly with it o Ice have lower density than water Possibly because the hydrogen bonding caused the water molecule to form an ordered, open arrangement which create the open cavities between water molecules and thus is further apart than liquid water Profoundly affects life on earth Ice float which allow aquatic life in cold weather Comparing Intermolecular Forces o Dispersion forces are found in all substances Dipole-dipole forces add to the effect of dispersion forces in polar molecules Hydrogen bond also add to the effect of dispersion forces o Hydrogen bond tend to be the strongest type of intermolecular attraction o Energy associated with the dispersion forces and dipole-dipole forces are in the range of 2-10 kj/mol, with hydrogen bond in the range of 5-25 kj/mol, while in the range of about 15 kj/mol for ion-dipole attractions o The effects of all these attractions are additive 11.3: Some Properties of Liquids Viscosity o Viscosity is the resistance of a liquid to flow It can be measured by measuring the amount of time for the liquid to flow through a thin tube under gravity, as well as measuring the rate at which steel balls fall through the liquid o Viscosity is related to the ease with which individual molecules of the liquid can move with respect to one another Viscosity increases with molecular weight Viscosity decreases with increasing temperature Surface Tension At higher temperature, the molecules have more energy to more easily overcomes the attractions between molecules
11.4: Phase Changes o The water molecules at the surface are pulled by the hydrogen bond from the interior, causing the molecules at the surface to be pack closely together This caused water droplets to be spherical because spheres have the smallest surface area o Surface tension the energy required to increase the surface area of a liquid by a unit amount Water has high surface tension because of its strong hydrogen bond Mercury has even higher surface tension because of its strong metallic bonds o Cohesive forces intermolecular forces that bind similar molecules to one another o Adhesive forces intermolecular forces that bind a substance to a surface Meniscus of water is curve downward because the adhesive forces is stronger than the cohesive forces Meniscus of mercury is curve upward because the cohesive forces is stronger than the adhesive forces o The rise of liquids up very narrow tubes is called capillary action Adhesive forces between the liquid and the walls of the tube tends to increase the surface area of the liquid Surface tension of the liquid tends to reduce the area, thereby pulling the liquid up the tube Use in plants to move water and nutrients upward In general, each state of matter can change into either of the other two states o These transformations are called either phase changes or changes of state Energy Changes Accompanying Phase Changes o The process of melting is called fusion Heat of fusion, or enthalpy of fusion (ΔH fus ), is the measure of the energy needed for the process o Vaporization is the process of turning into gas Heat of vaporization, or enthalpy of vaporization (ΔH vap ), is the measure of energy required for this process As temperature of the liquid increase, vapor pressure, which is the pressure exert by gas-phase molecules over the liquid, increases o ΔH vap values tend to be larger than ΔH fus because in the transition from liquid to the vapor state, the molecules must essentially sever all their
intermolecular attractions; whereas in melting, many of these attractions remain o Heat of sublimation (ΔH sub ) is the enthalpy change required for the molecules of a solid to be transformed directly into the gaseous state Equals to the sum of ΔH vap and ΔH fus o Heat of deposition is exothermic to the same degree that the heat of sublimation is endothermic o Heat of condensation is equal in magnitude to the neat of vaporization but opposite in sign o Heat of freezing is exothermic to the same degree that the heat of fusion is endothermic Heating Curves o During the process of changing in state, the temperature of the substance does not change even though the energy is being removed or added The energy is used to overcome the attraction rather than increasing the average kinetic energy o The slope of the line during the change in temperature depends on the specific heat of the substance in that state o Sometimes as we remove heat from a liquid, we can temporarily cool it below its freezing point without forming a solid This is called supercooling Results of a rapid removal of heat, so quick that the molecules literally have no time to assume the ordered structure of a solid Supercooled liquid is unstable; particles of dust entering the solution or gentle stirring is often sufficient to cause the substance to solidify Critical Temperature and Pressure o Critical temperature the highest temperature that a distinct liquid phase can form If the temperature is over the critical temperature, no amount of pressure can cause the gas to liquefies o Critical pressure the pressure required to bring about the liquefaction at the critical temperature o Above the critical temperature, the motional energies of the molecules are greater than the attractive forces that lead to the liquid state The greater the intermolecular forces, the greater the critical temperature of a substance
