Homework Problems Chapter 11 Homework Problems: 4, 14, 25, 34, 50, 52, 56, 59, 68, 74, 76, 78, 83, 84, 87, 104, 106, 108, 118 Skip sections 11-10, 11-11, and

CHAPTER 11 Liquids, Solids, and Intermolecular Forces

Solids, Liquids, Gases The three common states of matter are solids, liquids, and gases. Their properties can be summarized as follows stateproperties soliddefinite volume and definite shape liquiddefinite volume but indefinite shape gasindefinite volume and indefinite shape

Compressibility Compressibility refers to the decrease in the volume occupied by a substance when the pressure applied on the substance increases. Because the particles making up solids and liquids are in close contact with one another, these two phases are not easily compressed (to a first approximation solids and liquids are incompressible). Since gases are mostly empty space, they are highly compressible. This also means that solids and liquids will be much higher density than gases.

Effect of Temperature and Pressure on Phase The phase of a substance depends on both temperature and pressure. Generally speaking, substances go from solid to liquid to gas as temperature increases. It is often possible to condense a gas into a liquid or solid by increasing pressure. We will discuss this further later in the chapter.

Intermolecular Forces Intermolecular forces are the forces that exists between different molecules or particles. We are more concerned with long range attractive forces and will ignore short range repulsive forces. Ion-ion - The attractive force acting between cations and anions. These are strong, and are found in substances where ionic bonding occurs..

Dipole-Dipole Forces Dipole-dipole - The attractive force acting between polar molecules. The attraction is between the partial positive charge (  +) on one molecule and the partial negative charge (  -) on a different molecule. Generally speaking, the larger the partial positive and negative charges the stronger the dipole-dipole attraction.

Dipole-Dipole Forces and Boiling Point When molecules have strong intermolecular attractive forces it takes more energy to overcome those attractive forces. One way of seeing this is in the boiling point for a substance. Generally speaking, the stronger the dipole-dipole attraction between molecules the higher the boiling point, parti- cularly for substances with approximately the same molecular mass.

Dipole Moment and Miscibility Miscibility refers to the ability of one liquid to mix with another without forming two separate liquid phases. As a general rule (for reasons discussed in Chapter 12) polar liquids will mix well with polar liquids, and nonpolar liquids with nonpolar liquids. A polar and a nonpolar liquid will generally not mix to form a homogeneous solution.

Hydrogen Bonding Hydrogen bonding - A particularly strong form of dipole-dipole attractive force. It is the attractive force that exists between a hydrogen atom bonded to an N, O, or F atom and lone pair electrons on a different N, O, or F atom.

Evidence For Hydrogen Bonding One effect of hydrogen bonding is to raise the boiling point of a liquid. This occurs because it requires more energy (and so a higher temperature) to break apart strong attractive forces between molecules than it does to break apart weak attractive forces between molecules. substance boiling point hydrogen bonding? H 2 O100.0  Cyes H 2 S  C no H 2 Se  C no H 2 Te- 4.4  C no

London Dispersion Forces London dispersion forces - The attractive force that is due to the formation of instantaneous dipoles in a molecule. These instantaneous dipoles arise from the random motion of the electrons in the molecule.

Strength of London Dispersion Forces London dispersion forces are present in all molecules, but are the only intermolecular force present in nonpolar molecules. The strength of London dispersion forces is approximately proportional to the number of electrons, and so to the size of the molecule. The larger the molecule the stronger the London dispersion forces. substanceboiling point He  C Ne  C Ar  C Kr  C Xe  C

Ion-Dipole Forces Ion-dipole - The attractive force between an ion and a polar molecule. Responsible for the dissolution of some ionic substances in polar liquids such as water. Solvation - The close asso- ciation of solvent molecules with solute molecules or ions. Hydration - Solvation when the solvent is water.

To summarize the types of forces – Ion – ion. Forces between cations and anions. Dipole – dipole. Forces between molecules with a permanent dipole moment. This category includes hydrogen bonding, a particularly strong type of dipole – dipole force. London dispersion forces. Due to random movement of electrons. All particles have this type of force, but it is most important in molecules with no permanent dipole moment, Mixed forces: Ion – dipole is the most important. It is responsible for the solubility of some ionic compounds in water.