11.5: Vapor Pressure o Nonpolar, low molecular weight substances, which have weak intermolecular attractions have lower critical temperatures and pressures than those that are polar or of higher molecular weight o Critical temperatures and pressures of substances are often of considerable importance to engineers and other people working with gases Molecules can escape from the surface of a liquid into the gas phase by evaporation o E.g. Ethanol in an evacuated closed container Ethanol will quickly begin to evaporate, resulting in an increase of pressure exert by the vapor in the space above liquid After a short time the pressure of the vapor will attain a constant value, which we call the vapor pressure of the substance Explaining Vapor Pressure on the Molecular Level o At any instant, some of the molecules on the surface of the liquid possess sufficient kinetic energy to overcome the attractive forces of their neighbors and escape into the gas phase (distribution curve of kinetic energy) The weaker the attractive forces, the larger is the number of molecules that are able to escape and therefore the higher the vapor pressure o As the number of gas-phase molecules increases, the probability increases that a molecule in the gas phase will strike the liquid surface and be recaptured by the liquid Eventually, the number of molecules in the gas phase then reaches a steady value, and the pressure of the vapor at this stage becomes constant o The condition in which two opposing processes are occurring simultaneously at equal rate is called a dynamic equilibrium, usually referred to as equilibrium A liquid and its vapor are in dynamic equilibrium when evaporation and condensation occur at equal rates o The vapor pressure of a liquid is the pressure exerted by its vapor when the liquid and vapor states are in dynamic equilibrium Volatility, Vapor Pressure, and Temperature o When vaporization occurs in an open container, water spreads away from the liquid and was hardly recaptured at the surface of the liquid Equilibrium never occurs o Substances with high vapor pressure evaporate more quickly than substances with low vapor pressure o Liquids that evaporate readily are said to be volatile
o Vapor pressure increases with increasing temperature Vapor pressure in all cases increases nonlinearly with increasing temperature Vapor Pressure and Boiling Point o A liquid boils when its vapor pressure equals the external pressure acting on the surface of the liquid o The temperature at which a given liquid boils increases with increasing external pressure o Boiling point of a liquid at 1 atm is called its normal boiling point 11.6: Phase Diagrams Equilibrium other than the dynamic equilibrium can exist between states of matter Phase diagram is a graphic way to summarize the conditions under which equilibria exist between the different states of matter o The diagram is a two-dimensional graph, with pressure and temperature as the axes o Contain three important curves, each which represents the conditions of temperature and pressure at which the various phases can coexist at equilibrium o The curves may be described as follows: The line that separates the liquid and gas is the vapor-pressure curve of the liquid The vapor-pressure curve ends at the critical point, which is at the critical temperature and pressure of the substance Beyond the critical point, the liquid and gas phases become indistinguishable from each other, and the state of the substance is a supercritical fluid The line that separates the solid phase from the gas phase represents the change in the vapor pressure of the solid as it sublimes at different temperatures The line that separates the solid phase from the liquid phase corresponds to the change in melting point of the solid with increasing pressure An increase in pressure usually favors the more compact solid phase; thus, higher temperatures are required to melt the solid at higher pressures The melting point at 1 atm is the normal melting point o The point in the middle is known as the triple point
At this temperature and pressure, all three phases are in equilibrium o Any point on the curve represents equilibrium between two phases o Any point on the diagram that does not fall on a line corresponds to conditions which only one phase is present The Phase Diagrams of H 2 O and CO 2 o The solid-liquid equilibrium line of CO 2 follows the typical behavior, slanting to the right with increasing pressure, telling you that its melting point increases with increasing pressure o The melting point of H 2 O is atypical, slanting to the left with increasing pressure, indicating that for water the melting point decreases with increasing pressure o For CO 2 to exist as liquid, the pressure must exceed 5.11 atm Solid CO 2 does not melt but sublimes when heated at 1 atm 11.7: Structures of Solids In a crystalline solid the atoms, ions, or molecules are ordered in well-defined threedimensional arrangements o These solids usually have flat surfaces, or faces, the make a definite