Properties of Liquids There are several general properties of liquids. Viscosity - Resistance of a liquid to flow. Surface tension - Resistance of a liquid to spreading out. Generally speaking, the stronger the intermolecular forces, the larger the values for viscosity and surface tension.

Phase Transitions The conversion of a substance from one phase to another phase is called a phase transition. Transitions can be caused both by adding heat and by removing heat from a substance. adding heat (  H > 0)removing heat (  H < 0) s  fusion (melting)  s freezing  g vaporizationg  condensation s  g sublimationg  s deposition Recall that the enthalpy change for a phase transition is usually reported at the normal transition temperature, that is, the temperature at which the phase transition occurs when p = 1.00 atm. Since enthalpy is a state function:  H freez = -  H fus  H cond = -  H vap  H dep = -  H sub

Note  H sub   H fus +  H vap Enthalpy and Entropy Changes For Phase Transitions

Thermodynamics of Phase Transitions We can study the thermodynamics of phase transitions by finding the heating curve for a substance. This is simply a plot of temperature vs. amount of heat added, under conditions where the heat is added slowly enough to maintain equilibrium. Experimentally we ex- pect to see two regions in the heating curve. Normally the temperature of the substance will increase as heat is added. However, at the temperature where a phase transition occurs the added heat will be used to carry out the transi- tion, and so temperature will remain constant until the phase transition is complete.

Sample Heating Curve

Vapor Pressure The vapor pressure (p vap ) of a liquid or solid is equilibrium partial pressure of the substance in the gas phase above the liquid or solid. Experimentally, it is found that the vapor pressure of a liquid increases as temperature increases. This is not surprising, as the molecules in the liquid phase have a higher average energy as temperature increases.

Clausius-Clapeyron Equation The dependence of vapor pressure on temperature is well described by the Clausius-Clapeyron equation (which may be derived). ln(p vap ) = -  H vap + constant RT This may be rearranged to give ln(p 2 /p 1 )= -  H vap 1 1 R T 2 T 1 p 1 is the vapor pressure at T 1 p 2 is the vapor pressure at T 2 slope = -  H vap /R The Clausius-Clapeyron equation assumes that the value for  H vap is constant. In fact, the value decreases slightly with increasing temperature.

Example: The normal boiling point for water occurs at T =  C. The enthalpy of vaporization for water is  H vap = kJ/mol. Based on this information estimate the vapor pressure of water at T = 20.0  C.

Recall that one form of the Clausius-Clapyron equation is: ln(p 2 /p 1 )= -  H vap 1 _ 1 R T 2 T 1

Example: The normal boiling point for water occurs at T =  C. The enthalpy of vaporization for water is  H vap = kJ/mol. Based on this information estimate the vapor pressure of water at T = 20.0  C. Let T 1 =  C = 373. K ; p 1 = 1.00 atm T 2 = 20.0  C = 293. K ln(p 2 /p 1 ) = J/mol 1 1 = (8.314 J/mol. K) 293. K 373. K So (p 2 /p 1 ) = e = p 2 = (0.0279) p 1 = (0.0279) (1.00 atm) = atm = 21. torr

Phase Diagram A phase diagram is a diagram indicating which phase or phases are present at equilibrium as a function of pressure and temperature. H2OH2O

Important Features in a Phase Diagram Phase boundaries - Indicate where two phases can exist simul- taneously at equilibrium. Triple point - Indicates a point where three phases can exist simultaneously at equilibrium. Normal melting point - Solid-liquid equilibrium at p = 1.00 atm. Normal boiling point - Liquid-gas equilibrium at p = 1.00 atm. Normal sublimation point - Solid-gas equilibrium at p = 1.00 atm. (Note that substances will have either a normal melting and normal boiling point, or a normal sublimation point, but not both.) Critical point - Point below which a gas will undergo a phase transition (g  or g  s) when compressed reversibly at constant temperature. Above the critical point no such phase transition occurs. In this region of the phase diagram a supercritical fluid is present.