angles with one another o The orderly stacks of particles that produce these faces also cause the solids to have highly regular shapes An amorphous solid is a solid in which particles have no orderly structure o Lack well-defined faces and shapes o Do not stack together well o Intermolecular forces vary in strength throughout a sample Do not melt at specific temperatures Unit Cells o The repeating unit of a solid is known as the unit cell o The three-dimensional array of points called a crystal lattice Each point in the lattice is called a lattice point o Three kinds of cubic unit cells are primitive cubic, body-centered cubic, and face-centered cubic Primitive cube The lattice points are at the corners only Body-centered cubic When lattice points are at the corners and the center of the unit cell Face-centered cubic
When the cell has lattice points at the center of each face, as well as at each corner o The atoms on the corners and faces do not lie wholly within the unit cell o In an individual primitive cubic unit cell, each corner contains only one-eighth of an atom Because a cube has eight corners, each primitive cubic unit cell has total of 1/8 8 = 1 atom The Crystal Structure of Sodium Chloride o We can center either the Na + ions or the Cl - ions on the lattice points of a facecentered cubic unit cell The structure can be described as being face-centered cubic o The total cation-to-anion ratio of a unit cell must be the same as that for the entire crystal Within the unit cell of NaCl there must be an equal number of Na + and Cl - ions Similarly, the unit cell for CaCl 2 would have one Ca 2+ for every two Cl - Close Packing of Spheres o In many cases the particles that make up the solids are spherical or approximately so o The most efficient arrangement of a layer of equal-sized spheres is a sphere in the center, surrounded by six others in the layer Second layer of spheres can be placed in the depressions on top of the first layer If the third layer are placed in line with those of the first layer, the structure is known as hexagonal close packing If the third layer are placed so they do not sit above the spheres in the first layer, the structure is known as cubic close packing o In both of the close-packed structures, each sphere has 12 equidistant nearest neighbors We say that each sphere has a coordination number of 12 Coordination number is the number of particles immediately surrounding a particle in the crystal structure o When unequal-sized spheres are packed in a lattice, the larger particles sometimes assume one of the close-packed arrangements, with smaller particles occupying the holes between the large spheres 11.8: Bonding in Solids
The physical properties of crystalline solids, such as melting point and hardness, depend both on the arrangements or particles and on the attractive forces between them Molecular Solids o Consist of molecules held together by intermolecular forces o Because these forces are weak, molecular solids are soft o Have relatively low melting point o Properties of molecular solids depend on the strengths of the forces that exist between molecules and on the abilities of the molecules to pack efficiently in three dimensions E.g. Benzene Has higher melting point that toluene, a substitute of benzene because of the lower symmetry of toluene The boiling point of toluene is higher than that of benzene, indicating that the intermolecular attractions are larger in liquid toluene than in liquid benzene Covalent-Network Solids o Consists of atoms held together in a large networks or chains by covalent bonds o Much harder and have higher melting points than molecular solids o E.g. Diamond Each carbon atom is bonded tetrahedrally to four other carbon atoms The array of strong carbon-carbon single bonds that are sp 3 hybridized contributes to diamond s unusual hardness o E.g. Graphite Carbon are bonded in trigonal planar geometries, forming interconnected hexagonal rings Have delocalized π bonds extending over the layers Electrons are free to move through the delocalized orbitals, making graphite a good conductor of electricity along the layers Ionic Solids o Consist of ions held together by ionic bonds o The strength depends greatly on the charges of the ions o The structure adopted by an ionic solid depends largely on the charges and relative sizes of the ions o E.g. In NaCl, Na + ions have a coordination number of 6 while in CsCl, each Cs + ions have a coordination number of 8
Increase in the coordination number as the alkali metal ion changes from Na + to Cs + is a consequence of the larger size of Cs + compared to Na + Metallic Solids o Consist entirely of metal atoms o Usually have hexagonal close-packed, cubic close-packed, or body-centered cubic structures Thus, each atom typically is surrounded by eight or 12 adjacent atoms o Can be visualized as metal as an array of positive ions immersed in a sea of delocalized valence electrons This type of bonding is called metallic bonding o In general, the strength of the bonding increases as the number of electrons available for bonding increases o The mobility of the electrons explains why metals are good conductors of heat and electricity