Phase Diagram For CO 2 CO 2

Solids Solids can be divided into two general categories Crystalline solid - Has a regular arrangement of the particles making up the solid (a crystal structure). Four main types exist: ionic, molecular, covalent, and metallic solids. Amorphous solid - Does not have a regular arrangement of the parti- cles making up the solid (no regular crystal struct- ure).

Classification of Crystalline Solids Crystalline solids can be classified into three groups. 1) Ionic solids – Composed of cations and anions; held together by ionic bonding 2) Molecular solids – Composed of molecules; held together by weak intermolecular forces (dipole-dipole, hydrogen bonding, or dispersion) 3) Atomic solids – Composed of atoms a) Nonbonding – Composed of noble gas atoms. Held together by dispersion forces, b) Metallic – Composed of metal atoms, held together by metallic bonding. c) Network covalent – Composed of atoms where every atom is attached to other atoms in the solid through a network of covalent bonds.

Ionic solids An ionic solid is composed of cations and anions. Ionic solids are held together by the strong electrostatic force of attraction that exists between particles of opposite charge. Examples: NaCl, CaCO 3, Al 2 N 3, FeCl 3. Properties Hard and brittle High melting point High boiling point Poor conductors of heat and electricity in solid state Good conductors of electri- city when dissolved in water NaCl, T  fus = 801  C, T  vap = 1413  C

Molecular Solids A molecular solid is composed of molecules. Molecular solids are held together by the weak van der Waals attractive forces (dipole- dipole and London dispersion forces) that exist between molecules. Examples: H 2 O, Ar, CS 2, C 10 H 8 (naphthalene), C 6 H 12 O 6 (sugar). Properties Soft Low melting point Low boiling point Poor conductors of heat and electricity in solid state Poor conductors of electri- city when dissolved in water CS 2, T  fus =  C, T  vap = 46  C

Atomic Solids An atomic solid is a solid substance composed of individual atoms (as opposed to ions or molecules). There are three types of atomic solids Nonbonding atomic solid – An atomic solid where the atoms are held together by dispersion forces. The only nonbonding atomic solids are the solid forms of the noble gases (He, Ne, Ar, …) Metallic atomic solid – The solid form of a metal. Consists of individual metal atoms held together by a “sea” of electrons in metallic bonding. Network covalent atomic solid – A solid where every atom is connected to every other atom by a network of covalent bonds. In a sense a network covalent atomic solid is a single “supermolecule” diamond

Nonbonding Atomic Solid The only nonbonding atomic solids are the solid forms of the noble gases. Since these substances do not form covalent bonds with other noble gase atoms the forces holding nonbonding atomic solids together are weak dispersion forces. Properties Extremely low melting point Extremely low boiling point Poor conductors of heat and electricity in solid state Kr, T  fus =  C, T  vap =  C

Metallic Atomic Solids A metallic atomic solid represents the solid form for metals. Metallic solids can be thought of as metal cations immersed in a sea of loosely held valence electrons. Examples: K, Fe, Cu, Na, Pb. Properties Can be hard or soft Low to high melting point Low to high boiling point Good conductors of heat and electricity in solid state Insoluble in water Na, T  fus = 98  C, T  vap = 892  C

Network Covalent Atomic Solids A network covalent atomic solid is a “supermolecule” in which every atom is connected to every other atom through a network of covalent bonds. Because all of the atoms are connected by covalent bonds, these substances are generally among the hardest substances known. Exam-ples: C (diamond), C(graphite), SiO 2 (quartz). Properties Extremely hard Very high melting point Very high boiling point Poor conductors of heat and electricity in solid state Insoluble in water C, T  fus > 3550  C, T  vap = 4827  C diamond

End of Chapter 11 “Gibbs is perhaps the most brilliant person most people have never heard of. Modest to the point of near-invisibility, he passed virtually the whole of his life, apart from three years spent studying in Europe, within a three-block area bounded by his house and the Yale campus in New Haven, Connecticut. For his first ten years at Yale he didn't even bother to draw a salary. (He had independent means.) From 1871, when he joined the university as a professor, to his death in 1903, his courses attracted an average of slightly over one student a semester.” - Bill Bryson A Short History of Nearly Everything “Of all chemical bonds, hydrogen bonds are the weakest, the most important, the least understood, and the hardest to measure.” - John Emsley, “Science Watch” (2000)